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APPENDIX
USE OF GLASSWARE
Before beginning an experiment all glassware that will be used to perform the experiment must
be cleaned. Water sheets off of clean glassware, e.g. water does not bead on the surface of the
glass. Usually, most glassware needs a good rinse with tap water and then distilled water before
use. However, if oily or organic based solutions have been used, a scrub with soapy water is
needed, then a rinse with tap water followed by a rinse with distilled water.
The choice of which glassware to use is dictated by the experiment. It is important to ask how
precise and accurate does the experiment need to be in order to determine which glassware
would be most suitable. Sometimes a clue is given in the procedure, volumes to two decimals
(0.01 mL) require the use of volumetric ware and masses to 4 decimals (0.0001g) require use of
the analytical balance. Typically, in an analytical experiment, volumetric glassware and
equipment is used as the precision and accuracy needed is high. Pipets, (volumetric and
graduated), burets and volumetric flasks are calibrated in the factory so that the volume they
measure is both very accurate and precise.
Other glassware (graduated cylinders, beakers) found in the lab is also calibrated, however, the
calibration is rough and intended only to give an approximate amount, measuring cylinders being
more accurate than beakers. Often in synthetic and preparative labs, a top-loading balance and
measuring cylinders are adequate to determine correct quantities of reagents and solutions.
USE OF VOLUMETRIC WARE
Buret
Clean the buret by washing with detergent solution (if necessary) and warm water. Using a
beaker, thoroughly rinse the buret with tap water. (Do not try to fit the buret directly under the
tap). A final rinse with small portions of distilled water completes the washing. When liquid
drains from a buret and no droplets are left clinging to its inner walls, the buret is clean. Clamp
the buret upside down, stopcock open, to drain.
Before filling, the buret must be prerinsed three times with small amounts of the solution it will
deliver. Place - 5 mL of this solution in the buret; rotate it in a horizonal position so that all
parts of the body are wetted by the solution. Be sure to drain some solution through the
stopcock, discarding the remainder.
Fill the buret, with the aid of a funnel, to above the zero calibration mark, being sure the
stopcock and tip are filled and no air bubbles remain. Allow the solution to drain into a waste
container until the meniscus is just below zero. As all readings are taken by difference, it is a
waste of time to fill the buret to exactly zero. Remove the funnel from the buret before any
readings are taken as an extra drip unaccounted for will affect the results.
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In reading all the volumetric ware, the position of the meniscus must be accurately located. This
is affected by the position of the eye. To avoid this parallax error, it is important to read the
meniscus at eye level. A reading card is helpful to obtain reproducibility in buret readings.
Volumetric Pipet
A volumetric or transfer pipet is designed to deliver a particular fixed volume. Liquids are
usually drawn into pipets by application of a slight vacuum. The pipet is cleaned and drained
using the same method as for the buret.
Using a bulb, draw in a small quantity of the liquid to be sampled and thoroughly rinse the
interior surfaces of the pipet. Discard the rinsings. Repeat with two more portions.
To measure an aliquot, carefully fill the pipet somewhat above the calibration mark. Quickly
place a forefinger over the upper end of the pipet to arrest the outflow of liquid. Tilt the pipet
45B from the vertical and wipe the exterior free of adhering liquid with a Kimwipe (tissue).
Touch the tip of the pipet to the wall of a waste beaker and slowly allow the liquid level to drop
by partially releasing the forefinger. Halt further flow as the bottom of the meniscus coincides
exactly with the calibration mark. Place the tip of the pipet well into the receiving vessel and
allow the sample to drain. When free flow ceases, rest the tip against an inner wall for a full l0
seconds. Finally, withdraw the pipet with a rotating motion to remove any droplet adhering to
the tip. THE SMALL VOLUME REMAINING INSIDE THE TIP IS NOT TO BE BLOWN OR
RINSED INTO THE RECEIVING VESSEL.
Graduated Pipet
A graduated pipet is designed to deliver a volume of your choosing. The size of the pipet is
obviously chosen for the working range, but typically, you will encounter them in the 1 - 10 mL
size. Liquids are usually drawn into pipets by the application of a slight vacuum. Using a bulb,
draw in a small quantity of the liquid to be sampled and thoroughly rinse the interior surfaces of
the pipet. Discard the rinsings. Repeat with two more portions.
To measure an aliquot, carefully fill the pipet somewhat above one of the higher calibration
marks. Quickly place a forefinger over the upper end of the pipet to arrest the outflow of liquid.
Tilt the pipet 45B from the vertical and wipe the exterior free of adhering liquid with a Kimwipe.
Touch the tip of the pipet to the wall of a waste beaker and slowly allow the liquid level to drop
by partially releasing the forefinger. Halt further flow as the bottom of the meniscus coincides
exactly with the calibration mark. Place the tip of the pipet well into the receiving vessel and
allow the sample to drain slowly down to the lower calibration mark, thus delivering the desired
volume. Rest the tip against an inner wall for a full l0 seconds to allow the drip to drain off the
outer surface.
Volumetric Flask
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Volumetric flasks are calibrated to contain the specified volume when filled to the calibration
mark. They are used in the preparation of standard solutions, and in the dilution of samples to
known volumes prior to taking aliquot portions with a pipet or other means of sampling.
As with other volumetric ware, the flask must be cleaned before use. If used for dilution the
required volume of stock solution is transferred by pipet to the volumetric flask and distilled
water is added with a wash bottle to just below the calibration mark. The last few drops of
distilled water are carefully added with a medicine dropper until the bottom of the meniscus is
just resting on the calibration mark seen at eye level. To ensure uniform concentration, stopper
the flask and mix ten times by inversion.
When dissolving solids in a volumetric flask, it is recommended that all salt be dissolved by
swirling after about 2/3 of the required amount of distilled water has been added. After all the
solid has dissolved, the flask is topped up with the solvent as outlined above.
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LABORATORY TECHNIQUES
Titration
Carry out a titration as follows:
Ensure that the indicator used in the titration has been added before starting. Read and record the
initial buret reading to the nearest 0.02 mL.
Place the conical flask on a white tile under the buret. Controlling the flow of the solution from
the buret, allow solution from the buret to run into the Erlenmeyer while the flask is kept in
constant rotary motion. Continue to add titrant from the buret at a rapid drop rate as long as any
colour formed disappears immediately. As soon as the colour shows any tendency to linger,
slow down the rate of addition.
At this point, stop titrating and rinse down the inner walls of the Erlenmeyer flask with a small
spray of distilled water from a wash bottle. Continue slow dropwise addition of titrant until a
very faint colour persists throughout the whole solution for at least 30 seconds. This is the end
point. If the endpoint is uncertain, record the volume and continue until the endpoint is certain.
Read the buret (to the nearest 0.02 mL) and record the volume.
Use the value of the volume of titrant used as a guide in further titrations. Typically, the
procedure is repeated to give a total of three titrations.
Use of a Titration Table
A convenient and efficient way to display titration data is to use a titration table. Below is part of
a titration table summarising raw data that was obtained during the experiment and then a
continuation of rows gives the various calculations that were required to determine the final
result(s).
Table 1. Determination of the isolation and analysis of Y
Trial 1
Trial 2
Trial3
Final vol.of X ( mL)
25.9
26.14
25.86
Initial vol. of X (mL)
1.21
1.45
1.20
Amount of X added
(mL)
24.69
24.69
24.66
moles of X added
(moles)
0.01234
0.01234
0.01233
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moles of Y reacted
0.04936
0.04936
0.04932
mass of Y used up
(g)
19.55
19.55
19.54
average Y amount (g)
19.55
Gravity Filtration
At times it is necessary to remove solid particles from a solution. One simple technique for doing
this is gravity filtration. Filter paper is folded in such a manner that it can be placed in a glass
funnel. The glass funnel is usually placed in a conical flask. The solution is then introduced into
the folded filter paper and the solution slowly drips through with the solid particles being held
behind.
Vacuum Filtration and Washing isolated Solids
There are two convenient ways to separate a precipitated solid from
its mother liquor. The first and easiest method is to allow the solid to
settle and then pour off (decant) the upper liquid layer (supernatant
liquid). This can only be done if the solid is heavy and wellcoagulated. The remaining solid can then be washed, separated
(perhaps using a centrifuge) and dried.
The second, more common method is to filter the solid from the
mother liquor. In many cases, gravity provides sufficient force to
drive the liquid through a funnel lined with filter paper. Sometimes, a
vacuum applied to the lower end of the funnel will speed up that
process. This technique is known as a vacuum or Buchner filtration.
The source of vacuum is usually a water aspirator. The tap is always turned fully on. The hose is
attached to a Buchner flask (this is an Erlenmeyer flask made of thick glass, with a side-arm) and
flask is clamped to prevent tipping - a common problem. A rubber filtervac is placed between the
top of the flask and the porcelain funnel. Moistening the filtervac will help create the pressure
seal, but this can only be done if the water will not affect the filtrate. Filter paper is placed in the
funnel, and moistened to make the paper sit securely over the holes in the funnel. Consider what
would be a suitable solvent with which to moisten the filter paper. Using water is only suitable if
the residue or filtrate are not water sensitive. If they are, a second choice would be to use a little
of the solution to soak the filter paper. The apparatus is now ready for use.
Water aspirators have a habit of letting a little tap water run back into the flask - even if the
correct procedure is followed. (Note that if you turn off the water before breaking the vacuum
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seal at the filtervac, you will certainly see a lot of water flow back!) For this reason, gravity
filtration is usually recommended if the filtrate is the desired fraction. If a vacuum filtration is
absolutely necessary then it is usual to place another vessel between the flask and the water
aspirator to act as a water trap.
When washing a precipitate in a filter funnel, remove the filtrate first. (Sometimes the washing
liquid may react with the filtrate already collected. This is particularly relevant when nitric acid
is in the reaction mixture and the filtered residue is subsequently washed with ethanol. Ethanol
and concentrated nitric acid form an explosive mixture!)
Then, without applying suction, add to the filter funnel the wash solution/solvent. Using a glass
stirring rod, very gently stir the wash solution/solvent with the isolated solid. Ensure that the
filter paper does not get broken. Attempt to suspend and rinse the solid with the wash
solution/solvent. Then, apply suction to remove the wash solution/solvent. Repeat if and as
indicated.
Standardizing pH Meter for measuring pH of solutions
1.
Have function selector on STANDBY and lower combination electrode and thermometer
into buffer solution. (If using 610A or 910 model, set slope control at 100%).
2.
Wait about 2 minutes until electrodes and buffer solution reach thermal equilibrium, then
adjust TEMPERATURE control to temperature of buffer.
3.
Turn function selector to pH and adjust STANDARDIZE control so digital display reads
exact pH of buffer while swirling the solution.
4.
Turn function selector to STANDBY, raise combination electrode and thermometer and
rinse them with distilled water. Blot water droplets with kimwipe tissue.
Performing pH Determinations using a pH meter
1.
Lower combination electrode and thermometer into sample solution, wait until electrodes
and solution reach thermal equilibrium and adjust TEMPERATURE control to
temperature of sample solution.
2.
Turn function selector to pH. Read the indicated pH of sample all the while swirling the
solution.
3.
Turn back to STANDBY, raise electrode and thermometer from solution and rinse them
with distilled water. Blot them dry. Repeat steps 1 to 3 as required for further
determinations.
4.
When not measuring the pH of solutions, the combination electrode is to be kept in
distilled water.
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Preparation of Silver Electrode
Pick up a Ag+ electrode from the demonstration table. Gently sand the surface of the bottom of
the electrode when necessary. It is reactivated by polishing the surface with a very small
amount of AgCl/Ag2S powder. Use a small fire polished soft glass rod to do this. Continue
polishing until a smooth shiny surface is obtained and no graphite/teflon surface is visible
anywhere. This AgCl/Ag2S layer will function the same as a AgCl/Ag2S crystal. After
polishing, obtain a few millilitres of 0.1 M KCl solution and soak the electrode for 5 minutes.
Rinse the electrode with distilled water and leave it in distilled water for 5 minutes.
Zeroing the Meter for Use with the Selective Ion Electrode
The following procedure is to be carried out to check and adjust the zero reading on a Fisher pH
meter 610 & 610a when it is to be used as a digital millivolt meter.
With the meter on "Standby":
1) Disconnect the combination electrode from the meter, connect a shorting cable (one
end into the reference (REF) outlet and the other pin into a pin jack connector for
insertion into the INPUT jack).
2) Check the millivolt reading (turn knob from Standby to mV)..
3) Adjust the standardize control knob until the digital display indicates zero millivolts.
4) Turn knob from mV back to Standby.
5) Remove the shorting cable and the pin jack connector.
6) Connect the electrodes as required.
Operating Procedure for the LKB Novaspec Spectrometer
1.
Before turning on the spectrophotometer check that the sample compartment is
empty, then close the cover for the compartment.
2.
Set the ON/OFF rocker switch located at the back of the instrument to ON. The
instrument will run a self-diagnostic calibration program that takes about 90
seconds. The WAVELENGTH indicator on the display panel will read CAL
(calibration) and the MEASUREMENT indicator will count from 1 to 3. Do not
interrupt the program and do not open the sample compartment cover during
calibration. When the calibration program is complete, the WAVELENGTH
indicator will display the reading 360 nm. The instrument is now ready.
3.
Press the MODE key until the desired indicator on the display panel is
illuminated. (In most experiments, but not all, the solutions to be studied are
expected to have absorbance readings in the 0 to 1 range.)
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4.
Set the wavelength to the desired value by pressing and holding down either the
WAVELENGTH - or WAVELENGTH + key. The selected wavelength will be
indicated on the display panel.
5.
To zero the instrument, first wipe with a Kim-wipe, then insert the cuvette with
the 'blank' solution into the sample compartment and close the cover. Then press
the SET REF key (set reference). The MEASUREMENT indicator will display
0.00.
6.
Remove the cuvette containing the 'blank'. Insert a wiped cuvette containing the
sample to be studied into the sample compartment. Close the cover. Read and
record the absorbance reading displayed on the MEASUREMENT indicator on
the panel. Repeat this step with other samples as directed in the experimental
procedure.
Using a tlc Plate
Tlc plates, made with silica, are usually available from the demo bench. They are already cut to
the desired size. If there is a ragged short edge, it is important that it be used as the top edge of
the plate. There is a shiny and a matte side to the plate. The shiny side is the plastic backing that
holds the silica in place. The matte side is the side with the silica on it and is the side that is
spotted.
The solutions to be separated are spotted on the plate using a drawn out capillary tube. This
procedure will be demonstrated by your TA. When spotting, take extreme care not to break the
surface of the silica. Ideally, the capillary tube should not come in contact with the plate.
In the fumehood, the developing chamber is prepared by placing the developing solution into the
chamber along with filter paper wrapped partially around the inside of the chamber (usually a
beaker) and then covered.
After the spotted solutions have dried , the tlc plate is then very carefully placed end down into a
developing chamber. The filter paper which is now saturated with the solvent inside of the
developing chamber is used to keep the atmosphere of the chamber saturated with solvent vapour
which helps reduce the evaporation of the solvent from the plate. The chamber is covered and the
plate is allowed to develop. During the developing, the chamber is not to be moved as any
solution that splashes onto the plate can ruin the separation.
The plate is carefully monitored and once the solvent front if within about 1cm. of the top of the
plate, the plate is removed and the solvent front is immediately marked with a pencil. The plate
is then dried and treated as required for whatever separation was done.
Semi-Micro Qualitative techniques
In semi-micro qualitative analysis, for the most part, only drops of solution are used, and
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amounts of reagent and cations are in the weight range of l0 - l00 mg. In order to bring about
clean separation of compounds and to obtain satisfactory confirmatory tests, equipment must be
cleaned immediately before beginning work and immediately after use.
Most reactions are carried out in l0 x 75 mm test tubes and, since a given test tube will be used
over and over again during the analysis, each must be cleaned many times. If the test tube has
last had only solution in it, three rinsings of water from a wash bottle will be sufficient, but if
solid adheres to the walls cleaning with a brush/pipe cleaner, followed by rinsing with distilled
water will be necessary.
It is absolutely essential to handle all reagent bottles so that contaminants are not
introduced into the reagent solutions. If you make a mistake and contaminate a reagent bottle,
report this to the instructor at once, so that the work of other students is not invalidated.
Solutions are transferred by means of small pipets. As soon as you have finished using a pipet
for a specific transfer, clean it, allow distilled water to be drawn into the tube, expel it into an
empty beaker and repeat this three times. After the third rinse, again fill the pipet with distilled
water and store it in distilled water until required.
Equipment Set-up
Before you start the experimental procedure, set up your working space as follows:
Arrange your equipment in an orderly fashion, so that everything is at hand, but
not in the way of your working space. Hot plates are typically shared with other
students.
Put about 50 mL of tap water in a l00 mL beaker, set it on a hot plate and turn on
the switch to about the half way mark.
Clean a 250 mL beaker, fill it with distilled water, then place it near an empty 250
mL beaker. The distilled water will be used for washing pipets, the empty beaker
for collection of waste liquids.
Clean all equipment by rinsing three times with distilled water. Fill the wash
bottle with distilled water. Place a clean pipet and stirring rod in the beaker of
distilled water, and other clean equipment on a piece of paper or paper towel.
Washing a Precipitate
In semi-micro qualitative analysis, the amounts of solution and precipitate are too small to carry
out filtrations in the normal way; instead, the solid is collected in a compact mass at the bottom
of the test tube by centrifugation, and the solution with which the solid is in contact, the
supernatant solution, often called the supernatant, is removed by means of a pipet.
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To wash a precipitate in a test tube, add distilled water (or the solution specified) to the
precipitate. Break up the mass of solid with a clean stirring rod and mix so that the small
particles of solid make good contact with the solvent. Re-centrifuge the mixture. Discard the
supernatant unless otherwise directed.
Flame Test
Use a piece of nichrome wire one end of which has been fastened into a cork (the handle) and the
other end bent into a loop. Light the bunsen burner and set up a cool flame. Clean the loop by
first, dipping it into concentrated hydrochloric acid contained in a l0 x 75 mm test tube, and then
holding it in the flame just above the blue cone. Repeat the acid and flame treatment until the
flame shows no colour, then dip the clean loop into the solution of precipitate to be tested, and
hold it in the base of the flame. Several ions show a characteristic colour in the flame.
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LABORATORY REPORTS
The Laboratory Notebook
Documenting the experimental work is a crucial task of a scientist. Poor note keeping makes it
difficult to determine what really happened in a experiment and what legitimate conclusions can
be drawn from the results. A good set of notes provides enough information to allow the reader
to redo the experiment (by following the notes) and adequately compare the results. It is always
wise to document more than less.
Your laboratory notebook is to be a permanent record of the work that you did in the lab. It is to
be an accurate and unambiguous record of your lab experiments. On the first page, a Table of
Contents is to be prepared and kept up to date. The top of each page that has been used during
the experiment is to be filled out. During the lab period, all data and observations are to be
recorded in ink directly into your notebook. Only documentation written in pen will be accepted
. Any information not recorded in your notebook cannot be used in the writing of your report. At
the end of the experiment, each page is to be signed and dated by you and countersigned by your
TA. You are then to give the copy page(s) to the TA as it will be graded on the quality of your
note taking.
Writing the Report
At the end of each experiment, there is a summary as to what is to be included in the report. To
aid you in writing good lab reports, not all sections are required for each report, there being a
focus on specific parts.
The formal lab report is to be written on 8.5 x 11 inch paper and bound in a Duotang. Only that
which is written in ink or generated on a computer and submitted in a duotang will be read. If a
mistake is made, cross out, using one line, the incorrect value and write the new value next to it.
Do not use white-out. Correct scientific notation and formulations are to be used. The report is
to be written in the third person, impersonal, past tense, and passive voice, with all the pages
numbered.
The following example gives an experiment and then provides a general outline as to what is
expected in a full report for this experiment. There are many web sites available that give
information on report writing, The one below could be helpful.
http://odtl.dcu.ie/wp/1999/odtl-1999-03.html
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The experiment:
THE PREPARATION OF POTASSIUM DICHROMATE & STUDY OF THE OXIDATION
STATES OF CHROMIUM
Many transition metals can form compounds with varying oxidation states. In this preparation,
potassium dichromate, K2Cr2O7 will be made from chromium acetate, Cr(CH3COO)3AH2O and
then further reacted. The chart below provides information to help determine which oxidation
state the chromium is in and what species are in solution throughout the preparation and
subsequent reduction.
Table 1. Some oxidation Species of Chromium
Oxidation
number of
chromium
Ionic species present
Acidic solution
+6
Cr2O7-2 (orange)
+3
Cr(H2O)6+3 (violet)
0
Basic solution
CrO4-2 (yellow)
Cr(H2O)2(OH)4- (green)
Cr metal
Experimental Procedure
1. Weight out 2.47g of chromic acetate. Dissolve in 20 mL of water. Gently warm if necessary.
2. Collect 25 mL of 6M KOH in a 250 conical flask. Add the chromium acetate solution
(prepared above) slowly. Swirl the flask. Note what happens.
3. Add 20 mL of 3% hydrogen peroxide, add 2 boiling chips and heat the flask to boiling. If the
solution does not turn yellow in two minutes of boiling , add 10 mL of hydrogen peroxide.
Repeat , if necessary.
4. After the solution has turned yellow, gently boil the solution until the volume has been
reduced to 30 mL.
5. Allow the solution to cool to room temperature.
6. Add 18M CH3COOH, dropwise, until the solution becomes orange. Add 1 mL of 18 M
CH3COOH in excess.
7. Heat the solution to boiling and reduce the volume to 25 mL. Then, gradually cool the
solution, using an ice bath to crystallise K2Cr2O7.
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8. Filter the solid and wash with two portions of ethanol. Dry and weigh the crystals. Hand in a
sample to your TA.
9. Weigh 1.47g of K2Cr2O7 into a 250mL beaker and add 20mL of 3M H2SO4. Heat gently to
dissolve the potassium dichromate. Once all is dissolved, cool solution to room temperature.
10. Set up the SO2 generator as demonstrated by your TA. Follow the directions carefully as
outlines in the lab for the reduction. Ensure that the dichromate solution remains below 40EC
while the reduction ensues. Allow the reaction to continue until the solution turns from orange to
blue-green. (May take up to 15 minutes.)
The example report:
Experiment number, title, date& your name
Earl N. Meyer
July 3, 2008
Experiment 2. Preparation of Potassium Dichromate and the Study of
the Oxidation States of Chromium
Summary/ Abstract
Summarizes the content of the report. It defines the problem, describes the methods used in the
experiment, summarizes the results and gives the conclusion. It is typically short (100 words),
not more than a paragraph.
Potassium dichromate was prepared from a solution of chromic acetate that was
oxidised by H2O2, then acidified and cooled to precipitate orange K2Cr2O7. 0.345g
was isolated (23.0% yield). A sample of K2Cr2O7 was then allowed to react with
SO2,which reduced the chromium to form blue-green coloured KCr(SO4)2A12H2O.
Object
States the purpose of the experiment, usually contains an infinitive, ie. to + action verb, for
example to determine.
The object of this experiment was to prepare potassium dichromate and note the
various colors of the oxidation states of chromium throughout the preparation
and reaction of the product.
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Theory
Gives the concepts, ideas, and suppositions that are considered to explain the point of the
experiment. It introduces the reader to what the experiment is about. Equations are usually used
to describe what was done. It does not include the details of the procedure.
Chromium, a transition metal, has a several oxidation states(numbers). Oxidation
numbers are an arbitrary system of keeping track of electrons that are gained by
the substance reduced and lost by the substance oxidised 3. The typical oxidation
states and species(colour) of chromium are given below 1.
Table 1. Oxidation States of Chromium
Oxidation
number of
chromium
Ionic species present
Acidic solution
Basic solution
+6
Cr2O7-2 (orange)
CrO4-2 (yellow)
+3
Cr(H2O)6+3 (violet)
Cr(H2O)2(OH)4- (green)
0
Cr metal
Each chromium species can be easily identified by colour under acidic or basic
condition. Hydrogen peroxide is a suitable oxidising agent in basic solution. In basic
solution H2O2, which is a weak acid forms HO2- or O2-2. The half reactions1 , with
EEvalues, in basic solution are:
4OH- + Cr(OH)4- 6 CrO4-2 + 4H2O + 3e-
+ 0.13 volts
3OH- + 6 HO2- + H2O + 2e-
+ 0.88 volts
SO2 is also a suitable reducing agent as illustrated by the appropriate half
reactions.
SO2(g) + 2H2O 6 SO4-2 + 4H+ + 2e-
- 0.17 volts
2Cr+3 + 7H2O 6 Cr2O7-2 + 14H+ + 6e-
-1.33 volts
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Procedure
Usually refers to the appropriate pages in the lab manual, and includes any deviation from the
original procedure. It does not include a verbatim copy of the lab manual.
The experimental procedure followed was as described in the manual 2.
Data, Results
Usually consists of a tabular display of the data obtained during the experiment and the results
obtained through calculations.
Table 2. Preparation of potassium dichromate
Chemical
Amount
Cr(CH3COO)3AH2O
2.49 g
6M KOH
26 mL
3%H2O2
40 mL
boiling chips
2
18M CH3COOH
12 drops + 1 mL
ethanol
2 x 5 mL
K2Cr2O7 obtained
0.345 g
% yield of K2Cr2O7
23.0%
Table 3. Reduction of potassium dichromate
Chemical
Amount
K2Cr2O7
1.49 g
SO2
generated for 15 minutes
3M H2SO4
20 mL
6M H2SO4
5 x 10 mL portions
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Table 4. Description of Chemical Hazards
Reagents
Hazards 4
Cr(CH3COO)3AH2O
irritant
KOH
corrosive
H2O2
corrosive
CH3COOH
corrosive, lachrymator
ethanol
colourless flammable liquid, irritant
K2Cr2O7
cancer suspect agent, oxidiser
SO2
non flammable gas, corrosive
H2SO4
colourless liquid, oxidiser, corrosive
Calculations
Unless otherwise stated , all calculations are to be shown. Calculations are to be logically
comprehensible and clearly show how the result was obtained.
Calculation for the % yield of potassium dichromate.
In this experiment, it can be assumed that a large excess of hydrogen peroxide
was added to ensure compete oxidation of the chromium, therefore the maximum
amount of K2Cr2O7 that could be produced is based on the amount of
Cr(CH3COO)3AH2O that was used in the experiment. Therefore, Cr(CH3COO)3AH2O
is the limiting reagent.
Moles of Cr(CH3COO)3AH2O = 2.49/ 247.2 g/mole = 1.01 x 10-2 moles
The mole ratio is 2:1, ˆ 1.02 x10-2/2 = 5.10 x 10-3moles of K2Cr2O7.
The maximum number of grams of K2Cr2O7 that could be expected is :
5.10 x 10-3moles x 294.2g/mole = 1.50grams
Actual amount = 0.345 g, ˆ % yield = (0.345/1.50) x 100% = 23.0 %
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Discussion
Consists of discussion about your results and how they match up with what is known by either
theory or previous experimentation. Answer any questions posed in the manual (if there were
any). The limitations of the experimental measurements - and therefore results - are to be
mentioned in this section. Remember that most of the human errors are not to be mentioned in
this section but if necessary in the procedure section.
K2Cr2O7 was produced in a relatively low yield of 23.0%. This value was expected as
the starting material, Cr(CH3COO)3AH2O is difficult to purchase in high purity4.
Typically, it’s purity is around 25%. Given this, the % yield obtained is reasonable.
The preparation went as predicted with the green solution of the
Cr(CH3COO)3AH2O (+3 oxidation state) in base turning yellow (+6 oxidation state)
and then to an orange (+6 oxidation state) color in base.
In the reduction of K2Cr2O7, the orange color (+6 oxidation state) of the solution
readily turned a dark blue green (+3 oxidation state) as was expected.
Conclusion
Consists of a concise statement summarizing the results of the experiment. It relates back to the
object of the experiment. If an unknown was determined, then a complete identification of the
unknown would be included (e.g.Sample # 76, lot #42B3 contained...).
K2Cr2O7 was prepared with a yield of 23.0%.
References
This is a list of the sources of information used. The numbers relate back to the position (an
example might be this 1) in the text indicating from where information was obtained. If an
experiment is done in pairs, the name of the partner is given here. If a commercial product was
used, the specific details describing it would be written here, numerically relating back to when
it was first mentioned in the text.
1.
2.
3.
4.
L. Malm, Chemistry, An Experimental Science, pp.100-103. (W.H. Freeman:
San Francisco). 1963.
M. Reimer, Chem 123 Lab Manual, pp. 1-1-1-3. (University of Victoria:
Victoria, BC). Fall, 1918.
Brown, H.E LeMay, B.E. Bursten, J.R. Burdge, Chemistry, The Central
Science, 10th ed. p. 1017. (Pearson Education Inc.: Upper Saddle River, NJ).
2006.
Aldrich, Handbook of fine Chemicals and Laboratory Equipment ( Aldrich
:Oakville, Ontario). 2003-2004.
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Following are examples of formatting styles used for different references:
For books:
Author or Editor. Title, edition, volume, page(s) or chapter. (Publisher: city). Year.
1.
A. Author. The Moving Boundary and Me, p. 101. (Quick Draw Publishing:
Chemainus, B.C., Canada). 1983.
2.
D. Berry, ed. Chemistry 222 Laboratory Manual, pp. xx - yy. (University of
Victoria: Victoria, B.C.). Spring & Summer, 1998
For journals:
Author. Journal abbreviation year, volume, page(s).
A. Einstein. Ann. Physik 1905, 17, 891-921.
For a web site:
Harris, S. http://www.chemistry.uvic.ca/chemugrd.htm. University of Victoria, BC: last
accessed April 2002.
For a commercial product:
Commercial product. Manufacturer, manufacturer's address [city, province or state & postal
code], lot number and/or drug identification number.
Pep Up Vitamin Pills. Zing Pharmaceuticals, Victoria, B.C. W0W 0H0, Lot no. 54321.
For unpublished work. (can take various forms)
W. Pauli. Personal Communication, May 26. 2006
N. Bohr. Private Communication, May 26, 2006.
L. Pauling. Lab Partner, Chemistry 321, Section 94. (University of Victoria: Victoria,
B.C.). July 2006.
A. Einstein, Chemistry 999 Lecture Notes (Unpublished). University of Victoria:
Victoria, B.C.) Summer term, 2020.
Precision
A-19
The precision of the results might also be determined in some experiments. Precision describes
the reproducibility of the results. It refers to how close the individual results are to each other,
how well the results of each trial agree with one another.
To calculate the precision of the results, the standard deviation (s ) is calculated as given in the
following equation:
where xi = experimental result, '
x = average of experimental results, and n = number of trials.
To determine the relative precision (deviation), %RSD, the equation below is used:
Percent Yield Calculation
The % yield is the actual weight of product obtained expressed as a percentage of the maximum
weight obtainable. From the stoichiometric equation and the weights of the reactants used, the
maximum weight of product that can be expected to be obtained from the starting materials can
be calculated. Values less than 100% typically occur for two reasons: 1) because the reaction
does not go to completion, 2) or losses due to handling can occur.
(1)
(2)
(3)
(4)
(5)
Begin by writing a balanced chemical equation, and examine the ratios between the
reactants (e.g., how many moles of reactant are required for each mole of product?)
From the weights of the reactants actually used and their molar masses, calculate the
number of moles of each reactant that you began with.
Compare the ratio of the moles of reactants in (1) to the ratio in (2), and determine which
reactant is limiting and which is present in excess.
Calculate from the stoichiometric equation (i.e. mole ratio) how many moles of product
expected for each mole of the limiting reagent, and hence how many moles of product
you expect from your actual amount of limiting reagent.
Convert the expected number of moles of product in (4) to grams, and express the actual
weight of product in grams as a percentage of this. (i.e. experimental value/theoretical
value x 100)
Calculation of the Heat Capacity of the Calorimeter
A-20
1.
Calculate the heat lost by the warmer water, and the heat gained by the cooler water.
(Weight of water x ΔT K* x specific heat.) Use: the density of water = 1.00 g/mL, and
the specific heat = 4.18 Jg-1K-1.
2.
The difference, representing the joules lost from the water to the apparatus, divided by
the corresponding ΔT K, gives the heat capacity of the calorimeter in JK-1. Calculate a
value for each trial, and average your results. If the precision of your two results is not
reasonable, carry out a third trial.
An example will illustrate how the heat capacity may be determined, and later used as a
correction factor. Assume that the cooler water is in the calorimeter.
Heat Capacity from the calorimeter data:
Temperature of 50.0 mL of warmer water
Temperature of 50.0 mL of cooler water
Temperature after mixing
Heat lost by warm water:
(50.0 g x 8.3 K x 4.18 Jg-1K-1) =
Heat gained by cooler water:
(50.0 g x 6.7 K x 4.18 Jg-1K-1) =
Heat lost from water to apparatus
Heat capacity of this calorimeter:
(335 J/6.7 K)
*
35.7BC
20.7BC
27.4BC
1735 J
1400 J
335 J
50 JK-1
Note: 0BC = 273 K (Refer to Brown, Le May, page 15)
An example of the use of this heat capacity in the calculation of the heat of a reaction: During a
reaction in this calorimeter, 100 mL of a solution increases in temperature by 6.4 K.
Heat gained by the solution:
(102 g x 6.4 K x 4.02 Jg-1K-1)
Heat gained by the calorimeter:
(50 JK-1 x 6.4 K)
Total heat gained
Heat of the reaction, ΔH
=
2624 J
=
=
=
320 J
2944 J
- 2944 J
From the total heat produced (2944 J), the molar heat of reaction,( this could be heat of reaction
or heat of solution depending on which reaction is being studied) can be calculated by dividing
the total heat produced by either the number of moles of water formed or NaOH dissolved, as is
appropriate to the experimental data.
When doing these calculations, calculate to the nearest joule(J) and then in the final answer
round to the correct number of significant digits. Note that if the temperature in the calorimeter
increases, then heat is lost to the calorimeter and vice versa. Keep this in mind when using the
heat capacity.
Calculation of pH
A-21
Water
Water itself can dissociate according to the equilibrium:
HOH + HOH º H3O+ + OHThe equilibrium constant for this is defined as Kw = 1.0 x 10-14 at 25B
What this expression signifies is that in water, [H3O+] and [OH-] are intimately connected. If [H3O+]
goes up, then [OH-] must go down, and vice versa so that their product is always the same. In pure
water [H3O+] = [OH-], since the only source is the water itself.
Hence:
[H3O+] = [OH-] = 1.0 x 10-7 (because their product must be 1 x 10-14)
and so
pH = -1og10 10-7 = 7.00
If the pH = 4.00, then [H3O+] = 10-4, and [OH-] = 10-10 and so on.
An acid solution has a [H3O+] in excess of that of pure water; therefore, pH < 7. Similarly a basic
solution has a [OH-] in excess of pure water; therefore pH > 7. Strongly acid solutions have pH <
1; strongly basic solutions have pH >13.
Calculation of titration curve
If the titration is a neutralization between a strong acid and strong base, the following can be used.
In this case, both acid and base are assumed to dissociate completely, i.e.,
HCl + H2O 6 H3O+ + ClNaOH 6 Na+ + OHFor convenience consider two pH ranges:
(l) pH < 7, i.e., excess acid. It is assumed that any base added completely reacts with some of
the acid leaving some H3O+ unreacted and diluted by addition of NaOH.
pH = -log[H3O+]
Where VA = volume acid in mL; VB = volume base in mL; MA = concentration of acid;
A-22
MB = concentration of base.
(2) pH > 7, i.e., excess base. It is assumed that all H3O+ from acid is used up by some of the
OH- added and the OH- left is diluted.
[H3O+] = Kw/[OH-] = 1xl0-14/[OH-]
pH + pOH = 14.00
One Strategy for Problem Solving
There are many strategies for solving problems, here is one suggestion that you might find
helpful.
In general, for any problem or situation that needs solving, it must be first determined as to what
the problem is and what it is, that needs to be determined. If you do not know what the final
outcome that you want is, then it will be very difficult to get there.
Then determine what you know. Write all the initial information down, with units.
Now, you need to determine how to connect this two points. Determine what the relationships
are between what you already know and what you need to find out. These relationships will
allow to develop a bridge that gets you from one to the other. All problem solving also requires
you to think. A question like “Do I divide?” is not a demonstration of thinking and being
engaged in solving the problem. All problem solving also requires that you develop an agile
mind.
Once the answer is determined, it is always wise to check if the answer is reasonable.
Situation/Problem : Prepare 25 mL of a solution of 3.0 M NaOH.
1.We want to end up with 25 ml solution that is 3.0M in NaOH.
2. We know a volume (mL) and a concentration (moles/L).
3. However, to measure out amounts, we need to know a weight. The molecular mass (g/moles)
of the NaOH relates weight(g) to moles.
4. So, first, determine the number of moles needed for the solution.
Molarity(moles/L) = moles/ volume (L)
rearrange,
moles = Molarity (moles/L) x volume (L)
moles = 3.0moles/L x 0.025L
moles =0.075 moles of NaOH are needed.
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Now, this value can be related through the molecular weight to the amount that will need to be
weighed out.
molecular weight of NaOH is 39.99g/mol, therefore, 0.075 moles weighs 3.00g
A quick check could be that 3 moles in 1L would need about 120 g, however, only 25 mL are
needed, therefore about 1/40, (25mL/1000mL) would be needed. 1/40 of 120 which is 3.
Now, 3.00g of NaOH can be weighed out and quantitatively transferred to a beaker containing
25mL of water to give 25 mL of a solution that is 3.0M NaOH.
One book that you might find helpful is:
Robert S. Boikess, How to Solve General Chemistry Problems ,8th ed. . (Pearson Education,
Inc. Upper Saddle River, New Jersey 07458) 2009.
Writing a Procedure
Sometimes for a report the procedure is written out. It describes what was done in the
experiment, not what was given in the directions for the procedure in the lab manual. It gives all
the details of what was actually done, what was observed and quantities used and/or produced. It
is written in the past, passive, third person, impersonal.
Graphing Experimental Data
1. Use millimeter-ruled graph paper (10x10/cm). Quarter inch squared paper is not graph paper
and is therefore not suitable.
2. Draw and label axes. The dependent variable (that which is measured) is plotted on the
ordinate, the vertical (y) axis and the independent variable (that which you control in the
experiment) is plotted on the abscissa, the horizontal (x) axis.
3. Select a scale that reflects the precision of the measurements, and which allows the data points
to fill the graph paper.
4. Use ink for the axes, legends, titles, labels and any calculations that might be required on the
graph.
5. Use pencil for the data points, lines and curves. Circle each of the data points.
6. Draw the best possible smooth line through the data points, using a straight-edge ruler.
7. Give the graph a descriptive and informative title, one that does not restate the axes, and does
not include ‘vs.’, ‘as a function of’ and ‘relationship’. The title describes why or what is being
done. For example, “Calibration Curve for Platinum by Atomic Absorption Spectroscopy, λ =
299.8 nm.” instead of “Absorbance vs. Concentration of Pt.”
8. When data obtained from the graph are used for a further operation or as an answer show on
the graph how the values were obtained.
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Tables
The purpose of a table is to display data in a convenient and easily accessible manner.
1. Tables consist of rows and columns.
2. Table labelling includes numbering the table and giving the table a descriptive title. See page
A-12 to A-14.
3. The contents of the table are arranged such that there is minimum of repetition with all the
relevant information being clearly displayed.
4. Lines do not necessarily need to be drawn, but a column of data is not to snake behind the
previous column.
Significant Figures
Significant figures arise because of the uncertainty in measurements. A typical experimental
result contains only the digits known with certainty plus the first uncertain one.
1. For counting significant figures, it is easiest to convert the number into scientific notation and
count the digits.
2. For logarithms, the number of decimal places in the log value is equal to the number of
significant figures in the original measurement.
3. For multiplication and division operations, the number of significant digits in the result is the
same as the number of least number of significant figures.
4. For addition and subtraction operations, the number of decimal places in the result is the same
as the number of decimal places in the least precise measurement.
For a more detailed discussion of significant figures, see the text for this course, Chemistry The
Central Science by Theodore Brown, Eugene LeMay, Bruce Bursten and Julia Burdge.