Chemistry 317
Spring Quarter 2012
Inorganic Chemistry Laboratory
A. General Information
Introduction to Chemistry 317
Logistics and Schedule
Class Schedule
Assignments
Course Grading
Attendance/Punctuality Policy
Readings
Lab Notebook
Laboratory Safety
Guidelines for Writing Lab Reports
Techniques
B. Experiments
1. Chromous Acetate
2. The Chelate Effect
3. Phosphorous Acid
4. The Lewis Acid-Base Adduct BF3•NH3
5. Arene Molybdenum Tricarbonyl Chemistry
6. Linkage Isomers of Nitro-Pentammine-Cobalt(III)
7. Preparation of a Doped Phosphor, ZnS
p. 2
2
3
3
4
4
5
5
6
11
15
29
37
47
57
64
72
85
Instructor: Prof. Michael Heinekey
CHB 304 F, 543-7522, [email protected]
Lab Technician: Tom Leach, BAG 133D, 685-9466, [email protected]
Discussion Section/Lab Lecture: (first meeting March 27, 2012)
AA & AB sections: Tuesdays 8:30-9:20 AM in Bagley 108
BA & BB sections: Tuesdays 9:30-10:20 AM in Bagley 108
Laboratory Sections and Teaching Assistants:
AA & BA sections: TTh 12:30-4:20, Bagley 191
Amanda Weaver ([email protected])
Vlad Vlaskin ([email protected] )
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AB & BB sections: WF 12:30-4:20, Bagley 191
Eric Camp ([email protected])
Sophia Tran ([email protected])
Introduction to Chemistry 317
Welcome to Chemistry 317, Inorganic Chemistry Laboratory. Our goals are to make this
a stimulating, challenging and useful experience. You will be introduced to new techniques and
new kinds of chemicals and chemical reactivity. The class will tie together material you have
had in lecture courses, and will ask you to design and improve experiments. There are some lab
periods for which no instructions are provided; you must choose what you want to do and invent
a procedure to do it. We will help you, but we want you to bring your insights, enthusiasms,
questions, and skills to the course. Some of the material will be familiar, while other parts of the
class will be quite new. Some of the experiments work like a charm, others we are still
perfecting - and we hope you will help us make them better. We are eager to hear your
suggestions and comments.
Logistics and Schedule
Chemistry 317 consists of two laboratory periods and one “discussion” hour per week.
The experiments are designed for students to work in pairs, with a maximum of 20 students (10
pairs) in the laboratory at one time. In each lab period, half of the students will work on one
experiment and the other half will work on another. Those in the A section will be doing
experiments from the first column on the schedule (next page), and must come to the Tuesday
8:30 a.m. discussion hour. Those in the B section will work on the second column and must
come Tuesday at 9:30 a.m. In this way the discussions will be related to the experiments you are
doing.
The discussion hours will include some lecturing, to provide background and
understanding of the experiments. But primarily these hours will be forums for discussion of the
lab just completed, for instance how to analyze your spectra or numerical data. Please bring
your data, your questions, and your opinions! Most students have found these sessions quite
helpful. Attendance will be taken and is required.
The schedule for the class is given on the next page.
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Chem 317 Class Schedule:
Lab
period
Section AA, AB
1
2
3
4
5
6
7
8
9
10
11
12
13
14
15
16
17
18
19
20
Check in
Cr Acetate
Cr Acetate
Chelate Effect
Chelate Effect
Phosphorous Acid
Phosphorous Acid
BF3•NH3
(Arene)Mo(CO)3
Chelate Effect II
Linkage Isomers
Linkage Isomers
Linkage Isomers
Linkage Isomers
(Arene)Mo(CO)3 II
(Arene)Mo(CO)3 II
(Arene)Mo(CO)3 II
Linkage Isomers II
ZnS phosphor
Check out
Due*†
MSDS
Chromous(ST)
Chelate(AW)
Phos Acid(AW)
BF3•NH3(ST)
Arene (EC)
Chelate II(AW )
Linkage(VV )
Arene II(EC )
Linkage II(VV )
ZnS(ST )
Section BA, BB
Due*†
Date
Check in
Mar 27/28
Chelate Effect
MSDS
Mar 29/30
Chelate Effect
Apr 3/4
Chromous Acetate
Apr 5/6
Chromous Acetate
Chelate(AW)
Apr 10/11
ZnS phosphor
Apr 12/13
Chromous(ST )
Apr 17/18
(Arene)Mo(CO)3
Chelate Effect II
ZnS(ST)
Apr 19/20
Linkage Isomers
Arene (EC)
Apr 24/25
Linkage Isomers
Chelate II(AW)
Apr 26/27
Linkage Isomers
May 1/2
Linkage Isomers
May 3/4
(Arene)Mo(CO)3 II
May 8/9
(Arene)Mo(CO)3 II Linkage(VV )
May 10/11
(Arene)Mo(CO)3 II
May 15/16
Linkage Isomers II
May 17/18
Phosphorous Acid
Arene II( EC )
May22/23
Phosphorous Acid Linkage II( VV ) May 24/25
May29/30
BF3•NH3
Check out
Phos Acid(AW ) May31/June 1
June 5/6
BF3•NH3(ST )
* - For AA and BA sections, assignments listed as due Tuesdays can be turned in by 12:30 pm
on the following day (Wednesday). This will allow students to use the information presented at
the Tuesday morning discussion hour.
Warning: A lot of work piles up at the end of the quarter in this class.
Even with only five pairs of students working on a given experiment, there will
occasionally be times when you will have to wait to use a piece of equipment. Try to find
something else that needs to be done while you’re waiting. Making efficient use of your time is
a critical laboratory skill (and a skill you will be graded on). Within each experiment, you and
your lab partner will often be doing different things. You should try to follow what she or he is
doing, as the final lab write-up will require data from both of you.
Assignments
Each lab write-up in the lab manual ends with a description of the required assignment
for the lab. All assignments are due one week after completion of the experiment, as shown
on the schedule. Assignments must be typed, double spaced. Tables, graphs and equations
should be rendered digitally. For molecular structures, ChemDraw is suggested. UW has a site
license for ChemDraw, see http://depts.washington.edu/chem/facilserv/computing/software.html
A program called ChemSketch is available as freeware (http://www.acdlabs.com/download/).
You and your lab partner will have the same data, but you each need to write individual reports.
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This includes all text portions of the reports, such as the introduction and data analysis. Copies
of all spectra acquired should be turned in with the assignment. (One of you will have original
copies to turn in, and the other will have to make copies to turn in separately with his/her own
lab report.)
Two of the experiments require formal lab reports, as explained in their write-ups. The
most critical aspects of any lab report are clear thinking and maintaining your focus on the
important issues. Grading will be based not only on the science but also on the quality of the
writing. Mundane details like grammar and punctuation, when ignored, make it difficult for the
reader (grader) to appreciate the content of a report.
In addition to the lab write-ups, there will be a brief (50 min) written evaluation. This is
scheduled for Thursday, May 31, 2012 at 8:30 or 9:30 a.m. (your choice).. This is not a
scheduled course meeting time, so please reserve that date and time on your schedule now. If
you have a schedule conflict, contact Professor Heinekey as soon as possible.
Your grade will be based on the formal lab reports for two of the experiments, the shorter
assignments for the other labs, the laboratory evaluation, and on your overall ability in the
laboratory (as judged by your TAs). The point distribution is outlined below. The score for lab
skills given by the TAs is quite important. They will be looking at how prepared and punctual
you are, how well you use your lab time, your lab safety, lab awareness, and overall helpfulness
(especially to your lab partner), and the quality of your ideas, suggestions, and questions.
Course Grading
Lab Reports
Arene Molybdenum Tricarbonyl
Preliminary (5%) Proposal (5%) Final (15%)
Linkage Isomers
Other Experiments
Chelate Effect
Chromous Acetate
Phosphorous Acid
BF3•NH3
ZnS phosphor
25%
25%
9%
4%
4%
4%
4%
Lab Skills (TA input)
10%
Written Evaluation
(Thursday, May 31 at 8:30 or 9:30 a.m.)
15%
Attendance/Punctuality Policy
• You must be on time to each lab session. If you are more than ten minutes late, you will be
considered absent.
• If you have to miss a laboratory period due to illness or other circumstances, there is a clearly
defined procedure for an excusable absence. You will need to see Dr. Tracy Harvey to initiate
this process. Unexcused absences will cause an automatic 50% reduction in your score for that
experiment.
• Lab reports must be turned in according to the schedule. Work not turned in by the appropriate
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date and time will be assigned a grade of zero.
Readings
It is imperative that you carefully read the lab descriptions, and appropriate “techniques”
(eg, 1H NMR, gas handling, etc) section of the lab manual before entering the lab. Even more
important than reading, you must carefully think through what you will be doing. This is critical
for safety in the lab, and to manage your time efficiently. The TAs may take various steps to
insure that the reading is carefully done. You should also look up the Material Safety Data
Sheets (MSDS) of all compounds that you will be using and bring it with you to lab. There are
hazards to working with many of the chemicals involved in these experiments.
The lab descriptions contain occasional references to the “original literature,” the
scientific articles that originally reported the chemistry. These are given in the standard
American Chemical Society (ACS) reference format: authors, Journal Title year, volume #,
page. Literature can be searched on-line using the American Chemical Society (ACS)
publications website (http://pubs.acs.org/), the “Web of Science” (accessed via UW website:
http://www.lib.washington.edu/types/databases/). A powerful search engine is provided by
SciFinder(formerly Chemical Abstracts) (https://scifinder.cas.org/scifinder/) Relevant articles are
available in a folder in the 317 lab, along with various reference books. We will also try to have
books on reserve. These extra readings are optional, but may be quite useful and interesting. In
the arene-molybdenum-tricarbonyl experiment, for instance, many students have found these
papers very useful when they design their own procedures. The extra readings provide
background to help you understand your observations, and better interpret your data (both are
critical to good lab reports).
Lab Notebook
You must keep a good notebook in this laboratory (and in all scientific labs). Use a
bound book that pages cannot be removed from. Your notebook is your diary of what you did,
and it should be written as you are working. Do not make notes on scratch paper and transcribe
them into your notebook. The book should include numerical data (weights, volumes, voltages,
etc.), procedures (A was added to B dropwise over 20 minutes using an addition funnel), and
observations (it turned green after half the A was added). The most important features of a good
lab notebook are clarity and completeness. You should never remove a page or plan to go back
and fill in something later. If necessary, you can cross something out or recopy something for
clarity, just indicate why and make sure the original is still legible. The notebook is not a handy
piece of scratch paper. It should enable you to reconstruct what you did, including good and bad
aspects of the procedures. A TA will look at your lab book periodically and may collect it at
some point.
LABORATORY SAFETY
In this laboratory, as in any laboratory, there are a number of hazards. Learning how to
deal with hazardous situations safely is an important part of what you will learn in this class. If
chemistry majors cannot handle hazardous situations involving chemicals, then who in society
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can? It’s usually chemistry majors who write the rules for safe handling of chemicals. Safety is
an important focus of this class and we want you to think about safety as you read this lab
manual and, especially, as you work in the lab. Chemistry 317 has endured some scary incidents
in the past. Fortunately, no one was hurt and we will talk more in our discussion sections about
what went wrong and how these incidents could have been prevented.
The most important safety rule is to THINK! Good common sense will get you through
most situations. If there is anything that is unfamiliar or doesn’t seem right, stop what you are
doing and ask. There is at least one TA for every ten students in the lab, so there should always
be someone nearby to assist or explain. You will be using a number of expensive pieces of
equipment during the lab, so it is important that you understand how this equipment works.
Don’t just plow ahead if anything looks wrong. No one will be criticized for asking. It is,
however, critical that you arrive prepared for the laboratory, having worked out the procedures in
your own mind so you know what you’re going to do.
There are a few safety rules we will strictly enforce. Safety goggles and a lab coat must
be worn at all times in the lab. Eating in the lab is forbidden. Shoes must be worn at all times
(no sandals or open toed shoes), and no shorts or short skirts.
An increasingly important part of safety and safe handling of chemicals is their disposal.
The disposal of the solutions and products in each lab experiment is either described in the lab
write-up or your TA will explain the procedures. Note that there are waste bottles for each
experiment, as well as for acids and bases. Only chemicals go in these bottles. Syringe needles
go in the sharps bin, and glass products (including glass pipettes) go in the glass waste. Place all
waste in the appropriate container. If you aren’t sure where it goes, ask your TA. Leave all
“smelly” items in the hood with appropriate labels. When in doubt, put your waste in a bottle
and label it to indicate the contents. No potentially hazardous waste should be disposed of down
the drain or allowed to evaporate into the fume hood.
What follows is a detailed description of the safety rules for this class.
Safety goggles must be worn in the laboratory at all times. Any failure by a student to
observe this state health may result in removal of that student from the laboratory. Eyes are too
valuable to risk. Students will not be allowed to work in the laboratory without approved
standard laboratory goggles. Standard laboratory goggles that meet all state regulations may be
purchased from the bookstore or the stockroom. Safety glasses, etc. are not acceptable. If you
already have goggles, the stockroom personnel must first approve them before you can begin
working. Because of health regulations, goggles cannot be borrowed from the stockroom.
Dress appropriately for the laboratory. In order to work with chemicals safely, bare skin
should be kept to a minimum. This means socks, full shoes (no open toes), and long pants. In
addition, it is now a requirement to wear a lab coat in the lab at all times. Lab coats can be
purchased at Health Science Stores. FYI: cotton clothing (including denim) is particularly
susceptible to being eaten by acid solutions. The laboratory is not a good place to wear your
favorite clothes. Do not wear clothing so loose or bulky that it hampers your work (this is a good
way to break expensive glassware). Long hair should be tied back. Failure to dress
appropriately will result in a loss of available lab points, and you will be sent out of lab to
acquire the correct clothing. If you do not return in time to complete your work, the absence will
be unexcused.
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Do not eat, drink, or smoke in the laboratory. Do not even bring these materials into the
laboratory.
Wash hands often when working in lab, and always thoroughly before leaving. Do not taste
any chemicals. Do not put your hands, pens, or pencils in your mouth while working in the lab.
Keep coats, backpacks and other non-essential materials away from areas where people are
working. Lockers are available in the hallway.
Hall lockers: You may bring a lock from home and claim an empty hall locker for use during
the quarter. These are ideal for storing coats, backpacks, and other bulky items. Also, a locker is
the best place to store your goggles. Lockers must be emptied by the end of the quarter between quarters the locks will be cut off and the locker contents thrown away.
Cell phones and headphones may not be used in the lab.
Never attempt any unauthorized or unassigned experiments. Follow the experimental
procedures explicitly, checking and double-checking the identity of all reagents before you use
them. There are potentially hazardous combinations of chemicals present in the laboratory. If
you have an idea for further investigation, discuss it with your instructor.
Learn the location and operation of the safety showers, emergency eyewashes and fire
extinguishers in the laboratory. In the case of spills onto a person or clothing, the immediate
action should be water and lots of it. Do not hesitate to yell for help. Use the safety showers
and/or eyewashes and don't worry about the resulting mess. Don't use the safety showers for
non-emergencies since they are designed to deliver about 50 gallons of water before shutting off.
Report accidents to your instructor. All instructors have been certified to administer first aid. If
you are not familiar with the operation of the fire extinguishers ask your instructor to explain it
to you. The fire extinguishers should only be used for real emergencies since the chemicals they
contain can cause considerable damage. In any emergency that requires the assistance of the fire
department, aid car or police, send someone to the stockroom for assistance.
Become familiar with all of the exits from the laboratory. A repeating siren and flashing of
the FIRE indicator is the building evacuation signal. If this alarm goes off while you are in the
lab, turn off any open flames, grab your valuables, and leave the building as quickly as possible.
Don't sit in the hallways. Sitting in the hallways disrupts the natural flow of foot traffic and can
constitute a safety hazard. If you need to work on your lab report outside of the laboratory room,
please use the Chemistry Study Center (BAG 330).
WORKING WITH EQUIPMENT AND GLASSWARE
Do not leave a Bunsen burner or other heated apparatus unattended. The person working
next to you may not know what is involved with your setup and may be working with a
flammable material. Turn off open flames if you must leave your area. Make sure the gas taps
are completely off whenever the Bunsen burner is not lit.
Do not pick up hot objects with your bare hands. Be sure all apparatus is cool before picking
it up with your fingers.
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Chipped glassware and glass apparatus from your drawer may be traded for undamaged
items at the stockroom. We can fire-polish chipped glassware so it is usable, but we can’t fix
cut hands.
Do not adjust glass tubing connected to rubber stoppers. Severe cuts or puncture wounds
may result.
Lubricate rubber tubing. When slipping rubber tubing over connectors, such as filter flasks or
aspirators, lubricate with a drop of glycerin (hood) or liquid soap (in lab).
Do not force pipette bulbs onto pipettes. Apply just enough pressure to maintain a seal
between the pipette and the pipette bulb. Forcing the bulbs may cause the pipette to slip and
break, leading to severe cuts or puncture wounds.
WORKING WITH CHEMICALS
Reagents: Read the label (contents and hazards) before using reagents. Take only as much
reagent as you need - they are expensive and time consuming to prepare. When taking reagents,
transfer the amount you need to a clean beaker or other suitable container for taking the material
back to your desk. Replace the cap.
Never return unused reagents to their storage containers. If you take more than you need,
dispose of the excess in the appropriate manner.
Clean up spills immediately. The next person to come along has no way of knowing if the
clear liquid or white powder on the lab bench is innocuous or hazardous. Neutralize acid spills
with sodium bicarbonate (baking soda) before cleaning them up.
Hazard Identification. As part of the UW Laboratory Safety Manual, each laboratory has a
Chemical Hygiene Plan (CHP). This is available to all students in the lab at all times. As part of
the CHP, Material Safety Data Sheets (MSDS) must be readily accessible to all students. MSDS
are available through the campus computer network on the Lab Safety System (LSS). The
computers in the lab have a link to LSS.
Material Safety Data Sheets (MSDS) are provided by the manufacturer or vendor of a chemical.
They contain information about physical properties of the chemical and identify any hazards
associated with the chemical. They also identify any special handling precautions and protective
equipment needed when working with the chemical. You should be familiar with the MSDS
before working with any chemical.
Read chemical labels carefully. Chemicals are rated from 0 to 4 according to the hazard they
impose; with 0 representing no hazard and 4 representing high hazard. An example of a hazard
diamond label is shown below. Each chemical is rated for health, fire and reactivity. Special
warnings are reserved for the 4th diamond.
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Hexa ne
3
Fir e
2
0
Rea ct iv it y
Health
Flam ma b le
WASTE DISPOSAL
Dispose of chemical reagents and other materials properly. The proper disposal of chemical
wastes is essential to the health and safety of University faculty, staff, students and the
surrounding community. Chemical wastes must be managed and discarded in the most
responsible and environmentally sound method available. The University and Metro expect your
cooperation in taking care of the environment. Your laboratory manual will specify how to
dispose of chemicals used during the laboratory period. Do not put chemicals into glass boxes or
wastebaskets. Only specified non-hazardous water-soluble materials can be rinsed down the
drain. Waste containers for other materials will be provided. If you are unsure of how to dispose
of a particular material, ask your instructor.
1.
ALL NON-CHEMICAL SOLID WASTES used in this class go into the trash cans unless
otherwise noted. Paper towels, matches, pH paper, etc. should NOT be placed in the sinks.
2.
GLASS (whether broken or not) goes into the glass disposal box. Cleaning up broken
glass is greatly facilitated by using the broom and dustpan (located on the counter under the
windows). Custodial personnel will stop collecting trash after they find broken glass in the
trashcans!
3.
BROKEN THERMOMETERS should be taken to the UG stockroom (BAG 271) for
disposal after cleaning up any spilled mercury.
4.
ALL ACIDS AND BASES used in this class must be neutralized. Metro requires that
any solutions going down the drain be between pH 5.5 and 12. Acids may be easily
neutralized by the careful addition of sodium bicarbonate; bases may be neutralized by the
addition of dilute HCl and the final pH raised with sodium bicarbonate if necessary. The
final pH should be checked with pH paper before it is washed down the drain. Like other
solid wastes, pH paper should be placed in the trash.
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Suggested Procedure for Cleaning up Chemical Spills:
Solid Reagents: Wipe up small spills with a damp paper towel; rinse the reagent out of the
towel with water, then dispose of the towel in the trash cans. Clean up large spills using the
broom and dustpan (located on the counter under the windows) and dispose of the reagent in an
appropriate waste container. If glass is present in the spill, separate the glass from the reagent
before disposal. Do NOT place solid chemicals in either the trash cans or the glass box. Spills
in the weighing chamber of the balances should be immediately brushed out using the camel’s
hair brush provided; the reagent may then be disposed as above.
Liquid Reagents (Non-organics of near-neutral pH): Wipe up the spill using a damp paper
towel; rinse the reagent out of the towel with water, then dispose of the towel in the trash cans.
Acids (including phosphorus trichloride): Neutralize the acid by sprinkling solid sodium
bicarbonate over the area of the spill. Clean up the bicarbonate residue with either a damp towel
or the broom and dustpan, depending upon the amount used to neutralize the acid. Flush the
bicarbonate down the drain with an excess of water.
Organic liquids: Wipe up the liquid with paper towels. Do not rinse the paper towels or place
them in the trash. Instead, place them in a hood. Allow the liquid to evaporate and then dispose
of the paper towels in the trash cans.
Mercury: Obtain a “mercury sponge” from the stockroom. Moisten the sponge with water and
then rub it over the area of the spill (metal side down). The mercury should quickly become
amalgamated with the metal. When finished, place the sponge back into the plastic bag and
return it to the stockroom for disposal. During a mercury spill, small droplets may spatter a
surprising distance from the area of the spill, especially if the mercury falls from the bench to the
floor. Be sure to check a wide area around the spill to be sure that all the mercury has been
located and notify others in the lab to avoid the spill area. If you have a large spill, a special
mercury vacuum may be necessary; ask for assistance.
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Guidelines for Writing Lab Reports
1. Report Structure
Introduction:
The introduction of a lab report typically starts with an explanation of what you are trying
to accomplish or observe in the experiment. Include a brief discussion of why this is interesting.
You should introduce the compounds to be studied and/or the techniques to be used and their
interpretation. In the Linkage Isomers lab for instance, you might discuss the long history of the
cobalt complexes or the concept of linkage isomerism with reference to the nitro ligand. In the
Arene Molybdenum lab it would be appropriate to discuss arenes and CO as ligands in contrast
to typical (historical) inorganic coordination chemistry, and how 1H NMR and IR spectroscopies
can help characterize these compounds. This section should include references to help establish
the context of your work as a small piece in the general body of chemical knowledge.
Results:
This section should tell a story, describing what you did and what results you obtained. It
should describe the experiment as you performed it, not the one described in the manual.
Balanced chemical equations are strongly encouraged. Diagrams of the proposed and actual (if
different) structures of the complexes should be shown with labels on the atoms or groups of
interest. Data tables summarizing your characterizations should be included here, along with an
explanation of the data and corresponding assignments. For instance, in the Arene Molybdenum
lab you will record IR spectra for your starting material and products. Assign the peaks you see
in the IR spectra of the various compounds and how these spectra relate to each other.
When possible, numerical data should be presented in Tables rather than as part of the text.
Then the text can refer to the Table, as in “Voltage readings were obtained over the temperature
range 280 - 340 K and the data are given in Table 1.” All tables and graphs should be labeled as
well as referred to directly within the text. Any problems with data collection should be
explained. Whenever possible include the raw data, and show how the quantities of interest were
derived, giving any relevant equations. Always show sample calculations and use units. Error
analysis should be included in all calculations.
The Results section is a factual account only, interpretation should be left for the Discussion
section.
The grammatical tense in the Results section is generally past tense, but is not strictly adhered to.
Discussion:
The Discussion section should first of all provide your analysis of the results. Did you
make what you wanted? How do you know? If the expected product was not formed, discuss
why this might be the case and whether the actual product can be identified. Discuss what went
wrong, if anything did. In the Arene-Molybdenum lab, you’ll want to analyze your spectral data,
for example what the CO stretching frequencies indicate about what happened. Compare the
spectral data for the products to those of the starting materials and make comparisons. The
Linkage Isomers lab requires significant data and error analysis which go into the Results and
must be discussed in the Discussion section.
The Discussion section must include a clear description of your conclusions. You could
conclude that the error bars are too big for you to conclude anything or that the procedure didn’t
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work because .... But it’s important that you take a stand − no waffling. The Discussion should
return to the questions, goals, and issues raised in your Introduction. In a way, the Introduction
and Discussion are the bookends for the Results.
Experimental:
This section is a completely factual account (in past tense ONLY) of the experiments you
preformed and a formalized report of their characterization. The format in which this section
should be produced is strict and unwavering. This section is written in “cook book” manner, so
that future chemist could pick up your report and repeat your experiments exactly using only the
descriptions you provided. It should describe your actual procedure, not the one in the manual.
The first section should include general procedural methods although not standard procedures
(assume that we all know how to use a balance, glove box, IR spectrometer, etc.) Subsequently,
it should include descriptions of the reagents, the amounts of reagents used (grams, mL, eq., and
moles), your yields of the products obtained (both in grams and percent) for each product
obtained. A list containing all spectroscopic features and assignments (and specifics of how they
were obtained) and other methods of characterization should conclude the individual compound
sections. In general scientific writing, compounds already reported in the literature are referenced
only, however, you will include all synthesized compounds in your experimental.
Proper formatting is important in this section more than any other. There is an immense
amount of information which needs to be conveyed in a small amount of space, improper
formatting will only add to the confusion which often prevails in this section. The best way to
understand how to format this section properly is to look at an example from a published journal
displaying similar information.
References: This section should include a numerical tabulation of all the previously published
scientific literature which you cited in the above text. It should be in normal ACS (American
Chemical Society) format, authors, journal, year, volume, pages. For example…
1.
Kawamura, K.; Hartwig, J. F., J. Am. Chem. Soc., 2001, 123, 8422-8423.
2. Error Analysis:
It is important to understand and describe the uncertainties (the errors) in any numerical
values that you report. Errors are typically divided into two types: systematic errors and random
errors (and of course there are simple mistakes, like typing a number incorrectly into a
calculator). Random errors result from the fact that no measurement is perfect, for instance that
you can’t measure the temperature on a typical mercury thermometer without an uncertainty of,
maybe, ±0.2 ˚C. You should estimate your random errors when you make a measurement. For
example, if you measure an absorbance with a UV-vis spectrometer, measure the same sample
twice or three times to get a sense of how reproducible the values are. Systematic errors result
from problems that are not random, such as miscalibration of an instrument (e.g., your balance
doesn’t give 100.00 g for the standard 100 g weight). Perhaps your starting compound is only
95% pure, so your calculated concentrations are systematically 5% too high – or perhaps your
procedure leads to an inadvertent dilution of your samples. Systematic errors are more difficult
to identify and analyze but you should be on the lookout for them at all times. Error analysis is
discussed in more detail in the write-up for the Chelate Effect lab.
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3. Presentation and Style
The ability to write a clear and concise description of what you did and what it means is
an important skill, regardless of what job you find yourself in. This is a formal report, a single
document with a logical lay-out and flow. The arguments must flow within each paragraph and
from one paragraph to the next. Proper grammar, spelling and formatting must be used.
Scientific writing has its own peculiar styles. The experimental section is typically
written in the past passive voice. For instance: “The 0.1 M ammonia solution was added to the
solution of CuSO4 dropwise, with stirring. The solution rapidly turned dark purple....” You
should not imply that chemicals are active agents, avoiding, for example: “the ammonia turned
the copper ions purple.” Use the past tense when you are describing what you did at some time
in the past. In other places, try to use present tense: “Understanding the chelate effect is
important because ...” or “A plot of the ΔG values free energies versus temperature is shown in
Figure Z. ΔS is obtained from the slope of this plot .... The ΔS values imply that ...” Use
present tense because what you are describing is independent of time. You imply that anyone
looking at your data would derive the same ΔS and reach the same conclusion. For examples of
scientific writing, refer to the journal articles on reserve in the Chemistry Library or those
available in the laboratory.
You should avoid using personal pronouns, such as “I added ammonia to ....” The
statement “Addition of ammonia to a solution of Cu2+ causes a rapid change in color to dark
purple” is true whether you do it or someone else does it; it is true today and will be true in the
future. You can indicate that you’re not sure of something with phrases such as “it seems likely
that” or “it is possible that” or “perhaps.”
These are guidelines, not firm rules. “The voltages are/were converted to free energies (in
kJ/mol) using equation X (see Table Y).” could be present or past tense. The most important
features are clarity of thought and writing. Maintain your focus on the important issues and lead
the reader through your story and your arguments. Ask yourself “what am I really trying to say
here?”
4. Plagiarism
Reports which are plagiarized (in whole or in part, from published material or from other
students) will be automatically assigned a grade of zero, and will be referred to the Vice Provost
for Student Life in accordance with the University of Washington Student Conduct Code. See
http://depts.washington.edu/pswrite/plag.html
Below are some excerpts from the University of Washington web site on this issue.
1.
Using another writer's words without proper citation. If you use another writer's words,
you must place quotation marks around the quoted material and include a footnote or other
indication of the source of the quotation.
2. Using another writer's ideas without proper citation. When you use another author's
ideas, you must indicate with footnotes or other means where this information can be
found. Your instructors want to know which ideas and judgments are yours and which you
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3.
4.
5.
6.
arrived at by consulting other sources. Even if you arrived at the same judgment on your
own, you need to acknowledge that the writer you consulted also came up with the idea.
Citing your source but reproducing the exact words of a printed source without quotation
marks. This makes it appear that you have paraphrased rather than borrowed the author's
exact words.
Borrowing the structure of another author's phrases or sentences without crediting the
author from whom it came. This kind of plagiarism usually occurs out of laziness: it is
easier to replicate another writer's style than to think about what you have read and then
put it in your own words. The following example is from A Writer's Reference by Diana
Hacker (New York, 1989, p. 171).
o
Original: If the existence of a signing ape was unsettling for linguists, it was also
startling news for animal behaviorists.
o
Unacceptable borrowing of words: An ape who knew sign language unsettled
linguists and startled animal behaviorists.
o
Unacceptable borrowing of sentence structure: If the presence of a signlanguage-using chimp was disturbing for scientists studying language, it was also
surprising to scientists studying animal behavior.
o
Acceptable paraphrase: When they learned of an ape's ability to use sign
language, both linguists and animal behaviorists were taken by surprise.
Borrowing all or part of another student's paper or using someone else's outline to write
your own paper.
Using a paper writing "service" or having a friend write the paper for you. Regardless of
whether you pay a stranger or have a friend do it, it is a breach of academic honesty to
hand in work that is not your own or to use parts of another student's paper.
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Techniques
Many inorganic compounds are sensitive to oxygen, moisture, or both. Such compounds − you
will make several of these during this course − need to be protected from the ambient
atmosphere. You will become acquainted with a number of techniques to protect materials from
the atmosphere. In some ways, learning these techniques is a focal point of this class. By far the
best reference to the handling of air-sensitive compounds is the book The Manipulation of AirSensitive Compounds (2nd Ed.) by D. F. Shriver and M. A. Drezdzon, Wiley-Interscience, New
York (earlier editions with the same title are pretty similar).
The two primary ways you will handle air-sensitive materials in this class are with a
glove box or with a Schlenk line, as described below. The handling of gases is described later.
I. The Glove Box
Conceptually, the simplest way to keep things away from the oxygen and water in the
atmosphere is to work in a lab space where there is no oxygen or water. This is a little difficult,
however, since we humans need both to survive. But one could imagine a fully enclosed “bench
top,” containing an “inert atmosphere,” which one could reach into with gloves. Such a device is
called a “glove box” or a “dry box” if it has hard sides like a box. There are also cheap “glove
bags” that are what they sound like, bags you can fill with inert gas and reach into with attached
gloves.
The glove box that you will
use in this class is a fairly
sophisticated one, made by the
Vacuum Atmospheres Corp.
A
picture is shown at right, with the
gloves missing (they go on the big
circles in the middle). Despite the
name, this box does not have a
vacuum in it (the sides and gloves
would implode with the pressure).
Rather the glove box is kept full of
clean nitrogen gas. Clean in this
context means chemically clean,
containing as little oxygen and water
as possible.
The dry box has four important parts: (i) There is a large aluminum box, with a plastic
front window sprouting two gloves. This is the working area. Note that this is not a glass front
(too fragile), so use only water if you need to clean it - no organic solvents. (ii) There is an
antechamber (like a submarine or spaceship airlock) which is how things get in and out without
letting in air. This is the cylinder at right in the drawing above. (iii) The gas in the box is
constantly circulated over a scrubber (often called the “catalyst”) which removes any air or water
that has made its way into the enclosure. Since the catalyst is damaged by many kinds of
reactive chemicals (chlorinated solvents, sulfur compounds, etc.), we must be careful what we
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allow to evaporate into the box atmosphere. In the picture above, the catalyst is located in a
canister outside the left wall of the box. A fan inside the box circulates the box atmosphere
through the canister. (iv) Finally, the box must be able to regulate the pressure inside. If the
pressure gets too high the gloves will pop out or the front will break − and if the pressure gets too
low the gloves will suck in. Both are catastrophes. The device that regulates the pressure is
called the photohelic and it is located above the antechamber. You can recognize it because it
has a pressure gauge on it, calibrated in the unusual unit “inches of water” (more on pressure
units later). The box can take only a few inches of water positive or negative pressure, very little
change from one atmosphere. The photohelic automatically draws fresh nitrogen from a tank if
the pressure gets too low and automatically pumps nitrogen out if the pressure gets too high.
You can also do these things manually with the foot pedal, labeled R and L, for raise and lower
[pressure]. The instructor or a TA will demonstrate the operation of the box and the photohelic
in the lab. If the nitrogen cylinder that feeds the box runs out, tell a TA right away.
The drybox is an extremely useful piece of equipment but one that must be treated
carefully − one accident and you can contaminate the atmosphere and destroy everyone’s
chemicals. Some researchers in inorganic chemistry use a glove box as if it were a bench top,
and do all their work in there. However, this can be tedious, uncomfortable, and inconvenient. It
is difficult to maintain a good atmosphere in the box when you’re working with all sorts of
different solvents and reagents, it is hard to work with gases, and it is hard to do reactions that
require heating or cooling (remember, no water for reflux condensers). These boxes are also
expensive, costing well upwards of $20,000 each.
For these reasons, we only have two dry boxes in the lab, and they are used mostly for
storage, setup, and workup of experiments. You will do most of your chemistry out in the air,
using specialized glassware. In the aggregate, these strategies are called “Schlenk techniques.”
II. Schlenk Techniques
The centerpiece of this defense against atmospheric intrusion is the double manifold, or
Schlenk line. In the illustration below, the double manifold is in the middle, connected to
various equipment on both sides. The two long, horizontal tubes of the Schlenk line are called
“manifolds.” Each manifold can be filled with a gas, or “un-filled” with vacuum. Reactions and
manipulations are typically
done under an atmosphere
of an inert gas, usually
nitrogen, admitted into one
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of the manifolds (the “nitrogen side”). The other manifold (the “vacuum side”) is
connected to a vacuum pump through a liquid-nitrogen cooled trap. Any solvents that make it
into the vacuum line condense in the trap before they get to the pump, which protects the pump
and the pump oil. Glassware (see below) is connected to the Schlenk line via rubber hoses.
Then the glassware is exposed to the vacuum or the nitrogen manifold using the two-way
stopcocks. For safety reasons, the Schlenk lines will be set up in the hoods.
Be sure all stopcocks are lightly greased − if you forget this the glass pieces may stick
together and it can be a real pain getting them apart. The function of the thin film of grease is
just to allow the glass pieces to slip over each other. If you use too little grease the glass pieces
will bind. If you use too much grease, it will ooze into the holes and into your line and will
degrade more rapidly. When you’re done, the ground-glass part of the stopcock should look
clear, with no streaks in the grease. Your TA will help you get the hang of this.
**NOTE 1** We have replaced the mercury manometer, at left in the Figure above, with 2
pressure gauges: 1. A “Bourdon” diaphragm needle gauge and 2. a digital convection gauge.
These gauges attached to the vacuum manifold are your way of seeing what the pressure is in the
manifold. The pressure is measured in units of torr. Recall that one atmosphere is 760 torr. The
Bourdon gauge has a range of 1-760 torr and is not sensitive to the type of gas in the line. The
convection gauge has a range of 0.01-760 torr and is sensitive to the type of gas in the line and so
it is not as accurate. It is used to measure the low pressure range. These gauges are useful for
quantitative gas handling, as gases can be conveniently measured out by PV = nRT. It is also
critically useful to tell you qualitatively what is going on inside your Schlenk line. Whenever
you do anything on the Schlenk line, the pressure inside will change. Be sure you know what you
expect to happen whenever you open or close a valve on your line − and watch the gauge to be
sure that this is what happens. If anything surprising occurs, STOP, close the valve, and think
again. Note: go slowly and wait for the gauge to respond. These digital gauges are slower to
react than the mercury manometers.
** NOTE 2** We have replaced the nitrogen tank (see Figure above) with a link to the house
nitrogen supply located in the hoods.
The Nujol bubbler at the end of the nitrogen line is there to prevent air from getting in.
Be careful, if you pull a vacuum on the Nujol bubbler, you will suck Nujol back into your line
and make a big mess. Many groups do this at least once and then have to clean their line (not
fun). If you need to pull a vacuum on the nitrogen line, use a pinch clamp on the hose or squeeze
it with your fingers to prevent the Nujol from sucking back. You should also place a check valve
on the nitrogen line which will prevent the backflow of Nujol into the Schlenk line in case you
make a mistake. When you’re using the line, there should be a slow and steady flow of nitrogen
through the nitrogen manifold and out the bubbler. The stream of nitrogen bubbles tells you
what’s going on in this manifold.
You have an array of equipment to help you manipulate solutions without exposing them
to air. The lead players are pieces of glassware with sidearms attached, so-called Schlenk ware,
as illustrated below.
The Schlenk flask is an ordinary round-bottom flask with a sidearm with a stopcock (be
sure it’s greased!). You can connect this sidearm to the Schlenk line with thick rubber tubing and
use it to admit nitrogen to the flask or to evacuate it. The tubing needs to be thick so that it won’t
collapse under vacuum. You will put something in the neck of the flask, such as a glass stopper
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(greased) or another piece of apparatus such as a Schlenk addition funnel or a Schlenk filter.
Flasks
Addition funnel
Schlenk filter
The addition funnel is a convenient means for adding solutions to a Schlenk flask. With
the lower stopcock closed it is completely sealed off from the flask below it, so you can put
solutions in the funnel without contaminating the atmosphere in the bottom flask. To move a
solution from the funnel to the flask, first connect both to the nitrogen manifold via the sidearm
stopcocks. Then just open the valve at the bottom of the funnel. Connecting everything to
nitrogen is critical because this equalizes the pressure between the top and bottom; without it,
pressure would build up in the lower flask as the liquid flows down, and the flow would stop. In
all air sensitive work, you always have to worry about pressures − what the pressure is in each
piece of apparatus and how this will affect what you’re trying to do.
Another common way to cap a Schlenk flask is with a septum, an air-tight rubber
membrane. By piercing the septum with a variety of sharp instruments, you can add or remove
liquids from flasks without exposing them to the air. And septa (plural of septum) seal back up
(pretty well) when you remove a needle. Since septa will not usually give you a truly air-tight
seal, they should be swapped for a glass stopper if you need a flask to stay air-free for more than
an hour or so.
A syringe, which you’re probably familiar with, is one way of transferring liquids into or
out of a septum-capped container. If the container is under a pressure of nitrogen, you can
simply insert the needle into it, pull back the plunger to remove the desired amount of liquid,
then withdraw the needle and inject the liquid into another flask. In the time it takes to move
from one flask to another not much air can diffuse into the small bore of the needle. There are
two important hints about using syringes. First, the “dead space” in the needle and syringe body
will ordinarily have air in it. If this would be a problem, you should flush it out before using it.
Suck nitrogen into the barrel from a nitrogen-filled container (usually a flask attached to the
nitrogen line), then expel the nitrogen out into the air. Repeat this two or three times to purge air
from the syringe. Second, remember to always think about pressures when transferring via
syringe. If you try to suck a liquid out of a sealed container, you won’t get very far, since you’ll
build up a partial vacuum in the container. You can avoid this by making sure the flask from
which you are withdrawing is connected to a source of nitrogen, like your nitrogen manifold, or
by injecting a volume of nitrogen gas to compensate for the volume of liquid that you are going
to remove. Similar tricks can alleviate difficulties in injecting into closed flasks.
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A cannula is a hollow steel
needle with two sharp ends. It can
serve as a sort of express route for
transferring liquids when set up as
shown. If the pressure in the flask at
right is greater than that in the other
flask, the liquid will be pushed from the
right to the left flask.
To address the difficult conundrum of no-air filtrations, some twisted soul invented the
Schlenk filter (illustrated below). Its effective use requires some practice and a flair for
contortionism. The filter is placed on top of the flask with the material to be filtered, and on top
of it is placed a flask in which to catch the filtrate. The whole assembly is then inverted, and you
try to get as much of the solid as possible to run down on to the fritted glass disk. You can help
the solid down with the stir bar, which you can move around with a hand-held magnet on the
outside of the flask. Applying a touch of vacuum to the underside of the frit while the top is
under nitrogen will suck the filtrate through just like an ordinary suction filtration. It’s
considered tacky to pour the solution down one of the sidearms, so make sure you tip the setup
so the liquid runs down the other side (as shown). We encourage using more than two hands for
this operation − i.e., do it with your lab partner.
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III. Gas Handling: Tanks and Regulators
The handling of gases provides a number of challenges. First, all gases are “air
sensitive,” to the extent that if they get mixed with air they are no longer pure. (There are also
gases that undergo chemical change because of reaction with oxygen or water.) You will use the
vacuum manifold of your Schlenk line to work with gases.
Gases are purchased in thick (usually steel) tubes, either large cylinders or smaller
“lecture bottles.” The gas inside is present under pressure, typically pressures much greater than
one atmosphere. This represents the second challenge: a pressurized gas would love to get out
of its container and will do so with some force if allowed to. This is a serious hazard, and one
that must always be kept in mind while working with compressed gases. In the mid-90s, there
was a serious incident when someone opened a lecture bottle to their Schlenk line when the
stopcock on the line was closed. The high pressure gas was released to the rubber hose but had
no place to go. So what did it do? The gas blew open the thick hose, with sufficient force to
snap a piece of glass off the line. A fair amount of noxious gas was sprayed into the lab and
everyone had to be evacuated. Let’s not have a problem this quarter.
We use a series of valves and regulators to make sure that the flow of gases is controlled
(by us). The cylinder or lecture bottle that comes from the vendor has a valve on top, essentially
an on/off valve (there is a little bit of flow adjustment possible with this valve, but very little).
Lecture bottles are metal tubes, roughly 2” in diameter by 15” in length. Gas cylinders come in
all shapes and sizes, up to 5’ high and 300 pounds − and from there, gases can be purchased in
truck-loads or by the railroad tank car. The primary difference between a cylinder and a lecture
bottle (aside from the size) is that we the consumer buy the lecture bottle while the cylinder is the
property of the company and we pay demurrage, a “rent” of about 3¢ per day.
The valve on the tank is your first line of defense. If this is closed, no gas can get out.
But when that is open, the gas will come out at a high rate (flow): our nitrogen tanks come
pressurized to 2,000 psi (pounds per square inch) or ca. 150 atmospheres. If anything goes
wrong while you’re working with a gas, close off this main valve. If the student had done
this (in the mid-90s) when the hose blew, the incident would have been much less serious.
Note: you will be expected to know − perhaps on the laboratory evaluation − all the
various pressure units used in this class:
1 atm = 1.01 bar = 14.7 psi = 760 torr = 760 mm Hg = ~33 ft or ~400 in of water.
In this lab, you may use large cylinders of
nitrogen as your source of inert gas. To control the flow
of gas from a large tank, always use a regulator (such as
the one pictured at right). One other precaution: these
tanks should always be chained to a wall or a lab bench so
that they cannot tip over. If one tipped over and fell on
the regulator, it could snap off the valve on top of the
tank. This would turn your demure gas cylinder into a
steel torpedo, propelled by the high pressure gas inside.
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The regulators we use for large cylinders typically have two valves and two gauges⎯and
it is important that you know the function of all four things. The gauge closest to the cylinder
tells you the pressure in the cylinder; it will be calibrated from 0-3,000 psi or thereabouts (it will
also be calibrated in other units, too, just to be confusing). The valve that is in between the two
gauges is a diaphragm valve, which regulates pressure (not flow). This amazing device enables
you to set the pressure coming out of the regulator, which you can read on the second gauge, the
one farther way from the tank. Typically we use low pressures, 0-5 psi (above atmospheric).
Finally, the small valve leading out of the regulator is a needle valve that regulates flow, not
pressure. This valve is your last line of defense − open this valve slowly to bleed gas into the
line, then if you need more gas open it wider. The needle valve works as you would expect:
screwing it down will close the valve. The diaphragm valve, however, is opposite: it is closed
when screwed all the way out, and most open as you screw it in. Be sure that you understand
this difference.
Lecture bottles are typically used for gaseous reagents, such as the BF3 and
NH3 that you’ll use in Experiment 4. Lecture bottles contain much less gas and
typically have lower pressures than large cylinders. We typically use just a needle
valve (shown at right) on a lecture bottle to regulate the flow of gas out. Then you
cannot regulate the pressure coming out − only the flow. The procedure you should
use is to open the needle valve for an instant to let some gas out, then quickly close
it. This enables you to dispense just the amount you need into your Schlenk line.
Most importantly, be sure to look at the vacuum gauge when you are dispensing gas to be
sure the gas is going where you think it is. When you open the needle valve for an instant and
then close it, the vacuum gauge reading should rise. If it doesn’t do this, shut your gas off, think
about where the gas that you just let out of the lecture bottle went, and get your TA.
Whenever you first use a lecture bottle, you should assume that the hose and needle valve
are full of air, unless you know otherwise from the previous user. To get rid of the air, connect
the hose to your Schlenk line, make sure the main valve is closed, open the needle valve, and
pump out the hose and the needle valve. The main valve on a lecture bottle can look like a knob,
which is closed when screwed down. Or the top of a lecture bottle can look like two nuts, which
is closed when the nuts are tightened together. Your TA will show how to use two wrenches to
do this. To remove air from a regulator on an inert gas such as nitrogen, purge the regulator and
hose by running a good flow of gas through them. This avoids having to pump on the diaphragm
valve, which is not good for it.
You must be very careful not to let pressure build up in any piece of apparatus. Schlenk
techniques can tolerate pressures only slightly greater one atmosphere. If you have a pressure of
two atmospheres in a flask (external pressure plus one atm more) that’s 14.7 pounds on every
square inch of your apparatus. So a stopper with a one square inch opening will have 14.7
pounds pushing it open. This is equivalent to hanging a bowling ball off of it! Be
sure⎯whenever you work with gases⎯that you know what will happen anytime you open a
valve. Be sure you know where the gas is supposed to be going and where the gas will go if the
pressure by accident gets too high. Let’s have a safe and fun lab!
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IV. Spectroscopy
Spectroscopy deals with the interaction of light with matter. The energy of an atom or
molecule is quantized, that is only a limited number of energy levels are allowed.
Spectroscopies examine the gaps in energy between these energy levels, which can be quite
informative as to the nature and structure of the molecule. In Chem 317 you will use a number
of different kinds of spectroscopies, most if not all of which you have used in other classes.
Because of the variety of compounds and situations in this lab, it will be important for you to
understand what these spectroscopies are actually measuring and what problems could arise. An
overview is given here; addition information can be found in many of the individual
experiments.
A key issue to always keep in mind in almost all spectroscopic measurements is that your
sample is in a vessel and often in a solvent. The spectrometer measures – so your spectrum
shows – the entire sample, including the vessel and the solvent. For instance, the NaCl salt
plates that we use for IR spectra are not transparent over the whole IR region. Plastic cuvettes
absorb UV light. NMR solvents often have residual protons that will show up in a 1H NMR
spectrum (see table below). For instance chloroform–d is not 100% CDCl3; it probably is about
99% CDCl3 and 1% CHCl3. It is always a good idea to take a spectrum of just your solvent, or
just your cuvette, or just your salt plate. Then you’ll know what peaks correspond to your
sample and what to the other materials.
Solvent
Chemical Shift (ppm)
acetone-d6
2.05
benzene-d6
7.16
chloroform-d
7.26
deuterium oxide (D2O)
4.75 (varies)
A. NMR Spectroscopy
In this lab we will be mainly dealing with two types of NMR – phosphorus and proton.
Proton or 1H NMR is the most common type of NMR spectroscopy, which you may have studied
in organic chemistry. Recall that 1H nuclei have a spin of ½. In a magnetic field, these nuclei
have two energy levels, ms = +½ and ms = -½. Proton NMR looks at the energy of the transition
between these two levels. Phosphorus, on the other hand, consists exclusively of a single
isotope, phosphorus-31 (31P), which also has a nuclear spin of ½. So phosphorus or 31P NMR
spectra have chemical shifts and show coupling, just like proton NMR.
In the Phosphorous Acid lab, coupling is observed between two phosphorus nuclei and
between a phosphorus nucleus and a proton. This P-H coupling can be observed either in the 31P
or the 1H NMR. In that lab, we’ll also be looking at deuterium NMR (2H NMR). Similar to with
a proton, coupling can be observed between a phosphorus nucleus and a deuterium nucleus.
However, deuterium has a quantum spin of 1, so its nuclei can have 3 different orientations in a
magnetic field: +1, 0, or -1. This will cause the phosphorus coupling to look different than when
it’s split by a proton. When you get the deuterium (2H) spectrum from this experiment, it will
most likely look very similar to the proton NMR spectrum. The introduction to the Phosphorous
Acid lab contains more detailed instructions on how to look at and interpret these spectra.
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NMR Methods
To make an NMR sample, place about 10 mg of your sample in an NMR tube. Add
enough NMR solvent (use deuterated solvents unless you’re directed to do otherwise) to reach
about 1 ½ inches (“three fingers”) up from the bottom of your tube. Shake the tube until your
sample is dissolved. Note that if your sample does not dissolve, it will not show up in your
NMR spectrum. If necessary, try a new tube with a different solvent to see if you get better
dissolution. When your NMR tube is ready, label it with your name and the NMR solvent used
and give it to your TA.
B. Infrared (IR) Spectroscopy
Absorption of light in the infrared region of the spectrum typically corresponds to the
vibrations of the molecules in the sample, both stretching and bending vibrations. You probably
took a number of IR spectra in organic labs. IR spectra are usually reported in wavenumbers
(cm-1), the inverse of wavelength. Wavenumbers ( ν ) are a unit of energy, and are directly
proportional to frequency (ν): ν = ν × c where c is the speed of light (3 × 1010 cm s-1).
IR spectra are useful in the characterization of molecules because each bond type
typically has a characteristic range of frequencies (depending on the molecule). In organic
chemistry, there are a limited number of functional groups and the frequencies for each are
tabulated in standard texts, for instance Pavia et al. Introduction to Organic Laboratory
Techniques (Appendix 3) or Lambert et al. Organic Structural Spectroscopy (Chapters 8 and 9).
Inorganic compounds exhibit a much wider variety than organic ones – we deal with over 100
elements! – so one doesn’t find such tables. The best source is Nakamoto, Infrared and Raman
Spectra of Inorganic and Coordination Compounds, a copy of which should be found in the lab.
IR Methods
There are several different procedures for preparing a sample for IR spectroscopy. For
each compound you are asked to take an IR of, take a minute to think of which method might be
best for this particular compound.
I. Nujol Mull
A Nujol mull is probably the simplest method and is recommended if your product has
turned out to be too oily (but not wet). This method requires a mortar and pestle, Nujol (light
mineral oil), and two salt plates. Place a small amount of your solid (perhaps 20 mg) in the
mortar. Add a tiny amount of Nujol, not more than a couple drops. Use the pestle to grind the
Nujol and sample together into a thick and homogenous paste. Scoop up enough of this mull to
coat the center portion of a salt plate and add the other salt plate to make a sandwich. Don’t
squeeze the plates together too strongly or you’ll push the mull out (or worse, crack the salt
plates). These really are “salt” plates, made of NaCl, so you need to be gentle with them. They
crack very easily and dissolve quickly in water (they also cost more than $10 each). To clean the
salt plates, they can be wiped with a Kimwipe™, or rinsed with methylene chloride (CH2Cl2).
Never use water to clean the salt plates. Be sure to take a spectrum of Nujol by itself later to
account for these peaks in your mull spectra.
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II. KBr pellets
A KBr pellet can provide a much cleaner spectrum, free from pesky Nujol peaks, but may
require a little more effort. You may need to make several pellets to get a feel for the procedure
and get a decent IR spectrum. To make a KBr pellet, you will need a mortar and pestle
(provided in the lab), some dried KBr (which should be in the oven), and a KBr pellet press
(provided in the lab). To begin, place ~1-2 mg of your sample in the mortar. Add 100 times as
much KBr as your sample (i.e., 200 mg KBr for 2 mg of sample). Grind this mixture into a fine
powder, until no individual grains of KBr are visible. Scrape 100 to 150 mg of this powder into
the KBr pellet press. Consult your TA or the nearby directions if you are unfamiliar with the
setup or how to use the press. The pellet press can then be placed directly into the beam of the
IR spectrometer after removing the screws.
III. Solution cells
Another method involves the use of solution cells (cavity cells). This option is best if
you already have a product which will dissolve in an organic solvent (i.e., your Arene
Molybdenum Tricarbonyl II products). A solution cell is usually a block of NaCl with a cavity
of fixed path length created by a thin Teflon spacer. As with the salt plates, care must be taken
to keep the cells free of moisture (avoid solvents that would contain water or are aqueous – these
cells cost over $100 to repair). A dilute solution (1-2 mg/mL) of your solid in organic solvent is
put into the cell. Be sure to also take a solution cell IR spectrum of just your solvent. Empty the
cell, rinse with methylene chloride (CH2Cl2) and air-dry after using it. Place the cell back in the
desiccator when you are finished with it.
IV. Liquids
An alternate method to get spectra of solutions or liquids is to simply put a drop on a salt
plate and cover with another salt plate. If the solvent is volatile, like methylene chloride
(CH2Cl2), it may evaporate quickly and you’ll just get a spectrum of your sample. Again, don’t
use water solutions or highly polar solvents that could dissolve the salt plates (you don’t want to
pay for new salt plates). When you’re done, wipe the salt plates clean with a KimwipeTM and
rinse with CH2Cl2.
Using the FTIR Spectrophotometer (Perkin-Elmer 1600 series and similar)
First, notice that directly below the screen there is a row of grey keys which are not marked.
These keys are called “softkeys” and the operations or values they perform will be displayed
along the bottom of the screen itself (this will seem familiar to those of you who have used
graphing calculators).
1.
A background scan is not required (it doesn’t change much over time), but can be
performed if desired. Make sure the sample holder inside is empty. Press the gray
Backg key followed by the green Scan key, and then the key under the number “4”
displayed along the bottom of the screen. In 10-15 seconds, the machine should display
“Ready” instead of “Scanning” to indicate it has finished taking the background.
2.
Place your sample in the instrument and then press the Scan key followed by the key
under the number “4” displayed along the bottom of the screen. When scanning is
complete, your spectrum will appear on the screen.
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3.
At this point, several peaks should be visible, and these peaks need to be “marked” so
the wavenumber value will print out with your spectrum. First, press Shift and then the
Peaks key. A vertical line should appear on the screen which you can move from peak
to peak using the Peak and Vertical Cursor keys. To mark a peak (so the number will
print out), choose the Shift and then Mark keys. A small vertical dash should appear
below the peak, which indicates this peak has been marked. Repeat marking peaks
until all major peaks are marked.
4.
To print out your spectrum, press the green Plot key. Print out a copy for your lab
partner if you have time and there is not a line waiting to use the instrument. Record
the sample name, time/date taken, and method of preparation (i.e., KBr pellet) on your
printout.
5.
If the spectra becomes off-center on the screen while you are marking, first press Shift
and then Peaks to get rid of the cursor. Now you can use the Vertical Cursor and
Peaks keys to move the spectrum left or right. The Mark and Rerange keys expand or
shrink the spectrum along the x-axis (wavenumbers).
C. UV/Vis Spectroscopy
UV/Vis spectroscopy records how your sample absorbs light in both the visible and UV
regions (hence UV/Vis). This is a particularly useful technique than in inorganic chemistry
because many inorganic chemicals are colored. UV/Vis spectra are also known as electronic
spectra or optical spectra. For a review of how absorbance is measured and the caveats of Beer’s
Law, see the introduction to Experiment VI: Linkage Isomers. For more details on how to make
a UV/Vis sample and how specifically this method will be used, see the Linkage Isomers lab.
Using the UV/Vis Spectrophotometer (Ocean Optics USB2000)
The Ocean Optics UV/VIS Spectrophotometer (labeled USB2000) is connected by a fiber
optic cable to a cuvette sample holder which is connected by a fiber optic cable to a light source
(either model MINI-DT with a shutter or model MINI-D2T without a shutter). The light source
employs a combination of deuterium and tungsten-halogen lamps to give good signal coverage
over the wavelength range of 200 to 800 nm. The USB2000 is connected to a Dell laptop
computer via a USB cable. If you are using an Ocean Optics spectrophotometer which has a
cuvette holder/light source directly mounted on the spectrophotometer (and labeled USB-ISSUV/VIS), the instructions for setting up the instrument for the dark and reference spectra are
somewhat different from what follows, and appropriate instructions from the Laboratory
Technician or TA should be obtained.
1.
The light source for the instrument is powered by a 12V AC adapter. Be sure it is
plugged in and that the ON/OFF/REMOTE toggle switch is in the “ON” position. The
toggle switch is located next to where the power supply plugs into the light source.
2.
Start the program called OOIBase32 by double-clicking the Icon on the desktop of
the computer labeled "Ocean Optics". The program will start up with all the
parameters from the previous session except that the wavelength range will revert back
to the default settings (180-870nm). All the spectrophotometer functions needed for
this experiment are controlled by this program.
3.
At the top of the window are some toolbars. The four we will use are shown in
Figures 1 – 4.
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Figure 1. File Toolbar. Left to right, the six buttons are New Spectrum, Open, Save, Copy
Spectra, Print, and Help. (The ones in bold face are referred to later.)
Figure 2. Spectrum Controls Toolbar. Left to right, the six buttons are: Store Dark Spectrum,
Store Reference Spectrum, Snapshot, Single Exposure, Data Acquisition, Emergency Reset.
(The ones in bold face are referred to later.)
Figure 3. Part of the Spectral View Toolbar. The five buttons shown are Scope Mode,
Absorbance Mode, Transmission Mode, Irradiance Mode, and Specular Reflection Mode (the
ones in bold face are referred to later). Until both a dark and reference spectrum have been
stored, only Scope Mode can be used. Therefore, until both a dark and reference spectrum has
been stored, all the buttons except the S are grayed out.
Figure 4. Graph Scale Toolbar. Left to right the three buttons are Autoscale, Set Scale,
Unscale. (The ones in bold face are referred to later.)
4.
The instrument will start in Scope Mode which shows a graph of light intensity vs.
wavelength (Figure 5). You should see a number of peaks on the graph that are from
the deuterium/tungsten light source, including a major peak at 657nm. If all you see is
a flat line or a broad peak in 600-800 nm range, the lamp is not working properly.
Note: in the screen shots that follow, the display is of black lines and text on a white
background. The display you see may be of white lines and text on a black
background.
Figure 5. Ocean Optics software in Scope mode.
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5.
In Scope mode, the S icon (Figure 3) should be depressed. Press the S icon (or select
"Scope Mode" from the Spectrum drop down menu) whenever you need to return to
Scope mode (you should be in Scope mode when storing either a dark or reference
spectrum).
6.
Change the wavelength range to appropriate values for your experiment. To do this,
click on the Set Scale button (Figure 4, icon which has 2 perpendicular red arrows) and
enter the appropriate x-axis values. (no need to change the y-axis values). To help
select a wavelength range, you may want to place an empty cuvette into the cuvette
holder and observe the light intensity as a function of wavelength. Plastic cuvettes
absorb some UV light so you want to choose a wavelength range where there is
sufficient light reaching the detector. Otherwise, you will have noisy regions in your
absorbance spectrum.
7.
Place a reference cuvette filled 3/4 full with the reference solution (usually deionized
water unless otherwise specified) into the cuvette holder. Be sure that the cuvette is
aligned properly so that the light beam will pass through the clear sides of the cuvette.
Cover the cuvette with the cuvette cover and make sure that the cover is flush with the
base of the cuvette holder (light passes through the sample when either side of the
cover is flush with the base).
8.
If necessary, adjust the Integration Time until the highest peak in the light source
trace is between 3000-3500 Intensity units (usually a value between 8-20 ms is normal).
The Boxcar smoothing parameter should be set to 5 (this averages the 5 points before
and 5 points after each data point). The Average parameter should be set anywhere
between 10 and 40. More averaged scans give a better looking spectrum but slightly
slower response time.
9.
Calibrate the instrument by recording a reference and a dark spectrum (note that the
instrument should be allowed to warm up for at least 5 minutes before calibrating). To
record a reference spectrum, make sure the reference cuvette is in the cuvette holder.
Click on the yellow light bulb icon (Figure 2) or right-click on the screen and select
‘Store Reference’ from the menu that appears.
10.
To record a dark spectrum, slide the cuvette cover to the side so that it blocks all the
light getting to the detector, which you can see visually on the screen in Scope mode as
a flat baseline. For the light sources that have a shutter switch on the front (MINI-DT),
you could also move this shutter switch to the ‘OFF’ position. Click on the darkened
light bulb icon (Figure 2) or right-click on the screen and select ‘Store Dark’ from the
menu that appears.
11.
Select the Absorbance mode by touching the ‘A’ icon (Figure 3) or select
‘Absorbance Mode’ from the Spectrum menu.
12.
Place your sample cuvette into cuvette holder. Select ‘Autoscale’ by touching the
icon which looks like an up & down arrow (Figure 4) or right-clicking on the screen
and selecting ‘Autoscale’ from the menu that appears. The absorbance spectrum of
your sample should now show full screen.
13.
To determine an absorbance measurement, click on the screen at the desired
wavelength. A green cursor will be displayed on the screen. You can fine-tune the
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selected wavelength by using the green arrow icons or using the left and right arrow
keys on the keyboard. Three numbers appear in the legend at the top of the screen (if
displayed). The first number is the wavelength in nm, the second is the channel
number (not really useful), and the 3rd number is the absorbance (or the intensity if you
are in Scope Mode).
14.
To save a spectrum, choose File, Save, Processed from the pull down menus at the
top of the screen. Navigate either to the floppy disk drive, USB drive, or an appropriate
folder on the hard drive using the ‘Save In’ pull down box, give the spectrum a name,
and click ‘Save.’ The program will save the file as text and attach a suffix of
‘.Master.Absorbance’ to the end of your chosen filename. Note: Don’t use the other
save options (spectra, reference, etc.). If you don’t save your spectrum as a
processed file, you will not have any useable data!
15.
If you want to freeze the data acquisition before saving the file (particularly useful if
you are going to remove the sample before saving), click on the Snapshot option
(Camera icon in Figure 2). You will need to unclick this option to continue taking
measurements of other samples.
16.
To overlay spectra on top of one another on the screen, click on the Overlay pulldown menu and select an overlay channel. You will be prompted to choose a file
containing a previously saved spectrum to be imported into that channel. There are a
total of 8 overlay channels available so you may not be able to overlay all of your saved
spectra. To clear an overlay, simply click again on the channel number from the
Overlay menu.
17.
To print spectra, select File, Print Setup and choose ‘landscape’ orientation in the
bottom right corner (this will give a printout which fills more of the paper) and then
select File, Print. Landscape orientation remains the default until the program is exited.
18.
To manipulate a saved spectrum (outside of lab as lab computers do not have Excel),
open the file from Excel. Don't just double-click on it; the computer probably won't
know you want to open it in Excel since the filename doesn't end in .xls. Make sure
that in the ‘Files of Type’ box, you have chosen ‘All Files (*.*)’. Excel will recognize
the file as text and bring up a Text Import Wizard dialog box. Simply click ‘Finish’ to
open the file. You will see a column of wavelengths and a column of absorbances
starting on line 15. Select the data (both columns) corresponding to wavelengths of
interest and make a graph of the spectrum using the Chart Wizard.
19.
Shut down the Ocean Optics UV/VIS Spectrophotometer by selecting ‘Exit’ from the
File menu and then switching off the light source using the 2 lamp switches on the front
of the unit. It is very important to shut off the light source as the deuterium lamp
has a limited life. Leave the laptop running with the cover open. Don’t forget to take
the
cuvette
out
of
the
cuvette
holder.
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Chemistry 317 Lab Manual
Synthesis of Chromous Acetate
"The Schlenk Shuffle"
Background
The main purpose of this lab is to give you some practice manipulating the Schlenk
apparatus and using the glove box. It is important that you fully review the descriptions of these
in the Introduction above. But there is also some rather interesting chemistry involved in this
lab. You will carry out the following sequence of reactions:
Cr 3+ (aq)
Zn
HCl (aq)
Cr 2+ (aq)
NaOAc (aq)
Cr 2 (OAc) 4 • 2 H 2 O (s)
Cr3+ (aq) is a short hand way of writing Cr(H2O)63+, and Cr2+ (aq) similarly stands for the 2+
aquo ion; OAc is an abbreviation for acetate, CH3CO2-. Chromium(III) is the most common
oxidation state of chromium; its compounds are typically air-stable, octahedral complexes.
Chromium(III) can be reduced by strong reductants, like zinc metal, to chromium(II).
Chromium(III) compounds are referred to as chromic compounds, while chromous refers to
chromium(II) ⎯ so chromous acetate is chromium(II) acetate. This -ic/-ous nomenclature is
typically used to denote higher and lower common oxidation states. [The aqueous chemistry of
chromium is summarized in its pE/pH (Pourbaix) diagram (see Principles of Descriptive
Inorganic Chemistry G. Wulfsberg, p. 151).]
The chromium(II) aquo ion is bright blue and reacts rapidly with oxygen to regain the +3
oxidation state. This serves as a convenient visual check on your technique: if you let a solution
of chromium(II) come into contact with air, it will rapidly turn from bright blue to the forest
green color of Cr(H2O)63+ or, in the presence of ligands, the color of other chromium(III)-ligand
complexes.
Simple chromium(II) salts, like the chloride and sulfate, are water-soluble, bright blue,
and have magnetic susceptibilities that indicate four unpaired electrons per chromium, as
expected for a high-spin d4 electronic configuration. However, chromium(II) acetate, first
synthesized by Eugène Peligot in 1844 [Comptes Rendus Hebdomadaires des Séances de
L’Académie des Sciences, 1844, 19, 609] is unusual.
It is brick-red, rather insoluble in water (very rare for an acetate
salt) and was thought to be diamagnetic (no unpaired electrons).
It is also much more stable to air than the other chromium(II)
salts. In 1970, F. A. Cotton and co-workers [F. A. Cotton, B. G.
deBoer, M. D. LaPrade, J. R. Pipal, D. Ucko J. Amer. Chem. Soc.
1970, 92, 2926] re-determined its crystal structure by X-ray
diffraction. They found that it had the dimeric structure shown at
right with a short Cr-Cr distance of 2.362 (1) Å (R = CH3; L =
H2O; the (1) on the distance is an estimate of the error (1σ) in the
last decimal place; 1 Å = 100 pm).
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The short distance and the diamagnetism were explained by proposing that the two
chromium atoms form a quadruple bond, consisting of one σ bond, two π bonds, and one δ
bond. Quadruple bonds are now well established for a number of 4d and 5d transition metals
[see D. Shriver & P. Atkins Inorganic Chemistry, 3rd Ed., W.H. Freeman, New York, 1999, pp.
304-5.] . However, the presence of quadruple bonds in chromium compounds such as
Cr2(OAc)4 has been questioned and the compounds have been found to be slightly paramagnetic:
see S. Hao, S. Gambarotta, and C. Bensimon J. Am. Chem. Soc. 1992, 114, 3556-7 and F. A.
Cotton, H. Chen, L. M. Daniels, and X. Feng J. Am. Chem. Soc. 1992, 114, 8980-3. A
discussion can be found in: J. M. Mayer ChemTracts Inorg. Chem. 1992, 4, 209-213.
Method
This experiment covers two periods; you should work with your partner. The first day you
will prepare chromous acetate. You will reduce chromium(III) chloride to aqueous chromium(II)
with mossy zinc in acidic solution and then add the Cr2+ solution to a solution of sodium acetate,
which will result in the precipitation of brick-red hydrated chromous acetate, Cr2(OAc)4.2H2O.
You will isolate this somewhat air-sensitive solid by Schlenk filtration. The second day you will
take the solid into the glove box, weigh it, and characterize it. The air-sensitive products show the
presence of oxygen by changing color, so you will see if any air is leaking in.
Procedure
0. Get Acquainted with your Schlenk Line
Be sure you are familiar with the descriptions of Schlenk techniques and gas handling in
the Introduction (pp. 10-16). Take a moment to compare the actual Schlenk line with the picture
given. As if dissecting a frog in biology class, try to find the major systems: the line, the
bubblers, the vacuum pump, and the nitrogen tank.
Identify the nitrogen manifold. In your best imitation of Sherlock Holmes, trace along
the manifold in one direction to reach the nitrogen tank. In the other direction, you should find
the Nujol bubbler. Is the nitrogen flowing? We'll assume not, because it’s not good practice to
let the nitrogen flow 24 hours a day (you’re wasting it). You should leave the nitrogen on
through the lab period, and shut it off when you leave.
The source of nitrogen for the Schlenk lines is so-called “house” nitrogen, rather than a
nitrogen cylinder. This house nitrogen is plumbed into the fume hoods, and its flow is controlled
using the color coded valves to the left and right of the fume hoods. The house nitrogen comes
from boil-off from the departmental liquid nitrogen tank, which is at the south end of the loading
dock between Bagley and CHB – take a look at it next time you’re on the road behind Bagley.
A gas regulator and tank are available in the lab if you are interested in learning more
about their operation. For completeness, we include here the normal procedure for using a gas
regulator with a Schlenk line is as follows. Assume that the regulator is full of air, unless you
are sure of the contrary. Clear out the air by purging the regulator, passing a lot of gas through
it. First, close the needle valve at the end of the regulator ⎯ it screws in to close. Then open the
large valve on the top of the nitrogen tank. The pressure gauge nearer the tank should read
between 2200 and 200 psi ⎯ this is the pressure inside the tank. [If it's lower than this, tell your
TA.] Now cautiously open the little needle valve that sits between the regulator and the rubber
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tubing. Is any gas coming out? Try to do this without blowing Nujol all over your bench.
Adjust the diaphragm valve so that the pressure on the second gauge is a moderate 3-6 psi. This
valve gets more open as you screw it in. Use the needle valve to turn the gas flow down until it
barely bubbles out, then up to a moderate blast (again, watch the Nujol). Turn the nitrogen flow
back down to a slow bubble so you don't empty your tank. In general, the pressure should be set
so that gas will bubble pretty hard if the needle valve is way open. This way, when you need a
lot of nitrogen, you only have to play with one valve.
Now find the vacuum manifold. Trace out the path gas will take when it's pumped out,
starting from the digital vacuum gauge, through the line, through a stopcock into the cold trap, a
piece of vacuum tubing and ⎯ aha! ⎯ into the vacuum pump. Make sure all the stopcocks and
ground glass joints are greased (for instance, the cold trap). Close all the stopcocks on the
vacuum manifold and turn on the pump. You should be pumping only on the vacuum hoses and
the cold trap. The pump noise will increase, which should subside when the pump gets rid of the
air in the tubing and trap. The cold traps are now under dynamic vacuum; in other words, the
pump is actively sucking on them. Place a Dewar flask (these are just giant open Thermos
bottles) around each cold trap. The bottom of the Dewar must be firmly supported and you
should be careful not to drop them. The Dewars cost over $100 and shatter easily. Carefully fill
the Dewars with liquid nitrogen.
Liquid nitrogen presents two serious hazards in the lab. It is incredibly cold: 77 K = -196
°C = -321 °F! If your skin comes into contact with it for any length of time it will give you a
nasty cold burn. Fortunately, your skin (and everything else around you) is incredibly hot in
comparison to the liquid nitrogen, so the nitrogen evaporates like crazy. This is like drops of
water on a very hot griddle. So if you get a little onto your hand, you'll have a layer of N2 gas
between your skin and the liquid, insulating you. Just don't hold any liquid nitrogen in the same
spot on your hand for any length of time. The second hazard is scarier: liquid nitrogen traps are
perfectly willing to condense oxygen out of the air. A trap full of blue liquid O2 can explode if
there is any organic matter in the trap (as there often is), so blue liquid in a trap is one of the
most chilling sights anyone working at cryogenic [very low] temperatures could encounter.
Fortunately, our pumps are powerful and the vapor pressure of the liquid oxygen is fairly large;
oxygen will not stick around to give your TA a heart attack as long as the pump is on, or in other
words, when the system is under dynamic vacuum.1
Returning to your vacuum line, you should not be condensing any oxygen because your
system is under dynamic vacuum. You can go ahead and open the vacuum manifold to the pump
and wait until the pump sounds quieter. Watch what happens to the reading on the digital
vacuum gauge. When it stops changing (give it some time to do this), write down the reading.
Now close the stopcock to your cold trap so that the line is under static vacuum, without active
pumping. If the reading on the digital gauge rises, you have a leak in your line. Try to locate
this leak, checking your stopcocks and their grease, then look to make sure your hoses aren’t
cracked. Your TA should probably help with this.
1In
the unlikely event that your trap condenses some liquid oxygen, tell your TA, who will help you put the trap
back under dynamic vacuum (if the pump is not working, the Dewar flask is removed and the liquid oxygen allowed
to evaporate). While the liquid oxygen is evaporating, everybody evacuates the lab for a judicious period.
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The process you and your partner have just completed is called “bringing up the line;”
now you can impress your friends with lab slang. To take your line down, repeat the above steps
in reverse order.
Quick summary of “bringing up the line”:
1.) Turn on the N2 and adjust the needle valve so a gentle, constant stream of bubbles
forms in the bubbler.
2.) Close all stopcocks and make sure the trap is on.
3.) Turn on the pump and open the stopcock on the trap. Make sure the reading on
the vacuum gauge changes.
4.) Check for leaks.
5.) Fill the Dewar around the trap with liquid N2.
I. Generation of aqueous chromium(II)
In a hood, load a 125 mL Erlenmeyer flask with 5 mL distilled water, 20 mL
concentrated HCl (be sure to check the molarity), 5.0 g chromium(III) chloride hexahydrate, and
a magnetic stirbar. Stir the solution to dissolve the chromium(III) chloride. Pour the green
solution into a Schlenk addition funnel (check that the bottom stopcock is closed or the stuff will
pour out onto your shoes or onto the floor of the hood). Insert a lightly-greased ground-glass
stopper into the top of the addition funnel. Clip it on with a yellow plastic clip (these are called
‘keck’ clips) so that when you slightly pressurize the addition funnel with nitrogen the stopper
doesn’t depart for outer space.
You now have to get the oxygen out of the addition funnel, both the oxygen gas and the
oxygen dissolved in the solution. Connect the sidearm of your flask to the Schlenk line and open
the two-way stopcock on the Schlenk line to the vacuum manifold. Listen for the pump as you
open the stopcock and watch the vacuum gauge. Now slowly open the stopcock on the addition
funnel, while listening for the pump and perhaps for the gas whistling through the stopcock. The
solution will start bubbling a little as the water begins to boil at the reduced pressure. Let it
bubble for a second or two, then shut the Schlenk line stopcock.
Very slowly turn the Schlenk line stopcock 180° to admit nitrogen gas to the evacuated
addition funnel. BE CAREFUL. As you open the evacuated addition funnel to the nitrogen
manifold, the pressure in the manifold will drop dramatically, which will tend to cause the Nujol
in the bubbler to suck back into the tubing ⎯ very, very messy. Many groups wind up having to
clean oil out of their line at some point because they forget to worry about this. To minimize this
problem, turn up the nitrogen flow to “moderate blast” before you open the two-way stopcock,
and open the stopcock as slowly as possible. Putting your finger (or your partner’s finger) over
the exit tube on the Nujol bubbler also helps. Repeat this process of sucking out the gas and
refilling with nitrogen two more times. This is called pump-purge degassing: first you pump the
air away and then you purge (fill) with inert gas. Two or three pump-purge cycles are a standard
way to rid a vessel of oxygen; the atmosphere in the addition funnel is now essentially pure
nitrogen.
Add 9 g of mossy zinc (yes, this is really what it’s called; make sure the label says mossy
zinc) to a 250 mL Schlenk flask and clamp it to your aluminum rack. Place the addition funnel,
lightly greased, into the top of the Schlenk flask, and clamp it as well. Connect the flask to the
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Schlenk line with another one of the hoses. The addition funnel acts as a stopper. Degas the
flask with two pump-purge cycles. (Because there’s no solvent to dissolve gases, it’s easier to
degas than the first flask.) Once you have finished pumping and purging, turn the nitrogen flow
down to about 100 bubbles/min. Open both the addition funnel and the flask to the nitrogen line
and drip the acidic solution of CrCl3.6H2O in to the flask containing the mossy zinc. As both
heat and hydrogen are generated, the addition must be done drop by drop; if the flask with the
zinc feels quite hot, stop adding acid (a little hot is OK). You can also moderate the heat
generated by periodically applying an ice bath to the bottom Schlenk vessel. It is important that
both vessels be open to the nitrogen line to prevent a vacuum from developing in the addition
funnel and to allow the hydrogen gas evolved in the reaction to escape through the gas bubbler.
The chromium chloride/HCl solution will undergo a dramatic color change from green to robin’s
egg blue (almost sky blue) as it is reduced to Cr2+ (aq).
II. Addition of sodium acetate
You now have a solution of a highly air-sensitive intermediate, Cr2+, in a flask with an
addition funnel protruding from the neck. Turn off the stopcocks leading into the addition funnel
and unclamp it. Have a rubber septum handy. Turn up the nitrogen flow and make sure your
CrII flask is open to nitrogen. Cover the exit of the gas bubbler with a finger and remove the
addition funnel while nitrogen is blasting out of the flask with the chromium(II) solution.
NOTE: When you cover the exit of the nitrogen bubbler, you turn the nitrogen manifold into a
closed system with gas pressure building up inside it. If you do not promptly remove the
addition funnel to vent the system, it will create its own vent by popping open a part of your
apparatus. Also, keep a good grip on the addition funnel as you remove it, lest it leap onto the
benchtop. This is called doing something against a counterflow of nitrogen.
Immediately cap the flask with the rubber septum. Turn the two-way stopcock to open
the chromium flask to the vacuum and do a pump-purge cycle or two to help clear out any
oxygen that may have been inadvertently introduced. If air gets in, the blue Cr2+ will be
oxidized to green Cr3+, so you’ll be able to see if you’re doing it right. Any excess zinc that is
hanging around will reduce this unwelcome ion back to the Cr2+ species.
Pour 30 mL of saturated NaOAc solution and a stir bar into a 100 mL Schlenk flask, cap
with a septum, and hook it up to the Schlenk line. (Now you know why there are so many ports!)
Turn up the nitrogen and do 3 pump-purge cycles to degas the acetate solution. Place one end of
a cannula through the septum on the flask with the chromium solution, but leave the end of the
cannula above the level of the solution. Cover the outlet of the mineral oil bubbler on the
nitrogen line with your finger to increase the pressure of nitrogen; this should blow a good
stream of nitrogen out the cannula and thus purge it of air. Then place the other end through the
septum on the acetate flask. Place a small needle through the septum on the flask with the
acetate. This will serve as an exit for gases to escape. Use your finger to feel that nitrogen is
flowing through this needle. Close the stopcock on the flask with the acetate, take your finger
from the bubbler, and lower the cannula to the bottom of the flask with the chromium solution.
Your apparatus should look like the drawing of cannula transfer in the Introduction (p. 18), and
you are now ready to push the chromium solution over through the cannula.
By covering the bubbler outlet with your finger, you increase the pressure in the
chromium solution and push that solution through the cannula into the acetate flask. You can
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adjust the pressure by playing with the flow rate of nitrogen and how much you let out the
bubbler with your finger. The goal is to have the chromium solution be delivered slowly ⎯ you
should see individual drops fall off the end of the cannula, not a jet of material. Do not stir the
acetate solution. If the small pressure difference between the two flasks is not enough to make
the liquid flow, you can create a larger pressure difference by placing the acetate solution under a
partial vacuum. To do this, you need to remove the small needle so that you can pull a vacuum
on the acetate solution. Apply a slight vacuum to the acetate solution and see if the solution
flows through the cannula. Pull more vacuum as needed.
Once the solution is flowing through the cannula, you can just watch until it’s all done.
The solution will turn red-purple. When the addition is finished, lower the N2 pressure to normal
levels and remove the cannula and the needle (if still there). If you tilt the flask, you will see
small brick-red crystals of chromous acetate huddled at the bottom of the flask.
III. Filtration of the solid chromous acetate
The next step, a Schlenk filtration, will separate these crystals from their mother liquor.
It will also require three arms and your undivided attention. You should review the pictures of a
Schlenk filtration in the Introduction (p. 19). First, you have to get the Schlenk filter onto the
neck of the flask containing the chromous acetate in place of the septum that is currently keeping
out the air. Turn up the nitrogen to a good blast. Hook up both of the stopcocks on the Schlenk
filter to the nitrogen manifold and put a Schlenk or ordinary 100 mL round-bottomed (RB) flask
on one end to catch the filtrate. Be sure the flask has a stir bar in it. Clip the flask on with a
yellow plastic clip. Let the nitrogen run for a few minutes to flush out the filter apparatus ⎯ or
cap the open end and pump-purge it. While one hand blocks the nitrogen bubbler, remove the
septum on the chromous acetate flask with the other hand. Your third hand can now sneak
around and place the Schlenk filter (remember to grease the male joint) on the neck of the
chromous acetate flask. Secure it with a yellow clip. Close the stopcock on the flask and the
lowest stopcock on the filter; open the upper stopcock on the filter to the vacuum manifold. Do
two pump-purge cycles on the whole apparatus through this top stopcock to suck out any air that
may have ventured into the system. It’s better to use the top stopcock so that the air would be
sucked away from the reactive material, rather than over the solution.
Now comes the big moment. Open the stopcock on the flask to the nitrogen line. The
lower stopcock on the frit should still be closed, the upper one still open to nitrogen. Start
inverting the filter. Keep the side arms pointed up to avoid pouring solution into the hoses.
Untangle any hoses (if you reconnect any, pump-purge them twice.) Use a magnet to twiddle the
stir bar around and scrape the solid onto the filter. The solution will fill the top of the filter and
slowly come through. To accelerate the process, pump briefly on the apparatus below the fritted
glass disc. The liquid will flow into the RB flask at the bottom while the solid remains high and
dry on the frit. Repeat brief applications of vacuum to drain liquid, if necessary.
IV. Washing and Workup of the Chromous Acetate
Your solid is isolated, but it is wet. How can you weigh it out, get an IR spectrum, and
do other things with this air sensitive wet solid? You'll want to weigh it in the glove box, but
you can't take it in there wet, or you'll be bringing water into the box ⎯ an absolute no-no. No
water, no oxygen in the box! So first you need to dry the solid. You might have time to do this
in the first period. If not, empty the liquid from the Schlenk filter (think about how you might do
this!), seal the filter back up, pump/purge it a few times and leave it well sealed under nitrogen
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until the next period. If you can, cap and save the filtrate (the liquid that poured through the
filter).
To dry your solid, wash it while it is on the surface of the Schlenk filter, still under
nitrogen. First wash with ethanol, then with ether (twice). These organic solvents will wash
away any remaining water, and then the last bit of ether is easy to pump away. Volatile solvents
like ether (with large vapor pressures) can be pumped away while water evaporates too slowly to
use this technique. Choose how you want to do the washing, using a syringe, a cannula, or the
addition funnel. Notice the joints - you may need an adapter. When you've decided on a
procedure, talk it over with your TA. Remember, the solvents must be degassed before they are
added. The degassing of water solutions above was done by pump/purge cycles, which also
works for organic solvents. But with very volatile solvents like ether, pumping on them causes a
lot of evaporation. Pump/purge cycles still work ⎯ just make sure you pump for a very short
time. You can also freeze the ether in liquid nitrogen and pump on it; at this temperature none
will boil away. Then thaw the ether to let whatever gas was trapped in the solid escape, cool
again, and pump again. This is called freeze-pump-thaw degassing. [Don't do this with water,
since the expansion of water on freezing could crack your flask.]
When your solid is washed, remove the Schlenk flask with the solvents in it and cap the
apparatus with a 25 mL round bottom flask or something similar. Make sure you have a
counterflow of nitrogen whenever you take apart the apparatus. Pump on it for a few minutes.
Close all the stopcocks, disconnect the hoses, and send it into the glove box via the antechamber.
Note that if you're sending a sealed piece of equipment into the antechamber it must be under
vacuum. When the atmosphere in the box antechamber is pumped out, a sealed flask with gas in
it will likely pop open ⎯ the 14.7 psi inside will not be balanced by any pressure outside.
Alternatively, you can leave your flask under nitrogen and open a stopcock immediately prior to
placing it into the antechamber. The open stopcock will allow the pressure to be equal both
inside and outside the flask while entering the drybox. No matter which method you choose,
black electrical tape can be used as a safeguard against exploding flasks. This slightly stretchy
tape will allow an improperly prepared flask to expand in the antechamber, relieving pressure,
without the vessel exploding.
V. Weighing and Exploring Chromous Acetate; Working in the Glove Box.
Working in the drybox seems simple ⎯ just like a benchtop ⎯ but it takes some getting
used to. Think through exactly what you're going to do before you send your stuff into the box.
How are you going to scrape out your product? (Don't scrape too hard as this damages the
filter.) How are you going to weigh it? What are you going to weigh it in? Is it labeled? In
sum, make sure you bring in all of the things you'll need in the box with you. Note also that the
rubber gloves are pretty thick (to prevent diffusion of air through them) so you'll be clumsy.
Don't plan anything too complicated. Weigh your solid so you can calculate your yield.
Take at least two small samples of your product out of the drybox. Put two samples on
the benchtop, and wet one sample with water. How fast do the two samples react with air? Do
you have any sense of how fast the Cr2+ (aq) solution reacts with air? Use your other small
samples to explore other reactions of Cr2(OAc)4, whatever you think would be interesting (if you
have time). Can you protonate off the acetate ligands with HCl to reform Cr2+ (aq)?
Make up a Nujol mull of your product for an IR spectrum. Does this have to be done in
the drybox? You'll find a description of how to make up a Nujol mull in the Phosphoric Acid
experiment, along with an introduction to IR spectroscopy. You should also get IR spectra of
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Chromous Acetate
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plain Nujol and dry NaOAc (dry, not the aqueous solution - NEVER put water on salt plates) for
comparison. Try to get all of your spectra on the same horizontal (frequency/wavelength) scale.
Safety and Waste Disposal
All of the waste from this experiment (except zinc) should go in a 4-L waste bottle
labeled appropriately, “Chromous Acetate Waste”. In this and all other waste issues, the TA’s
will instruct you.
Be sure to clean your Schlenk equipment well as others will be using it shortly.
Lab Write-up
A report template (see page 94) will be distributed to you electronically. Fill in the blanks with
the requested information and submit a printed copy one week after the completion of the
experiment.
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Chelate Effect
Chemistry 317 Lab Manual
The Chelate Effect
Background
Transition metal cations are Lewis acids; they want to share electron pairs. The
molecules or ions that supply these electron pairs are called ligands, and they coordinate or bind
to the metal center. Previous chemistry classes (165, 312, or 416), may have dealt primarily with
water as the ligand, but essentially anything that can act as a base can act as a ligand to a metal
center. The addition, loss, and exchange of ligands at metal ions form the core of coordination
chemistry.
Because of the importance of ligand-exchange reactions, it is convenient to have a way of
assessing the relative thermodynamic stability of various complexes. In aqueous chemistry
(chemistry in water), the stability of complexes is typically measured relative to the aquo ion of
the metal in a given oxidation state. The aquo ion is the complex ion where all ligands are water.
In the example shown, addition of ammonia displaces water from the aquo ion to make the
tetraammine complex. The more stable a given complex is, the larger its Keq will be.
[Cu(H2O)n]2+ + 4 NH3 → [Cu(NH3)4]2+ + n H2O
Keq =
(1)
[Cu(NH3)42+]
[Cu(H2O)n2+][NH3]4
(2)
It has long been known that “all ligands are not created equal.” Often the better
something is as a base, the better a ligand it will be. Hard/soft acid-base effects are also
important: for example, cobalt(III) has a marked affinity for nitrogen donors while the softer
mercury(II) is partial to sulfur. But consider the following reaction:
H 3N
NH 3
Cu
H 3N
2+
2+
NH 2 NH 2
+ 2
NH 3
Cu
NH 2 NH 2
NH 2 NH 2
+ 4 NH
3
(3)
Ethylenediamine is just about the same as ammonia, as far as the lone pair on nitrogen is
concerned. The major difference between the two is than the nitrogen in ethylenediamine (en) is
tethered to a second nitrogen which can also bind to copper. Ethylenediamine is termed a
chelating ligand because it can bind to a metal through more than one site; it is a bidentate ligand
because it binds in two places. Measurements (such as the ones you will make in this lab) show
that, given the same set of lone-pair donors, a metal will take bidentate or polydentate ligands
over monodentate ones. This is a general phenomenon known as the chelate effect. For a good
discussion of the chelate effect, see F. Basolo, R. G. Pearson Mechanisms of Inorganic Reactions
2nd Ed. Wiley 1967, pp. 27ff, 223ff.
The chelate effect is of enormous importance in coordination chemistry and in bioinorganic chemistry. Nature almost invariably uses a chelating ligand to bind a metal. A classic
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Chelate Effect
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example is the heme group in hemoglobin, which strongly binds iron. Chelating ligands are also
widely used as reagents in the laboratory and even as drugs. For instance, ethylenediaminetetraacetate (EDTA), the hexadentate chelate shown below, is used (as its calcium complex) to
treat lead and other heavy metal poisoning. The heavy metal binds to EDTA and the complex is
less toxic and is more readily excreted from the body.
O
n-4
O
-
O2 C
N
O2 C
EDTA
N
4-
CO 2
-
CO 2
-
N
O
M n+
N
O
O
O
O
EDTAmetal
complex
O
EDTA is also frequently added to foods to retard spoilage (e.g. mayonnaise⎯check your bottle
at home). EDTA binds metals (often iron) that would otherwise catalyze the air oxidation
(termed autoxidation) of foods.
The chelate effect is quite general and also works in main group chemistry. For example,
the British developed BAL (British anti-Lewisite) in World War I to combat the arseniccontaining war gas, Lewisite. As a bidentate sulfur donor, BAL binds tightly to the soft arsenic
center. It is still used to treat some kinds of heavy metal poisoning.
SH
Cl
As
Cl
Lewisite
S
As
SH
+
S
SO 3 -
+ 2 HCl
SO 3 -
BAL
Both EDTA and BAL fight heavy metal poisoning because the chelate effect enables them to
bind strongly to the undesired metals and the complexes formed are small and charged, hence
water-soluble and easily excreted from the body.
In this lab you will investigate in more detail why the chelate effect is observed (see page
99). In the reaction of tetraamminecopper(II) with ethylenediamine (equation 3 above), three
particles are converted to five particles. This means more translational disorder in the products,
and the reaction should have a favorable entropy (ΔS > 0). This has been the traditional
explanation of the chelate effect. However, some chemists have suggested that enthalpy may also
play a role. You will address this question by measuring ΔH and ΔS of this chelate-formonodentate exchange reaction.
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Consider the following electrochemical cell:
Ammonia forms a rather stable complex with copper(II). Another way of saying this is to say
that the copper(II) would like to move from the left-hand beaker, where it is surrounded by mere
water molecules, to the right-hand beaker, where it could enjoy the attentions of the more basic
ammonia molecules. The copper ions can’t make this leap directly, but indirectly the same
movement is accomplished by moving electrons from the right-hand beaker to the left-hand
beaker. Simultaneous with this, copper metal is converted to copper(II) in the ammoniacal
solution and, in the water solution, copper(II) is plated out onto the copper electrode. Inert ions
flow through the salt bridge to complete the circuit. Thus:
Left-hand beaker (cathode):
Right-hand beaker (anode):
Cu2+ (aq) + 2e- → Cu (s)
Cu (s) + 4 NH3 (aq) → [Cu(NH3)4]2+ (aq) + 2e______________________________________________
Overall reaction:
Cu2+ (aq) + 4 NH3 (aq) → [Cu(NH3)4]2+ (aq)
Note that the overall reaction is not a redox reaction at all, but a Lewis acid-base reaction!
Nevertheless, the voltage which can be read on the voltmeter between the two cells is a measure
of how desperately the reaction wants to proceed; the larger the voltage, the more spontaneous
the reaction. Put more precisely, the difference in electrical potential (E) between the two cells
is proportional to the difference in chemical potential (a fancy synonym for the change in free
energy, ΔG) between the two sides of the net reaction (see the Data Analysis section below).
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Method
This experiment is in two parts, in which you will work with your lab partner. This lab
manual describes the first part: a two-period lab experiment. In the first period you will
synthesize and isolate [Cu(NH3)4]SO4.H2O and [Cu(en)2]SO4.2H2O. In the second period you
will use the compounds you have synthesized to make up half-cells and measure the potential
between them as a function of temperature. You will then write a lab report describing what you
have done and deriving the thermodynamic parameters ΔH and ΔS for each of the three
reactions.
There is a common complaint about lab courses that they are “too cookbook" and don't
represent what scientists really do. A lot of what scientists do is to try something and find that it
doesn't work too well⎯that they overlooked some source of error or that the method wasn't great
or that they just messed it up or whatever. By thinking about the problems encountered, we
devise what we hope is an improved method and then we try it. In your lab report based on the
first two lab periods, you will propose how to make this experiment better. You can decide to
collect the same data, with an improved procedure. Sometimes, just doing the same procedure a
second time is enough to make it work better. Or you can decide that your data from the first
attempt is fine, and propose an extension of this experiment, perhaps using a different complex
or different reaction conditions.
After you get your report back you will talk about your ideas with your TA. Then you
will be given a lab period (called Chelate Effect II in the schedule) to put your ideas into
practice. No write-up is provided here for this "extra" lab period⎯you have to devise your own
approach. You will then write a revised lab report, taking into account the comments on the first
version and describing your work in the Chelate Effect II lab period. You will be graded both on
the quality of the report and on quality of your use of this “extra” lab period.
Experimental
I.
Synthesis of tetraamminecopper(II) sulfate monohydrate, [Cu(NH3)4]SO4.H2O, and
bis(ethylenediamine)copper(II) sulfate dihydrate, [Cu(en)2]SO4.2H2O.
In each pair of students, one should prepare the tetraammine complex, the other the
ethylenediamine complex. Each person should follow the procedure below, with their respective
ligand. Work should be done in the hoods instead of on the bench to avoid stinking up the lab.
Dissolve 10.0 g copper(II) sulfate pentahydrate (CuSO4.5H2O; 0.040 mol) in about 10
mL of water (heating and stirring will hasten its dissolution). Allow the solution to cool
somewhat. To make the tetraammine complex, add 18 mL conc. ammonium hydroxide to the
copper solution and stir vigorously. To make the ethylenediamine complex add dropwise 5.35
mL (4.8 g, 0.080 mol) ethylenediamine with vigorous stirring. The basic amines can cause light
blue Cu(OH)2 to precipitate, but this usually dissolves after a few seconds with the formation of
deeply colored complex ions. If the light blue solid persists, break it up with a glass rod or a
spatula; warming the suspensions may also help to get rid of the light blue solid. Adding more
ammonia or ethylenediamine may help as well. If there is still light blue solid, filter the solution.
Allow the solution to cool. Add about 20 mL of ethanol to precipitate the complex, and
cool further in an ice/water bath. Isolate the solids by suction filtration on a Büchner funnel;
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Chemistry 317 Lab Manual
wash thoroughly with 5 mL ethanol [20 mL in the case of the bis(en) complex], followed by 5
mL diethyl ether, then air-dry. The bis(ethylenediamine) complex precipitates as the 4.5-hydrate
but is converted to the dihydrate on washing with ethanol; the tetraammine complex is isolated
as the monohydrate. Dry the products in the oven at < 100 deg C, placing them on labeled watch
glasses. You may want to stir periodically to ensure even drying. Check your compounds often:
if the oven is too hot or you leave them in for too long, light blue solid will form around the
edges of the dark blue/purple product you want. As always, determine your yield. Store the two
compounds in labeled vials until the next period.
II. Preparation of electrochemical cells
The second lab period of this experiment can be a long one. Before you get to lab, think about
how you can be most efficient. When you’re in the lab, get right to it.
Make up approximately 600 mL of electrolyte solution by dissolving 12 g KCl in 600 mL
DI water (0.27 M). On the lab benches, there are set up six constant temperature baths ranging
from 0 to 45 degrees C. Fill the appropriate number of electrochemical cells (the clear vials with
stoppers glued to the bottom) about 80% full. Keep in mind that the level of the KCl solution
must be higher that the level of the solution in the electrodes that will be constructed (below).
Place one of the electrolyte cups in each of the water baths and let it sit there to come to thermal
equilibrium (at least 30 min!). [How will you know when everything is at the same temperature
(thermal equilibrium)?] Note the temperature in your electrochemical cell before you make any
electrochemical measurements. Does it stay constant throughout your measurements?
To construct the electrodes, dissolve 2 g gelatin and 1 g KCl in 50 mL boiling water in a
small beaker (the water must be boiling prior to addition of the gelatin, although the KCl can be
added earlier). Once all of the gelatin has dissolved (forming Agar), remove the beaker from the
hot plate and dip the tips of 5 special cut plastic pipettes into the Agar as instructed by your TA
(you only need 4 pipettes but one can be used as a spare). Allow the Agar to cool and harden.
Use volumetric glassware and deionized water to prepare three solutions:
A: 10 mL of a solution 0.50 M in CuSO4.5 H2O
B: 10 mL of a solution 0.50 M in [Cu(NH3)4]SO4.H2O and 1.00 M in ammonia
C: 10 mL of a solution 0.50 M in [Cu(en)2]SO4.2H2O and 1.00 M in ethylenediamine
Calculate how much of each reagent you will need (do this prior to coming to lab) and
compare your results with your lab partner. The density of neat ethylenediamine is 0.90 g/mL;
concentrated ammonia is 14.8 M. You can measure out the ethylenediamine by volume or by
weight (which do you think is more accurate?). An easy way to make up the solutions is to
dissolve the appropriate amounts of the complexes in roughly 3-4 mL of water, with stirring and
perhaps heating. After everything is dissolved and cooled back to room temperature, add the
solutions to 10 mL volumetric flasks, rinse the beakers with a few mL of water and add the
washings to the volumetric flask, and finally make up to exactly 10 mL. You will need
approximately 2-3 mL of each solution to prepare an electrode.
Remove the plastic pipettes from the beaker after the Agar has hardened and place them in a
special electrochemical cell (see page 102) that is provided in the lab. The electrochemical cell
consists of a plastic vial with a plastic cap with holes punched in the top where the pipettes sit.
The vial is glued to a large rubber stopper so that it can be placed in a water bath without tipping.
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Pour each of the copper solutions into a separate pipette. Place a copper wire electrode into each
pipette. Use a rubber septum at the top of each copper electrode to adjust the height of the
electrode such that it is immersed in the solution but does not touch the Agar plug at the bottom
of the pipette. Label your electrodes!! The level of the copper solutions within your electrodes
should be below the level of the KCl solution. Finally, fill the fourth electrode with KCl solution
and place a digital thermometer inside. This will be used to determine when the temperature
inside the electrodes has come to thermal equilibrium with the KCl solution.
Use a voltmeter (your TA will give you instructions for the specific brand you have) to
measure the potential between A-B, B-C, and A-C (see page 102). Be sure to record which lead
of the voltmeter was connected to which electrode so you can figure out the signs. The voltage
readings may drift a bit initially, but they should settle down in a few minutes. Check and
double check that you have good connections between the copper and the voltmeter. Read the
voltages to one more decimal place than you think is meaningful. The potential across A–C
should be equal to A–B plus B–C (once you get your signs right). In other words, the values
should be internally consistent. If they are not, think about what might cause a deviation, and
repeat your measurements. We suspect that the quality of the copper may influence the results,
so try to use a clean piece of wire. You can also roughen the surface with sand paper. You
might want to change the wire during the experiment to see if there’s any effect. Periodically
check electrodes for blue/black discoloration. Replace corroded electrodes, or polish them with
sand paper. If they don’t look shiny and metallic, they may not work correctly.
Note: if your voltmeter is not giving reasonable readings, ensure that there is a direct connection
between the agar and the electrolyte solution. If air bubbles are trapped inside the end of the
pipette tip there will not be an electrical connection.
You should make these measurements at 6 temperatures (at least) over as wide a
temperature range as possible. Only move your electrodes and the white cap between baths! The
more measurements and the larger the temperature range, the more reliable your data will be.
However, keep the solutions below 45 degrees or weird things might happen⎯i.e., your agar
plug (salt bridge) could melt or the amines might evaporate [how would you know?]. If you
have time left, you might repeat measurement(s), in order to check the reproducibility of your
data.
In addition to recording your data in your lab notebook, we will ask you to record it in a
computer spreadsheet provided in the lab. We will use this pooled data during the lecture section
in our discussions on error.
Safety and Waste Disposal
All of the copper waste from this experiment (including the solutions in the cells) should
go in a 4-L waste bottle labeled “Chelate Effect Waste”. The KCl electrolyte solutions can be
poured down the drain (unless a significant amount of copper has leached into it). Used copper
wire goes into a special collection box. Agar can be disposed of in liquid waste.
REMEMBER TO SAVE YOUR LEFT OVER COMPOUNDS TO USE IN THE
CHELATE EFFECT II THAT YOU WILL DO IN A FEW WEEKS.
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Data Analysis
Your goal is to determine the thermodynamic quantities ΔH, ΔS, and ΔG, for each of the three
reactions below (equations 1 & 3 are repeated from above). Also determine Keq at 298 K for
Cu(NH3)42+(aq) and for Cu(en)22+(aq).
Cu2+(aq) + 4 NH3(aq)
Cu2+(aq) + 2 en(aq)
Cu(NH3)42+(aq) + 2 en(aq)
Cu(NH3)42+(aq)
Cu(en)22+(aq)
(1)
(4)
Cu(en)22+(aq) + 4 NH3(aq)
(3)
To obtain these values from the measured voltages, you need the basic thermodynamic relations
in equations 5-7. If these don’t look familiar, you should review the relevant sections on
thermodynamics and electrochemistry in your first-year chemistry text (you might look in the
index for free energy or Faraday constant).
ΔG = -nFΔE
(5)
where n is the number of electrons involved in the reaction (in
this case, 2) and F is Faraday's constant, 96.5 kJ/mol V.
ΔG = ΔH - TΔS
ΔG = -RT lnKeq
(6)
(7)
Equating the expressions for ΔG in equations 5 and 6,
-nFΔE = ΔH - TΔS,
-ΔH
ΔS
ΔE = nF + T nF
or:
(8)
Thus a graph of ΔE versus temperature (in Kelvin) should have a slope of ΔS/nF and an intercept
of -ΔH/nF. Make such plots for each of the three reactions. Determine ΔH and ΔS, and then ΔG
and Kf at 298 K from equations 6 and 7. Make a Table with your results. Are the values you
obtain ΔH˚ and ΔS˚ or just ΔΗ and ΔS? Keep in mind that we are using electrochemistry, so you
need to think about the definition of standard state conditions for electrochemistry.
Note that if you reverse the direction of reaction 1 and add it to reaction 4, you get
reaction 3. Compare the values of Keq, ΔH, and ΔS that you obtain in this manner for reaction 3
with those you calculated from the direct measurement. (Watch out for sign flips! Ask your TA
if you aren’t sure about the math.) If there are differences, are they significant? What does the
presence or absence of differences mean?
Estimate the uncertainties in the values you report. We’re looking for a qualitative feel
for how reliable you believe your values are⎯there’s quite a difference between ΔS = 3 ± 2
J/mol.K and ΔS = 3 ± 200 J/mol.K. A good way to get this feel is by inspection of your graph.
What are the maximum and minimum values that the slope (or intercept) could reasonably be,
given the data that you measured?
Think about the reasons for your observed uncertainties. If there is markedly more
uncertainty in the enthalpies than the entropies, or vice versa, say why. What experimental
values go into these calculations (concentrations, temperatures, etc.) and how accurate are these
values? How would small errors in these values affect your measured voltages? Use the Nernst
equation (eq 9), where Q, the reaction quotient, has the same functional form as the equilibrium
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Chemistry 317 Lab Manual
constant (as before, check your freshman text if this looks unfamiliar).
ΔE = ΔE° - lnQ
(9)
How would errors in your measured ΔE’s affect your calculated enthalpies and entropies? Do
you think your errors are random errors, caused by uncertainties in measurement, or systematic
errors that result from your particular experimental method and equipment? Systematic errors
might include your thermometer being a few degrees off or your not accounting for some other
voltage difference in your cells. How might any systematic errors you find cause your calculated
thermodynamic values to deviate from the “true,” accepted values?
Many 317 students have found these data and error analyses a little tricky. So we will
devote most of the lab discussion hour the Tuesday morning following the lab to these issues.
Please make a first try at the analysis before this class period. Bring your data and your plots to
the discussion hour, so we can talk about them. This way you can get advice and you can see the
results from a number of groups. Most people find that this is very helpful when you’re writing
your lab report.
Error Propagation
The discussion of errors above is qualitative, not quantitative. You should look hard at
your data and your plots and decide for yourself what you believe, what you can confidently say.
How different could the slope of the line be, given the error bars on each point? Where do the
errors mostly come from and how could they be minimized? This qualitative, almost intuitive
approach is what we really want you to do in Chem 317.
Some students would like to go farther and do a more mathematical analysis of how an
error in a measurement will affect the final result, an exercise called error propagation. This is
an option for 317. An overview of this topic is given here but the interested reader should
consult books on this topic (the following should be on reserve at the Chemistry Library: Philip
R. Bevington, D. Keith Robinson Data Reduction and Error Analysis for the Physical Sciences
Boston : McGraw-Hill, 2003). Essentially, error propagation enables the calculation of the
uncertainty in a quantity A when A is a function of a number of inputs x, y, z, … [A =
f(x,y,z,…)], given the uncertainties in the various inputs. This is only applicable to random
errors, not systematic errors.
In general, measuring a parameter x multiple times will give a distribution of results
centered about a mean (average) value
x which is typically taken to be the
“true” value.
The width of the
distribution is defined by the standard
deviation σ (listed as “s” in the figure).
Out of all the measurements, 68% will
lie within one standard deviation of the
mean, 95% will lie within 2σ, and more
than 99.5% within 3σ.
So the
uncertainty in A is given as the standard
deviation σ(A). Rather than giving the full derivation (in terms of partial derivatives ∂A/∂x etc.),
we will just give the results here for three common cases.
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• When A is a multiple of one parameter:
A = 3x
then σ(A) = 3 σ(x)
• When A is the sum of other parameters:
A = 3x + 8y – z
then σ(A) = {[3σ(x)]2 + [8σ(y)]2 + [σ(z)]2}1/2
• When A is the product of other parameters:
—(x) 2
—(z) 2 1/2
—(y) 2
A = 4xy/z
then σ(A) = A{[
] +[
] +[
]}
z
y
x
—(x) 2
—(z) 2 1/2
—(y) 2
or
σ(A) = 4xyz/v{[
] +[
] +[
]}
x
z
y
• When A is the logarithm of a parameter:
—(x)
A = ln(x)
then σ(A) =
x
• When A is the exponential of a parameter:
A = ex
then σ(A) = e x σ(x)
Lab Report (first two lab periods)
A report template will be distributed to you electronically (see page 104). Fill in the blanks and
submit a printed copy one week after the completion of the experiment.
Please think about the following questions, the answers do not need to be submitted, but will
prove useful in writing the final report.
1) Write balanced equations for all of the syntheses of Cu complexes you performed in this lab.
2) Write out balanced half reactions for the redox reactions you measured in this lab.
3) Make a table showing all of your electrochemical measurements. See example below.
temperature (K)
275
298
Cu(NH3)42+ vs. Cu2+
0.00 V
0.00 V
Cu(en)22+ vs. Cu2+
0.00 V
0.00 V
Cu(NH3)42+ vs. Cu(en)22+
0.00 V
0.00 V
4) Construct plots of ΔE vs. T. Calculate ΔH, ΔS, ΔG and Keq for each plot. It may also be
helpful to organize these data in a table. Are these values ΔH°, ΔS°, and ΔG° or just ΔH, ΔS, and
ΔG? In other words, what does the (°) mean? What is the difference between Q and Keq?
5) Every reported value (ΔE, ΔH, ΔS, ΔG, etc.) should include an estimate of error (e.g. ΔH = 10
± 1 kcal mol-1). Were there fluctuations in temperature or readings from the voltmeter? Are there
any other sources of error? Some error many not be easily quantifiable (e.g. decomposition of
copper complexes at high temperature). Are the errors random or systematic? Explain your
reasoning.
Plot these fluctuations as error bars on the above plots. Using these error bars estimate the
maximum and minimum slope and intercept of the fits to the data. Translate these fits into error
bars for ΔH, ΔS, ΔG and Keq.
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6) Do the error bars account for the scatter observed in the ΔE vs. T data? Do the error bars on
ΔH, ΔS, ΔG and Keq seem appropriate? Do the errors seem too large? Too small? In other words,
based on observed experimental errors do the propagated error bars make sense? Explain your
reasoning.
7) Does the data support the chelate effect? Is the chelate effect enthalpy or entropy driven?
Explain the logic used to make this conclusion. How does the above error propagation effect the
conclusion? The correct answer may be that the errors are so big that a conclusion cannot be
reached.
8) Describe the plan for the final lab period (in a couple of weeks). This should be no longer
than one double spaced page of text! Based on the above questions, particularly the discussion
of errors, describe the experimental procedure you plan to pursue during the “extra” lab period.
Be sure to address how your approach will address any problems you found. Feel free to discuss
this with your TA. This section must be submitted or we can’t let you work in the lab during the
Chelate II lab period. Be certain to discuss your plan with your lab partner!
Final Lab Report
The final lab report should be written in the style of the discussion/conclusion section of a formal
lab report (3 pages of text maximum, NOT including plots and data tables). DO NOT write a full
lab report (Introduction, Experimental, etc.). Data should be worked up in the same manner as
the first report (tables and plots where appropriate). You do not need to re-answer the questions
from the first report. They should be used as an outline/guideline for the topics to be addressed
in your final report. Include data from both the first and second days of the lab. It may be useful
to compare and contrast these data. Did your plan for the final lab period decrease the overall
error in the data? You might choose to omit data from the first set of measurements if you think
the values from the “extra” lab period are better; explain your reasoning in the report.
Be sure to concisely address the following points in your report:
Analysis of the ΔE versus T data. Discuss your graphs and your values of ΔG, ΔH, and ΔS
and Keq for each reaction. What are the errors on these numbers? If you have not
discussed how you derived the errors in the results section, do so now.
Did you observe the chelate effect? How do you know?
Is the chelate effect enthalpy- or entropy-driven? How do you know?
How does uncertainty affect your conclusions? Specifically address the magnitude of the
errors in the calculated values.
Think about the magnitude and sign of each of your results. Does the sign of ΔS for reaction
AB match your prediction? Was the most favorable reaction the one you thought it would
be?
How do your calculated values (with error) compare to the pooled data you discussed in
lecture section? Describe the differences between your values and “pooled” value and
how these differences arise. What are the advantages/disadvantages of increased data
collection?
Explain any anomalies or discuss any other items you think are noteworthy.
Include a short (~1 paragraph) conclusion, summarizing your findings.
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Phosphorous Acid
Background
Phosphorus lies below nitrogen in the periodic table, and one might expect its chemistry
to resemble that of nitrogen. However, phosphorus is able to expand its coordination sphere in
ways of which nitrogen would never dream. The classic example of this is the violation of the
octet rule found in phosphorus pentafluoride, PF5. This is often explained by invoking the five
empty 3d orbitals on phosphorus but recent research indicates that the d orbitals are not involved.
The ability of phosphorus to exceed the octet rule appears to be due to its larger size (five
fluorine atoms just don’t fit around nitrogen) and to its lower electronegativity.
The lower electronegativity of phosphorous also leads to a very high affinity for oxygen.
You may have learned about the Wittig reaction in organic chemistry because it is a versatile
carbon-carbon bond forming reaction. In it, a phosphorus atom exchanges a “:CHR” fragment
for the oxygen atom of an aldehyde (or ketone):
Phosphorus forms numerous compounds with oxygen and hydrogen and these compounds
display a wide variety of structures and states of aggregation. For example, there are three kinds
of oxygen-containing acid with only one P atom: phosphoric acid, H3PO4; phosphorous acid,
H3PO3; and hypophosphorus acid, H3PO2. [Note the spelling: the element is phosphorus, but the
acids follow the inorganic nomenclature –ic for the highest common oxidation state and –ous for
the next lower; thus the acid you will prepare and study here is phosphorous acid.]
In this lab, you will investigate the structure of phosphorous acid, H3PO3. This acid can
be made by adding water to phosphorus trichloride:
PCl3 + 3 H2O
→ H3PO3
+ 3 HCl
The phosphorus acid has two plausible structures, as illustrated below (the dots in the left
structure indicate a lone pair of electrons):
Modern spectroscopic techniques are of tremendous importance in figuring out the
structures of inorganic molecules, which span a wider range than is found in organic chemistry.
This is primarily a spectroscopy experiment; you will take NMR and IR spectra and interpret
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Chemistry 317 Lab Manual
them. You will use a technique called isotopic substitution to confirm the spectroscopic
assignments. Hydrogen and deuterium are isotopes of the element hydrogen (abbreviated H and
D or 1H and 2H). H and D are chemically very similar but they show up quite differently in
NMR and IR spectra.
To understand your spectra, you will need some basic theory of NMR and IR
spectroscopies (which you have likely seen something about in other chemistry classes; see also
the end of the Introduction to this manual). Understanding the basic principles is particularly
important in inorganic chemistry because of the wide range of elements, bonds, and structures
found. Some reference materials can be found in the lab; we encourage you to read them
carefully and ask questions. It is also important that you learn how to prepare samples and
operate the spectrometers. Finally, you will have to interpret your data. A good way to learn to
interpret IR and NMR spectra of unfamiliar molecules is to look at lots of spectra of similar
molecules. The rest of this section will walk you through the spectra of some phosphorus
compounds.
NMR Spectroscopy
The most common type of NMR spectroscopy is proton or 1H NMR, which you have
studied in organic chemistry. 1H nuclei have a spin of ½. In a magnetic field, these nuclei have
two energy levels, ms = +½ and ms = -½. Proton NMR looks at the energy of the transition
between these two levels. Phosphorus consists exclusively of a single isotope, phosphorus-31
(31P), which also has a nuclear spin of ½. So it has NMR spectra, just like the proton NMR, at
least in theory. Phosphorus NMR spectra have chemical shifts and show coupling, just like
proton NMR. Coupling is observed between two phosphorus nuclei and between a phosphorus
nucleus and a proton. This P-H coupling can be observed either in the 31P NMR or the 1H NMR.
First we’ll consider 1H NMR spectroscopy of phosphorus-containing molecules.
Consider the spectrum of dimethyl phosphite, P(O)H(OMe)2, shown above. The horizontal scale
at the bottom is given in chemical shifts (1, 2, … in ppm). This spectrum was taken on a 90
MHz NMR spectrometer. Since the spectrometer frequency is 90 MHz (90 x 106 Hz), 1 part per
million (1 ppm) = 90 Hz. Recall that Hz is a unit of frequency, seconds-1. In this spectrum, there
is a peak outside the normal range of 0-10 ppm, occurring at 10.6 ppm. On older NMR
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instruments you would access this peak by shifting the region the spectrometer looks at by 90
MHz (10 ppm); such a peak would appear at the top of the spectrum, at 0.6 ppm, labeled 90 MHz
shift. On the much newer Fourier-Transform (FT) NMR instruments we use today, you’ll be
able to set the horizontal scale to whatever you want so you won’t need such a shift.
There are only two kinds of protons in P(O)H(OMe)2, the OCH3 and the PH. There are
six methoxy protons and one PH proton, so two peaks in a 6:1 ratio are expected. But wait –
there are four peaks visible! The two big peaks at 3.7 and 3.9 ppm are in fact a doublet centered
at 3.8 ppm. The 0.2 ppm splitting corresponds to an 18 Hz coupling constant (1 ppm = 90 Hz).
This splitting is the coupling between the phosphorus nucleus and the methoxy protons. The
symbol for coupling is J and this coupling is described as 3JPH. The 3 refers to the fact that the
phosphorus is three bonds away from the methoxy protons: P-O-C-H. The PH refers to the fact
that this is coupling between P and H. The remaining two signals are the small peaks at 2.8 ppm
and at 10.6 ppm. These peaks are actually a doublet, centered at 6.7 ppm (halfway between 2.8
and 10.6), with the enormous coupling constant of about 700 Hz! (1JPH = 700 Hz). This doublet,
due to the PH proton, has 1/6th the area of the methoxy doublet. The one-bond coupling is much
larger than the three-bond coupling because the interaction is larger when the nuclei are closer.
Take a look at the next spectrum (below), which shows phenylphosphinic acid
(C6H5)P(O)H(OH). The aromatic protons appear, as is typical, in a mess between 7 and 8 ppm.
Now there are two peaks beyond δ = 10 ppm; one at 12.8 roughly twice the integral of one at
10.8. The one at 10.8 has the same integral as the one at 4.2, indicating this to be the doublet due
to the PH proton (what are the chemical shift, δ, and the coupling constant 1JPH?). The large peak
at 12.8 ppm is a typical POH peak. It is not split into a doublet by the phosphorus two bonds
away because the proton is very acidic and exchanges very rapidly with other POH protons with
the effect that the coupling is washed out. At very low temperatures, you might be able to slow
the exchange rate down, but at room temperature the POH proton is rather like a carboxylic acid
proton.
Now let’s focus on phosphorus NMR (31P NMR), where we look at the phosphorus
nuclei. The chemical shift range for phosphorus nuclei in different magnetic environments is
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much larger than that for protons, ranging over several hundred ppm (proton chemical shifts
usually range from 0 to 15 ppm). The reference material for 31P NMR is 70% H3PO4; its single
peak is arbitrarily called “0 ppm” (TMS, Me4Si, is assigned as zero for proton NMR). TMS
would of course not give a signal, as it has no P atoms!
Just as we saw P-H coupling in the proton NMR, we also see it in the 31P NMR spectrum
as well. The origin of this coupling is that the “spinning” proton is a little magnet. As noted
above, quantum mechanics tells us that the spin ½ proton can have two possible orientations (ms
values): +½ or –½. The magnetic field felt by the phosphorus atom will depend on the spin state
of the neighboring protons, so the place where you see the signal will be different – the
difference between the two signals is the coupling constant J, which is the energy of the
interaction. Since a proton can have an ms of +½ or -½, the phosphorus signal is split into two
peaks, a doublet. The coupling between P and H that you see in the phosphorus NMR is exactly
the same as the coupling between H and P that you see in the proton NMR because it is the same
interaction energy. Seeing the same coupling constant in both the 1H and 31P spectra confirms
that the coupling is a JPH.
Now what happens if we use deuterium instead of hydrogen? A deuterium nucleus, called
a deuteron, has a nuclear spin quantum number of 1. When it is attached to a phosphorus atom,
the 31P NMR really gets interesting. (At the same time, the proton NMR gets pretty dull, as there
are no protons!) A nucleus with a spin of 1 can have 3 orientations in a magnetic field: +1, 0, or
–1. Any phosphorus nucleus nearby will feel one of three effective magnetic fields. So the 31P
signal will appear as a 1:1:1 triplet. The coupling constant (JPD) is the distance between two
neighbors in this triplet. The magnitude of the coupling constant between deuterium and any
other nucleus is less than the equivalent hydrogen coupling constant, a factor of 6.51 smaller
(this is the ratio of the nuclear gyromagnetic ratios for H and D). If the P-H coupling constant is
651 Hz, the P-D coupling constant will be close to 100 Hz.
Infrared (IR) Spectroscopy
IR spectroscopy involves vibrations of a molecule, rather than the nuclear spin flips seen
by NMR. The theory of IR spectroscopy (actually the theory of any vibration) says that the
frequency of vibration (ν) for a bond (or spring) connecting atoms (or objects) a and b is related
to the stiffness of the spring and the weight of the objects, according to the following equation:
The stiffness of the bond is represented as kab, the force constant. μ is the reduced mass, a sort
of averaged mass. For a system of two particles with masses ma and mb:
With this in mind, consider the spectra of H2O and D2O below.
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The spectra look similar, just shifted. The highest frequency band (at the left) is due to O-H (or
O-D) stretches. H2O and D2O are chemically very similar, so the force constants are essentially
identical (kOH = kOD). The big difference is in the reduced mass, since D is twice as heavy as H:
For O-H,
while for O-D,
The reduced mass approximately doubles, so the observed bands of D2O should shift to
lower frequency by a factor of about 1/√2 or roughly 0.7. The stretch involving the heavier
isotope occurs more slowly – at lower frequency. This is just as you would expect with balls
vibrating on a spring; heavier balls will move more slowly. Check the spectra of normal and
heavy water to see if the bands shift by the factor of 0.7. The lower frequency band is a bending
mode, corresponding to a scissors motion of the two H atoms.
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Consider now the spectrum of dicyclohexyl phosphine shown below.
The C-H stretches come at around 3000 cm-1, like all C-H stretches. The P-H bond has
roughly the same reduced mass (check this) but it is weaker so it has a smaller force constant (its
spring is less stiff). Thus P-H stretching bands come at lower frequency than νCH, usually
between 2300 and 2400 cm-1.
The spectrum on the next page shows a strong P=O band at about 1250 cm-1. Can you
spot the less intense P-H stretch? Even though the P=O bond is stronger than the P-H (higher
force constant), it has a much larger reduced mass so it appears at a lower frequency. The OH
group attached to a phosphorus atom shows up in about the same place as other OH groups (e.g.,
in water or alcohols), at about 3400 cm-1. The band is often very broad due to hydrogenbonding.
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IR spectra are often quite crowded below 1000 cm-1, as many different kinds of
vibrations show up there, including bending and rocking motions. So you cannot identify the
P-O single bond stretches in this spectrum, which also occur in this region. Assigning bands in
this region for large molecules is very difficult and it not usually attempted without extensive
study.
Method
You will do this lab in groups of two. One member of each pair will synthesize ordinary
phosphorous acid, H3PO3, while the other partner will synthesize phosphorous acid in which the
protons are substituted by deuterium, D3PO3. You will then collect and interpret IR, 1H NMR,
2
H NMR and 31P NMR spectra of the normal and deuterated acid to distinguish between the two
possible structures of H3PO3 shown on page 48.
This experiment will take two laboratory periods. In the first period you will synthesize
the phosphorous acid. In the second period you will obtain the IR and NMR spectra (see page
108).
Safety and Waste Disposal
PCl3 (phosphorus trichloride) is a strong Lewis acid. When it encounters moisture it
gives off acidic fumes (HCl) which can burn your lungs and eyes. You are, of course, wearing
your goggles, aren’t you? Do you remember where the eye wash station is? Refresh your
memory before handling PCl3. Small spills (a few drops) and the residue left in your syringe
should be washed away with sodium bicarbonate solution (NaHCO3 aq); when it stops fizzing,
there is no more acid to worry about and you can flush it down the sink or soak it up with paper
towels. If you spill a lot, alert your TA.
Needles and other metal “sharps” require special disposal. Please rinse the needles with
bicarbonate solution and then put them in the Red “Sharps” container (not the glass waste!). The
plastic syringes are disposable, after rinsing. No glass or needles should ever go in the trash!
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Experimental
Synthesis of Phosphorous Acid or Deutero-phosphorous Acid
Connect a 50 mL Schlenk flask to the Schlenk line. Add a KOH trap after the Nujol
bubbler at the end of the nitrogen manifold, as shown below. Use ~ 10 pellets of KOH and add
enough water to cover the entry tube. This will catch most of the HCl generated. This trap is not
perfect, however, so you should vent it to the back of the fume hood. Then the HCl that makes it
through the trap won’t get into the lab atmosphere that you have to breathe.
Add a magnetic stirring bar and 20 mL CH2Cl2 to the flask. Flush the flask with nitrogen
for a few seconds and close it with a rubber septum. Purge a syringe by pulling nitrogen into it
through the septum. This should be done with the flask open to the nitrogen line. After two or
three cycles, leave about 2 mL of N2 in the syringe.
While working in the hood inject the nitrogen in the syringe into the septum-sealed
container of PCl3 and withdraw 1.60 mL of the liquid (2.52 g, 18.4 mmol). Inject the PCl3 into
the dichloromethane in the Schlenk flask through the septum. In another syringe, obtain 1.0 mL
of distilled water, or D2O if you’re making the deuterated phosphorous acid. Note that you need
to treat D2O as an air sensitive material, because it will mix with H2O in the air, diluting the
isotopic purity. With the Schlenk flask containing PCl3 open to the nitrogen line, pierce the
septum with the syringe and then slowly drip the water into the stirred, chilled solution of
phosphorus trichloride. HCl will be evolved and will flow out into your nitrogen line, where it
will be swept through the Nujol bubbler and mostly trapped by the KOH. Once the addition is
complete and the reaction has moderated, remove the ice bath and stir the reaction at room
temperature for about half an hour.
Now turn the stopcock on the Schlenk line to the vacuum, carefully evacuate the flask,
and pull off all the volatile compounds. Be sure your trap is full of liquid nitrogen so none of the
corrosive HCl gas gets to your vacuum pump. If the solution is not too thick, you should stir it
(with the magnetic stir bar) to avoid bumping. You should eventually get a solid. If you get an
oil, try adding some fresh CH2Cl2 and pumping it off. Once the dichloromethane has been
removed, place the flask in a warm water bath (~60 °C) and continue pumping on it for the rest
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of the period, to remove most of the remaining water. When the end of the period draws near,
fill the flask with nitrogen and replace the septum with a glass stopper, against a counterflow of
nitrogen. Remember to grease the stopper so it won’t get stuck. Shut the stopcock on the
sidearm and remove the flask from the Schlenk line. Make sure you have your name somewhere
on the flask, as it will be placed in a vacuum desiccator along with many other flasks. Put the
flask in the desiccator and open it to the dynamic vacuum so that it can be pumped on until the
next class. This is to get the phosphorous acid product dry as possible.
Spectroscopy
Phosphorous acid and deuteron-phosphorous acid are hygroscopic – they suck water out
of the air. If you have time, prepare your IR samples in the drybox. Previous 317 students have
said that making the IR samples out in the air is fine if you work quickly to avoid contamination
with water; make your own decisions. Assemble a dry agate mortar and pestle, Nujol, and two
salt plates (for a Nujol mull). Pulverize some of the phosphorous acid in the mortar, maybe 20
mg; if it is an oil rather than a solid, smear some into the mortar with a clean glass rod or pipette.
Mix with a tiny amount of Nujol oil to prepare a Nujol mull; this should be a thick and
homogenous paste. In this experiment, you should make this mull as concentrated as possible in
order to see the bands of interest, which are weak. Once the acid is mixed with the Nujol it is
fairly stable to atmospheric water. Grind well, for as much as 10 min (!) according to some
previous students, to get the best resolution. Scoop up enough of the mull to coat the center
portion of a salt plate and add the other salt plate to make a sandwich. Don’t squeeze the plates
together too strongly or you’ll push the mull out (or worse, crack the salt plates). These really
are “salt” plates, made of NaCl, so you need to be gentle with them. They crack very easily and
dissolve quickly in water (they also cost more than $10 each).
Record the IR spectrum from 4000-500 cm-1 or thereabouts. Also get an IR spectrum of
pure Nujol, so that you can identify the peaks in your spectrum just due to the Nujol. Plot all
your IR spectra in this lab on the same horizontal scale (e.g., 4000-500 cm-1), so that you can just
lay one over the other to identify which bands are different. It would probably be useful to get
expansions of your spectra, for instance 1400-700 cm-1; just be sure to plot the same expansion
for all the different spectra.
Make up NMR samples of H3PO3 in D2O, and D3PO3 in H2O (label your tubes). About
10 mg of sample is all you need for each. Talk with your TA about obtaining 1H, 2H, and 31P
NMR spectra (see page 108).
Data Analysis/Lab report
Make sure your lab partner has copies of all your spectra; each of you should analyze
both compounds. You should turn in, one week after completion of the lab, a packet containing
your and your partner’s spectra and tables of the important bands with your assignments. The
only writing required is a paragraph or two stating your conclusion as to what is the structure of
phosphorous acid, and your evidence for the structure assignment. Be sure to site specific
spectral evidence which either supports or excludes either of the two structures. We’ll discuss
the spectra and their interpretation in the Tuesday morning lab discussion hour, so make sure to
bring your spectra!
IR Spectra: Assign the major bands due to phosphorous acid in both IR spectra (maybe three or
four bands in each spectrum). You don’t need to interpret every wiggle. Use the model spectra
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included above to help you with the assignments. In particular, you should be able to figure out
which peaks in the spectra involve a bond to a proton or a deuteron. These will be the modes
that are very different between the spectra of the H3PO3 and D3PO3 - roughly a factor of 0.7 in
frequency. Can you figure out from the spectra whether the samples were wet? If so where did
the water come from? Make a table with the assignments and the frequencies of the bands for
both H3PO3 and D3PO3, such as shown below.
Observed bands (cm-1)
H3PO3
D3PO3
Assignment
xxxx
yyyy
O-H/O-D stretch
etc.
…
…
NMR Spectra: List all the resonances and coupling constants for each spectrum, and assign all
the peaks. The proper way to do this is to first indicate the chemical shift in units of ppm, then
multiplicity of the peak – a singlet, doublet, triplet, quartet, and so on – and then, if it is a
multiplet, the coupling constant in units of Hz. The coupling constant is separation of two
neighboring peaks in Hz. You may need to get the spectrometer frequencies for the 1H and 31P
spectra from your TA, if they’re not printed on the spectrum. If the frequency is 50 MHz (50 x
106 Hz), then 1 ppm (1 part per million) in the phosphorus spectrum equals 50 Hz. Another good
way to calculate the coupling constant, J, is to use the chemical shift difference between the
peaks (Δδ, in ppm) and the frequency at which the spectrometer is set (f, in MHz):
J = Δδ × f
Note that f will differ depending on the type of NMR (1H, 31P, etc.) and the spectrometer used.
Processing NMR data with MestReC
MestReC is available in the organic study center, the computers across the hall from the NMR
spectrometer, and you can download a temporary copy for free online. There are manuals with
information on processing your data in the organic laboratory workstation. This information
presented here is “extra,” so we are skipping a lot of instructions which you can easily get
elsewhere.
Your data is stored at ftp://phoenix.chem.washington.edu/DPX200/
You will need to navigate through a few folders to find the exact location of your data.
When you are processing your NMR data the entire folder(s) of data must be stored on the local
computer (So, you have to copy the folders from the phoenix server onto your computer before
you can process them).
The important file for processing is the fid or FID file. You can open this from within
MestReC or you can drag & drop it into MestReC. Below are example FIDs from 1H and 2H
spectra:
56
Things to notice about “fid” data:
•
•
“FID” stands for “Free Induction Decay.” The data shown in the above graphs is the current produced inside the detector after your sample was excited. The horizontal axis is Time. We need to manipulate this data in some way to turn it into a more familiar representation of NMR data. How can we do that? The Fourier Transform (or FT) is an operation which converts data from the time domain to the
frequency domain. In other words, we can use the FT to turn our data with time units (a FID)
into data with frequency units (a spectrum). To turn you FID into a spectrum click the
following commands inside MestreC:
“Process” Æ ”Fourier Transform” Æ ”Apply along T1”
Your FID should now look like an NMR spectrum. For example, the FIDs from above now look
like this:
For the phosphorous acid experiment you will be acquiring 1H, 2H, and two sets of 31P NMR
spectra. They may all contain coupling and if you get mixed up about which nucleus you are
probing in a given experiment it will be difficult (impossible) to extract valuable information
from the spectrum. So, remember that if we know the strength of the magnetic field and the
magnetogyric ratio of the pertinent nuclei we can determine the resonant frequency of the
nucleus in question. NMR spectrometers are commonly referred to by their resonant frequency
for the 1H nucleus. The NMR frequencies of 2H and 31P nuclei are different than the NMR
frequency of the 1H nucleus. This is why we need to manually adjust the instrument in between
experiments that probe different nuclei. Below is a table of some important values for NMR
experiments.
Nucleus
Spin
Natural
Magnetogyric
Relative NMR
Abundance
Ratio
frequency
1
H
0.5
99.985 %
26.7519
100
2
H
1
0.015 %
4.1066
15.4
31
P
0.5
100 %
10.841
40.5
The values above will be useful for interpreting your NMR data. If you are ever unsure which
NMR spectrum you are looking at (“Is this data a 2H NMR spectrum or a 31P NMR spectrum?
OH NO I don’t remember!”) you aren’t doomed. Inside MestReC you can determine the
spectrometer frequency (and therefore the nuclear spin transitions being observed) using the
following commands:
“Options” Æ “Properties” Æ “Parameters” tab
You will now see a table containing descriptions and accompanying numerical data. One row
says “Spectrometer Frequency” and then has a number. This is the frequency of the nucleus
being detected in this experiment. Problem solved! (Keep in mind that the DPX200 which we
used to acquire your data is a “200 MHz NMR Spectrometer”).
If your spectrum has a series of periodic undulating “peaks” on both sides of all your signals, like
this:
Then your spectrum has “sinc wiggles.” Sinc wiggles are an artifact of applying the FT. They
can be removed (at the expense of signal intensity) using the following commands, and starting
over from the FID:
Option 1:
“Process” Æ ”Fourier Transform” Æ ”Set Functions” Æ check the box next to
”Exponential” and adjust the numerical value (You want to use the minimum value which
eliminates the sinc wiggles) Æ “Ok” Æ ”Apply along T1”
Option 2:
Perform the FT as described at the beginning of this document, then, “Process”
Æ ”Smooth” Æ ”Ok”
BF3ŠNH3
Chemistry 317 Lab Manual
The Lewis Acid-Base Adduct BF3.NH3
Background
Boron trifluoride, BF3, is a canonical (prototypical) Lewis acid. In it, boron forms only
three bonds, leaving a vacant orbital that it could use to share a pair of electrons donated by
another molecule. Ammonia, NH3, is likewise a canonical Lewis base, with a pair of electrons
available for sharing. Not surprisingly, they combine to form a stable adduct which is
paradigmatic of the Lewis acid-base interaction:
H
F
F
B
F
+
:
N
F
H
H
F
F
H
B: N
H
H
In this lab you will prepare this adduct simply by mixing boron trifluoride and ammonia.
There is one unusual technical twist involved here: the reagents BF3 and NH3 are gases at
ambient temperatures and pressures. They are also both rather hygroscopic and need to be
protected from atmospheric moisture. (The adduct, though, is an air- and water-stable solid.)
The Schlenk techniques you have learned for protecting compounds from air and moisture need
to be adapted to work with gases⎯in particular it does no good to have an inert gas around. You
will instead use a set of gas-handling strategies known as vacuum-line techniques.
Schlenk methods use differences in pressure to move liquids around, for instance pushing
solutions through a frit or a cannula. In contrast, vacuum-line methods use differences in
temperature to move gases around. Say flask A shown below was filled with BF3 and you
wanted to transfer the gas into flask B, and both flasks were attached to the vacuum manifold of
your Schlenk line.
Schlenk vacuum manifold
Flask A
Flask B
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First you evacuate flask B and the vacuum manifold by opening them to the vacuum pump.
Then close the stopcock between the vacuum manifold and the trap, to isolate the manifold and
leave both the manifold and flask B under static vacuum. Static vacuum means that nothing is
pumping on them (the opposite, dynamic vacuum, means that it's open to the pump). You need
to isolate the vacuum manifold so that you wouldn't lose any gas in the system by pumping it
away.
Now when you open the stopcock on flask A, the BF3 gas in there will simply expand
until it fills the manifold and both flasks at some pressure⎯recall PV = nRT. How can you get
the gas to all go into flask B? The only way is to stop it from being a gas, to condense it.
Cooling flask B in a liquid nitrogen bath at 77 K (-196 °C) will freeze the BF3 in flask B, turning
it into a solid. As the BF3 condenses, the pressure in the flask drops and more BF3 flows in, and
within a short time all the BF3 has condensed in the flask and⎯voila!⎯manifold and flask A
have no more gas. You can follow the progress of this transfer if you keep a close eye on the
vacuum gauge (or mercury manometer) on the vacuum manifold.
You have successfully "vacuum transferred" the gas from flask A to flask B, using a
temperature gradient. Effectively, this vacuum transfer is a reduced-pressure distillation through
the vacuum line, and it works well for substances with a substantial vapor pressure at room
temperature (gases and volatile liquids) that can be condensed using a convenient cryogen
(typically dry ice/acetone at -77° C or liquid nitrogen). The whole while, the chemicals are
protected from reaction with oxygen or water by an inert "atmosphere"⎯not nitrogen gas, but
vacuum!
You can now close the stopcock of flask B and let it warm to room temperature. One
warning: you must be sure that flask B is large enough to hold the gas without building up a
pressure greater than one atmosphere inside. If you do build up pressure, something will pop.
Letting the flask warm up is one example of heating a closed system, which should be done only
with great caution. This applies to heating a gas or solution at room temperature, or (as in this
case) letting a flask or trap warm from low temperature.
You will characterize the BF3.NH3 adduct by IR spectroscopy. When the adduct is
formed, the empty p orbital on the boron atom receives a dose of electron density from the
nitrogen atom’s lone pair, which has two major effects on the BF3 unit. First, the three fluorine
atoms bend away from the nitrogen; this changes the number of boron-fluorine vibrations that
quantum mechanical selection rules allow us to see with IR spectroscopy. Secondly, the boronfluorine bond strength changes, so the observed frequency of B-F vibration shifts. One can
compare IR spectra of boron trifluoride, ammonia, and the adduct and see all these spectral
changes.
The comparison between the IR spectra is complicated by the different appearance of IR
spectra of gases vs. condensed phase samples (solids or liquids). The difference has to do with
their rotational motion: molecules in gases rotate freely unlike molecules in solids and liquids
which bang into each other. Quantum mechanics says that a rotating molecule cannot assume
just any old rotational angular momentum⎯it has more-or-less regularly spaced rotational
energy levels. The energy of a bond vibration might be around 1000 cm-1, while the energy
required to boost a molecule from one rotational energy level to the next is much smaller, on the
order of 0.1 - 10 cm-1 (recall that the average energy at room temperature is 207 cm-1). The net
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result of all of this quantum mechanical stuff (that you may have seen in Pchem) is that a
vibration of a gas phase molecule will appear in the IR spectrum not as a single band but as a set
of lines⎯a broad envelope of more-or-less regularly spaced spikes. Each spike results from a
transition that has both a vibrational and a rotational component (e.g., a molecule in the ground
vibrational state and the tenth rotational state goes to the first vibrational excited state and the
ninth rotational level).
Two such sets of lines are sketched below, one with a central spike (b) and one without
(a). You measure the vibrational frequency at the spike or the center of the open area. The
presence of the central spike (called a Q branch) depends on the symmetry of the molecule.
Spectrum (a) is plotted as an absorption spectrum (like a UV/Vis spectrum) while (b) is plotted
as % transmittance through the sample, as is more typical for IR spectra. In (a), the spacing of
the lines and the resolution of the spectrometer is such that individual lines are observed, while
in (b) only the overall envelope is seen because the lines are so close together. The distance
between the lines is determined by the moment of inertia of the molecule about a rotational axis,
such that widely spaced lines (case (a)) are observed only for small molecules with low moments
of inertia. NH3 has a low moment of inertia because only the light hydrogens move on spinning
about the nitrogen, while in BF3 the much heavier fluorine atoms must move on rotation.
(a)
(b)
IR spectra of (a) HCl gas at high resolution, plotted in absorbance (peaks go up) and (b) CH3Br
gas at low resolution, plotted in % transmittance (peaks go down).
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Method
You will work in pairs for this experiment, which will take one period. You will fill two
Schlenk flasks with BF3 and NH3 gases from gas cylinders; the pressures you measure on the
digital vacuum gauges and the volumes of the flasks will allow you to determine the quantity of
gases. You will also take gas-phase IR spectra of these two reagents. You will make solid
BF3.NH3 by vacuum transferring NH3 gas into a flask containing BF3 at liquid nitrogen
temperature, and allowing the BF3/NH3 mixture to warm to room temperature. Comparison of
the IR spectrum of the adduct with that of the starting materials may shed some light on the
changes in bonding that occur when the adduct forms.
Procedure
0. SAFETY!
Please review the section in the Introduction of this manual on gas handling (pp. 20-21).
Do not open the lecture bottles unless the TA is working with you. It is imperative that you deal
with these gases safely, as both BF3 and NH3 are toxic. There have been several incidents in this
lab that vented a significant amount of BF3 into the air and the lab had to be evacuated.
The key to safe handling of gases is to be sure that pressure does not build up in any
portion of your apparatus. Before you let any gas in, trace the path the gas will take. Make sure
that the gas can flow through to the manometer and vacuum gauge (i.e., check that all the
stopcocks that you think are open actually are open). One of the incidents noted above involved
opening the BF3 when the stopcock on the Schlenk line was closed. Pressure built up in the
rubber hose until the hose ruptured. Another incident also resulted from too much pressure in
the Schlenk line. In that case, the students did not wait long enough for the digital gauge to
respond before adding more gas. To avoid such problems, you should let only a small amount of
gas in at first. Know where the gas is going, and check that you see the gas entered your line by
looking for a response on the gauge. Be sure to wait for the gauge to respond.
Both of the gases have built-in leak detection characteristics. BF3 reacts on contact with
the moisture in air, forming a white fog of boron oxides and HF (quite toxic); this is not a good
thing to inhale. NH3 doesn’t leave visual evidence of a leak, but it has a characteristic odor with
which you should be familiar. If you see white fog seeping from some part of your apparatus or
smell a strong ammonia smell, close the main valve on the lecture bottle and alert a TA.
I. Transfer of gases from cylinders to flasks
*The procedure for filling a flask and the IR cell with NH3 is identical to that for BF3. If the
BF3 cylinder is in use, by all means do the ammonia first. Remember to use separate Schlenk
flasks for each gas, though! Also, use a slight excess (in pressure) of ammonia over boron
trifluoride. At the end of Part I, you should have two closed, gas-filled flasks.
Before you begin, your Schlenk line must be brought up to specifications. A good static
vacuum must be demonstrated to your TA and approved before you begin using them to
move these highly toxic gases.
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To begin, connect the cylinder of BF3, a 250 mL Schlenk flask with a glass stopper, and a
gas-phase IR cell to the Schlenk line, as shown on the following page. Be sure to use the IR cells
with the proper gas written on the label. The cells can contain traces of gas from previous users
and if you combine the two different gases, they will react. You will then have the unenviable
task of cleaning the IR cell. The hose connecting the BF3 to the line should be clamped on, to
help prevent leaks. Also, all of the stopcocks should be clamped if possible⎯we have had
several stopcocks that were not clamped become airborne during the course of this lab.
Schlenk vacuum manifold
BF 3 cylinder
gas phase IR cell
Flask A
We have two BF3 cylinders, two NH3 cylinders, and four gas phase IR cells in the lab (two for
each gas). Try to be flexible with your procedures as you will need to share these reagents and
gas cells. Stay in touch with the TA(s).
Open the vacuum manifold to the pump, then open all the stopcocks in the setup: the
three on the vacuum manifold, the one on the IR cell, and the one on the Schlenk flask. Make
sure the main valve on the cylinder is closed but that the needle valve is open. Always use the
vacuum manifold of the Schlenk line for reagent gases, because that manifold has the vacuum
gauge on it. Let the system pump out for a few minutes, to reach to lowest pressure inside that
the vacuum pump can attain. Now test to see if there are any leaks. Close the stopcock that
connects the vacuum manifold to the trap and pump. Watch the level on the vacuum gauge: if it
increases at an appreciable rate, it means that air is leaking into the system from somewhere.
You can close various stopcocks to see if you can isolate the trouble; poor hose connections are
the most likely sources of leaks. If you need help finding or fixing your leaks, ask a TA. If the
gauge is steady with the manifold cut off from the pump, then your system is not leaky and you
are ready to admit the BF3. Write down the reading on the vacuum gauge that corresponds to
full vacuum.
CAUTION: Boron trifluoride will corrode your rubber hoses if it is left too long in contact with
them, so work as quickly as possible. (It's also corrosive to lungs.)
With the vacuum manifold and flasks etc. all under vacuum but not open to the vacuum
pump, you're ready to admit the BF3 gas to the manifold+flasks system. Close the needle valve
and then open the main valve on the lecture bottle. (Do this carefully⎯every so often you will
have a leak and spew gas into the lab. If there are any problems, close the main valve, close your
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hood sash and get your TA.) Check to be sure that the valve on the Schlenk line is open so that
the gas has a place to go. Open the needle valve a little for an instant, and then close it. This
should let a small pulse of gas out. Watch the vacuum gauge on the Schlenk line when you're
fiddling with the valve. The reading on the gauge should go up when you let gas in, then stop
(you have no leaks, remember?). Make sure to wait for a response from the gauge. This can
take a little time. If what you anticipate does not happen, shut the valve off and think through
what you are doing. If you can't quickly figure out what is wrong, get your TA. If you're getting
gas, you're ready for business.
Let in the gas in small puffs, opening the needle valve briefly and then closing it. Wait in
between puffs to see a change on the digital vacuum gauge. One partner can work the needle
valve while the other (both?) watch the vacuum gauge. Let in roughly 400 torr of gas. It doesn’t
matter if you get exactly this amount in the system, as long as you record the exact amount you
do let in. At this point, close the stopcocks to the gas IR cell and the Schlenk flask to imprison
the BF3 therein. Then open the stopcock to the pump, which sweeps out the unused BF3 away
from the rubber hosing into the liquid nitrogen-cooled trap. Keep the trap full of liquid nitrogen
so that the BF3 is efficiently trapped and not pulled through to corrode the pump. Once all the
BF3 has been pumped away from both the line and the rubber tubing, you can detach the cylinder
and let someone else use it.
The amount of BF3 in the IR cell is higher than is desired for best resolution. Close all
the stopcocks on the manifold except the one leading to the IR cell. Close the stopcock
connecting the manifold to the pump and open the IR cell to the manifold. We think that the best
resolution and intensity in the IR spectrum is obtained for pressures in the cell of roughly 30-50
torr. You may need to pump a little of the BF3 away to get to this pressure range (don't pump it
all away!). Re-close the IR cell stopcock and pump the excess gas in the vacuum line into the
cold trap. After all the BF3 has been removed, you can disconnect the cell from the line and take
the IR spectrum. When you’re finished, reconnect the cell to the line and pump out all the BF3
in it. Please be quick with the IR cell so others can use it.
Repeat the above procedure using NH3.
II. Reaction of boron trifluoride with ammonia
At this point you should have two closed gas-filled flasks connected to the vacuum line
as shown below. Keep the stopcocks on the Schlenk flasks closed until you want to actually
transfer the gases (or you’ll pump away your material). Evacuate the system and test for leaks as
before. Once you’re satisfied everything is leak-free, put the manifold and hoses under dynamic
vacuum and cool the flask with the BF3 in a liquid nitrogen bath. You should see a white solid
condense on the walls of the flask (what is it?). When the flask is cold, open the stopcock on the
BF3 flask to dynamic vacuum as well, to remove any air that might have leaked in (why doesn’t
this pump away the BF3?). Once vacuum is established within the flask and manifold, close the
vacuum manifold off from the pump. The manifold and hoses are now under static vacuum.
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Schlenk vacuum manifold
BF 3
NH 3
You are now ready to transfer the NH3. Open the stopcock on the cold BF3 flask, then
the stopcock on the NH3 flask. The pressure in the line should rise as NH3 flows into it, then fall
as it condenses in the cold BF3 flask. Once you have judged the transfer to be complete, pump
any gases remaining in the line into the cold trap. You can leave the stopcock on the cold flask
open for this; these condensed gases have a negligible vapor pressure at -196 °C and won't go
anywhere. Close the stopcock on the cold flask⎯be sure the stopcock is clamped⎯and remove
the liquid nitrogen bath. Let it warm slowly to room temperature; record your observations.
Once the flask has come to room temperature, open it to the vacuum line and pump out any
unreacted gases. Then scrape out the solid, weigh it, and take an IR spectrum (preferably as a
KBr pellet but a Nujol mull is OK).
CAUTION: Your trap will contain both BF3 and NH3 at the end of the experiment. Be sure to
keep the Dewar full of liquid nitrogen. Let your TA shut down your Schlenk line; he or she will
remove the trap while still very cold, and let it warm up in the back of the fume hood.
Questions
Please turn in your spectra and answers to the questions below one week after completion of the
experiment.
1. Calculate the number of moles of the reagents you used. The ideal gas law is appropriate for
this calculation. When you measure out solids or liquids for a reaction you have to divide the
number of grams you used by the molecular weight to get moles, but in this experiment you
just use roughly the same pressure of both reagents. Why?
2. What was your yield of product? Where did the rest of it go?
3. Assign your spectra and prepare a table of IR data with peak positions and assignments. You
will likely want to consult with the excellent reference book by Nakamoto, Infrared and
Raman Spectra of Inorganic and Coordination Compounds. Spectra for both BF3 and NH3
appear in here, in the section on XY3 molecules in the earlier portion of the book. What do
the IR spectra tell you about BF3.NH3?
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Arene Molybdenum Tricarbonyl Chemistry
Background
The metal compounds you have dealt with in this class so far⎯and in Chem 165 or 312
or 416⎯all formally contain metal cations that act as typical Lewis acids. The metals bind
typical Lewis bases such as water, ammonia, chloride, and acetate. Such metal complexes have
been a staple of inorganic chemistry for well over a century, with some of the compounds known
since the days of the alchemists. A new field known as organometallic chemistry, has developed
in the last 50 years into one of the major areas of chemistry. As the name suggests,
organometallic chemistry is the chemistry of metals bound to organic groups, mostly compounds
with metal-carbon bonds. These organic ligands include alkenes and alkynes, arenes and other
aromatic compounds, alkyl and aryl groups, and carbon monoxide.
The organometallic field has blossomed both because it is new and interesting chemistry,
and because it can be very useful, in a practical sense. In both the organic laboratory and the
industrial plant, organometallic reagents and catalysts have led to a number of new processes.
For example, acetic acid, a major industrial chemical, is now made in large part by carbonylation
of methanol using an organometallic rhodium catalyst. This process (in the reaction shown
below) was developed by what was then called the Monsanto Corporation.
O
CH 3 OH + CO
CH 3 C
[Rh(CO) 2 I 2 ]OH
I - co-catalyst
A number of polymers are made using organometallic catalysts, for instance polyethylene from
ethylene, so called Ziegler-Natta polymerization:
n H2C = CH2
n-2
TiCl2 , AlEt 3
The pentagons with ovals in them stand for the cyclopentadienyl anion, C5H5 , which is aromatic
just like benzene (and is a very good ligand for most transition metals). Billions of dollars are
currently being invested in alkene polymerization plants using the newest transition metal
catalysts. You can learn more about organometallic chemistry in the course on this topic, Chem
417.
Many of the organic ligands found in organometallic chemistry are not classic Lewis
bases, for instance the arene and carbon monoxide (CO) ligands you will use in this experiment.
Even carbon ligands that are formally anionic, such as CH3– form quite covalent bonds with most
metals and are not well described by M+←CH3–. So the bonding and the nature of
organometallic compounds is in some ways quite different from the bonding in "classical"
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coordination compounds. Because the compounds are fairly covalent, they resemble organic and
main group compounds. Instead of the octet rule for main group compounds, organometallic
compounds usually obey an "eighteen electron rule." Just as the octet rule arises from the filling
of the four available orbitals (one s + three p orbitals), metals which have d orbitals want to have
a total of 18 valence electrons in nine orbitals (one s + five d + three p orbitals).
The molybdenum compounds you will work with in this lab are good examples of the
eighteen electron rule. Your starting material molybdenum hexacarbonyl, Mo(CO)6, has six
carbon monoxide (CO) ligands bound to the molybdenum. Each CO has a lone pair of electrons
that it uses to bind to the molybdenum, in other words each CO contributes two electrons toward
the molybdenum's goal of eighteen. The molybdenum itself brings six valence electrons, so the
total electron count is
6 from Mo + 6 × (2 per CO) = 18 total electrons about Mo.
The reaction you will do in the first part of the lab is displacing three of the carbon monoxides by
an arene ligand, mesitylene:
O
C
OC
OC
CO
Mo CO
+
reflux
C
O
Mo
O CC
O
+ 3 CO
CO
The product has a so-called "three-legged piano stool" structure as drawn above. The fact that all
six aromatic carbon atoms are bound to the molybdenum is indicated by inserting η6 in the
name: (η6-C6H3Me3)Mo(CO)3 [η is the Greek letter eta, which rhymes with beta]. The bonding
within the arene is delocalized, with equal C–C distances. But we can think of the arene as being
composed of three double bonds, each of which has two π electrons. Thus an arene is sort of a
chelating tris-alkene ligand, which can donate six electrons to the molybdenum. Therefore, one
arene displaces three CO ligands and the compound is still eighteen electrons:
6 from Mo + 6 from the arene + 3 × (2 per CO) = 18 total electrons about Mo.
As implied above, arenes and CO act as electron pair donors (Lewis bases) in their
bonding to a metal such as molybdenum. This σ bonding interaction is shown as A on the next
page: the lone pair on carbon monoxide donates into an empty d orbital on the molybdenum.
This σ bond is just like the donation of water or ammonia to a metal center. But there is an
additional interaction, made possible because CO has an empty π-antibonding orbital located
mostly on the carbon (B). The molybdenum can donate its electrons back to the CO, into this
orbital. This empty orbital on CO acts as a Lewis acid⎯it accepts a pair of electrons⎯in a π
bond. For this reason, CO is called a "π-acid" ligand. Arenes and alkenes and many other
ligands have similar interactions.
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A.
B.
C
O
C
Mo
O
Mo
π backbonding
π donation from Mo to CO
σ bonding
σ donation from CO to Mo
This π-backbonding has a number of critical effects, as explained in Chem 417. One
feature is that it sucks electron density off of the metal, and therefore stabilizes low metal
oxidation states. In the compounds in this experiment, molybdenum is formally in a zero
oxidation state! With ligands that cannot π-backbond, like water or ammonia, molybdenum is
only found in oxidation states +2 and higher.
Another key feature of this π-backbonding is its effect on the carbon monoxide. The
π-backbonding interaction puts some electron density into a CO antibonding orbital, which
weakens the C–O bond. This is most evident in the CO stretching frequency. The C≡O stretch
in free carbon monoxide gas occurs at 2143 cm-1, while in Mo(CO)6 this IR band is at lower
frequency (you'll determine the value in the lab). When three CO ligands are displaced by the
arene, the amount of backbonding changes and the CO frequency shifts. Since these shifts are
substantial and the stretches are quite intense, the CO bands in the IR spectrum provide an
important handle on what is going on. The development of metal carbonyl (M-CO) chemistry
has relied heavily on IR spectroscopy. Mesitylene is a poorer π acid and a better σ donor than
CO (almost all ligands are), so the reaction you're doing makes the molybdenum more electron
rich. This should cause the stretching frequencies for the remaining CO ligands to shift to lower
frequency.
The conversion of Mo(CO)6 to (η6-C6H3Me3)Mo(CO)3 also changes the symmetry at
the molybdenum center, from octahedral to three-fold symmetry. The symmetry at the metal
determines, via quantum mechanical selection rules, how many CO stretches you will see in the
IR spectrum. Although Mo(CO)6 has six carbonyl ligands, there is only one CO stretching mode
in the IR. The three CO ligands in (C6H3Me3)Mo(CO)3 should give rise to two IR bands. Thus
both the position and the number of CO bands in the IR provide useful information.
Method
Four lab periods have been allocated for this experiment, in two blocks. In the first
block, you and your lab partner will use one lab period to react Mo(CO)6 with mesitylene and
isolate the product. You will also characterize the product by IR spectroscopy and by NMR. In
the second block, three periods have been left open for you and your partner to choose what you
want to do, based on your interests, the papers in the 317 folder and additional inspiration you
find in the library. You might want to explore the reactivity of the (η6-C6H3Me3)Mo(CO)3 you
made or you might make something else from Mo(CO)6. We expect that you and your partner
will do two reactions (of 4 proposed) of your own development (or more if you have time and
interest).
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Procedure
0. Safety
Carbon monoxide is a highly toxic, odorless and colorless gas. Since reactions of metal
carbonyls often give off CO, they should be done with care and only in a fume hood or on a
well-vented Schlenk line. Metal carbonyl compounds themselves are often highly toxic,
especially volatile compounds such as Ni(CO)4 (boiling point 43 ˚C). Mo(CO)6 is not so volatile
at room temperature and therefore does not present a serious danger, but it should be handled
with care in a hood.
1. Preparation of Mesitylene Molybdenum Tricarbonyl, (C6H3Me3)Mo(CO)3.
Adapted from Angelici, R. J. Synthesis and Technique in Inorganic
Chemistry University Science Books, Mill Valley, CA 1986 pp.
129-146.
Put 2.0 g of Mo(CO)6 (0.0072 mol), 10 mL mesitylene and a
stir bar in a 100-mL round-bottom Schlenk flask. Attach a
condenser to the flask and connect to a mineral oil bubbler, as
shown in the drawing at right. The bubbler should set up into the
hood so that you will know that the CO gas evolved is making it to
the hood. Connect the side arm of the Schlenk flask to your
Schlenk line.
Mo(CO)6 is volatile enough that it will sublime out of the
reaction flask upon heating. You should let the mesitylene vapors
reflux high into the condenser to wash the Mo(CO)6 down into the
reaction flask. Therefore, don't cool the condenser with water⎯the
room temperature will be adequate to condense the mesitylene,
since it boils at such a high temperature (b.p. 165 ˚C). Place the
reaction apparatus in a heating mantle appropriate for the size
Schlenk flask you are using.
Mo(CO)6 and (C6H3Me3)Mo(CO)3 are both pretty stable to
air at ambient temperatures, but would react with atmospheric
oxygen and/or water at elevated temperatures. For this reason you
need to conduct the reaction under an inert atmosphere. This could be done by preparing an all
sealed system, by evacuating a flask and closing all the stopcocks. But great care should be
taken before ever heating a closed system⎯the system has a tendency to build up pressure and
blow up. (In this reaction, you're liberating three equivalents of CO gas, in addition to the
normal increase in vapor pressure of the solvent and the gas expansion on heating from 25 to 165
˚C.) So it's much better to do this reaction open to your oil bubbler, which keeps everything at 1
atmosphere pressure. First you need to flush the apparatus with a moderate to fast flow of N2 for
ca. 5 minutes. To get the nitrogen to flow through your apparatus, you will likely need to
prevent the nitrogen from flowing out the bubbler at the end of the nitrogen line. Double check
to make sure that nitrogen is flowing through your apparatus to your new bubbler. Once the
setup is flushed, turn off the nitrogen, close the stopcock (why in that order?), and turn on the
heating mantle with the variac.
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Reflux at a moderate boil for about 30 minutes. Note what is happening in the reaction as
heating occurs. Remove the heating mantle and quickly turn on a slow nitrogen flow to prevent
the mineral oil from being sucked back into the apparatus as the mesitylene vapors condense.
The nitrogen flush may also carry unreacted Mo(CO)6 away from your product to deposit in the
condenser. This is good because it purifies your product, but too much of a flow will carry the
Mo(CO)6 into your vent line, which you don’t want. When the solution has cooled to room
temperature, turn off the nitrogen flow and take the apparatus apart. Add 15 mL of hexane to
complete precipitation of the yellow product (which is much less soluble in hexane). Isolate the
product by suction filtration and rinse with 5 mL hexane. Dry the product by pulling air through
the material on the frit. Weigh the material to obtain a yield.
If your product is not spectroscopically pure, you may want to purify it. You will need a
pure product for accurate analysis of the IR and NMR spectra. You can recrystallize it by
dissolution in a minimum of CH2Cl2 (~15 mL), filtration, and reprecipitation by adding 25 mL
hexane. If necessary, a second batch of product can be obtained from the initial filtrate by
pumping away some of the solvent on the Schlenk line and addition of more hexane. Some 317
students reported that this recrystallization can be done out in the air, while others took
precautions to exclude air. Other workers have found that recrystallization made their product
worse. Use your judgment and let us know what you thought was necessary. The product can be
further purified by sublimation at approximately 120 ˚C. Solid (C6H3Me3)Mo(CO)3 will
decompose over a period of weeks in the presence of light and air, so it should be stored in the
dark in the drybox (perhaps in a vial wrapped in aluminum foil). Characterize your product by
taking its IR spectrum and its proton NMR spectrum (see Section 3, below).
2. Planning your next step
There is roughly a three-week break before you have three “open” lab periods to do at
least two other reactions in this area, reactions of your own choosing. Some guidance and
suggestions are given below. After the first lab period, write up a preliminary lab report (a
template will be provided (see page 110) + see below) AND a proposal (see page 113) for at
least two additional experiments. (Two from each individual student for a total of 4 proposed
reactions for a pair of students.) If you don’t let us know what you want to work with, we can’t
get it for you or insure that what you want to do is safe.
To guide you in your choices, we have put some papers in the Chem 317 folder in the lab. You
will need to read some of these papers and you will also need to do reference searching in the
Chemistry library or online. Pick something that you think is interesting and reasonable within
the constraints of this course, especially the time available. We recommend the Journal of
Chemical Education as an excellent resource for experiments that are of appropriate difficulty.
We can often find unusual chemicals and equipment; ask your TA or the instructor. The TA’s
will not, however, make choices for you. The Schlenk line and the drybox will be available for
the first two periods.
Examples of the kind of reactions we have in mind are:
(a) Preparation of a compound similar to (C6H3Me3)Mo(CO)3.
(b) Explore the reactivity of (C6H3Me3)Mo(CO)3 with other ligands.
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(c) Explore the reactivity of (C6H3Me3)Mo(CO)3 with an oxidizing agent such as Br2 or I2.
(d) React Mo(CO)6 with a different type of ligand.
(e) Do similar reactions with other metal carbonyl compounds.
A very valuable class of ligands for organometallic chemistry are phosphorus ligands. They are
soft donors with some π-acid character, although much poorer π-acids than CO. We will have
available in the lab triphenylphosphine, PPh3, and trimethyl phosphite, P(OMe)3, in case you
want to use them. Apparatus for various kinds of chromatography can also be made available if
you need to separate product(s) from your reactions.
Before you choose what to do, think about the questions in part II of the lab report description on
the next page. Most importantly, how will you know what happened? This is a key question in
any synthetic procedure, and relates to how you will characterize your products, as described in
the next section.
3. Characterization of (C6H3Me3)Mo(CO)3 and the products of your other reactions.
You should get IR spectra of (i) your starting material Mo(CO)6, (ii) the mesitylene
complex (C6H3Me3)Mo(CO)3, and (iii) whatever materials you obtain in your other reactions.
Please refer to the introduction to IR spectroscopy in the techniques section (pp. 23-25) if you
haven’t done that lab already. This section describes, for instance, how to prepare a sample for
IR as a Nujol mull. IR spectra of molecules in solution are also very useful, as solution spectra
are often better resolved. Solution IR cells will be available (be careful⎯they're very expensive,
fragile, and made of salt, so no water!). Dissolve a few crystals in 2 mL heptane or hexane or
chloroform. Get a blank of the solution IR cell with solvent in it, and then replace that solvent
with your sample solution using a syringe. You should also get spectra of just the solvent using
the empty cell as a blank. Wherever the solvent absorbs strongly you’re not going to get any
information about your compound (why not?). Wash the solution cell by pouring the solution
out (into a waste bottle), and filling it with fresh solvent and decanting at least twice. Dry the
cell with a rapid flush of nitrogen and return it to a desiccator. If any sample gets on the outside
surface of the cell window, rinse it off immediately with chloroform or methylene chloride.
You should also get proton NMR spectra of mesitylene, (C6H3Me3)Mo(CO)3, and your
products from the other reactions. Your TA will advise you on how to make up the NMR
samples⎯something like 10 mg of sample dissolved in CDCl3 (chloroform-d). (Note that
Mo(CO)6 does not give a proton NMR spectrum as there are no protons.) If you use phosphorus
ligands, you will likely want to get 31P NMR spectra as well (see the Phosphorous Acid
experiment, #3, for instructions on obtaining these spectra).
You and your lab partner should divide up the tasks you need to do. One can stay and set
up a reaction while the other is taking NMR spectra. IR spectra can be taken during a slow time
in the lab, for instance while reactions are refluxing. Do a preliminary analysis of your spectra
right after you get them, as this will help guide you in what to do next.
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Lab Report
Preliminary Report for Part I - due one week after the conclusion of the first part of the lab
The basic elements of an Introduction, Results, Discussion, Experimental, and Reference
sections should all be included (see page 111). Follow the formatting guidelines provided for you
in the introduction section of this manual and use the templates from chromous acetate and the
chelate effect as examples. Be sure to include the following items in the appropriate sections of
your report:
•
Background and purpose of the synthesis and subsequent characterization. Include a
minimum of 2 literature references (the 317 manual is NOT literature).
•
Give and discuss your percent yield.
•
Prepare tables summarizing your IR and NMR spectra, including any assignments you
can make. (In addition to appropriate characterization in the Experimental section.)
•
Compare and contrast the complexes you made and/or starting materials. Is there any
starting material in your product? Does it look clean (by IR, NMR, and visual
appearance)?
Proposal for Part II - due one week after the conclusion of the first part of the lab
•
A template for this assignment will be distributed electronically (see page 110).
Please complete this template and attach the requested supporting documents. Note that
this assignment should NOT be attached to your part I report as it may be sent to a
different grader.
Final Report Parts I and II
Return your graded reports and grade sheet from the first part of this lab, in addition to the final
report. Do not include the proposal documents. Refer to the guidelines given in the introduction
of this lab manual for the correct formal report formatting and use the templates from chromous
acetate and the chelate effect as examples.
The format of this report should remain the same as the preliminary report, but will
reflect the syntheses and characterization done in all four lab periods.
Use comments received on the preliminary report and incorporate the following ideas
from part II to form a comprehensive final lab report:
Describe the experiments that you designed. What did you hope to accomplish and did
you succeed? Provide references for any sources you used. Answer, to the best of your
knowledge, the key questions for any synthesis:
i) Was a single product formed?
ii) Is there any starting material left?
iii) What did you make?
iv) How do you know?
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v) Compare and contrast the characterization you obtained with those reported in the
literature. (Remember to appropriately cite this literature in your report.)
A full analysis can probably be done in roughly three to five pages of text, not
including attached spectra and other data. But clarity and completeness are most important, so if
you need to write a longer report you should do so. You will be graded on how well you
approached and analyzed your synthesis, and how clearly you wrote it up. Please don’t do
something just because you think it will be easy. We want you to be imaginative, adventurous,
and to have some fun here.
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Linkage Isomers of Nitro-Pentammine-Cobalt(III)
Background
Coordination chemistry is the study of how ligands (bases) coordinate to (bind to) metal
cations (Lewis acids). Common ligands include water, as in aquo ions [M(H2O)x]n+, ammonia
(NH3), and pyridine ( N
). The nitrogen bases are in general much better ligands than water
because they are better bases. You have already seen in Experiment 2 that adding ammonia to a
light blue aqueous solution of Cu2+ causes rapid formation of the dark purple copper ammonia
complex. A complex with ammonia is called an ammine complex.
Ammine complexes were the first coordination complexes studied in detail, back at the
turn of the 20th century. Alfred Werner and others made a large number of ammine complexes
of cobalt(III), such as [Co(NH3)6]3+. Studies of the chloro-ammine complexes showed that
[Co(NH3)5Cl]2+ exists as one isomer, [Co(NH3)4Cl2]+ has two isomers, and Co(NH3)3Cl3 has
two isomers. This proved (or at least strongly suggested) that these compounds have octahedral
structures. The structures of the two geometrical isomers of Co(NH3)3Cl3 (facial and
meridional) are shown below. [You can find translations of some of Werner’s classic papers in
Classics in Coordination Chemistry, G. B. Kaufman, Ed., Dover, New York, 1968.]
Cl
Cl
H3 N
H3 N
Co
Cl
Cl
H3 N
H3 N
Co
Cl
NH 3
Cl
NH 3
meridonal
facial
Another form of isomerism common in coordination chemistry is linkage isomerism.
This can occur when a ligand has two different basic sites, so it can bind in two different ways.
The oldest recognized example of linkage isomerism involves the nitrite anion, NO2¯. In 1894,
Jørgensen reacted [Co(NH3)5Cl]Cl2 with sodium nitrite, NaNO2, to give a red solution from
which he isolated red crystals. If the red solution was boiled with acid, it changed color and
deposited yellow crystals on cooling. No problem, thought Jørgensen, who doubtless thought he
had made some other cobalt complex. However, the red and yellow complexes had identical
elemental compositions. Jørgensen and Werner proposed, based on the colors of other cobalt
ammine complexes with nitrogen and oxygen donor ligands, that the red compound contains a
“nitrito” ligand, bound to cobalt through an oxygen atom while the yellow compound has a
“nitro” ligand, bound through its nitrogen atom. The isomers [Co(NH3)5ONO]Cl2 (nitrito) and
[Co(NH3)5NO2]Cl2 (nitro) are drawn below. The formula [Co(NH3)5NO2]Cl2 means that there is
a [Co(NH3)5NO2]2+ dication, in which the cobalt is coordinated to (bound to) five ammonia
ligands and one nitro group, and that there are two chloride counterions to make the compound
electrically neutral.
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The nitro ligand can bind either through N or through O because it has lone pairs on both atoms,
as shown by the two resonance forms of NO2-:
In this lab, you will synthesize these linkage isomers the way Jørgensen did, by substitution of
the single chloride ligand in [Co(NH3)5Cl]Cl2. The nitrito complex is the kinetic product of this
substitution⎯it is formed first⎯while the nitro isomer is the thermodynamic product (the more
stable one). You isolate the nitrito and then apply heat to convert it to the nitro complex. You
will follow this interconversion the way Jørgensen did, using the colors of the compounds. The
beautiful colors of many transition metal compounds arise because the compounds absorb some
frequencies of light and transmit others. This is a result of the partial occupancy of d orbitals, as
is described in Chemistry 416. Your eye is a phenomenally sensitive instrument and can
distinguish colors well, but it is not a quantitative instrument. In the laboratory, you will
measure the visible spectrum (also sometimes called the optical spectrum) with a
spectrophotometer. The spectrometer also records spectra in the ultraviolet (UV) region (higher
energy than visible light), so these spectra are sometimes called UV/visible spectra (UV/Vis for
short). They are also termed electronic spectra because the transitions involve changes in
electron configuration. You use a different spectrometer to measure vibrational spectra in the
infrared (an IR spectrometer), on the low energy side of the visible region.
The visible spectra of these compounds are recorded as the absorbance at each
wavelength in the visible region. The spectrophotometer measures the intensity of the light, both
the light going into the sample (Io) and the intensity that comes out (I). Absorbance of light (A)
is measured on a logarithmic scale:
I
A = log o
I
Thus an absorbance of 1 corresponds to 90% of the light being absorbed and A = 2 means that
99% of the light is absorbed. The absorbance is proportional to the concentration of absorbing
molecules and to the distance that the light travels through the solution. The mathematical
statement of these facts, with C standing for the concentration and l the path length (usually 1
cm), is called Beer’s law:
A = εCl
ε is the extinction coefficient and is a property of the molecule, indicating how efficiently the
molecule absorbs a given wavelength of light. If you know the ε at a certain wavelength, you
can measure the concentration by measuring A at that wavelength. Molecules have different
extinction coefficients at each wavelength; your UV/Vis spectrum is a plot of ε vs. wavelength.
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This experiment probes the mechanism of the isomerization of the nitrito to the nitro
compound, by studying the kinetics of the reaction. One can never prove a mechanism, but one
can distinguish between possible options. For instance, consider the two mechanisms drawn
below. Does the cobalt slide across the nitrito ligand from O to N (path A below), or does the
nitrogen atom of the nitrito ligand reach out and grab another cobalt before the oxygen atom lets
go of the first cobalt (path B)? In these drawings, the square brackets indicate intermediates that
are not observed or transition states, marked with ‡.
A goal of any mechanistic study is to figure out what is the rate-determining or ratelimiting step in the reaction. In the one step mechanism (A), the rate limiting step is as shown,
while in (B) the rate limiting step could be loss of nitrite to generate Co(NH3)53+ or attack of the
second cobalt complex on the nitrite nitrogen. One common way to address these questions is to
study the kinetics of the reaction.
Reaction kinetics are usually interpreted using a model called transition state theory. In
this theory, a chemical reaction occurs when the reactant(s) come together with and have
sufficient energy to climb an energy barrier. The top of this barrier is called the transition state,
the point at which the reactants cross the barrier to the product side. The transition state is like a
pass connecting one valley to another (reactants to products). The lifetime of the transition state
is very short⎯it is not an intermediate and it is never present in any significant concentration. In
a multistep reaction, the rate limiting step is the one that goes over the highest energy barrier.
The first information you get from the kinetics is the order of the reaction: is it first order
in nitrito complex or second order? First order would mean that the transition state contains only
one nitrito complex, while second order means that two nitrito complexes (or two species
derived from the nitro complex) have to come together to do the reaction. Nitrite dissociation,
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analogous to an organic SN1 mechanism, would show first order kinetics (that’s what the “1” in
SN1 means).
However, if two cobalt complexes are involved in the rate limiting step, as in the middle
structure in Path B above, the reaction could show second order kinetics. Different reaction
orders can be distinguished by looking at how fast the nitrito complex disappears over time, as
described below in the Data Analysis section. Intuitively, if two nitrito complexes have to come
together, that's going to get pretty slow when the concentration of the nitrito gets very low
toward the end of the reaction. Another way to tell this⎯often more convincing⎯is to examine
the kinetics of the same reaction at different starting cobalt concentrations. If you triple the
starting concentration, a first-order reaction will go three times faster but a second-order reaction
will go nine times faster, a pretty big difference.
The temperature dependence of the rates of reaction are interpreted as changes in the free
energy of the transition state, ΔG‡, according to the equation ΔG‡ = ΔH‡ - TΔS‡ (see below).
The enthalpy of activation, ΔH‡, and the entropy of activation ΔS‡, are called the activation
parameters for the reaction and they provide information about the transition state. The symbols
ΔH‡ and ΔS‡ have a double dagger to indicate them as kinetic parameters, to distinguish them
from regular ΔH and ΔS that are thermodynamic parameters describing the ground state enthalpy
and entropy changes in the reaction. Note the difference between ΔG‡ and ΔG in the figure
above. [In fact, activation parameters are not strictly enthalpies and entropies, but we treat them
the same way.] If the transition state is more disordered than the reactants, as you might expect
for a dissociative process where two particles are made from one, then ΔS‡ should be a positive
number or perhaps a small negative number (more than about -80 J mol-1 K-1). An associative
process, however, should give a large negative ΔS‡ (more ordered) because two particles are
coming together. On the other hand, ΔH‡ (which is essentially always positive) should be large
for a dissociative process that involves bond breaking but smaller for an associative reaction in
which there is bond formation in the transition state.
Now it’s equation time. When Eyring created transition state theory he derived the
famous Eyring rate equation:
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k =
κkBT -ΔG‡/RT
h e
(1)
k is the measured rate of reaction, T is the temperature (in degrees Kelvin, K), and kB, h, and R
are the Boltzmann, Plank, and gas constants:
kB = 1.381 × 10-23 J K-1;
h = 6.626 × 10-34 J s;
R = 8.314 J K-1 mol-1
κ (Greek letter kappa) is the transmission coefficient, the fraction of the time that species at the
transition state really make it over to products as opposed to falling back to the starting materials
(you can assume it’s 1). As mentioned above,
ΔG‡ = ΔH‡ - TΔS‡
(2)
Substituting into equation (1) gives
k =
κkBT -(ΔH‡ - TΔS‡)/RT
e
h
or:
k h
TΔS‡/RT] e-ΔH‡/RT
T κkB = [e
or:
k h
ΔS‡/R] e-ΔH‡/RT
=
[e
T κkB
(3)
(4)
The right side of equation (4) has an initial term, in square brackets, that is independent of
temperature and a second term that is temperature dependent. So with a batch of rate constants
determined at different temperatures you can determine ΔH‡ and ΔS‡, as described in the Data
Analysis section below.
Method
This experiment is in two parts. This write-up describes the first part while in the second
part, a couple of weeks later, you will execute your own plan to extend this study. This is the
same pattern that we followed in the Chelate Effect experiment.
The part described here is a four-period lab in which you will work in pairs. You will
synthesize the cobalt(III) amine complex [Co(NH3)5Cl]Cl2 and convert it into the nitrito and
nitro linkage isomers. You will then study the kinetics of the isomerization of the nitrito to the
more stable nitro complex to determine the reaction order and the rate constant, k, at various
temperatures. To follow the progress of the reaction in your kinetic studies you will use UV/Vis
spectroscopy.
Procedure
I. Synthesis of [Co(NH3)5Cl]Cl2
References: W. L. Jolly, The Synthesis and Characterization of Inorganic Compounds, 1970 p.
461-462. G. Schlessinger Inorganic Syntheses 1967, 9, 160. [Inorganic Syntheses is a
wonderful set of volumes containing syntheses that really work⎯not only do they work for the
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researchers that submitted them for publication but they have been checked by another group.]
The original references are S. M. Jørgensen Z. Anorg. Chem. 1894, 5, 147 and 1898, 17, 455.
In the hood, dissolve 10.0 g of ammonium chloride, NH4Cl, in 60 mL of concentrated
aqueous ammonia in a 500 mL Erlenmeyer flask with a magnetic stir bar. While magnetically
stirring the solution, slowly add 20.0 g of finely powdered CoCl2•6H2O. You will end up with a
slurry, not a solution. Add 16 mL of 30% hydrogen peroxide to the stirred slurry, drop by drop.
Both heat and gas are evolved, so take it slowly! When the mixture stops fizzing, add drop by
drop 60 mL of concentrated HCl. Still stirring, heat the contents of the flask to about 85 ˚C for
20 minutes, then cool the flask to room temperature. Note the color changes in your reaction
mixture as the oxidation state of the cobalt and its ligands change.
Vacuum filter the precipitated [Co(NH3)5Cl]Cl2, using a fritted Büchner funnel, with no
filter paper. Wash the precipitate with 2×20 mL of ice-cold 6 M HCl. Pull air through the
powder to dry it partially, then put the solid in a labeled beaker in the 100 ˚C oven to dry for
about an hour. If the material does not look dry, or if you run out of time, you can leave the solid
in the oven until your next class. The material is air and water stable (clearly, from its
preparation) and may be stored in a vial.
Obtain an IR spectrum of the dry product (KBr pellet if possible; Nujol OK). This is a
good way to find out if the stuff is dry. In the third lab period you will also determine the visible
spectrum of [Co(NH3)5Cl]Cl2 (see part III); if you have time and there’s a spectrophotometer
available you could start on this now.
II. Preparation of [Co(NH3)5ONO]Cl2 and [Co(NH3)5NO2]Cl2
In the hood, heat 160 mL of water to about 80 ˚C in a large beaker with a stir bar. Add
15 mL concentrated ammonia and dissolve 10.0 g [Co(NH3)5Cl]Cl2 in the solution. A slight
precipitate of cobalt oxide may form; vacuum filter the solution to remove it. Cool the solution
to about 10 ˚C and put it in a cool water bath (cold tap water is OK). Slowly add 2 M HCl with
stirring until the solution is neutral to pH paper. Add 10.0 g of sodium nitrite and stir until it
dissolves; then add 10 mL of 6 M HCl to precipitate the [Co(NH3)5ONO]Cl2. Put the flask on
ice for half an hour to ensure complete precipitation. Then collect the precipitated crystals by
filtration on a Büchner funnel. Wash them with 2×25 mL ice water, 2×25 mL ice-cold ethanol,
25 mL of cold ether, and then suck air through the crystals to dry them. The dry solid may be
stored indefinitely in a labeled vial in the freezer; if you leave it standing at room temperature
some will isomerize. Be sure your vial is labeled with what it is, your and your lab partners
name, and the date you made it (including the year).
To make [Co(NH3)5NO2]Cl2, put about a gram of the nitrito complex into a labeled
beaker or small glass vial and put it in the 80 ˚C oven for an hour or so. You can follow the
progress of the isomerization by taking IR spectra⎯make sure you have reasonable spectra of
both the nitrito and nitro isomers. These spectra are reported in R. B. Penland, T. J. Lane, J. V.
Quagliano J. Amer. Chem. Soc. 1956, 78, 887. Some of the peaks are simply due to the
ammonia ligands and are found in most metal ammine complexes, such as the pentamminechloro starting material (compare with your spectrum taken above). There should also be two
additional bands for the nitrite ligand in both the nitrito and nitro complexes, between roughly
800 and 1500 cm-1. The nitrite ligand bands should be different for the two isomers, since the
bonding in the nitrite ligand is different in the two isomers.
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III. Kinetics and Activation Parameters for the Isomerization Reaction
Plan of action:
In the third lab period, you will obtain UV/Vis spectra and perform one kinetic run. In
the last lab period, you will obtain kinetic data at three other temperatures, measuring the
absorbance of the reaction mixture at just one wavelength.
A. UV/Vis Spectra
Your TA will give you an introduction to the UV/Vis spectrophotometer (see also the
Techniques section earlier in this manual). Before you can do any kinetics, you need to find a
concentration of the nitrito complex that gives a maximum absorbance somewhere in the range
0.5-2 for the band of interest. Spectrophotometers measure absorbance well up to about A = 2 or
so, but much above that there's just too little light getting through to the detector. Previous
students have found that the best kinetic data is obtained by looking at the changes in the UV
bands rather than the visible bands, but you might feel differently. For the UV absorption, you
might start at ~10-3 M concentration. Note that your eye cannot judge absorption in the UV.
Use deionized (DI) water as your solvent. Record the wavelengths of whatever peaks you see
(λmax) and calculate ε values for these peaks; for example, λmax = 500 nm, ε = 50 M-1cm-1).
Once you have a reasonable concentration, start the kinetics run as described in part B.
Make sure you obtain and save UV/Vis spectra of the nitrito, nitro, and chloro
complexes. Compare the spectra of the nitrito and the nitro isomer at the same concentration and
determine at what wavelength there is the largest difference in absorbance. This will be the
wavelength that you will monitor for your kinetics. You should discuss your wavelength
choice with your TA. Make sure you record the absorbance of each isomer at this wavelength.
You should be able to finish the kinetics run in this period and you may have time to take other
spectra while the kinetics are running. But some students have found that if they're slow at the
beginning this can be a long lab. Come prepared!
B. Kinetics Run at One Temperature
Make up 100.0 mL of a buffer solution that is 0.1 M in both ammonia and ammonium
chloride. (The buffer minimizes the formation of cobalt hydroxide when you heat the nitrito
isomer, in part by inhibiting decomposition via dissociation of ammonia. It’s generally a good
idea to have a buffer present in aqueous reactions to control the pH.) Use volumetric glassware
and an analytical balance here and in all solutions for kinetic studies. Pour your solution into a
labeled 250 mL Erlenmeyer flask and put the flask in one of the hot water baths (set to 40˚C) that
have been set up in the lab. Generally, 40˚C works well but if you're running late, you can use a
slightly higher temperature so the reaction will go faster. Stopper your flask to prevent
evaporation of ammonia.
Make 20 mL or so of an aqueous stock solution of [Co(NH3)5ONO]Cl2 at a concentration
ten times the concentration you need for UV/Vis spectra. It should be ten times as concentrated
because you will dilute it roughly ten-fold with the buffer solution. This solution is unstable at
room temperature and must be stored in an ice bath. Use a volumetric pipette to add 10.00 mL
of the cold stock solution to 90 mL of the preheated buffer solution. Swirl the flask to mix the
contents thoroughly and record the time. Using a plastic pipette, quickly squirt an aliquot (a
small portion) from the flask into a cuvette and cap it. [Note that you have the option to check
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out a quartz cuvette from the stockroom but the special UV plastic cuvettes available in the lab
work just as good for this experiment.] Leave the flask, stoppered, in the water bath, put the
cuvette in the spectrophotometer and record the spectrum. Be sure to save the spectrum at each
timepoint. Don’t change any spectrometer settings once the run has started as this will invalidate
your data. In your notebook, record the absorbance at the peak and the time of this spectrum.
Use a stopwatch if one is available. What should you use as tinitial for the reaction, the time of
mixing or the time of the first spectrum or ...?
Pour the contents of the cuvette into the cobalt waste container, then rinse the cuvettes
with distilled water and dry with compressed air. Repeat the process of taking aliquots and
measuring their spectra at noted times. At the beginning, the absorbance will change rapidly and
you should try to take a measurement as often as possible. Later in the reaction the absorbance
will change at a more sedate pace and you can take a measurement every quarter hour or so.
You can judge the progress of the reaction by overlaying the saved spectra on the screen of the
spectrometer.
When you have time during the kinetics, view some of the spectra using the overlay
function. Only eight spectra can be plotted at once so you should pick the first and last spectrum,
as well as six in the middle. Make sure you get a plot like this when you’re all done. As you
look at your set of spectra on the screen, are there any wavelengths at which the absorbance
doesn’t change over the course of the reaction? These are called isosbestic points, wavelengths
where the starting material and the product have the same extinction coefficients (ε’s). Their
presence indicates that there are only two species with absorbance in the visible region in the
reaction mixture in any appreciable concentration. Do your spectra go precisely through the
isosbestic point(s)? Or is there some scatter or a systematic deviation, such as the isosbestic
point appearing to drift up or down?
The data analysis (see below) requires a final absorbance for the reaction, Af. Your
reaction may be fast enough that it is complete by the end of the lab period (the absorbance stops
changing significantly). But in most cases, you will need to help the reaction along to
completion. About half an hour before the end of the period, heat your flask to 80˚ (not hotter!)
for a couple of minutes to fully convert the nitrito to the nitro complex. You can monitor this by
taking aliquots, as above. Spectra of the final solution, cooled to room temperature, give the Af
value. Boiling or overheating the solution could change the concentration of the nitrito product
or cause decomposition and thus give an incorrect Af.
Be sure that all spectra are preceded by a blank spectrum. You can use the same blank
for multiple spectra only if you're sure no-one else has used the spectrometer in between.
Because your samples are 10 mL complex in DI water plus 90 mL buffer, your blank should be a
1:9 mixture of DI water and buffer.
C. Three Kinetics Runs at Once
In the last lab period, you will follow the kinetics of three reactions simultaneously on
one spectrophotometer, measuring the absorbance at the same wavelength as the first kinetics
run.
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As in part B, make up a buffer solution 0.1 M in both NH3 and NH4Cl, enough to put
100.0 mL of buffer solution in each of three labeled 250 mL Erlenmeyer flasks. (You might
want to make 500.0 mL of buffer solution just because you have a 500 mL volumetric flask).
Preheat the stoppered flasks in water baths at three temperatures that are different from the
temperature you ran in part B. Find three UV/Vis cuvettes and label them something catchy like
25˚, 30˚, and 35˚ (don’t put labels on the optical windows!). As above, make 50 mL of a stock
solution of [Co(NH3)5ONO]Cl2 and store it in an ice bath. Add 10.00 mL of the stock solution to
each of the preheated flasks with 90.0 mL of buffer solution. Swirl to mix, record the time, and
take an aliquot from each flask into the appropriate cuvette. Leaving the flasks in their
respective water baths, quickly measure the absorbance at your chosen wavelength, recording the
value and the time in your notebook. Collecting data for three concurrent kinetic runs is hectic at
the beginning when the measurements are taken most frequently. You can monitor the rate of
change by making a rough plot of absorbance vs. time in your notebook as you collect the data.
As before, toward the end of the period you should heat each of your flasks to 80˚ for a couple of
minutes to finish the reactions and obtain final absorbance values.
Data Analysis
1. Determination of the reaction order and the rate constant (k)
You measured absorbance (A) as a function of time. You convert the absorbance to
concentration of the nitrito isomer as follows. At time t, the measured absorbance At is the sum
of the absorbance due to the nitrito complex and that due to the nitro complex. [CoONO]t and
[CoNO2]t are the concentrations of the complexes at time t, and εx is the extinction coefficient of
complex X at the wavelength chosen. Square brackets are used to indicate concentrations and,
for convenience, the formulas are written omitting the five amine ligands and the charge.
At = εCoONO [CoONO]t l + εCoNO2 [CoNO2]t l
(5)
Assume that the nitrito starting material was 100% pure so that the initial absorbance, Ai is:
Ai = εCoONO [CoONO]i l
If the nitrito cleanly converts to the nitro isomer, then by mass balance the total amount of cobalt
complexes at any time is equal to the initial concentration of nitrito, [CoONO]i, which is equal to
the final concentration of the nitro, [CoNO2]f.
[CoONO]t + [CoNO2]t = [CoONO]i = [CoNO2]f
or:
[CoNO2]t =
[CoNO2]f
– [CoONO]t
Substituting eq 6 (the ‘mass-balance’ equation) into eq 5,
At = εCoONO [CoONO]t l + εCoNO2 ([CoNO2]f – [CoONO]t) l
At = εCoONO [CoONO]t l + εCoNO2 [CoNO2]f l – εCoNO2 [CoONO]t l
The middle term is simply the final absorbance, Af:
At = εCoONO [CoONO]t l + Af – εCoNO2 [CoONO]t l
Af - At = (εCoNO2 – εCoONO) [CoONO]t l
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Multiplying both sides by [CoONO]i,
[CoONO]i (Af - At) = (εCoNO2 [CoONO]i l – εCoONO [CoONO]i l) [CoONO]t
Since [CoONO]i = [CoNO2]f, the terms inside the parentheses on the right are Af and Ai:
[CoONO]i (Af - At) = (Af - Ai) [CoONO]t
or:
[CoONO]t =
Af - At
Af - Ai [CoONO]i
(7)
Check this equation to make sure it looks right: does it give the right answers at t = tinitial and t =
tfinal?
Think about the assumptions that went into this equation. Is your nitrito complex really
100% pure? What effect would it have on the equations below if it isn’t? Does the reaction
really proceed in 100% yield? (The tightness of your isosbestic points might help you answer
this.) It turns out that the most critical assumption in using eq 7 is that you know the final
absorbance, Af. Make a plot of your absorbances vs. time (your raw data) and estimate where
this would level off. Calculate the Af you would expect based on the εCoNO2 at this wavelength
from your UV/Vis spectrum of the nitro complex (Af = εCoNO2 [CoNO2]f l = εCoNO2 [CoONO]i
l). Do these estimates of Af agree with the one you got from heating the solution? Compare the
ε’s you get from the kinetic studies with those from your UV/Vis spectra of the solids. Which do
you think are more accurate? With absorbances and ε’s at two wavelengths, you can calculate
the concentrations of both species without the mass balance assumption (eq 6). Can you see
how?
Now onto the kinetic analysis. If the reaction is first order in [CoONO], then
d[CoONO]t
dt
Integrating:
(k has units of s-1)
= -k [CoONO]t
ln[CoONO]t = -kt + ln[CoONO]i
(8)
Using eq 7 to recast eq 8 in terms of the direct observable, At, gives:
⎛Af - At
⎞
ln⎜A - A [CoONO]i⎟ = -kt + ln[CoONO]i
i
⎝ f
⎠
or:
⎛Af - At⎞
ln⎜A - A ⎟ = -kt
i⎠
⎝ f
(9)
If the reaction is second order in [CoONO], then
d[CoONO]t
dt
= -k [CoONO]t2
1
[CoONO]t
=
(so k has units of M-1 s-1). Integrating:
1
kt + [CoONO]
i
(10)
As above, use eq 7 to recast eq 10 to give an equation that relates At to t. In these equations, k is
the rate constant for the reaction at a particular temperature.
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Pick a temperature at which you think you have good data and make some plots. If the
reaction is first order, a plot of ln[(Af - At)/(Af - Ai)] vs. t should be linear, with a slope of -k (eq
9). Work out what plot should be linear for a second-order reaction following your recast
version of eq 9. See which best approximates a straight line. The one that is best will have some
scatter above and below the best line, but will not look like a curve. If you are unsure of your Af
values (see above), try making plots with the various values. Alternatively, you can have a
program like Microsoft Excel fit your raw At vs. t data. If the reaction is first order, it should fit
the exponential form of eq 9:
At = a e-kt + b
where a = Ai - Af and b = Af
(11)
You can derive the equivalent expression for a second order process. These direct fits have the
advantage that they don’t require you to input Af or Ai values, but it’s harder to tell whether the
fit is good or not because the program is adjusting three parameters. Make sure that the Af or Ai
values you get are reasonable.
Repeat the first-order vs. second-order analysis for another temperature. The order of the
reaction should not change with temperature (unless there is a change in mechanism which is
quite unlikely). The distinction between first and second order will be most clear for data sets
that extend farthest toward the end of the reaction. If you only have one half life of data (the last
point corresponds to only 50% of starting material converted), it'll be very difficult to tell first
order from second order. The plots will be much clearer if you have two or three half lives (75%
or 87.5% completion). From your plots determine k at each of the four temperatures. Be sure
(always!) to indicate the units of k.
2. Determination of the enthalpy and entropy of activation
In the background section, we “derived” equation (4):
k h
‡
‡
ΔS /R] e-ΔH /RT
=
[e
T kk B
(4)
Taking the natural logarithms of both sides, we obtain
k
h
ln T + ln
=
κk B
ΔS‡
R
ΔH‡
– RT
A plot of ln(k/T) vs. 1/T⎯an Eyring plot⎯should be linear (you should have four data points).
The y intercept is ΔS‡/R - ln(h/κkB) and the slope is -ΔH‡/R. Watch out for units when you do
this plot (in fact, whenever you use an equation): T must be in degrees Kelvin, K, and the rate
constant k should be in s-1 (for a first order rate constant; s = second) or M-1s-1 (for second
order). If you get confused or want to check your constants, there's a nice summary in The
Chemists Companion A. J. Gordon, R. A. Ford, 1972, Wiley, New York.
Estimate the error in each rate constant (see page 12). Considering your various values of
Af, how much does k depend on which value you chose? This provides another measure of error
in k. Put these error bars on your Eyring plot and crudely estimate your errors in ΔS‡ and ΔH‡.
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Suggest reasons for your observed uncertainties. Is there more uncertainty in the
enthalpies than the entropies, or vice versa? Why? Think about all the experimental values that
go into these calculations⎯concentrations, temperatures, absorbances, etc.⎯and how uncertain
they are. How would small errors in these values affect your calculated enthalpies and
entropies? Do you think your errors are random errors, caused by uncertainties in measurement,
or systematic errors that result from your particular experimental method and equipment?
Lab Report
A lab report on this experiment is due a week after the experiment is completed. It
should have the long report format described in the beginning of the lab manual, focused on the
kinetics portion. You will revise this lab report after the "extra" lab period. Try to be concise
and to the point, but completeness is more important than staying within the suggested guidelines
for the length of each section.
I. Introduction (several paragraphs).
A. describe why this experiment was done and why it is of interest.
appropriate literature references.
Use
II. Experimental (try to keep to 1-2 pages).
A. briefly describe the syntheses of all synthesized products, with yields and
colors. Include:
1. IR stretching frequency, with assignments.
2. UV/vis λmax and ε values.
B. describe how the kinetics was done.
1. describe how the was raw data collected.
2. determination of Af
III. Results
A. Tables summarizing IR and UV/Vis characterization (attach original UV/Vis and IR
spectra for all synthesized products at the end of your report).
B. An overlay UV/vis plot of one kinetic run (pick a good one) illustrating an
isosbestic point
C. Plots of raw data (At vs. t) for all temperatures.
D. Plots for 1st and 2nd order fits for all temperatures, showing linear fits.
E. Eyring plot, showing error bars.
a. calculations for determination of ΔH‡ and ΔS‡
IV. Discussion (aim for ~5 pages of text).
A. analysis of the kinetic data.
1. determination of reaction order, rate constants.
B. analysis of the rate constants: determination of ΔH‡ and ΔS‡
C. Discussion of errors.
D.Interpretation of reaction order, rate constants, and activation parameters
in terms of the reaction mechanism (the most important part).
V. Plan for the final lab period (~1 page)
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Describe an experimental procedure for what you will do in the additional lab period,
with a clear rationale for your choice. Consult the discussion below.
The "Extra" Lab Period
As in the Chelate Effect and (Arene)molybdenum experiments, you plan your own
procedure for an "extra" lab period. We suggest that you extend this study, exploring the
mechanism of the reaction by doing kinetics under different conditions. You should plan on
running at least two kinetic runs in the extra period, one under your new conditions and one with
the same solutions but under the “old” conditions. This is the best way to tell if the change you
made had any effect, comparing two runs on the same solutions in the same temperature bath. It
has the additional value that you will see how reproducible your data are from day to day.
Suggestions:
You might vary the concentration of the starting cobalt complex, to confirm your first
order/second order conclusion. If the reaction is first order, its half life (the time for half of the
material to go away) will be independent of concentration. Conversely, a second order reaction
will go faster at higher concentrations. Note that the concentration you choose must be
appropriate for your experimental conditions. For instance, proposing to increase the
concentration ten times will not work because your Co complexes will not be soluble at this
concentration.
You could determine the order of the reaction with respect to nitrite anion, by doing kinetics runs
in the presence of added NaNO2. If the mechanism requires an initial equilibrium dissociation of
nitrite, then adding nitrite should slow the reaction down by shifting this pre-equilibrium to the
left and reducing the amount of reactive stuff. (It's called a pre-equilibrium because it comes
before the rate limiting step in the reaction.)
[Co(NH3)5(ONO)]2+
[Co(NH3)53+] + NO2–
You could examine the effect of pH on the reaction, by changing the ratio of the components of
the buffer. Be sure to check the pH of your reactions with a pH meter. [You'll also need to
worry about pH if you add NO2-, as this is a weak base.] Note that one unit change in pH
corresponds to a factor of ten change in the concentration of H+ and OH-.
You could add excess Cl–, I–, or some other ligand to try to trap [Co(NH3)53+].
Or you could find your own questions (usually the best way). Whatever approach you take, be
sure to discuss it with your TA.
One week after the extra period, turn in your revised lab report. This should take
into account the comments on the first version and incorporate what you did with the extra
period. You should describe what you hoped to learn from the experiments you planned and
whether you actually did learn those things. You will turn in this revised lab report, along
with your first lab report and grade sheet.
Questions for thought (no need to turn in answers):
•
•
If you had some isotopically labeled nitrite, say Na15NO2 (too expensive to give you), what
experiments might you do test the hypotheses described above?
You could also try to make other cobalt complexes of the form [Co(NH3)5X]Cl2. What X
groups might you want to use? Would you expect linkage isomers from nitrate (NO3–)?
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What about cyanate (OCN–)? How might you do these syntheses? Do you think you could
tell from the visible spectra what the structures of the complexes are?
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Preparation of a Doped Phosphor: ZnS
Background
The emission of light from matter is a fascinating and very useful phenomenon. All
matter glows when it is heated to high temperatures, which is incandescence (think of an
incandescent light bulb or the heating elements in an electric oven that turn orange when hot).
Incandescence is an example of black body radiation and it occurs over a very broad continuum
of wavelengths (it has a broad spectrum). Emission occurs in the visible region only at
temperatures above about 1000 °C. Certain minerals and insects, on the other hand, emit light at
room temperature; this phenomenon puzzled natural philosophers hundreds of years ago. In
1603, an Italian shoemaker heated the mineral barite (BaSO4) with charcoal and made material
that glowed in the dark after exposure to sunlight; he called this a “phosphor”, which means
“light bearer” in Greek. Today, luminescent materials of all kinds, including phosphorescent
ones, are used in TV screens, X-ray films, fluorescent lights, light emitting diodes (LEDs),
semiconductor lasers for fiber optics, invisible hand stamps, anti-counterfeiting inks for
currency, and many other applications.
In order to get light out of a luminescent material, you have to put energy in. There are a
myriad of sources of energy: sunlight, for those “glow-in-the-dark” articles found in cereal
boxes; beams of electrons, for the luminescent coating which produces the picture on your TV or
computer screen (unless you have a flat panel screen); high-energy chemicals, such as the ATP
with which fireflies fuel their glow. The phosphor which you will make in this lab is stimulated
by photons in the ultraviolet part of the spectrum. A UV photon comes in, excites the material,
and a photon of lower energy (visible light) is emitted; the difference in energy is lost to the
material as heat.
The photon serves to boost the luminescent system up to an excited state. This takes
about 10-18 seconds (this is the time it takes a photon, at the speed of light, to travel 300 pm!).
The excited state may go back to the ground state by releasing thermal energy (heat) or by
rapidly re-emitting light of roughly the same frequency (this takes 10-12 to 10-6 seconds). This
fast re-emission of light is called fluorescence; it is exemplified by the coatings on fluorescent
light bulbs. Another option for a molecule in an excited state is to decay to another excited state.
This sounds pointless, but sometimes it has dramatic effects. The interconversion of photons,
excited states, and the ground state is called the photophysics of the problem. The excited state
can also use its energy to break bonds or to react with something, in other words to do
photochemistry. You might use or have used UV light to do a reaction of a molybdenum
carbonyl compound in experiment 6.
A sodium atom provides a simple illustration of photophysics. The ground state of a
sodium atom is described by the electronic configuration [Ne]3s1. Excitation of a sodium atom
by a photon in the yellow region of the visible spectrum gives the lowest energy excited state,
configuration [Ne]3p1. In the energy level diagram below, the most stable level is at the bottom
and the highest energy one is at the top. Events that involve a photon are drawn with wavy
arrows while the transfer of heat energy is depicted with straight arrows. Excitation with a
photon of energy hν1 gives the excited state which then either emits the light back out by
fluorescence or decays to the ground state without emitting a photon. Decay without emission is
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called radiationless decay, and the energy of the excited state is just dumped as heat (perhaps in a
collision with another sodium atom).
This isn't a very exciting system: you put light in and you get light of the same energy
out, just in lower yield. But let's say you could excite the sodium atom to the next lowest excited
state, [Ne]4s1, perhaps by collision with an electron. This is shown in another energy level
diagram on the next page. The [Ne]4s1 excited state has a very low probability of emitting to the
ground state, because of quantum mechanical selection rules. So it waits around until it is hit by
another atom or something else happens to allow it to decay to the [Ne]3p1 excited state. This is
now a much more interesting photophysical system: the energy of the electron (or whatever is
used for excitation) is converted into a photon of wavelength ν1. This is the principle of sodium
vapor lamps, the yellow lights you see in tunnels and elsewhere around town. Another example
of interesting photophysics, the ruby laser, is described in the Appendix to this experiment.
The system you will work with in this experiment is zinc sulfide, commonly employed in
the electronics industry as a coating for cathode ray tubes (such as TV tubes or computer
screens⎯a “cathode ray” is a beam of electrons). ZnS (the minerals zinc blend and wurtzite) can
be thought of as an ionic solid (Zn2+S2–), but based on its optical and electronic properties it is a
semiconductor. Solids such as ZnS have energy levels, just as molecules do. In a molecule, a
relatively small number of atomic orbitals interact to give molecular orbitals of specific energies.
In a solid, ca. 1023 atomic orbitals interact and give rise to ca. 1023 molecular orbitals. These
orbitals typically fall into groups with similar energies, which are called bands of orbitals. The
electrons, of course, occupy orbitals, so in a solid, electrons occupy bands of orbitals (see the
drawing on page 79). In solids that are electrically conducting such as metals, the highest energy
band is only partially filled, leaving the electrons mobile. In insulating solids, there is a large
energy gap between the highest filled band (the valence band) and the lowest unoccupied band
(the conduction band). [This is like the typical gap in a molecule between bonding (typically
filled) orbitals and antibonding (empty) orbitals.] In a semiconductor such as silicon, there is a
small band gap.
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Defects in a crystal bring their own electrons and orbitals. Often the energies of these
orbitals fall within the band gap and therefore influence the optical and electronic properties.
Defects can be lattice vacancies (e.g., a missing zinc ion in ZnS), added atoms in the holes within
the lattice, or impurities that are substituted for lattice atoms, for instance a Cu2+ replacing Zn2+.
This is how the electronic structure of silicon is varied to make transistors: if a silicon atom is
replaced by a phosphorus atom, there is an extra electron in the solid but if Si is replaced by Ga,
there is one too few electrons (a hole). These are n-doped and p-doped silicon (negative and
positive), respectively.
Zinc sulfide is a semiconductor with a fairly large band gap between its valence and
conduction bands. An ordinary photon of visible light doesn’t have enough energy to excite an
electron across this gap (that's why the solid is white or yellow). However, crystal defects
introduce energy levels that lie within the energy gap. Absorption of a UV photon promotes an
electron to the conduction band, leaving a positive hole in the valence band. If this electron
relaxes to a defect energy level, light can be emitted at a lower frequency than the excitatory
light. This photophysics is similar to that in the sodium case above. In the ZnS phosphor, UV
light goes in and visible light comes out. (The case drawn shows defect energy levels that are at
least partially empty so electrons can be transferred into them; it is also possible to have filled
defect levels into which the hole relaxes up.) Photons promoting electrons to the conduction
band is the essence of “photoconductivity” in semiconductor photocells.
So how would one go about making a doped solid? One can make ZnS solid by mixing
aqueous solutions of Zn2+ (aq) and HS- (aq) but the product still has water and hydroxide ions in
it and is not very crystalline. Thus it has lots of defects and is not very pure, so it is not usable
for semiconductor applications. (The formation of pure, crystalline solids under such mild
conditions is a challenge, and a topic of much current research.) The problem with making a
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unfilled
(empty)
hν
(UV)
Defect energy levels
hν
(visible)
filled
Pure Semiconductor
(no defects)
Semiconductor
with Defects
solid at room temperature is that if the Zn2+ or S2- ions don’t fall in exactly the right place
initially, it’s hard to move them into their proper crystal locations. Diffusion in a solid is
typically very slow at room temperature.
The synthesis of solids is typically done at high temperatures, to facilitate solid state
diffusion. The classic way to make ZnS is to mix solid zinc and solid sulfur in a sealed tube
under vacuum and heat to about 450 °C, the boiling point of sulfur, making this a gas/solid
reaction. To make crystalline samples, you might then heat the crude solid to 1000 °C to get
solid diffusion to occur with any speed. High temperatures are especially needed in cases where
both reactants are solids even at 1000 °C, as reaction then can take place only where the particles
touch. In this case, finely powdered reagents are often used to maximize particle contact.
This experiment starts with commercial "electronics grade" zinc sulfide and a dopant
such as Cu2+ is introduced. Cu2+ can substitute for Zn2+ in the lattice and it has a partially filled
d shell. The dopant is added to the ZnS powder as a solution, to disperse the cationic dopant
fairly evenly over the surface of the ZnS particles. Imagine the problems if you just placed one
crystal of CuSO4 next to all the ZnS particles! You would never get an even distribution.
Dispersing the dopant over the particles helps achieve a random distribution of defect sites.
Initially the dopants are only on the surface of the ZnS particles, so the solid must be heated to
900 °C to allow the dopants to move (diffuse) through the lattice.
Even at 900 °C, this reaction is slow so you will use a flux, the solid state chemist’s
equivalent of a solvent. You will use NaCl as your flux. While most chemists think of NaCl as a
white solid, it is a liquid at 900 °C. A flux can act like a typical solvent and dissolve all the
reactants, or, more commonly, it dissolves only a small surface layer but this is enough to assist
the reaction. NaCl is a nice flux because when everything has cooled back to room temperature,
it can be easily washed away from the product with water.
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Method
This is a one-period lab in which you should work individually, but coordinating with the
rest of your lab section. Choose a cation as a dopant, checking with your TA to see what’s
available. Each person should try to choose a different dopant, as well as one person who will
not use any dopant at all, also consider using iron as a dopant. You grind together the dopant,
the NaCl flux, and the ZnS into a fine powder to get the particles as well mixed as possible.
Then you load the powder into a small “boat” for heating in the tube furnace at 900 °C. Once it
cools, you wash away the flux with water. You place the doped ZnS under a UV light to see if
the product phosphoresces (“glows in the dark”), or if your dopant quenches (prevents) emission
of light.
Experimental
First, choose a dopant to use. This experiment was originally written using Cu2+ as the
dopant, but ZnS can be doped with a variety of metal ions. You might use salts of manganese,
silver, gold, chromium, nickel, uranium, lead, tungsten, etc. Luminescence is claimed to be
completely quenched by iron and nickel, even in concentrations as low as 10-6 mole per mole of
ZnS (one part per million, or 1 ppm). You could test whether this quenching really happens. It
is also reported that when ZnS is heated with no dopant, it loses some sulfur and becomes a nonstoichiometric compound (ZnS1-x), which might reasonably have lots of crystal defects and
luminesce.
Clean everything that will touch the phosphor materials with 1 M HCl. Rinse well with
deionized water, and then with acetone. Remember that even traces of iron are reported to
quench the emission.
Prepare a dopant solution of molarity roughly 10-5 M in water (there may be some
already made). If you’re making one up, mark the concentration on the bottle and leave it in case
someone else wants to use it (remember, there are lots of lab sections). Weigh out about 1 g
NaCl (don’t use a metal spatula!) and put it in a clean mortar. Add 1.0 mL of your dopant
solution using a clean 1 mL pipette. Put the wet salt into an oven at about 120 °C until it is dry,
about 1/2 hour. [Why do you think the Na+ from the flux does not dope the ZnS, but divalent
cations and Ag+ will?]
Weigh out about 1 g of electronics grade ZnS. Grind it and the doped salt together in the
mortar until you have a fine powder. Add a few drops of ethanol and continue grinding until you
get a smooth paste. Pack this paste into a clean vitreous boat and dry in the 120˚ oven for about
20 minutes.
Note which boat is yours before you put it in the tube furnace (how?⎯you can’t use a
marker or tape as this will just burn off at 900 °C). Let your phosphor heat in the furnace under a
constant stream of dry nitrogen; your TA will set this up. The nitrogen is critical because any
oxygen present at 900° would convert the ZnS to ZnO, which doesn’t emit much (it has a very
large band gap). Allow the boats to cool in the nitrogen stream.
Check the crude product for luminescence in the UV light box. You might also examine
some of the starting material, ZnS. Check out the samples others have made.
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Lab Report:
Observe which dopants effect phosphorescence using the UV light box. Record this for
all doped ZnS samples made in your lab period. Write a short (approx. 1 single-spaced page)
news article, suitable for the science column of a newspaper or a science-oriented magazine.
Your article should include enough background to enable a scientifically literate reader to
understand it, and it should communicate the latest research in the field undertaken by students in
Chem 317. The dry scientific writing you have done in previous reports is not appropriate to
reach a more general audience, as we are asking you to do here. Part of the challenge for this
report is providing enough background to bring your readers up to speed, while leaving enough
space to explain the actual “news story,” aka, what your class achieved in the lab period.
Although the format is different, be sure to discuss relevant details of your experiment. Be
succinct and imaginative - have some fun with this one.
Questions for you to think about (but not turn in.):
Why is it that the ZnS must be heated under nitrogen, yet the products may be handled in the air?
Why should you grind your starting material to a fine powder?
Why are high temperatures needed?
Why should you not use a metal spatula? (What metal do you think it is made of?)
Why type of glass is used in the tube furnace? (Yes, there are different types.) We normally use
Pyrex® glass in the lab. What is the maximum temperature that Pyrex remains useful? How can
you tell different types of glass apart in the laboratory?
Appendix:
Another Example of Photophysics: The Ruby Laser.
Ruby contains chromium(III) ions doped into a crystal of Al2O3 (so the correct formula is
Al2-xCrxO3). The Cr3+ ions are isolated from each other and therefore interact with light as
isolated ions. We can therefore discuss what happens to only one ion and we’ll have most of the
picture. The Cr3+ ions are octahedral holes of the oxide lattice so that the d orbitals split into the
familiar “2 above 3” pattern. The ground state of octahedral Cr3+ is labeled 4A2 in the symmetry
language used for such problems; the 4 stands for 3 unpaired electrons plus 1 (2S + 1 in quantum
mechanics-ese) and the A2 refers the orbital symmetry of the state. A photon of light can excite
a chromium atom to an excited state (4T2), also with three unpaired electrons, as shown in the
diagram below. As above, events that involve a photon are drawn with wavy arrows while the
transfer of heat energy is depicted with straight arrows. The absorption of light (4A2 → 4T2) is
what gives the aquo chromium(III) ion, Cr(H2O)63+, its forest green color (remember, in the
Chromous Acetate experiment?). When the chromium is bound to oxide ions in a ruby lattice,
instead of water molecules in Cr(H2O)63+, the energy gap between the 4A2 and 4T2 states is such
that blue/green light is absorbed and rubies look red.
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4
T 2 Excited State
thermal decay
"intersystem crossing"
2
excitation
E=hv 1
E Excited State
Energy
reemision
E=hv 1
emission
E=hv 2
thermal
decay
Ground State
The 4T2 chromium excited state in ruby rapidly decays predominantly by intersystem
crossing to another excited state, where one of the electrons has flipped its spin and there is now
only one net unpaired electron (2E). To get back to the ground state, this electron has to flip its
spin back again, but since photons are not trained to help out with changes in electron spin,
emission of light is a difficult process; it is said to be “spin-forbidden.” (This quantum
mechanical “forbiddenness” is why incoming light does not excite ground state Cr3+ (4A2)
directly to the 2E excited state.) “Forbiddenness” doesn’t stop emission, but it slows it down,
leading to lifetimes from 10-6 seconds to several seconds or even minutes. This slow glow,
exemplified by “glow-in-the-dark” stickers and suchlike, is called phosphorescence. In ruby,
decay is slow enough that by “pumping” with the right frequency, most of the chromium atoms
can be put into this excited state. This is called a "population inversion," when there are more
atoms in the excited state than in the ground state, and it is a prerequisite for a laser. Then a
photon of the right energy (hν2 in the drawing) comes along and stimulates emission from the
excited states. This “stimulated emission” gives light all of the same frequency and with the
same phase⎯the special characteristics of a laser⎯because it was started by a single photon.
The word laser comes from light amplification by stimulated emission of radiation.
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Chromous Acetate Lab Report Template
Name________________________
Student ID Number_____________
Partner_______________________
Section_______________________
Synthesis, Characterization, and Reactivity of Chromous Acetate
(The goal of this activity is to assess your understanding of the chromous acetate lab and to
provide some guidance on what is expected when writing a formal lab report as you will do later
in the quarter. As part of this assignment it is expected that you will READ carefully the section
titled “Guidelines for Writing Lab Reports” in the lab manual.)
Introduction
Chromium acetate has been known for over 100 years and is still of interest to chemists
because the nature of the metal-metal bond is not well understood despite its having been studied
by a variety of techniques including _____.i (read this reference and choose one of the
techniques) One proposed structure for chromous acetate is shown in Figure 1.
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see chem. 317 website for directions to download ChemDraw. You must create your own image,
copy/paste from anywhere is not acceptable.
Figure 1. Structure of chromous acetate
This structure is based on…. (summarize the data presented on page 29-30 of the lab manual.
Read references as needed.)
Results
The synthesis was carried out according to the balanced equation below.
In this space, put a balanced equation with subscripts and superscripts as appropriate
In the first step chromium (oxidation state?) starting material was reduced to chromium
(oxidation state?) using ____ as the reductant as shown in the balanced redox reaction below.
In this space, write a balanced redox equation.
This (exo/endo)thermic reaction was performed under acidic conditions and was accompanied by
a color change from ____ to _____. Once the chromium(oxidation state?) was generated it was
reacted with a sodium acetate solution forming the chromous acetate complex. This reaction was
accompanied by a color change from ____ to _______.
The product was characterized by IR spectroscopy. The results are reported in Table 1
along with characterization of the sodium acetate starting material and the Nujol.
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Table 1. Assignments of IR spectra for chromous acetate, sodium acetate, and nujol
assignment type of vibration sample sodium acetate chromous acetate nujol acetate N/a water N/a Nujol N/a A sample of chromous acetate was found to react with air in an amount of time becoming
______. When water was added to a sample of chromous acetate in air, it reacted in an amount
of time becoming ________.
Discussion
The synthesis of chromous acetate was carried out using Schlenk techniques because
chromium (ox state?) reacts with the______in air. If water is present this reaction occurs
faster/slower based on the comparison of samples of wet and dry chromous acetate being reacted
with air.
IR spectroscopy was used to characterize the product and determine whether the
chromous acetate is Cr2(OAc)4 or the hydrate, Cr2(OAc)4.2H2O. From the spectrum of just Nujol
it was found that the vibrations at ______ result from the Nujol and therefore can be ignored in
other spectra. The spectrum of sodium acetate reveals vibrations at _____ which can be
associated with free acetate. In the spectrum of chromous acetate, a vibration at _____ was
assigned as the _______ of the metal-bound acetate, consistent with published work that suggests
that metal-acetate complexes show IR vibrations around ______.ii The IR spectrum of chromous
acetate (does/does not) show bands for water suggesting that the formula of the product is
(anhydrous/hydrated). Based on this molecular weight, the yield of chromous acetate is ___%.
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Conclusions
Chromous acetate,(type formula here), was synthesized and characterized by IR
spectroscopy. Its reactivity with oxygen was observed and it was found to react with air
(faster/slower) in the presence of water
Experimental
General Considerations:
All manipulations were performed using standard Schlenk techniques or in a nitrogen-filled
glovebox unless otherwise specified. Solvents were degassed by _________prior to use. All
other reagents were used as provided. IR spectra were recorded in Nujol.
Synthesis of____________:
In air, chromium (III) chloride hexahydrate (__g, ___ mol) was dissolved in water (___ mL) and
concentrated HCl (___ mL). The green solution was then degassed and added dropwise to
mossy zinc (___g, ___mol). When the reaction was complete, the blue solution was added to
saturated sodium acetate solution (___mL) by cannula. The resulting precipitate was collected
by filtration on a frit and washed with ethanol ( ___ mL) and diethyl ether (___mL). The red
solid was dried under vacuum (____g, ____mols, ____%). IR: ____νC-O, ___ νH2O. 1H NMR
(acetone-d6, 200.1 MHz) δ=1.68 (s, 12H).
Reaction of ________ _________ with air:
A sample of______________________(___g, ___mol) was opened to air. A color change was
observed after an amount of time.
Reaction of aqueous ___________________ with air:
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A sample of ____________________(___g, ___mol) was opened to air and water (___mL) was
added. A color change from ____to ____was observed after an amount of time.
Supplemental Information
Clearly labeled spectra are attached
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Diagram of Electrochemical Cells for Chelate Effect Expt
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Chelate Effect Lab Report Template
Name________________________
Student ID Number______________
Partner_______________________
Electrochemical Comparison of [Cu(NH3)4]2+ and [Cu(ethylenediamine)2]2+
Introduction
The chelate effect is a phenomenon that explains why
.1 Read this
reference and explain, in 1-2 sentences, what the chelate effect is. The chelate effect is the
fundamental principle behind
. Pick one example from the lab manual or from
your own experience that uses the chelate effect and briefly explain it. Find a reference for this
example that is not the lab manual.
Ammonia and ethylenediamine complexes of copper (II) are studied by electrochemical
techniques to obtain values for
. What thermodynamic values were the chief results
from your study? These values support/don't support the established theory that the chelate
effect is primarily an (enthalpic/entropic) phenomenon.
Results
[Cu(NH3)4]2+ and [Cu(en)2]2+ were synthesized by the following reactions:
(1) BALANCED reaction for synthesis of [Cu(NH3)4]2+ (use superscripts and subscripts where
appropriate)
(2) BALANCED reaction for synthesis of [Cu(en)2]2+ (use superscripts and subscripts where
appropriate)
Electrochemical cells were constructed to study the redox reactions:
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What were the half reactions and complete balanced reactions that you analyzed?
Electrochemical potential measurements were measured between the different electrodes at
various temperatures and are reported in Table 1.
Table 1. Electrochemical potentials for Cu(NH3)42+ and Cu(en)22+ versus different CuII electrodes
Temperature (K) Cu(NH3)42+ vs. Cu2+ (V) Cu(en)22+ vs. Cu2+ (V)
Cu(NH3)42+ vs. Cu(en2)2+ (V)
Fill in this chart. The whole thing should Move it around if it's not
be on one page.
From electrochemical measurements, thermodynamic data can be calculated from Equation 3
which equates the two fundamental thermodynamic equations, 4 and 5.
(3) Put the equation here
(4) one equation here and
(5) one equation here
The resulting calculated thermodynamic data are tabulated in Table 2.
Table 2. Thermodynamic data and equilibrium constants derived for Cu(NH3)42+ and Cu(en)22+.
ΔΗ (kJ/mol)
2+
Cu(NH3)4 vs. Cu
2+
ΔS (J/mol K)
Fill this table out
and include errors
Cu(en)22+ vs. Cu2+
Cu(NH3)42+ vs. Cu(en2)2+
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ΔG (kJ/mol)
Keq (give units)
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Insert plots with appropriate captions here. Include error bars. All axes should be clearly labeled.
If you are using Excel, don’t connect the points, show fit lines and R values, and get rid of the
horizontal lines.
Discussion
Our results confirm/conflict with the theoretical principle. The chelate effect is
. Explain how your data
ethalpically/entropically driven, and our data
agrees or disagrees. Evaluate the uncertainty in your data and the effect this has on your results.
which could be prevented by
The error was a result of
.
In an upcoming study, we will seek to improve our methodology. We plan to
. Explain what you're doing and why you think it will improve your results.
Conclusions
Thermodynamic data for
were obtained from plots derived from
Equation 3. These results confirm/conflict with the theory behind the chelate effect because
. Write a one sentence explanation of your error.
Experimental
Include any additional observations! Note color changes, gas evolution, temperature changes
(exo/endothermic reaction), precipitates, etc.
General Considerations:
The starting materials, __ List chemicals _, were obtained from XYZ chemical co.
Electrochemical measurements were performed using a
(brand)
.
voltmeter.
Synthesis of [Cu(NH3)4]2+.H2O:
Copper(II) sulfate pentahydrate (
The solution was cooled for
g,
mol) was dissolved in
mL of water at ____ °C.
minutes until it reached _____°C. After cooling,
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mL of
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___M (give molarity) ammonium hydroxide, _____ (give chemical formula and state), was
added to the solution with vigorous stirring. The solution was cooled for
min., then
mL of EtOH was added. The solution was cooled to _____°C. The precipitate was collected by
vacuum filtration and washed
mL EtOH and
min and in an oven at
dried in air for
g(
times with
mol;
°C for
mL Et2O. The solid was
min.
%) was collected.
Synthesis of [Cu(en)2]2+.2H2O:
Using the same outline as the synthesis of Cu(NH3)42+.2H2O write your own synthesis of
Cu(en)22+ .2H2O (DO NOT just copy and paste the above section)
Electrochemical Analysis of Copper Solutions:
Electrodes were constructed from a solution of gelatin (
water heated to _____°C . Three solutions were prepared:
of [Cu(NH3)4]2+.H2O, and
g) and KCl (
g ) in
mL
M of CuSO4.5H2O,
M
M of [Cu(en)2]2+. 2H2O. For each measurement, the solutions
were equilibrated at a given temperature using a heat bath.
. Explain
how you determined the solution was equilibrated. The potential was measured at each
temperature between
at
temperatures. Which potentials did you measure and
at what temperatures?
References
1. Basolo, F.; Pearson, R. G.; Mechanisms of Inorganic Reactions, 2nd Ed. Wiley 1967, 27ff,
223ff.
i
Cotton, F. A.; Chen, H.; Daniels, L. M.; Teng, X. J. Am. Chem. Soc. 1992, 114, 8980.
ii
Nakamoto, K. Infrared and Raman Spectroscopy of Inorganic and Coordination Compounds.
Part A: Theory and Applications in Inorganic Chemistry, .Fifth edition. Wiley, NY, 1997.
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Arene Proposal Template
Name________________________
Student ID Number_____________
Partner_______________________
Section_______________________
Arene Molybdenum Tricarbonyl Chemistry- Part II Proposals
Proposal 1:
Reaction of interest:
(Write a balanced equation for a reaction you intend to perform, draw the structures of any
starting materials or proposed products. SciFinder is a great resource for finding reactions and
some papers are in a blue binder located by the computer in lab.)
Scientific Rationale:
(What principle is being investigated? What will performing this reaction/ analyzing this
product tell you about organometallic chemistry? How does this experiment build on your
investigations from part I?- This section must be written in narrative form using appropriate
scientific language. Approximately 1 paragraph will typically be sufficient.)
Brief Procedure:
(This may be written as bullet points but must include details such as volume/mass of material,
reaction times, reaction temperature, any hazards associated with a specific step- including
strong odors, carcinogenic/toxic reagents, and whether air free techniques will be used for some
or all of the procedure)
Timeline:
(You and your partner will have four experiments to choose from, two provided by each partner.
The timeline will help you decide how long each of the reactions will take and may be useful as
you decide which to perform. This should include not just reaction times but also estimates of
how long it will take to assemble the apparatus, perform the work up, and characterize the
product. Indicate times when one or both partners can be working on another reaction.)
Materials List:
(Compile a list of all chemicals that will be used during this lab including reagents, solvents,
NMR solvents, filtering agents, etc. Your list should include the quantity of each material.
Review the MSDS sheets for each material that you plan to use. If any seem especially
hazardous attach a copy of the MSDS for the TA to review. Mark any chemical for which you
are attaching an MSDS with a “ #”. Place an “*” after any material that is not on the chemical
list located in the Chem317 lab so that we may procure it for you. Make sure that all of your
starting materials not already available in the lab are available from a commercial supplier in
quantities > 5 g.)
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Special Equipment (optional):
(List any equipment that you will need that is not listed on the materials list or stockroom list in
the Chem317 lab.)
Reference:
(Cite at least one reference where the procedure for this synthesis or a related synthesis can be
found. Use proper ACS citation style.)
Supporting Information:
A reference detailing the procedure for this reaction or a related compound is attached and
labeled “Proposal 1”. (If the procedure is for a similar molecule include a brief statement of why
you think this procedure is appropriate for the reaction you plan to perform. If you are applying
a procedure in a new way use evidence from a textbook or other journal articles to show why you
think the procedure will work. Do not submit references which use your product as the starting
material but cite the original synthesis elsewhere. You will not receive credit – or be allowed to
perform a synthesis – without the original synthesis reference.)
Proposal 2:
Reaction of interest:
(Write a balanced equation for a reaction you intend to perform, draw the structures of any
starting materials or proposed products. SciFinder is a great resource for finding reactions and
some papers are in a blue binder located by the computer in lab.)
Scientific Rationale:
(What principle is being investigated? What will performing this reaction/ analyzing this
product tell you about organometallic chemistry? How does this experiment build on your
investigations from part I?- This section must be written in narrative form using appropriate
scientific language. Approximately 1 paragraph will typically be sufficient.)
Brief Procedure:
(This may be written as bullet points but must include details such as volume/mass of material,
reaction times, reaction temperature, any hazards associated with a specific step- including
strong odors, carcinogenic/toxic reagents, and whether air free techniques will be used for some
or all of the procedure)
Timeline:
(You and your partner will have four experiments to choose from, two provided by each partner.
The timeline will help you decide how long each of the reactions will take and may be useful as
you decide which to perform. This should include not just reaction times but also estimates of
how long it will take to assemble the apparatus, perform the work up, and characterize the
product. Indicate times when one or both partners can be working on another reaction.)
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Materials List:
(Compile a list of all chemicals that will be used during this lab including reagents, solvents,
NMR solvents, filtering agents, etc. Your list should include the quantity of each material.
Review the MSDS sheets for each material that you plan to use. If any seem especially
hazardous attach a copy of the MSDS for the TA to review. Mark any chemical for which you
are attaching an MSDS with a “ #”. Place an “*” after any material that is not on the chemical
list located in the Chem317 lab so that we may procure it for you. Make sure that all of your
starting materials not already available in the lab are available from a commercial supplier in
quantities > 5 g.)
Special Equipment (optional):
(List any equipment that you will need that is not listed on the materials list or stockroom list in
the Chem317 lab.)
Reference:
(Cite at least one reference where the procedure for this synthesis or a related synthesis can be
found. Use proper ACS citation style.)
Supporting Information:
A reference detailing the procedure for this reaction or a related compound is attached and
labeled “Proposal 1”. (If the procedure is for a similar molecule include a brief statement of why
you think this procedure is appropriate for the reaction you plan to perform. If you are applying
a procedure in a new way use evidence from a textbook or other journal articles to show why you
think the procedure will work. Do not submit references which use your product as the starting
material but cite the original synthesis elsewhere. You will not receive credit – or be allowed to
perform a synthesis – without the original synthesis reference.)
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Notes on the Arene Molybdenum Tricarbonyl (Part I) Lab Report
The following are a list of common formatting issues, “problems,” and points of confusion that
were noticed amongst the first draft of “Arene” reports. The items that follow are included here
because they occurred in a lot of reports, so it behooves you to read this document carefully
before writing your final draft of the Arene lab report.
The Introduction:
Do not simply change the order of words in the lab manual. If you mention specific
chemical examples that are given in the lab manual introduction to the arene lab (such as
Cp2TiCl2 or [Rh(CO)2I2]-) then you should elaborate on those examples in some way that shows
off the extra research that you have done on those compounds.
The idea behind this class is for you to become more familiar with accessing the
chemical literature. Find some articles, read them (or at the least, skim them), and then
synthesize your newly acquired knowledge into an introduction. Review articles are probably
your best bet since they cover a wider breadth of topics than most journal articles.
This section should acquaint your reader with organometallic chemistry, perhaps give
some specific examples of interesting compounds or reactions, and eventually focus on the
electronic properties of carbon monoxide ligands. By the end of the introduction your reader
should have a rough idea of which experiments you are going to do and should understand why
you decided to do those specific experiments.
The Experimental Section:
This section of a lab report or journal article is very specialized. The purpose of this
section is to clearly relate to the reader exactly what you did. It should also contain a concise list
of all spectroscopic observations and other data (such as percent yield). Other observations
which could be helpful to someone trying to duplicate your work (such as color changes) should
also be included. This section should be separated into paragraphs for each compound studied
(such as (η6-mesitylene)Mo(CO)3).
This section contains “Just The Facts.” Any data interpretation, explanations, or suspicions
should be stated elsewhere.
The experimental section should:
• be written in the past tense
• use specific, descriptive wording
• specify which instruments were used to acquire data (model numbers or other
specifications)
Infrared Spectroscopy
Within the experimental section, stretches are reported in the following format:
IR (method of sample preparation): frequency of stretch (strength of stretch, νatoms involved).
For example:
IR (KBr): 3011 cm-1 (m, νCH), 1810 cm-1 (s, νCO)
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-s, m, w, br are common abbreviations used to describe the shape of IR signals. These
abbreviations stand for strong, medium, weak, and broad.
-Common methods for preparing infrared samples are as KBr pellets (KBr), nujol mulls (nujol),
solutions (_____ solution).
NMR Spectroscopy
Spectra should be reported as follows (in the experimental section):
Nucleus NMR (NMR frequency, solvent): δ chemical shift (multiplicity, coupling constant (if
applicable), integration, peak assignment). For example:
1
H NMR (500.0 MHz, acetone-d6): δ 7.02 (t, JHH = 7.5 Hz, 1H, Ar-H para), 6.86 (d, JHH = 7.5
Hz, 2H, Ar-H meta), etc.
When interpreting and/or reporting NMR spectra, keep in mind:
• If your spectrum contains an internal standard then the spectrum should be referenced to
that signal. Referencing involves manually setting a signal with a known chemical shift
(such as TMS) to it’s known chemical shift value. This is a way of ensuring that signals
have the “correct” chemical shift and can be compared between labs across the globe.
• If you suspect that one or more signals come from an impurity, do your best to identify
that impurity. You can compare the observed chemical shift to the published chemical
shift (a helpful reference with many chemical shifts was distributed earlier entitled “NMR
chemical shifts of common laboratory solvents as trace impurities”). Alternatively (or
additionally) you can add a small quantity of the suspected impurity to your sample tube
and take another spectrum. If the impurity peak intensity increases and the rest of the
spectrum is unchanged then you have correctly identified the impurity! This is called a
“spike.”
• When labeling your NMR spectra, use labels that are as descriptive as possible. For
example, your labels for the 1H NMR spectrum of mesitylene could be Ar-H and Ar-Me
(these are more informative than labeling the same peaks a and b).
Helpful Hints
• Some of you may have noticed that samples of the same compound often have slightly
different infrared stretching frequencies when they are prepared in different ways (such
as a KBr pellet vs. hexane solution). So, if you want to compare IR spectra of two
different compounds, try to acquire spectra using the same method (ex: both samples
prepared as KBr pellets). Then the only thing different between the two samples is the
identity of the compounds you are studying.
• Referring to compounds as “the product,” “the reactant,” or other similar names is not
effective. Essentially every reaction that you will conduct in this class will have multiple
substances that could be “the product” or “the reactant” for any particular reaction. Use
wording that is as specific as possible, or, better yet, write out the chemical formula.
Using formulas is often faster and takes up less space than other options
• Within Microsoft Word: a highlighted character may be turned into a superscript or a
subscript character by pressing the following buttons simultaneously:
Ctrl, Shift, and +(superscript) orCtrl, and =(subscript)
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…Hopefully this makes typing chemical formulas a more attractive prospect!
When listing IR or NMR data for a compound you should exclude signals attributed to
impurities. If you know that a signal comes from something other than your compound,
then it doesn’t make sense to include it, for example, in a table entitled “1H NMR peak
assignments for (mesitylene)Mo(CO)3.” Including spectral data for impurities in the
same place as spectral data for “desired compounds” is an easy way to confuse your
reader.
• Discussions of electron density are more meaningful if they are accompanied with a
specific atom, orbital, or other location.
ex: “This ligand substitution reaction resulted in a compound with less electron density in the d
orbitals.”
•
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