LABORATORY
MANUAL FOR
CHEMISTRY 102
Prepared by
Department of Chemistry and Physics
Los Angeles Valley College
This Lab Book Belongs To:
Copyright © 2017 by the Department of Chemistry and Physics, Los Angeles Valley College.
All rights reserved. No part of this publication may be reproduced or distributed in any form
or by any means, electronic or otherwise, or stored in a database or retrieval system, without
written permission of the copyright holder.
2
TABLE OF CONTENTS
Contents
LABORATORY SAFETY RULES ....................................................................................................... 3
LAB NOTEBOOK ................................................................................................................................ 7
Ksp, G, H, AND S OF POTASSIUM NITRATE DISSOLVING IN WATER ....................... 9
LE CHÂTELIER'S PRINCIPLE ....................................................................................................... 15
WEAK ACIDS AND BASES.............................................................................................................. 23
DETERMINATION OF Ka BY pH TITRATION ............................................................................ 30
BUFFERS AND pH ........................................................................................................................... 39
ACID-BASE EQUILIBRIUM PROBLEMS ..................................................................................... 43
A SOLUBILITY INVESTIGATION ................................................................................................. 48
SOLUBILITY AND Ksp DETERMINATION .................................................................................. 52
DETERMINATION OF Kf BY SPECTROPHOTOMETRIC METHODS .................................... 56
DETERMINATION OF PERCENT OXALATE BY OXIDATION-REDUCTION TITRATION 61
ELECTROCHEMISTRY ................................................................................................................... 67
ELECTROLYTIC DETERMINATION OF THE MOLAR MASS OF LEAD ............................... 76
FACTORS AFFECTING THE RATE OF A REACTION ............................................................... 79
CHEMICAL KINETICS .................................................................................................................... 84
DETERMINATION OF THE HALF-LIFE OF POTASSIUM-40 ................................................. 91
EQUILIBRIUM BETWEEN TWO COMPLEX IONS OF Co2+ IN SOLUTION ......................... 96
SYNTHESIS AND ANALYSIS OF A NICKEL COMPLEX ......................................................... 104
MOLECULAR MODELS OF TRANSITION METAL COMPLEXES ........................................ 110
CHECK OUT INSTRUCTIONS ..................................................................................................... 114
APPENDIX A ................................................................................................................................... 116
APPENDIX B ................................................................................................................................... 117
APPENDIX C ................................................................................................................................... 127
3
LABORATORY SAFETY RULES
Note: Failure to follow safety rules will result in expulsion from this course.
1.
Wear approved safety goggles at all times in the laboratory.
2.
It is not advisable to wear contact lenses during lab.
3.
Do not wear loose clothing to lab. It is a fire hazard.
4.
Tie back long hair. It too is a fire hazard.
5.
Wear closed shoes to lab.
6.
Never put anything into your mouth while in the lab.
7.
Immediately wash off any chemicals spilled on your skin or clothes.
8.
Keep the lab neat. Return reagent containers and equipment to proper locations. Put any
belongings not needed for experimental work on the shelves provided.
9.
Clean up all chemical spills or broken glass immediately.
10.
Think about how much chemical you will need before you take it from a stock (reagent) bottle.
Never return unused chemicals to stock bottles. Never dip into a reagent bottle with anything
(spatula, dropper, pipet, etc.)!
11.
Dispose of waste chemicals only as instructed.
12.
Behave in a responsible manner.
13.
You should be aware of the location and use of laboratory safety equipment.
14.
Immediately report accidents and injuries to your professor.
15.
Do not perform unauthorized experiments.
16.
Thoroughly wash your hands any time you leave the lab.
17.
No smoking on the Los Angeles Valley College campus.
I have carefully read all of the safety precautions summarized above and recognize that it is my
responsibility to observe them throughout this course.
Chemistry 102
Date
Section Number
Printed Name
Signature
4
5
LABORATORY SAFETY RULES
Note: Failure to follow safety rules will result in expulsion from this course.
1.
Wear approved safety goggles at all times in the laboratory.
2.
It is not advisable to wear contact lenses during lab.
3.
Do not wear loose clothing to lab. It is a fire hazard.
4.
Tie back long hair. It too is a fire hazard.
5.
Wear closed shoes to lab.
6.
Never put anything into your mouth while in the lab.
7.
Immediately wash off any chemicals spilled on your skin or clothes.
8.
Keep the lab neat. Return reagent containers and equipment to proper locations. Put any
belongings not needed for experimental work on the shelves provided.
9.
Clean up all chemical spills or broken glass immediately.
10.
Think about how much chemical you will need before you take it from a stock (reagent) bottle.
Never return unused chemicals to stock bottles. Never dip into a reagent bottle with anything
(spatula, dropper, pipet, etc.)!
11.
Dispose of waste chemicals only as instructed.
12.
Behave in a responsible manner.
13.
You should be aware of the location and use of laboratory safety equipment.
14.
Immediately report accidents and injuries to your professor.
15.
Do not perform unauthorized experiments.
16.
Thoroughly wash your hands any time you leave the lab.
17.
No smoking on the Los Angeles Valley College campus.
Come to lab prepared!! Carefully read the experiment before coming to lab.
6
Quantity
Description
2
Beaker, 20 mL
2
Beaker, 50 mL
2
Beaker, 100 mL
2
Beaker, 150 mL
2
Beaker, 250 mL
1
Beaker, 400 mL
1
Beaker, 600 mL
1
Bottle, 500 mL, Screw Cap
1
Bulb, Pipet
1
Clamp, Buret
1
Cylinder, Graduated, 10 mL
1
Cylinder, Graduated, 50 mL
3
Flask, Erlenmeyer, 250 mL
1
Forceps
1
Funnel, Small, 45 mm
1
Holder, Test Tube
1
Microspatula
1
Pipet, Graduated 1.0 mL
2
Pipet, Graduated 5.0 mL or 10.0 mL
1
Pipet, Volumetric 10 mL
1
Pipet, Volumetric 25 mL
2
Rack, Test tube
4
Shell Vials
10
Test Tube, 10 mm x 75 mm
10
Test Tube, 13 mm x 100 mm
1
Test Tube Brush, 12 mm x 62 mm
1
Thermometer, -20°C to 110°C
1
Tongs
1
Wash Bottle, Polyethylene, 250 mL
1
Watch Glass, 75 mm
7
LAB NOTEBOOK
A record of all experiments you perform in the laboratory will be kept in a notebook that is bound and
has page numbers. This laboratory notebook is as important as the actual experiments you perform
and constitutes a permanent record of your experimentation. Therefore, all entries are to be made in
ink, and mistakes are to be crossed out with a single line (no white out, no erasures). Use the first page
of your notebook as table of contents that is kept up-to-date. An experiment should always be started
on a new page of the notebook. If you have to start another experiment before finishing the previous
one, leave 2 or 3 extra pages at the end of the experiment, so that all of the work for a given experiment
will be on consecutive pages. All work should be done in the notebook and not on separate sheets of
loose paper. For example, additional questions or mechanisms that you are asked to address in the
discussion should still be included in the notebook after the results section. Use professional language
throughout the notebook; avoid first and second person pronouns like I, my, you, etc. You should
number, sign and date all pages of an experiment. Each experiment should have the following format:
I.
Title
II. Purpose: A brief yet complete summary of the goals of the lab. In the context of these goals,
briefly mention which basic techniques are to be used and the role that those techniques serve (for
example, "isolated by extraction, purified by distillation, and analyzed by GC"). It takes practice to
write a good purpose statement. You may want to leave a blank space and write the purpose after
you completed sections III-VI, to ensure that you really understand why a particular experiment is
being done.
III. Balanced reaction(s): Use line structures or Lewis structures, not abbreviations; do not
include mechanisms; do include possible side reactions; where necessary.
IV. References
V. Table of reagents and products: List all chemicals (name and structure) to be encountered
in this experiment – all reactants, reagents, solvents, and products. Include molecular weights and
relevant physical properties (e.g. mp, bp, density, solubility, concentration) for all entries. Note: An
incredibly useful website for finding data for various compounds is chemfinder.com. You are
encouraged to register and use it!
compound name and structure
MW
other properties
VI. Procedure: Start writing the procedure on a new page of
the notebook. The stepwise listing of operations is to be written
using the left column of a page. In general it is a good idea to
leave some space between steps (to allow ample room for
accompanying observations), to sketch pictures of an
apparatus the first time it is used, and to write instructions in
your own words, grouping various operations according to how
you would actually perform them in lab.
P rocedure
D ata and O bs erv at ions
1. T o 1-but yl alcohol (15. 42 g, W hen the sul furic acid was
0.208 m ol es ) i n a 50 m L round added, t he m i xture rem ai ned
bott om fl as k was added i n 3
clear and col orl ess .
port ions 15 m L chi l led
concentrated H
..
.
S
O
4
2
2.
VII. Observations: Use the right column to record raw data and accompanying observations for
each step of the procedure. You should include enough detail so that another person could use
your notebook to perform a lab and he/she would not encounter any unexpected results. It is most
important that data and observations be recorded directly in the notebook immediately at the time
of measurement.
8
VIII. Results: Start writing results on a new page. Return to using the full width of the notebook
(only procedure and observations are written in two column format). All calculations go in this
section, including calculation of percent recovery, or theoretical and percent yield. Show all work
for your calculations. This section should always include a boxed final table that summarizes all of
the pertinent results of the experiment, e.g. unknown identification, composition of mixtures, yields,
etc.
IX.
Discussion and Conclusions: First answer the question: “Did you accomplish the goal of
the experiment?” The discussion is a succinct analysis of the meaning of your results and will often
be guided by questions/statements provided by the instructor. When possible, compare results to
literature values. Answer any assigned questions in this section.
9
Ksp, G, H, AND S OF POTASSIUM NITRATE DISSOLVING
IN WATER
INTRODUCTION
Solubility Equilibrium
When potassium nitrate (KNO3) dissolves in water, it dissociates into potassium ion (K+) and
nitrate ions (NO3−). Once sufficient quantities of K+ and NO3− are in solution, however, the
ions recombine into solid KNO3. Eventually, for every pair of ions that forms, another pair
recombines. As a result, the concentrations of the ions remain constant; we say the reaction
is at equilibrium. This solubility equilibrium of KNO3 is shown in Equation 1,
KNO3  K+ + NO3−
(Eq. 1)
where the opposing arrows indicate that the reaction is reversible. We call this system, where
undissolved solid is in equilibrium with its dissolved ions, a saturated solution.
We can describe the saturated solution with its fixed concentrations of ions with an
equilibrium constant expression. Equation 2 defines the equilibrium constant, Ksp, for KNO3
dissolved in water.
K sp = K +  NO3− 
(Eq. 2)
The sp stands for solubility product and the square brackets around the ions symbolize molar
concentration (M or mol/L). The equation serves as a reminder that the equilibrium constant
not only is concerned with solubility but also is expressed as a product of the ions’ molarities.
The value for Ksp can be large, greater than 1, for the very soluble KNO3, or small, less than 1010, for an insoluble compound such as silver chloride. In addition, because the solubility of a
compound changes with the temperature, its Ksp is likewise a function of the temperature.
Thermodynamics
We use thermodynamics to understand how and why KNO3 dissolves in water. The enthalpy
change, H, for KNO3 dissolving in water provides the difference in energy between solid
KNO3 and its dissolved ions. If H is positive, heat must be added for KNO3 to dissolve. On
the other hand, if H is negative, dissolving KNO3 ion water gives off heat. The entropy
change, S, for KNO3 dissolving in water indicates the higher number of possible energy
states being occupies by the dissolved ions with respect to the lower number of energy states
occupied by the solid KNO3. We expect ΔS for solid KNO3 dissolving in water to be positive
because the two ions on the product side of Equation 1 can occupy more possible energy states
than the KNO3 crystal lattice can as a reactant. Finally, the free energy change, ΔG, for KNO3
dissolving in water indicates whether this process occurs spontaneously. If ΔG is negative,
solid KNO3 spontaneously dissolves in water.
We relate the equilibrium constant to the standard free energy change by Equation 3,
∆G =
−RT ln K sp
(Eq. 3)
10
where R is the ideal gas constant, 8.314 J K−1 mol−1, T is the temperature in Kelvin, and ln Ksp
is the natural logarithm of the equilibrium constant. Like Ksp, the free energy change for a
reaction also changes with temperature.
We also relate the standard free energy change to standard enthalpy and standard entropy
changes by the Gibbs–Helmholtz equation, Equation 4.
∆G =∆H  − T ∆S 
(Eq. 4)
Substituting Equation 3 into Equation 4 yields Equation 5.
−RT ln K sp =∆H  − T ∆S 
(Eq. 5)
Using algebra, we rearrange the equation into the form for a straight line, y = mx + b
∆H   1  ∆S 
ln K sp =
−
+
R  T  R
(Eq. 6)
so that a plot of ln Ksp on the y-axis, versus 1/T on the x-axis, is linear with a slope, m, of
–ΔH°/R and a y-intercept, b, of ΔS°/R. One assumption in this derivation is that ΔH° and
ΔS° are constant, independent of the temperature.
PROCEDURE
1. Prepare a hot water bath by placing a 400-mL beaker half-filled with tap water on a hot
plate.
2. On a balance, weigh about 20 g of KNO3 on a tared piece of weighing paper. Record the
exact mass (to ±0.0001 g) of KNO3 on your report sheet. Transfer the KNO3 to a clean
25×200-mm test tube.
3. Using a graduated cylinder, add 15 mL of distilled or deionized water to the test tube
containing the KNO3. Clamp the test tube in the beaker. Heat the test tube in the
assembled hot-water bath. Stir the mixture with a thermometer until all of the KNO3
dissolves.
4. Determine the volume of the KNO3 solution by filling another 25 × 200-mm test tube with
tap water until the volumes in both test tubes are the same. Measure the volume in the test
tube filled with tap water by pouring this water into a graduated cylinder. Record this
volume on your report sheet.
5. Remove the test tube with the KNO3 solution from the hot-water bath and allow it to cool
while slowly and carefully stirring the solution with your thermometer.
6. Record the temperature when crystals first appear. This is the temperature at which the
solution is just saturated with potassium nitrate (the very small amount of solid is assumed
to be in equilibrium with the ions in solution).
11
7. Add 5 mL of distilled water to the test tube containing the KNO3 solution. Warm and stir
the mixture in the hot-water bath until the solid has completely redissolved. Using the
same method as in Step 4, determine and record on your report sheet the new solution
volume.
8. Remove the test tube containing the KNO3 solution from the hot-water bath. Allow it to
cool slowly. Record on your report sheet the temperature at which crystals first appear.
9. Repeat Steps 7 and 8 for a total of 6 determinations. Record all volume and temperature
measurements on your report sheet.
10. Pour the contents of your test tube containing KNO3 into the container labeled “Discarded
KNO3 Solution”.
11. Use the mass of the KNO3 to calculate the number of moles of KNO3 present.
12. Use the number of moles of KNO3 and the volumes you determined at each temperature
to calculate the molar concentration of KNO3 in the solution at each temperature.
Because, with only a very small amount of solid present, nearly all the KNO3 is still in
solution, its molar concentration equals the molar concentrations of K+ and of NO3- in the
saturated solution
13. Use Equation 2 to calculate the equilibrium constant, Ksp, for dissolving KNO3 in water at
each temperature.
14. Convert the temperatures in degrees Celsius (°C) to Kelvin (K).
15. Determine the natural logarithm of Ksp (ln Ksp) at each temperature.
16. Use Equation 3 to calculate ΔG° at each temperature.
17. Calculate the reciprocal of each Kelvin temperature, 1/T (K-1).
18. Using the graph paper provided at the end of this lab manual or a computer spreadsheet
or graphing program, construct a graph with the y-axis as ln Ksp and the x-axis as 1/T (K1).
19. Determine the slope of the resulting straight line on this graph by choosing two widely
separated points on the line that are not data points.
20. Calculate ΔH° for the reaction. Remember that the slope of the straight line in the ln Ksp
versus 1/T plot equals –ΔH°/R, according to Equation 6.
21. Calculate ΔS° at each temperature using Equation 4. Determine the average ΔS.
22. Calculate S° from the y-intercept from your graph ( b =
the average S° from step 21.
∆S 
R
) and compare this value to
12
13
QUESTIONS FOR Ksp, G, H, and S of KNO3 NAME_______________________
1. (a) Is the process of KNO3 dissolving in water spontaneous at all temperatures studied?
Briefly explain.
(b) Is the reaction in (a) one that gives off heat or requires heat? Briefly explain.
(c) Is your value of ΔS° consistent with the expected change in disorder for the reaction in
Equation 1? Briefly explain.
2. A few compounds exist whose solubility decreases as the temperature increases. How
would the values for ΔG°, ΔH°, and ΔS° for these reactions be different from those values
observed for the solubility of KNO3? Briefly explain.
14
QUESTIONS FOR Ksp, G, H, and S of KNO3 NAME______________________
3. (a) Why must the temperature be measured when only a small amount of solid has been
formed?
(b) What could not be calculated if the temperature was measured after a large quantity of
crystals precipitated?
(c) If you calculated G using temperatures when a large amount of solid had been
formed, disregarding the error of doing so, how would the result impact G’s value?
Would it be higher or lower? Explain why.
15
LE CHÂTELIER'S PRINCIPLE
INTRODUCTION
When a chemical system at equilibrium is disturbed by a change in a component's
concentration/pressure or by a temperature change, the system must shift to counteract the
perturbation while simultaneously attempting to reestablish equilibrium. It is the description
of this "shifting" process that is referred to as Le Châtelier's Principle.
The equilibria to be studied in this experiment involve the formation of transition metal
complex ions. (You can refer to the text for more complete information.) Complex ions formed
in this experiment are made from transition metal ions with Lewis bases (called ligands)
attached to the central metal ion through coordinate covalent bonds.
In general, complex ion formation equilibria can be described by the following equation:
Mx+ + y :LBn  [M(:LB)y](x+yn)
Once a formation equilibrium is established, a change in temperature; in the concentration of
the metal ion, (Mx+); in the concentration of a ligand, (:LBn), (which may or may not carry an
overall charge, n); or in the concentration of the complex ion itself, would disturb the system.
Students will establish and perturb three different complex ion formation equilibria in this
experiment, and will observe each system's response to these perturbations.
PROCEDURE
A.
Fe3+, SCN−, and [Fe(SCN)]2+
1.
Work in groups of 2 or 3 students. Trays containing dropper bottles of the reagents will
be provided.
2.
Clean a 10 mL graduated cylinder, four test tubes (all of them must be the same size and
hold at least 4 mL), and a 100 mL beaker. Use labeling tape to label the test tubes as 1, 2,
3, and 4.
3.
Note the color of the reagents prior to mixing.
4.
Add 20 mL of distilled water from a graduated cylinder to the 100 mL beaker. Next, add
10 drops of the iron(III) nitrate solution and 10 drops of the potassium thiocyanate
solution to the beaker. Stir the mixture thoroughly. The color in the beaker will be due to
the formation of the complex ion, Fe(SCN)2+. Record your observations.
5.
Using a 10 mL graduated cylinder, add 3 mL of the solution prepared in step 4 to each of
the first three test tubes. Add 3.5 mL of the solution prepared in step 4 to the fourth test
tube.
6.
Add 10 drops of the 0.1-M iron(III) nitrate solution to test tube 1. Stir the contents of this
test tube.
7.
Add 10 drops of the 0.1-M potassium thiocyanate solution to test tube 2. Stir the contents
of this test tube.
16
8.
Add 10 drops of distilled water to test tube 3. Stir the contents of this test tube.
9.
Compare the color of the solutions in test tubes 1, 2, and 3 with the color of the solution
in tube 4. (For ease of comparisons, view each test tube's contents down its length against
a white background.) Because the depth of solution and the final volume in all four test
tubes are the same, the intensity of each solution's color is directly proportional to the
complex ion's concentration. (The difference between tubes 3 and 4 may be difficult to
see.)
10. Record your observations and determine which tube(s) contain(s) the highest
concentration of the complex ion.
B.
Ni2+, NH3 and [Ni(NH3)6]2+
1.
Clean a test tube.
2.
Observe the 6-M aqueous ammonia, the 6-M hydrochloric acid and the 0.1-M nickel(II)
nitrate. Record the colors of the reagents.
3.
Place 10 drops of 0.1-M nickel(II) nitrate in the test tube.
4.
Add 6-M aqueous ammonia (also known as aqueous ammonia) one drop at a time to the
test tube from step 3 with stirring after each addition until there is a definite color change.
Remember that aqueous ammonia is primarily ammonia, with ammonium and
hydroxide ions in equilibrium with the ammonia molecules. The ammonia molecules
react with nickel(II) ions to form the colored complex ion, [Ni(NH3)6]2+. Record your
observations.
5.
To the solution from step 4, add 6-M hydrochloric acid (not 12-M HCl) drop wise with
stirring until the color changes once again. (The acid reacts with the basic molecules of
ammonia to form ammonium ions. Ammonium ions have no lone pairs of electrons and
therefore cannot act as Lewis bases.) Record your observations.
C.
Co2+, Cl− and [CoCl4]2−
1. Place a small beaker containing tap water on the hot plate and heat to a gentle boil.
2. Place 5 drops of 0.1-M cobalt(II) nitrate in a clean test tube. Record the color of this
reagent.
Do not remove the concentrated hydrochloric acid from the fume hood!
neutralize and clean up any spills!!
Immediately
3. In a fume hood, add 8 drops of 12-M hydrochloric acid (not 6-M HCl) to the solution in the
test tube from step 2. Stir the mixture and record the color. (This color is characteristic of
the complex ion [CoCl4]2−.)
4. Add 5 drops of distilled water to the contents of the test tube from step 3. Stir to mix.
Record the color. (There may or may not be a color change in this step.)
5. Place the test tube from step 4 in the hot water bath and wait a few minutes for a color
17
change. Record the color. What has been formed (as evidenced by the color change)?
6. Cool the test tube from step 5 in an ice-water bath until the color changes once more.
Record the color. (Think about what has occurred that caused this color change.)
7. The next experiment requires clean, dry glassware. Always put your glassware away clean
so that it will be dry by the next lab period. You will waste valuable lab time if you have to
wash and dry glassware.
18
QUESTIONS FOR LE CHATELIER’S EXP.
1.
NAME _______________________
a. Write a balanced net ionic equation for the equilibrium reaction in Part A, formation
of [Fe(SCN)]2+.
b. For each of the changes in Part A, give the immediate effect of each perturbation on
the value for Q (increase, decrease, or no change). Do the color changes you
observed agree with the shift predicted by the change in Q? Explain your answers.
i. additional iron(III) is added
ii. additional thiocyanate is added
iii. additional water is added
19
2.
a. Write a balanced net ionic equation for the equilibrium reaction in Part B, formation
of [Ni(NH3)6]2+
b. Select which component from the equilibrium mixture reacts with HCl and then
write a net ionic equation for that reaction (NOT AN EQUILIBRIUM!).
c. The addition of hydrochloric acid impacts one of the components in the equilibrium
reaction shown in 2a. Determine the immediate effect on the value of Q due to the
addition of HCl.
d. Do the color changes you observed agree with the shift predicted by the change in
Q? Explain your answer.
20
QUESTIONS FOR LE CHATELIER’S EXP.
3.
NAME ______________________
a. Write a balanced net ionic equation for the (equilibrium) formation of the
tetrachloridocobaltate(II) complex.
b. Based on your observations of color changes in Part C, did heating the reaction
mixture cause a shift in equilibrium? Which direction? Explain your answer based on
the color changes you observed.
c. Is the formation of tetrachlorocobaltate(II) complex ion exothermic or
endothermic? Explain your answer based on the shifts in equilibrium caused by
heating and cooling the reaction mixture.
d. What is the effect of an increase in temperature on the value for the equilibrium
constant? (Increase, decrease or no change)
21
QUESTIONS FOR LE CHATELIER’S EXP. NAME_________________________
4. Consider the hypothetical equilibrium: A + B  2 C + D
H<0
a. Write the equilibrium expression for this reaction.
b. Suppose a change is made to the system. Fill in the following table—answer using
one of the symbols given in each question. Note: NC means no change, and NS means
no shift.
For each change given at the
top of a column, answer the
questions below
What will be the immediate
effect on Qc (↑, ↓, or NC)?
What will be the effect on Kc
(↑, ↓, or NC)?
In comparing the values
from above how does the
size of Qc compare to Kc
(Q = K, Q < K, or Q > K)?
Which way will the change
cause the reaction shift to
re-establish equilibrium,
right (→), left (←) or NS?
When the new
A
equilibrium has been
established, is the
B
amount of each substance
present greater (), less (),
or unchanged (NC)
C
from what before the it was
change?
D
Changes
C is added A is added
D is
removed
A catalyst
is added
Temperature
is decreased
22
23
WEAK ACIDS AND BASES
INTRODUCTION
One method of measuring the acidity or basicity of a solution is to use a pH meter. A pH meter
is a voltmeter that measures the potential of an electrical current flowing through a solution
that is in contact with both a pH sensitive glass electrode (the measuring electrode) and a
constant voltage (reference) electrode. In many pH meters, these two electrodes are fused
together into one "combination" electrode. These electrodes feed their signals into a voltmeter
that is calibrated so that the overall voltage is converted directly to pH units.
In this experiment, a pH meter will be used to study acid-base equilibria of a weak acid, acetic
acid, and a weak base, aqueous ammonia. Because
pH =
− log H3O + 
and pOH =
14.00 − pH
− log OH −  =
pH measurements can be used in the calculation of the equilibrium hydrogen ion and
hydroxide ion concentrations in any aqueous solution. If the initial concentration of a weak
acid is known and the hydrogen ion concentration at equilibrium is calculated from the pH,
then the percent ionization (dissociation) of the weak acid in solution can be determined. For
example, let HA represent any monoprotic weak acid. Then
Initial
Change
Equilibrium
HA + H2O  A− + H3O+
Y
0
~0
−x
+x
+x
Y−x
x
x
where y is the initial concentration of the acid. The percent ionization (dissociation) for the
monoprotic acid is (x/Y) times 100.
The effect of weak bases on pH is due to the ionization (hydrolysis) of water. If B represents
any weak base and y is the initial concentration of that base, then
Initial
Change
Equilibrium
B + H2O  HB+ + OH−
Y
0
~0
+x
+x
−x
x
x
Y−x
and the percent of the weak base involved in ionization is (x/Y) times 100.
24
PROCEDURE
A.
Calibration of the pH meter
The number of groups will be limited by the number of pH meters available. Follow your
professor's instructions as to the number of students per group.
1.
Instructions for standardizing the UB-5 pH meter.
a. Immerse the electrode in a standard buffer solution. Stir gently. Allow the electrode
to reach a stable value.
b. If necessary, press and release the mode button until the display indicates pH mode.
c. Clear existing buffers when performing a new standardization. Use the setup and
enter buttons to clear existing buffers.
d. Press standardize. The meter flashes the current buffer set and detects the flashing
buffer. When the signal is stable, or when you press enter, the buffer’s pH is stored.
e. The meter displays the percent slope of the electrode as 100.0% on the first buffer. On
entering a second or third buffer, the meter performs a diagnostic check on the
electrode and displays the slope.
f. To enter a second buffer, rinse the electrode with deionized water, gently dry it with a
chemwipe and place the electrode in the second buffer solution. Stir and allow time
for the electrode to stabilize, and press standardize again. The meter detects the
buffer and when the signal is stable, or when you press enter, the buffer’s pH is stored.
g. Next, the meter performs a diagnostic test of the electrode. The display indicates
electrode’s condition. The meter displays the % slope obtained from the values read
by the electrode.
h. If Error displayed with the Slope symbol this indicates that your electrode is not
working properly. The electrode response must be between 90 and 105% slope.
Measurements causing Slope Error are not accepted, used or stored by the meter.
Press enter to continue.
i. To enter a third standard, clean the electrode as before and place the electrode in the
third buffer solution, stir, allow it to stabilize, and press standardize. The results will
be the same as in steps g and h.
j. After entering each buffer, the Standardizing symbol goes off and the Measuring or
Stable symbol appears on the display to indicate that the meter has returned to
Measuring operation.
k. Standardize your meter and electrode using at least two buffers with pH values above
and below the expected pH of your samples.
25
B.
The Effect of Dilution on the pH of a Weak Acid Solution
1.
Clean a shell vial. When used in the experiment, the vial can be wet but should be well
drained. If you have not already done so, standardize your pH meter with the pH 4, 7 and
10 buffer solutions. Remember to rinse the pH meter’s probe well with distilled water
between the measurements.
2.
In a clean 10 mL graduated cylinder, obtain 4 to 5 mL of 1.0-M acetic acid. Pour it into
the clean shell vial. Measure and record the pH of the acetic acid. (The solution must
cover the tip of the probe while the measurement is made.)
3.
Pour exactly 1.0 mL of the acetic acid back into the 10 mL graduated cylinder. Discard
the remaining acid. Add distilled water to the acid in the graduated cylinder until the
total volume is exactly 10.0 mL (you have just made a 1 to 10 dilution). Mix well by
carefully pouring the solution back and forth between the vial (the one from which you
discarded the excess acid) and the graduated cylinder. Pour sufficient diluted acid into
the shell vial to allow you to measure and record the pH of this diluted solution.
4.
Save exactly 1.0 mL of diluted acid from step 3 in the cylinder and discard the remainder.
Again add distilled water to the cylinder until the volume is 10.0 mL. You have now made
a second 1 to 10 dilution. (What is the overall dilution?) Mix well and record the pH of
this solution.
5.
Again, save exactly 1.0 mL of the diluted acid from step 4 in the cylinder and discard the
remainder. Again add distilled water to the cylinder until the volume is 10.0 mL. (What
is the overall dilution now?) Mix well and record the pH of this solution.
C.
The Effect of Dilution on the pH of a Weak Base
1.
Clean a shell vial and repeat steps 2 through 5 for Part B above except use 4 to 5 mL of
1.0-M aqueous ammonia for the initial solution. (Note: aqueous ammonia is also known
as aqueous ammonia and thus 1.0-M NH3 would also be an appropriate label for this
solution.)
2.
Rinse the probe well. If the probe had a cap and no storage solution is available, put a
small amount of tap water into the cap before gently sliding it onto the probe. If your
probe did not have a cap, leave the tip of the probe dipped into a beaker containing tap
water. Save your standardization buffers for pH experiments that will be completed on
other lab days.
D.
Calculation of Equilibrium Constants
1.
Calculate the initial molarity (before dissociation or hydrolysis) of the acid or base for
each of the diluted solutions.
2.
From the pH readings, calculate the hydronium ion concentration (in molarity) of each
acidic solution and the hydroxide ion concentration for each basic solution. Use these
data to calculate the percent dissociation for each acetic acid solution and percent
hydrolysis for each aqueous ammonia solution.
3.
Calculate the Ka for each of the acetic acid solutions, and the average Ka. Calculate the
26
percent relative average deviation (see Appendix A at the end of the lab manual) for the
four Ka’s. Use the Ka value for acetic acid given in your textbook as the accepted value
and calculate your percent error (see Appendix A at the end of the lab manual). Because
you are calculating very small numbers and because this experiment was done at nonstandard conditions, your experimental error may be quite large. Use analogous
calculations to calculate percent relative average deviation and percent error for your Kb
for aqueous ammonia.
27
QUESTIONS FOR WEAK ACIDS & BASES EXP.
1.
NAME _______________________
Examine the data for acetic acid and discuss the effects of dilution on the percent
dissociation of this weak acid.
a. What immediate effect did dilution have on Q?
b. Did K change? Should it have changed? Why or why not?
c. Which way did any changes cause the equilibrium to shift? Why?
d. How did the shift affect the percent dissociation?
2.
Should the effects of dilution on % dissociation for a weak acid be any different than %
hydrolysis of a weak base undergoing dilution?
28
QUESTIONS FOR WEAK ACIDS & BASES EXP.
NAME ______________________
3.
A 0.0150 M solution of a weak monoprotic acid is found to be 6.8% ionized. What is the
pH of this acid solution? What is the Ka for this weak acid?
4.
A weak base has a Kb of 6.3×10-3. Calculate the percent hydrolysis of the base and the pH
of the solution if the initial concentration of the weak base is 0.25 M.
29
30
DETERMINATION OF Ka BY pH TITRATION
INTRODUCTION
From previous chemistry lab work students should already be familiar with acid-base titration
techniques. Those experiments probably used a pH indicator (such as phenolphthalein) to
determine the "endpoint" of the titration—the point at which a stoichiometrically equivalent
amount of base had been added to the acid (or acid to base). In such a titration, the only data
collected are the mass or volume of acid and base that have been added to the titration flask
when the equivalence point is reached. However, to construct an acid-base pH titration curve,
both pH and buret readings must be recorded after each addition of reagent from the buret.
From the volume and molarity of the reagent added, the moles of reagent added can be
calculated and then this is plotted against pH.
Acid-base titration curves for monoprotic acids have a characteristic shape. The titration
curve shown below is typical of one obtained when a strong base is added to a weak acid.
mol NaOH added
At the beginning of a titration, pH changes slowly as base is added. Acid is in excess and only
a small percentage of the acid is neutralized after each addition of base. As more base is added,
the ratio of the conjugate base formed to the remaining (unreacted) weak acid in the titration
flask continues to increase. However, as the equivalence point is approached, very little acid
remains and, as base continues to be added, there is a sudden excess of base. It is at this point
in a titration that the pH changes very rapidly. After passing this rapid pH change region the
pH becomes dependent only on the gradually increasing concentration of excess strong base
and again changes slowly.
Acid-base pH titrations can provide information that titration to an indicator endpoint cannot.
Both methods will identify the equivalence point, but the pH titration provides information
31
which allows the pKa and Ka for the acid being titrated to be determined. One method of doing
this is to plot pH as a function of the moles of base added. After the titration curve has been
constructed, two straight lines can be drawn through the data that is almost horizontal (see
the diagram on the previous page). A vertical line that is parallel to the y-axis is drawn between
the two "horizontal" lines. The midpoint of the vertical line (1/2 the distance between the
horizontal lines) is the approximate equivalence point (moles of original Hydrogen ion equal
to moles of Hydroxide ion added). Note: in a titration between a weak acid and strong base,
at the equivalence point all the weak acid has been converted to its conjugate, weak base. An
alternative method of determining the equivalence point is to construct, on the same graph,
the first derivative curve. The first derivative shows how the pH changes for the amount of
base added. The change in pH will be relatively constant at first and then start to increase as
we approach the equivalence point. After the equivalence point the change in the pH will start
to decrease and then become relatively constant again.
Now the pKa and ultimately the Ka of the acid can be calculated. Remember, the equivalence
point is the point at which the acid has been completely neutralized by the strong base. The
weak acid has been converted completely to its conjugate base and water. To use the
Henderson-Hasselbalch equation:
conjugate base 
=
pH pK a + log 
 weak acid
you must determine from the graph at what point half of the acid was neutralized. It is only
at this point that half the acid has been converted to its conjugate base and thus the
concentration of the two are equal; when pH = pKa.
Alternatively, we can think of the Henderson-Hasselbalch equation as:
pH pK a + log
=
Vb
Ve − Vb
Where Vb is the volume of base added and Ve is the equivalence point volume. If we plot pH
on the y-axis and log(Vb/(Ve−Vb) on the x-axis for volumes from about 20% to 80% of the
equivalence point (because this is the region where it’s a buffer so the Henderson-Hasselbalch
equation applies), we will get a straight line. The point on the pH scale where this line crosses
0 on the x-axis is the point at which the pH is equal to the pKa.
The experimental value for Ka can then be determined from the equation:
pK a = −log K a
So,
K a = 10 − pK a
In this experiment, an acid-base pH titration curve will be constructed for potassium hydrogen
phthalate (KHC8H4O4 or KHP). KHP is a monoprotic acid having a structural formula of:
32
O
H
H
C
C
C
O
C
C
H
C
C
O
+
K
H
C
H
O
An experimental Ka for KHP will be determined in this experiment.
PROCEDURE
1.
Each group should obtain a buret, a Vernier LabQuest, a pH probe and a Drop Counter.
2.
Plug the pH Probe into the port on the top of the LabQuest
labelled “CH 1” (on the top) and the Drop Counter in the port
labelled “DIG 1” (on the side). The display should look like the
image to the right.
3.
Attach the Drop Counter and a buret clamp to a ring stand such
that the Drop Counter is below the buret clamp.
4.
Obtain about 100 mL of standardized (approximately 0.1-M)
NaOH in a clean dry beaker. Record the exact molarity of the
NaOH from the bottle.
5.
Using the same techniques learned in previous titration
experiments, clean the buret, rinse and flush it with 1 to 2 mL
of the NaOH solution, discard the rinsings and fill the buret
with the NaOH solution.
6.
You need to make sure that the Drop Counter can “see” each drop that passes through it.
Place the buret filled with the NaOH solution in the buret clamp so that the tip of the
buret is just above and approximately centered over the slot in the Drop Counter. Place
a waste beaker under the Drop Counter to collect the solution. Turn on the LabQuest.
When it has
started you should see that both probes are connected. Press
the
“Collect”
button. Open the stopcock on the buret such that the NaOH
solution comes out one drop at a time (about 1 drop every second or two). If it is aligned
correctly, you should see the volume increase incrementally on the screen. If it is not,
adjust the buret side-to-side until the LabQuest shows the volume changing. When
everything is aligned correctly close the stop flow of the solution and press the “Collect”
button to stop data collection.
7.
Calculate the approximate mass of KHP (FM = 204.23) that would be required to
neutralize about 25 mL of 0.1-M NaOH.
8.
Clean and label two 250 mL or 400 mL beakers (they can be wet). From your professor,
obtain a small amount of KHP in a dry shell vial and take the KHP and titration beakers
to the analytical balance room. Use the "weighing by difference" technique to place the
approximate mass of KHP determined in step 4 into each of the two beakers. Record the
33
mass of KHP in each beaker (±0.0001 g). Obtain approximately 50 mL of the supplied
NaOH solution. Be sure to record its molarity.
9.
Add approximately 50 mL of distilled water to each beaker and swirl until the KHP is
dissolved.
10. Fill a small beaker with distilled water and stand the pH probe in the beaker. The probe
should be free from the holder so that it can be moved easily between the beaker and
titration beaker.
11. Add a magnetic stir bar to titration beaker 1 and place
the beaker under the Drop Counter on top of the
magnetic stir plate. Lower the tip of the pH probe
through the hole in the Drop Counter into the solution
of beaker 1. Turn on the magnetic stir plate and set
the speed to the maximum setting.
12. Press “Collect” button on the LabQuest. Open the
stopcock on the buret such that the solution flows out
one drop at a time at a rate of no more than about 1
drop per second.
13. When pH reaches 11 to 12 and you have added about
2 to 3 mL of solution at that pH you can close the
stopcock on the buret and stop the run by pressing the
“Collect” button. At this time the contents of beaker 1
can be discarded.
14. Refill the buret with the NaOH solution.
15. Repeat Steps 11 through 13 for beaker 2. Before pressing the “Collect” button, click on
the file cabinet icon to add another run to the data collection (Run 2).
16. Attach a USB flash drive (it must be USB 2, USB 3 will not work) to the USB-A port on
the top of the LabQuest. It may take a few seconds for the device to recognize the USB
drive. Click on “File” and then “Export.” Click on the USB icon and save the data as a
text file onto the flash drive (both runs will be in the file). Give the file a meaningful
name. Make sure that the data (including both runs) is on the drive and that all members
of the group have a copy of the data. The data is saved as a tab-delimited ASCII file (.txt).
17. Using Excel, or another graphing program (i.e., Google Sheets, Origin, or Numbers on a
Mac), open the file from your USB flash drive. It will recognize that it is a tab-delimited
ASCII file. Just click on “Next” to choose the default option for everything. The
spreadsheet should look like this:
18. In Excel, create a two new columns for
your data. The first column should be
labelled pH/V, the second column is
explained in Step 22. In this column
starting with row containing the first
34
data point enter the formula “=(B9−B8)/(A9−A8)” (without the
quotation marks). Here we used B8, B9, A8 and A9 because the
data starts in row 8. Press ENTER. Put the cursor at the bottom
right corner of the cell containing that formula (it should change
into a bold cross) and click and drag it down to the last row of
data for Run 1 to copy it down that column (there are likely about
600 to 800 data points so be careful). Do this for both runs.
19. Construct two titration graphs, one for each run. Plot pH
(vertical axis) as a function of the volume of NaOH added
(horizontal axis) and pH/V (vertical axis) as a function of the
volume of NaOH added on the same graph. You can do this by
highlighting all three columns of data and selecting
Insert/Chart/Scatter/Scatter with Smooth Lines. Be sure to properly title and
label your graphs. Also be sure that each graph shows the correct precision. See
Appendix B for a review on graphing. Do this for both runs.
20. You can more easily see where the peak of the pH/V vs. volume graph is by plotting
the pH/V data on a secondary axis. The point on the x-axis (volume) where this line
has its maximum value is the equivalence point. Do this for both runs.
21. Determine the pKa from each of your pH vs. volume graphs. Show on the graph how you
determined the pKa value.
22. The second column should be labelled log(Vb/(Ve−Vb)). Starting at a volume that is about
20% of the equivalence point enter the formula “=log(AXX/(yy.yy−AXX))” (again without
the quotation marks). XX indicated the row number you are starting at and yy.yy is the
equivalence point volume you determined in step 20. Copy this formula down to about
80% of the equivalence point. Do this for both runs.
23. Create a new graph (Scatter X-Y, Line) and plot pH on the y-axis and the new column of
data on the x-axis. You should get a straight line. Where this line crosses the y-axis (x=0)
is the point where the pH=pKa. Do this for both runs.
24. Read the pKa from each graph. Mark each graph to show how you got the pKa.
25. Average the two pKa values. Calculate the experimental Ka for KHP from the average pKa
value.
26. Attach printouts of all 4 graphs to your lab report.
35
36
QUESTION FOR DETERMINATION OF KA …
NAME ______________________
1.
Should the mass of KHP used for the pH titration change the experimental value for the
Ka? Explain your answer.
2.
Calculate the experimental Kb for the phthalate ion, C8H4O42−, from the average
experimental Ka for KHP.
3.
Using the experimental Kb in Question 2, calculate the pH of a 0.83-M K2C8H4O2 solution.
37
QUESTION FOR DETERMINATION OF Ka …
NAME _______________________
4.
Recall that an optimum buffer is one that contains equal (or close to equal)
concentrations of a weak acid and its conjugate base. At approximately what pH reached
during the titration would the solutions in titration flasks 1 and 2 meet the criterion for
an "optimum" buffer? Explain your answer.
5.
Using duplicate calculations (one for each graph), use the equivalence point on each
graph to determine an experimental molar mass of KHP (remember it is monoprotic).
Average your results. Now use the true molar mass of KHP (204.23 g mol−1) and calculate
the percent error for this experiment (see appendix A of this lab manual). This is a
measure of the accuracy of your work in this procedure. Show calculations.
38
39
BUFFERS AND pH
INTRODUCTION
An acid-base buffer is a solution that resists change in pH when small amounts of acid or base
are added. This type of buffer contains two species, a weak acid and its conjugate base. The
weak acid reacts with and partially removes from solution added base, and the weak acid’s
conjugate base reacts with added acid. If hydrogen ion is removed from solution, the buffer’s
weak acid dissociates to partially replace the hydrogen ion that was removed. If hydroxide is
removed from the system, it is partially replaced through hydrolysis of water by the weak
conjugate base. These processes are examples of Le Châtelier’s principle. The original buffer
solution is at equilibrium. Added material temporarily disturbs this equilibrium, and the
system shifts to restore equilibrium. Thus, concentrations of hydrogen ion and hydroxide ion
are “buffered” and the pH of the solution remains relatively constant.
To be a pH buffer, both a weak acid or base and its conjugate base or acid must be initially
present. In other words, both must be present before dissociation by the weak acid or
hydrolysis by the weak base can be considered. An “optimum” buffer, which has equal capacity
to neutralize either added acid or added base, is created when the concentrations of the
conjugate acid/base pair in the buffer solution are equal. However, a solution does not have
to contain equal amounts of the pair to be considered a buffer. The equation
H +   A− 
Ka =    
HA 
can be rearranged into the Henderson-Hasselbalch equation:
 A− 
=
pH pK a + log
HA 
From the equation above, it can be seen that the ratio of the conjugate acid-base pair can be
varied to create a buffer solution with a desired pH so long as that pH is close to the pKa of the
acid form of the weak pair. The buffer solution does not have to be made by combining the
weak acid and its conjugate base directly. It also can be created by partial neutralization of a
weak acid by a strong base, or by partial neutralization of a weak base using a strong acid. For
example, if a weak acid (HA) is neutralized by a strong base, the net ionic equation for the
reaction would be:
HA + OH−  A− + H2O
If the hydroxide ion from the strong base were the limiting reactant, some weak acid, HA,
would remain in solution after reaction was complete. The HA remaining in solution, along
with its conjugate base, A−, (formed in the partial neutralization) would create the buffer. An
analogous approach would be to use an excess of weak base with a limited amount of strong
acid.
40
In this experiment, various solutions will be prepared and studied. A pH meter will be used
to determine the experimental equilibrium concentration of hydrogen ion in each solution.
Using the Henderson-Hasselbalch equation, the theoretical pH and hydrogen ion
concentration can be determined from the Ka and the mole to mole ratio of the conjugate
acid/base pair for each solution studied.
PROCEDURE
1. Due to the limited number of pH meters, students will work in groups. Follow your
professor’s instructions regarding the number of students per group.
2. Each group will need to obtain pipet pump.
3. Prepare your meter’s electrode for use and standardize the pH meter (refer to the
instructions provided in the “Weak Acids and Bases” experiment).
4. Obtain approximately 40 mL each of 0.20-M acetic acid and 0.20-M sodium acetate
solutions in separate clean, dry 50 mL beakers. Obtain approximately 15 mL each of 0.10M hydrochloric acid and 0.10-M sodium hydroxide solutions in separate clean, dry 20 mL
beakers.
5. Pour enough of the 0.20-M acetic acid solution into a clean, dry shell vial so that you can
measure its pH. Record the pH.
6. Clean your pipet and use the solution remaining in the shell vial to rinse the pipet. Discard
the solution used for rinsing.
Do not pipet by mouth; use a bulb or a pipet-pump!
7. Using the pipet, measure 25.0 mL of the acetic acid solution into a clean, dry 150 mL
beaker. Save the acetic acid solution remaining in the 50 mL beaker.
8. Repeat steps 5 and 6 using the sodium acetate solution.
9. Using the freshly rinsed pipet, measure 25.0 mL of the sodium acetate solution and add it
to the 150 mL beaker containing the acetic acid solution (Step 7) and mix well. This is the
combined solution that will be referred to throughout this experiment. Save the sodium
acetate solution remaining in the 50 mL beaker.
10. Pour enough of the combined solution into a clean, dry shell vial to measure its pH, and
record. Do not discard the remaining combined solution in the 150 mL beaker.
11. In all subsequent steps, you may use a clean shell vial that has been rinsed with deionized
water and well-drained.
41
12. Use clean 10 mL graduated cylinders and follow the chart to carefully measure the volume
of each reagent indicated into separate, well-drained shell vials.
Shell Vial
Number
Combined
Solution
1
2
3
4
5
6
7
8
9
10
11
3.0 mL
4.0 mL
4.0 mL
3.0 mL
3.0 mL
4.0 mL
4.0 mL
0.20-M
acetic acid
2.0 mL
0.20-M
sodium
acetate
2.0 mL
0.10-M
HCl
3.0 mL
4.0 mL
H2O
3.0 mL
3.0 mL
2.0 mL
3.0 mL
4.0 mL
0.10-M
NaOH
3.0 mL
2.0 mL
3.0 mL
2.0 mL
3.0 mL
2.0 mL
13. Cover each shell vial with Parafilm™. Mix the contents well by inversion, then measure
and record the pH of each solution. Rinse the probe well with deionized water between
every measurement.
14. Rinse the probe well. If the probe had a cap, put a small amount of storage solution or tap
water into the cap before gently sliding it onto the probe. If your probe did not have a cap,
leave the tip of the probe dipped into a beaker containing tap water.
15. Be sure to clean the pipet and rinse it with deionized water. Return the pipet pump if you
borrowed one.
16. Calculate the initial molarity (after dilution but before any shift to achieve equilibrium) of
the acetate and the acetic acid in the combined solution.
17. Calculate the initial moles (due to the combination of solutions or after any neutralization
reaction but before any shift to achieve equilibrium) of the acetate ion and the acetic acid
present in tubes 1 through 11. In some of the solutions these species come from more than
one reagent. In others, acid-base neutralization calculations must be completed before the
initial moles can be determined.
18. Determine the ratio of moles of acetate ion to moles of acetic acid for tubes 1 through 11.
Express your ratios as 1:1, 1:3, 2:1, etc.
42
QUESTIONS FOR BUFFERS AND pH EXP.
NAME ______________________
1. In theory, which of the 14 solutions tested should have similar pH’s? Why? Use results
from your calculations for step 18 of the Procedure to help explain your answer for each
solution.
2. The experimental pH values for the solutions should be in fairly good agreement with the
theoretical pH values for each of the solutions tested. Why? What are some things that
could cause the experimental pH to be different than the theoretical pH?
3. Which of the 14 solutions tested are buffers? Identify any of the solutions that would be
considered “optimum” buffers (have the same number of moles of weak acid and conjugate
weak base present).
43
ACID-BASE EQUILIBRIUM PROBLEMS
1. Calculate the pH of a solution that contains 0.15 M oxalic acid. Calculate the concentration
of the oxalate ion in this solution.
2. Calculate the pH of a 0.0035 M solution of methylamine.
44
3. 65.3 mL of 0.156 M hydrochloric acid is added to 145.3 mL of 0.078 M aniline solution.
What is the approximate pH of the resulting solution?
4. Out of the following, which is the best acid/base to use to prepare a buffer with a pH of
8.00?
a. sodium cyanate
b. sodium lactate
c. hydrazine
What ratio of masses of the weak acid/base and its conjugate should you use to make the
buffer of the required pH? Use the sodium salt of the conjugate base if you chose a weak
acid or the chloride salt of the conjugate acid if you chose a weak base.
45
5. Calculate the pH at the equivalence point when 25.00 mL of 0.10 M iodic acid is titrated
with 0.080 M barium hydroxide solution.
6. What is the pH of a solution obtained by adding 100.0 g of sodium benzoate to enough
water to make 1.50 L of solution?
46
7. 25.00 mL of 0.15 M hydroxylamine is titrated with 0.20 M hydrochloric acid. When 12.56
mL of the acid have been added what should the approximate pH be?
8. If Kw at 40.0°C is 2.916×10−14, what is the pH of pure water at this temperature?
47
9. The pH of a 0.15 M solution of butanoic acid is 2.82. What is the Kb of the butanoate ion?
10. Ethanolammmonium ion has pKa of 9.498. What is the pH of a 0.050 M solution of
ethanolamine?
48
A SOLUBILITY INVESTIGATION
INTRODUCTION
Most metal ions are soluble when mixed with most anions. There are some exceptions as
delineated in the solubility rules in your textbook. In this experiment we are going to examine
some of these insoluble salts and the circumstances that affect their solubility.
One factor that can affect solubility is the pH of a solution. If the anion in the insoluble salt is
the conjugate base of a weak acid, the salt will become more soluble as the pH decreases. For
example, barium sulfate is an insoluble salt with an equilibrium reaction shown as
BaSO4  Ba2+ + SO42−
Eq. 1
and the solid will become more soluble as the pH decreases because sulfate ion will react with
hydronium ions in an acid/base equilibrium
H3O+ + SO42−  HSO4− + H2O
Eq. 2
As Equation 2 proceeds to the right it effectively removes sulfate ion from the first equilibrium
(Eq. 1) causing the first reaction to shift to the right, (i.e., more barium sulfate dissolves), to
re-establish equilibrium.
Increasing pH can also affect an equilibrium if the pH is raised in the correct manner. Silver
chloride is an insoluble salt with an equilibrium reaction of
AgCl  Ag+ + Cl−
Eq. 3
If we increase the pH of the mixture shown in Equation 3 by adding aqueous ammonia, the
ammonia forms a complex with the silver ion
Ag+ + 2 NH3  [Ag(NH3)2]+
Eq. 4
which removes silver ion from Eq. 3 causing more of the silver chloride to dissolve as
equilibrium is re-established.
Another example involves amphoteric hydroxides. Amphoteric hydroxides are compounds
that can react with either acids or bases. Aluminum hydroxide is an amphoteric hydroxide. If
we have aluminum ion in solution and we start to increase the pH by adding a strong base, we
initially produce an insoluble compound
Al3+ + 3 OH− Al(OH)3
Eq. 5
Continued addition of hydroxide allows another equilibrium to occur in which a complex ion
is formed between the aluminum ion and the hydroxide
Al(OH)3 + OH-  [Al(OH)4]−
Eq. 6
which results in the solid Al(OH)3 dissolving. But, addition of a strong acid would also dissolve
solid Al(OH)3:
3 H3O+ + Al(OH)3  Al3+ + 6 H2O
Eq. 7
49
PROCEDURE
A. Effect of lowering the pH on the solubility of an insoluble salt.
1.
Obtain 5.0 mL each of 1.0 M calcium chloride and 0.25 M sodium oxalate solutions
2.
Pour both solutions into a 50 mL beaker (mixing well). What reaction has occurred?
3.
Add approximately 10 mL of 6 M nitric acid and stir. What reaction has occurred? What
do you observe? Dispose of the solution in the appropriate waste receptacle and
thoroughly clean the beaker.
B.
Effect of raising the pH on the solubility of an insoluble salt
1.
Obtain 15.0 mL each of silver nitrate and sodium chloride solutions
2.
Pour both solutions into a 100 mL beaker (mixing well). Allow the mixture to sit for 15
minutes and observe if any noticeable amount of silver chloride has precipitated.
3.
Add approximately 25 mL of 6 M aqueous ammonia and stir. Allow the mixture to sit for
15 minutes and observe if any noticeable amount of silver chloride has dissolved.
4.
Add approximately 25 mL of 6 M nitric acid and stir. Observe any changes that occur in
the beaker. Dispose of the solution from step 4 in the appropriate waste receptacle and
thoroughly clean the beaker.
C.
Effect of adding a strong acid base to an amphoteric hydroxide
1.
Obtain approximately 20.0 mL of 1.0 M zinc nitrate solution and place it into a 150 mL
beaker.
2.
Add 6 M sodium hydroxide (with mixing), in a drop-wise fashion, until a reasonable
amount of solid appears.
3.
Divide the mixture from step 2 into approximately two equal portions. (This mixture
contains the amphoteric hydroxide.)
4.
To one of the two portions, continue to add 6 M sodium hydroxide (with mixing) until you
see a distinct change in the mixture. Note how much sodium hydroxide solution was
added. (Recall: 20 drops  1 mL)
5.
To the other portion, add 6 M nitric acid (with mixing) until you see a distinct change in
the mixture. Note how much nitric acid solution was added. Dispose of the solutions in
the appropriate waste receptacle.
50
QUESTIONS FOR A SOLUBILITY INVESTIGATION
NAME__________________
1. Write the equilibrium reaction for the mixture in the beaker in Part A, step 2.
Write the net ionic equation for the reaction (which involves one of the species in the
reaction that you’ve just written) that occurs when nitric acid is added to the beaker in Part
A.
Examine the two reactions shown above for part A and explain, using Le Châtelier’s
principle, why the changes occurred in the beaker after adding nitric acid.
2. Write the equilibrium reaction for the mixture in the beaker in Part B, step 2.
As in question 1, write the net ionic equation for the reaction (which involves one of the
species in the reaction that you’ve just written) that occurs when aqueous ammonia is
added to the beaker in Part B.
Again, according to Le Châtelier’s principle, why does the precipitate dissolve upon
addition of ammonia?
As above, write the net ionic equation for the reaction that occurs when nitric acid is added
to the beaker in part B.
As previously, explain why the precipitate reappears upon addition of nitric acid.
51
QUESTIONS FOR A SOLUBILITY INVESTIGATION
NAME _________________
3. Write the net ionic equation for the reaction that initially occurs when aqueous sodium
hydroxide is added to the zinc nitrate solution. (Formation of the amphoteric hydroxide.)
Write the net ionic equation for the reaction that occurs when an excess of sodium
hydroxide is added to the amphoteric hydroxide. (Step 4)
Write the net ionic equation for the reaction that occurs when nitric acid is added to the
amphoteric hydroxide. (Step 5)
In the context of part C of this experiment, explain what an amphoteric hydroxide can do
that:
•
acetic acid can’t do
•
aqueous ammonia can’t do
•
sodium chloride can’t do
52
SOLUBILITY AND Ksp DETERMINATION
INTRODUCTION
Calcium iodate is an ionic compound that is only slightly soluble in water.
solution, an equilibrium forms between the solid salt and its ions:
In aqueous
Ca(IO3)2  Ca2+ + 2 IO3−
The solubility of calcium iodate can be determined by measuring the concentration of either
the calcium ion or the iodate ion in a saturated solution. In this experiment the concentration
of the iodate ion will be determined.
This analysis involves two reactions. First, the saturated solution of calcium iodate is acidified
and reacted with excess potassium iodide, converting all the iodate ions into molecular iodine.
IO3− + 5 I− + 6 H+ → 3 I2 + 3 H2O
The molecular iodine formed is then titrated with standardized sodium thiosulfate.
I2 + 2 S2O32− → 2 I− + S4O62−
The titration uses as indicators, the brown color of the molecular iodine (the iodate and iodide
ions are colorless) and the dark blue color of an iodine-starch complex, (seen in the chemical
kinetics experiment).
PROCEDURE
1. Each group will need one buret and a pipet pump. In this experiment you will need a
clean, dry shell vial, a clean, dry 10 mL graduated cylinder, three clean and dry filter
funnels, and seven clean, 100 to 250 mL beakers (they don't need to be the same size).
Four of the beakers must be dry, the other three can be wet. If your group does not have
these available, clean them and put them in the oven now (remove any plastic parts from
the graduated cylinders BEFORE putting them in the oven).
2. Label three beakers (the ones that can be wet) A-1, B-1, and C-1. Put about 50 mL of
distilled water into each beaker. Bring a clean, dry shell vial to your instructor to obtain
about 4.5 g of calcium iodate. Using the balances in the lab (NOT the analytical balances),
weigh out approximately 1.0 g of calcium iodate and place it in beaker A-1. Again using
the balances in the lab, weigh out approximately 1.5 g of calcium iodate and place it in
beaker B-1. Weigh out approximately 2.0 g of calcium iodate and put it in beaker C-1.
3. Stir the contents of each beaker with a separate clean stir rod. Allow the solutions to sit
for at least 20 minutes, stirring every few minutes. Calcium iodate is only slightly soluble
and the saturated solution forms slowly. Use this time to prepare for titration (Steps 4-7).
If you have put cylinders, beakers, and/or funnels into the oven, remove them now and
allow them to cool.
53
4. Using one of the cooled, clean, and dry beakers, obtain about 100 mL of standardized
sodium thiosulfate. Record the exact molarity of the sodium thiosulfate solution.
5. Using the clean, dry graduated cylinder that you prepared, measure out three separate
samples of about 1 cm3 (1 mL) each of solid KI. (A paper funnel might help you pour the
KI into the cylinder without spills. Clean up any spilled KI!) Set these aside for use in step
14.
6. In each of three clean test tubes (they can be wet) put about 2.0 mL (40 drops) of 1% starch
solution. You will need this indicator solution later in the titrations.
7. Clean the buret, rinse and fill it with the sodium thiosulfate solution.
8. Set up the three clean, dry funnels you have prepared using buret clamps. Place the three
clean, dry beakers labeled A-2, B-2, and C-2 under these funnels. Put dry filter paper cones
into each funnel.
9. After allowing the calcium iodate mixtures to come equilibrium (it takes at least 20
minutes) pour each solution through its own filter cone, catching each filtrate in its own
dry beaker. Do not add water. Any precipitate remaining in the beakers can be
discarded. It is the solution filtering into the beakers that you will be titrating and you do
not want to change its concentration by adding rinse water. After the solution has filtered
through, discard the filter papers and precipitates.
10. Set up three clean 125 mL titration flasks (they can be wet) labeled A-3, B-3, and C-3.
11. Clean the 10.0 mL volumetric pipet. Shake as much water as possible from the pipet and
then rinse the pipet twice (each time with 1 to 2 mL) with the filtered solution from beaker
A-2. Pipetting carefully (do not pipet by mouth; use a pipet pump or a bulb!), transfer
exactly a 10.0 mL sample (aliquot) of the solution from beaker A-2 to flask A-3. Rinse the
pipet twice with the filtered solution from beaker B-2 and transfer 10.0 mL of solution
from beaker B-2 to flask B-3. Repeat the procedure for beaker C-2/flask C-3.
12. Add about 20 mL of water to each flask. (Think about why is it okay to add water now.)
13. Add about 8 drops of 6 M HCl to each flask.
14. Add about 1 cm3 (1 mL) of solid KI to the flask that you are now ready to titrate. (As KI is
added to the flask, it reacts with iodate to form brown I2.) Swirl each flask until the KI is
dissolved.
15. Record the initial buret reading. Set flask A under the buret (a white piece of paper under
the flask will help you see color changes).
16. Start adding sodium thiosulfate from the buret into the flask. Add about 1 mL at a time
and swirl well after each addition. The sodium thiosulfate will react with the brown Iodine
and will convert it to colorless iodide ions. When the color of the solution has faded to
pale yellow, add one of the 2.0 mL aliquots of starch solution to the titration flask. The
starch will react with the remaining iodine in the flask to produce a dark blue complex. (If
the starch had been added at the beginning of the titration, the very large amount of iodine
present would create numerous complex ions with the starch that would make it much
more difficult to titrate.)
54
17. Continue to titrate slowly. The blue-black color will start to fade. The endpoint is when
one drop of sodium thiosulfate causes the solution to become colorless. (Note: It is more
difficult to titrate from a colored to a colorless solution than vice versa.)
18. At the endpoint, record the final buret reading and calculate the volume of sodium
thiosulfate used.
19. Refill the buret and repeat steps 14 - 18 for flask B and then again for flask C.
20. Calculate the molarity of the iodate ion which was in the saturated solutions (in beakers
A, B, and C).
Note: You must use the mole to mole relationships from both of the chemical reactions
provided in the Introduction to go from moles of thiosulfate to moles of iodate.
21. Determine the average molarity of the iodate ion in the saturated solutions.
22. Calculate the molar solubility, the solubility in g/100 mL, and the solubility product
constant, Ksp, for calcium iodate. Include appropriate units.
55
QUESTIONS FOR SOLUBILITY…
NAME_______________________
1.
How do the molarities of the iodate ion in each of the saturated solutions compare?
Should they be the same? Explain.
2.
How would adding water in step 9 to wash the solid calcium iodate precipitate onto the
filter paper change the Ksp value which was determined experimentally? Would the
calculated value for the constant be higher, lower, or unchanged if extra water had been
used in this step? Explain.
3.
In step 12, extra water is added to the titration flask. This added water does not alter the
value obtained for the Ksp. Explain why.
4.
Should a precipitate form when 10.00 mL of 0.1500 M silver nitrate is added to 20.00 mL
of 0.1864 M potassium acetate?
5.
If a 1.00 M potassium chloride solution is added dropwise (no significant volume change)
to a solution containing both 0.010 M silver nitrate and 0.020 M mercury(I) nitrate, which
insoluble chloride starts to precipitate first?
What percent of the cation that precipitated first remains in the solution just as the other
cation reaches its saturation point with the chloride?
6.
What is the molar solubility of barium fluoride in a solution that contains 3.25 M acetic
acid and 3.25 M sodium acetate? Hint: Combine two equilibria reactions to determine the
Kc for:
BaF2 + 2 H+  Ba2+ + 2 HF
and then solve for the molar solubility using the approximation method. Be sure to
validate! Think about what the total concentration of fluoride ion must be, both as the
free F- and as HF.
56
DETERMINATION OF Kf BY SPECTROPHOTOMETRIC
METHODS
INTRODUCTION
This experiment will use a spectrophotometer to obtain the data needed to calculate an
equilibrium constant, Kf, for the formation of a complex ion from iron(III) and thiocyanate
ions. The net ionic equation for the reaction is:
Fe3+ + SCN−  [Fe(SCN)]2+
and the equilibrium constant is given by the expression
 Fe ( SCN )  2+ 
 

K f =  3+
Fe  SCN − 
To determine the Kf value, one must be able to measure or calculate the equilibrium
concentrations of the three ions that appear in the equilibrium expression. Because the
reactants, iron(III) ion and thiocyanate ion, are colorless, and the complex ion product is red,
spectrophotometry can be used to determine the equilibrium concentration of the complex
ion. This data can then be used to calculate the equilibrium concentrations of iron(III) ion
and thiocyanate ion, assuming the starting concentrations of those ions are known.
Spectrophotometry is based on the principle that the light absorbed by a solution is directly
proportional to the concentration of a component of that solution. The relationship between
absorbance and concentration is given by the equation:
A = ε bc
where A represents absorbance, c represents concentration, and ε and b are constants.
A spectrophotometer operates by separating light into its component wavelengths and
selectively measuring the intensity of a given wavelength of light before and after it passes
through a solution. The absorbance (A) is then calculated (by the spectrophotometer) using
the relationship
A = − log
I
I0
where I0 is the intensity of the light entering the solution and I is the intensity of the light that
has passed through the solution. It is customary to "zero" the spectrophotometer using the
solvent that will be used for the test solution. This "zeroing" process accounts for light that is
absorbed by the solvent or is scattered by the cuvet (a special test tube made of optically
uniform glass).
Before determining the concentration of a particular solute in a solution, a “standard curve”
for the solute must be prepared. The standard curve (which is actually a straight line) is
prepared by measuring the absorbances of solutions having known concentrations of the
solute. The absorbances of the known solutions are plotted as a function of their
concentrations. The unknown's concentration is then obtained from that solution's
57
absorbance and the “standard curve.”
PROCEDURE
A. Preparation of a Standard Curve
1. Each group will need one 1.00 mL pipet and three 5.00 or 10.00 mL graduated pipets, a
pipet pump and a cuvet. Also obtain some Parafilm™.
2. In separate clean, dry labeled beakers or shell vials, obtain about 25 mL of 0.000075 M
Fe(NO3)3, about 40 mL of 1 M KSCN, and about 45 mL of 0.1 M HNO3. Be sure to use the
correct concentrations. Different concentrations are used in part B.
Do not take more reagent than you need. If you do not have clean, dry beakers or shell
vials (you should always put your glassware away clean) you will need to rinse the clean, wet
beaker and/or vials with 2 to 3 mL of the reagent that you are obtaining. Discard the rinse
solution.
3. Set up 12 large test tubes (they should each hold at least 8 mL). They should be clean and
dry. If you need to wash them, rinse them with 1 to 2 mL of 0.1 M HNO3 and discard rinse.
4. Label the 1 mL pipet for Fe(NO3)3. Use this pipet for the volumes of Fe(NO3)3 that are 1.0
mL or smaller. Label the three 5 or 10 mL pipets, one for Fe(NO3)3 (to use for volumes
greater than 1.0 mL), one for KSCN, and one for HNO3. Rinse each pipet with 0.5 to 1.0
mL of the reagent for which it will be used.
5. Using the appropriate pipet for each reagent, add the amount of each reagent to each tube
that is shown on the chart below. Pipet carefully using the bulb!
Tube
No.
1
2
3
4
5
6
7
8
9
10
11
12
0.000075 M Fe(NO3)3
(mL)
4.00
3.50
3.00
2.50
2.00
1.50
1.00
0.80
0.60
0.40
0.30
0.20
1 M KSCN
(mL)
3.00
3.00
3.00
3.00
3.00
3.00
3.00
3.00
3.00
3.00
3.00
3.00
6. Cover each tube with parafilm and mix well.
0.1 M HNO3
(mL)
1.00
1.50
2.00
2.50
3.00
3.50
4.00
4.20
4.40
4.60
4.70
4.80
58
7. The spectrophotometer must be “zeroed.” This means that the light absorbed by the
solvent (aqueous Nitric acid in this experiment) and light scattered by the cuvet must be
blanked out so that it does not register on the display. The procedure below must be
followed each time you “zero” the spectrophotometer.
GENESYS 20 INSTRUCTIONS
a. Set the wavelength to 450 nm.
b. Rinse a cuvet, first with distilled water, and then with 0.5 to 1 mL of 0.1-M HNO3. Discard
the rinse and fill the cuvet about three-fourths full with 0.1-M HNO3. Wipe any
fingerprints off the cuvet with a chem wipe.
Cuvets are made of special optically uniform plastic that needs to be protected against
scratches. Use only chem wipes to clean them and do not allow chemicals to stand in them.
Rinse them well with distilled water immediately after using them.
c. Place the cuvet containing the HNO3 into the cuvet holder. Position the cuvet so the light
passes through the clear walls. Close the lid.
d. Press the A/T/C button to select absorbance (A) mode. Press the 0 Abs/100%T button.
After a few seconds, 0.000 should be displayed.
e. The spectrophotometer has now been “zeroed.” Discard the HNO3 in the cuvet, rinse the
cuvet with 1 to 2 mL of the well-mixed solution in tube 1 and discard the rinse. Fill the
cuvet about three-fourths full with the solution from tube 1, wipe the cuvet with a chem
wipe and place it in the cuvet holder with the label on the cuvet facing forward and close
the lid. Read and record the absorbance reading from the display.
f.
Repeat this procedure for each of the other solutions in tubes 2 through 12. (You do not
have to re-zero the machine; just continue with the next solution, rinsing the cuvet as
before.)
g. Clean all your glassware and return the cuvet and pipet pump.
h. Calculate the molarity of complex ion, Fe(SCN)2+, that was present at equilibrium for tubes
1 through 12. Note that the concentration of KSCN was very large compared to the
concentration of iron(III) nitrate. This resulted in the reaction for the formation of the
complex ion being driven essentially to completion and you can assume that all the
iron(III) ion was converted to the complex ion. However, to calculate the concentrations
you must take dilution into account. (Note: the final volume for all 12 tubes was 8.00 mL.)
i.
Construct a graph (using the graph paper provided after Appendix D) plotting absorbance
(vertical axis) as a function of concentration (horizontal axis). Be sure to properly label
and title your graph. Draw a best fit straight line through the data. (See Appendix C for a
review of graphing.) Include 0.00, 0.00 as a data point.
59
B. Determination of Kf for the Fe(SCN)2+ complex ion.
1. Each group will need a cuvet, a 1.00 mL, a 5.00 mL and a 10.00 mL graduated pipet and
a pipet pump. If you must wash beakers or test tubes, follow the same rinse procedure as
for Part A.
2. In separate clean, dry beakers obtain about 15 mL 0.0025 M Fe(NO3)3, about 20 mL of
0.0025 M KSCN, and about 60 mL of 0.1 M HNO3. Be sure to use the correct
concentrations. Different concentrations were used in part A. Do not take more reagent
than you need.
3. Label the 1 mL pipet for Fe(NO3)3, the 5 mL pipet for KSCN and the 10 mL pipet for HNO3.
Prior to using, rinse each pipet with between 0.5 and 1.0 mL the reagent to be used in that
pipet.
4. Set up 10 large clean dry test tubes. Pipetting carefully, using the pipets you have just
prepared, transfer the following amounts of reagent to each tube.
Tube No.
1
2
3
4
5
6
7
8
9
10
0.0025 M Fe(NO3)3
(mL)
0.50
0.50
0.50
0.50
0.50
1.00
1.00
1.00
1.00
1.00
0.0025 M KSCN
(mL)
0.50
1.00
1.50
2.00
2.50
0.50
1.00
1.50
2.00
2.50
0.1 M HNO3 (mL)
7.00
6.50
6.00
5.50
5.00
6.50
6.00
5.50
5.00
4.50
5. Cover each tube with Parafilm™ and mix well.
6. Following the procedure from Part A, set your spectrophotometer on 450 nm, zero the
machine using 0.1 M Nitric acid solution and determine the absorbances of the solutions
in the tubes 1 through 10.
7. Clean and return borrowed items.
8. Using your standard curve prepared in Part A, determine the equilibrium concentration
of the complex ion in tubes 1 through 10.
9. Calculate the initial concentration of Fe3+ and SCN- in each tube. You must account for
dilution.
10. Calculate the equilibrium concentrations of Fe3+ and SCN-.
11. For each tube, calculate the experimental equilibrium constant, Kf, for the formation of
the complex ion. Calculate the average Kf and the percent relative average deviation (see
Appendix A).
60
QUESTIONS FOR DETERMINATION OF Kf …
NAME ______________________
1. What most affects the precision for this experiment?
2. What most affects the accuracy for this experiment?
3. Would you expect all the Kf values determined in this experiment to be the same (or nearly
so)? Explain your answer.
4. The reaction:
Co2+ + 6 NH3  [Co(NH3)6]2+
has an equilibrium constant of 5.0  104. Solutions were mixed so that the initial
concentration of the cobaltous ion was 0.250 M and the ammonia was 2.00 M. What are
the equilibrium concentrations of all three species in the reaction?
61
DETERMINATION OF PERCENT OXALATE BY OXIDATIONREDUCTION TITRATION
INTRODUCTION
The fundamental event in an oxidation-reduction reaction is electron transfer. Balancing an
oxidation-reduction equation requires that the quantity of electrons lost by the reducing agent
be equivalent to those gained by the oxidizing agent. Determination of the equivalence point
in an oxidation-reduction reaction can be accomplished by using titration techniques.
In this experiment each student will work alone and:
a. prepare an approximately 0.02 M potassium permanganate solution.
b. standardize the permanganate solution against pure, solid sodium oxalate.
c. determine the percent by mass of sodium oxalate in an impure sample.
Heat is required to catalyze (speed up) the permanganate-oxalate ion reaction. Manganese(II)
sulfate will also be present to catalyze the reaction. A catalyst speeds up a reaction but is not
consumed in the reaction.
Notice that when permanganate reacts with oxalate in the presence of an acid, two of the
products are manganese(II) ion and carbon dioxide. If acid is not present in sufficient
quantity, the permanganate ion will instead react according to:
4 MnO4− + 2 H2O → 4 MnO2 + 3 O2 + 4 OH−
Evidence that this reaction has occurred and contaminated the titration is the appearance of
a muddy brown color due to manganese(IV) oxide. This reaction must be avoided!
Should a brown color persist in a flask during a titration, then that trial must be discarded. A
momentary brownish discoloration, which completely disappears, is nothing to worry about.
(however, you should add the permanganate solution more slowly during subsequent
titrations to avoid the chance of permanent contamination by manganese(IV) oxide.)
62
PROCEDURE
A. Preparation of the Potassium Permanganate solution
1. Thoroughly clean one of your large screw cap bottles. Also clean its cap. (If the bottle or
cap has been stained brown see your professor for cleaning instructions.)
2. Calculate the mass of solid potassium permanganate required in the preparation of
approximately 300 mL of a 0.02 M solution.
3. Potassium permanganate is very corrosive to metal and can destroy analytical balances.
Use the beam balances in the lab room to weigh out the approximate mass of
potassium permanganate needed and place it in the clean bottle. Add ten drops of 0.001M sulfuric acid solution and add approximately 100 mL of deionized water.
4. Cap the bottle and mix well. After all solid potassium permanganate has dissolved; fill the
bottle to the 300 mL mark with deionized water. Mix well.
5. Put your name and/or locker number on the bottle.
Note: The permanganate solution must "rest" for several days before it can be
standardized. It must be shaken before each use to insure uniformity. Also, because a
potassium permanganate solution will degrade if exposed to light for extended periods of
time, you should store the solution in your dark lab drawer as much as possible.
B. Standardization of the Potassium Permanganate solution
1. Obtain and clean the buret assigned to your locker number (see signs posted in the lab).
2. Rinse the buret three times with 2 to 3 mL of your potassium permanganate solution. If
you wish to use a beaker or funnel to help fill your buret you must also clean them and
then rinse them with your potassium permanganate solution prior to their use.
3. Fill the buret with the permanganate solution, drain the solution through the buret tip to
eliminate air bubbles, and note the initial buret reading. Due to the intense color of the
permanganate solution it will be easier to read the volume at the top of the meniscus.
4. Obtain a shell vial with pure sodium oxalate standard (Formula Mass = 134.00 g mol−1).
Label this vial and keep it capped when it is not in use.
5. Clean a 250 mL Erlenmeyer flask (it does not have to be dry).
6. Calculate the mass of pure sodium oxalate that will require about 25 mL of approximately
0.02 M potassium permanganate solution for complete reaction with the sodium oxalate,
you must have the balanced chemical equation for this step.
7. Take the vial of pure sodium oxalate, your clean flask, and the data sheet to the analytical
balance room and measure sodium oxalate into the flask. Use the approximate mass (+/0.05 g) calculated in step 6 as a guide. Place the vial of sodium oxalate on the analytical
balance pan and zero the balance while the vial remains on the pan. Carefully avoiding
spills, transfer some of the sodium oxalate into the flask and then place the vial back on
the balance pan (without re-zeroing). The negative mass displayed is the mass of sodium
63
oxalate dispensed into the flask. Record the exact mass of sodium oxalate dispensed into
the flask to the nearest 0.0001 g.
8. Add about 25 mL of distilled water to the flask.
9. Add about 25 mL of the 3 M sulfuric acid solution (that contains some manganese(II)
sulfate catalyst) to the flask. The oxidation-reduction reaction is catalyzed by
manganese(II) sulfate but the catalyst in no way affects experimental results.
10. Set up a water bath using about 200 mL of tap water in your largest beaker on a hot plate.
Heat the water to a gentle boil and then place the Erlenmeyer flask containing the acidified
oxalate solution in this bath for 1-2 minutes (until the flask feels quite warm).
11. Titrate the warm, acidic oxalate solution in the flask with the permanganate solution,
swirling constantly. When the flask starts to cool, return it to the hot water bath for a
minute or two and then continue the titration. The end point is the persistence of the
(diluted) permanganate color (a very light purple). If you are not extremely careful as you
approach the end point you will add too much permanganate solution, the end point color
will be too dark, and you will have to discard that trial.
12. Record the final buret reading and calculate the total volume of permanganate used for
the titration. The contents of the flask can now be discarded. Save your oxalate for
additional trials.
Note: If your titration volume was at least 10.00 mL (4 sig. figs.) this titration can be
included in your calculations. However, a larger titration volume (closer to 25 mL) will
give better precision. On the other hand, an unnecessarily large titration volume (more
than 25 mL) is time consuming. The volume of potassium permanganate solution
required is directly proportional to the mass of oxalate titrated. If the volume for your
first your titration was not between 20 and 25 mL, adjust the mass used for the rest of your
trials.
13. Clean three Erlenmeyer flasks (they can be wet) and label them #1, #2, and #3.
14. Take your oxalate, the flasks, and your report sheet to the balance room and weight out a
sample of oxalate into each flask using the method in Step 7. Record the mass of oxalate
transferred to each flask.
Note: You should refill the buret for each titration. The potassium permanganate
solution remaining in your buret at the end of each lab session should be saved in a clean
dry beaker and used for rinsing the buret at the next lab session. Never put any unused
solution back into your stock bottle. You risk contaminating your solution.
15. Titrate each flask following the procedures in Steps 8-12.
16. From the volume of the potassium permanganate solution used in each titration and the
mass of sodium oxalate in each flask, calculate the molarity of the potassium
permanganate solution (at least three values are necessary for a good average).
64
17. Using at least three molarity values calculate your percent relative average deviation (see
Appendix A at the end of this lab manual). (Note: Percent relative average deviation is a
measure of precision and at least 3 trials are required for the calculation to be
meaningful.). If your average deviation is less than 2%, it means that the data you have
collected show good precision and you have performed enough trials. If it is greater than
2%, then additional trials are needed.
18. Determine and record the average molarity for your potassium permanganate solution.
This solution will be used to analyze your unknown. Take good care of it!
C. Determination of percent sodium oxalate in an impure sample
1. Obtain a shell vial containing an impure sodium oxalate unknown. Record your
unknown number and label your unknown vial. Keep this unknown in your locker until
you have received your graded lab report.
Note: The shell vial of impure sodium oxalate (unknown) contains enough sample for at least
six trials. No additional unknown will be provided! Should you spill your unknown, a
different unknown will be obtained and you will start the unknown’s analysis from the
beginning.
2. Check out and clean the buret assigned you your locker number. Rinse and fill the buret
with your permanganate solution, as before.
3. Do one titration with the unknown using two to three times as much mass as was used for
pure sodium oxalate (record the exact mass of unknown used). Use the same titration
procedure as was used in steps 7-12 for Part B.
4. You will need to do at least two more trials. If the total volume of potassium
permanganate solution used in your first titration was less than 20 mL use a little more
unknown for your subsequent titrations. If your titration volume was greater than 25 mL
use a little less unknown. (The mass of impure sodium oxalate and the volume of
potassium permanganate solution used in the titration are directly proportional.)
5. Using the average molarity of your potassium permanganate from part B in your
calculations, determine the percent by mass of sodium oxalate in your impure sample for
each trial.
6. Using the percent by mass calculated for each trial, determine the percent relative average
deviation. If it is greater than 2%, do additional trials.
7. Report the average percent by mass of sodium oxalate for your unknown.
65
QUESTIONS FOR RE-DOX EXP.
NAME ______________________
1. Dichromate and ferrous ions react in acidic solution to form chromic and ferric ions,
respectively. If 1.285 grams of iron(II) bisulfate dissolved in sulfuric acid solution requires
35.78 mL of sodium dichromate solution for complete titration, what is the molarity of the
sodium dichromate solution? Show the balanced net ionic and complete chemical
equations for this reaction.
2. The sodium dichromate solution from problem 1 was used to titrate a solution made by
dissolving 3.500 g of a pure ferrous salt in sulfuric acid. The titration required 47.22 mL
of the sodium dichromate solution. Calculate the percent by mass of iron in the pure salt.
The net ionic equation here is the same as in problem 1. You will not be able to write a
balanced molecular equation for this because the anion in the ferrous salt was not specified
in this problem.
66
67
ELECTROCHEMISTRY
INTRODUCTION
Electrochemistry involves the transfer of electrons from a reducing agent to an oxidizing
agent. For the electrons involved in the transfer to be used in a productive fashion (e.g.
electroplating flatware, starting a car, etc.), an electrochemical cell is usually set up. An
electrochemical cell is a device that converts the energy of a chemical reaction into electrical
energy. In such a cell, the reaction proceeds by the transfer of electrons, producing an electric
current. A reaction involving the transfer of electrons is called an oxidation-reduction
reaction. If the oxidized and reduced species are separated from each other in different
containers but are allowed to maintain contact through a salt bridge or porous cup, the
electron transfer can be made to occur through a wire which is in contact with the oxidized
and reduced species. The flow of electrons through the wire, called the current, can be used
to produce electrical work. The common dry cell, for example, is an electrochemical cell.
When the terminals of the dry cell are connected to a motor, electrons flow from the cell
through the motor, producing work. An electrochemical cell can only function when there is
a complete electric circuit. In a cell in which there are two half-cells, a salt bridge must be
used to maintain electrical neutrality.
An electrolytic cell uses current from an outside source (a battery or other power supply) to
cause a reaction to run in the direction that is “non-spontaneous.” In this laboratory exercise,
the electrolysis of aqueous potassium iodide will be studied.
Students will record observations and information, and then perform calculations pertaining
to the electrolytic cell. (Examples of common electrolytic processes are recharging “dead”
batteries and anodizing metals such as aluminum.)
If a current spontaneously flows when an electrochemical cell's circuit is complete, then the
cell is referred to as a voltaic or galvanic cell. (Examples of these cells are cell phone and
laptop batteries.) This experiment will include the study of voltaic cells formed from half-cells
involving pairs of the following half-reactions:
Cu2+ + 2 e−  Cu
Fe2+ + 2 e−  Fe
Fe3+ + e−  Fe2+
Zn2+ + 2 e−  Zn
If a solid metal is a component in a half-reaction, then that metal will be used as the electrode
for that half-cell. A graphite rod will serve as the electrode in a half-cell which involves no solid
metal.
The voltaic cells in this experiment will not be “standard” cells. In standard cells all molarities
are 1-M, all partial pressures of reactant and product gases are 1 bar. In addition, the “ideal”
standard cell would be constructed with perfect electrical connections and zero resistance
electrical leads and utilizes circuits that draw no current. The imperfect voltages obtained
from the non-standard cells in this exercise will be compared to standard potentials for that
type of cell.
The next exercise in this experiment will be to construct a concentration cell. This cell will
measure the potential generated by a difference in copper(II) concentrations in
copper/copper(II) half cells.
68
Lastly, you will construct an electrochemical cell to determine the solubility product constant
of copper(II) carbonate. This can be accomplished by measuring the potential of a cell which
has a saturated solution of copper(II) carbonate in one of the half cells. This potential,
compared to the cell potential of a standard cell allows us to determine the copper(II) ion
concentration and the Ksp.
PROCEDURE
Due to equipment limitations, students will work in groups during this experiment.
Each group will need a voltmeter, a porous cup, a 100-mL graduated cylinder, and 2 copper
electrodes.
Day 1 – A.
Voltaic Cells
1. Often, the metal electrodes are stored in oil, which must be removed before use. If so,
pour a small amount of acetone on to a paper towel and wipe the metal electrodes well.
Rinse the electrodes with tap water and then distilled water.
DO NOT COMPLETE A CELL'S CIRCUIT UNTIL YOU ARE READY TO MEASURE
ITS VOLTAGE.
2. Obtain a volt meter and insert the red plug into the “V” connector on the meter and the
black plug into the black “COM” connector. Press the button in the center of the dial and
turn
the
voltmeter’s dial to “V.” You should hear a beep from the meter. Set the meter display to
read three places after the decimal. The meter will run through an internal self- check and
will be ready for use when the display reads approximately 0.000 VDC. (The meter should
read zero if you clip the leads together.)
3. Take the volt meter to the various cell set-ups and measure the voltage of each cell.
4. For example, go to the Cu|Cu2+ Fe|Fe2+ cell. Attach one of the lead's alligator clips to the
top of the copper electrode. Attach the other lead's alligator clip to the iron electrode. If
the meter indicates a negative voltage, then it has been hooked up backwards. (This meter
is designed to yield positive voltages when its black lead is connected to the anode.) Swap
the positions of the alligator clips. Record the voltage for this cell. As soon as the voltage
is read, remove one of the alligator clips to break the circuit and stop the current flow.
5. Determine and record the half reactions for the cell. Write the cathode reaction as a
reduction and the anode reaction as an oxidation. Write the overall chemical reaction.
6. Using the half-reaction potentials in your text, calculate the standard voltage potential for
this cell.
7. Write the shorthand cell notation for each cell that you tested. Remember that these
cells were not standard cells.
8. Repeat steps 4 through 7, for each of the half-cell combinations listed on the report sheet.
69
B.
Concentration Cell
1. Measure 1.0 mL of 0.10 M copper(II) nitrate in a clean 10 mL graduated cylinder. Transfer
this solution to a clean 100 mL graduated cylinder and add distilled water to bring the
volume to the 100 mL mark. (Use some of the water to rinse out the smaller cylinder. Add
the risings to the larger cylinder, and then finish the dilution by adding water directly to
the larger cylinder.) Pour this solution into a clean 250-mL beaker. Stir to mix well.
2. Use some of the diluted solution to rinse the 10 mL graduated cylinder before measuring
1.0 mL of the diluted copper(II) nitrate solution. Transfer the 1.0 mL of diluted solution
to the well rinsed 100 mL graduated cylinder and make a second dilution by adding
distilled water to bring the total volume to 100 mL as before. Mix well.
3. Transfer ~50 mL of solution from the second dilution into a clean 150 mL beaker.
4. Place a clean copper electrode in the diluted solution to form a half-cell.
5. Place ~30 mL of 0.10 M copper(II) nitrate solution into a porous cup and carefully place
the porous cup into the 150 mL beaker from step 3. Place a copper electrode into the
porous cup.
6. Record the temperature of the solution in the beaker.
7. Set the voltmeter to 300 mV. Connect the clips to the electrodes and measure the
concentration cell's voltage. The display will be in mV.
8. Add 10 drops of 0.10 M copper(II) nitrate to the more dilute solution in the cell apparatus.
Stir the mixture (you can use the electrode to stir the solution).
9. Measure the concentration cell's new voltage in mV.
10. Rinse the electrodes with distilled water. Dry the electrodes with paper towels.
11. Calculate the concentration of the copper(II) nitrate in the diluted solution taking both
dilutions into account (you have to calculate each dilution separately). Use the data to do
the calculations and answer the questions in the lab report.
12. Return all borrowed equipment.
70
For Day 2, each group will need a voltmeter, a porous cup, a timer, a set of electrodes with
transformer, an electrode holder, and the following electrodes: 1 copper, 1 zinc.
Day 2 – C.
Electrolysis of Aqueous Potassium Iodide
1. Clean the 100-mL beaker and place 50.0 mL of distilled water into it.
2. Weigh out 1.00 g of potassium iodide and put it into the water in the beaker. Stir until the
potassium iodide is completely dissolved.
3. Take the aqueous KI solution to the pH meter that has been set up for use by the class.
Measure and record the initial pH of the solution.
Be careful of the platinum electrodes because they can be easily damaged! Do not twist or
bend them.
4. Insert the ends of the platinum electrodes into the glass tubes of the electrode holder and
place the entire assembly in the beaker containing the aqueous KI solution (see
diagram). (If necessary, the cork electrode holder can rest on top of the
beaker.) The wires should exit through the spout of the beaker and the
electrodes should rest on the bottom of the beaker. Adjust the glass tubes
of the electrode holder up or down so that the glass holds just the very tip
of each electrode. The purpose of the holder is to make sure the
electrodes do not touch each other during the electrolysis. Do not plug in
the transformer until you have the electrodes in the proper orientation!
5. Start the timer as you plug in the transformer. Observe and record what
is happening at each electrode, initially and several times during the
electrolysis.
6. Allow the electrolysis to proceed for 20-25 minutes. (You can start Part D of the
experiment during this time.) Do not move the electrodes until you have unplugged the
transformer! Stop the timer as you unplug the transformer. Record the exact amount of
time elapsed.
7. Remove the electrode assembly. Measure and record the pH of the solution after
electrolysis.
8. Dip the electrodes into the sodium thiosulfate cleaning solution provided in a container in
the hood. Gently swirl the electrodes in the solution for about 30 seconds to remove any
iodine adhering to the electrodes.
9. Carefully rinse the electrodes 2-3 times with distilled water and gently blot them with a
paper towel to dry them. Rinse and dry the glass tubes in the electrode holder.
10. The data that was collected in this experiment will used in calculations in the Report and
Questions section of the lab.
71
D.
Determination of the solubility product constant of copper(II) carbonate.
1. Obtain a volt meter and insert the red plug into the “V” connector on the meter and the
black plug into the black “COM” connector. Push the button in the center of the dial and
turn the voltmeter’s dial to “V.” The meter will beep. Set the meter display to read three
places after the decimal. The meter will run through an internal self-check and will be
ready for use when the display reads approximately 0.000 VDC. (The meter should read
zero if you clip the leads together.)
2. Place ~50 mL of 1.0 M sodium carbonate and a clean copper strip into a 150-mL beaker.
Add 5 drops of 1.0 M copper(II) nitrate solution to form a precipitate (stir the solution).
Record the temperature of the solution in the beaker.
3. Place ~30 mL of 1.0 M zinc nitrate solution into the porous cup. Place a clean zinc
electrode into the porous cup.
4. Carefully place the porous cup into the 150 mL beaker from step 2.
5. Connect the volt meter to the metal strips and record the voltage. Switch the connections
if you get a negative voltage.
6. Return all borrowed equipment.
72
73
QUESTIONS FOR ELECTROCHEMISTRY EXP.
NAME __________________
1. Use the standard reduction potential tables to answer the following questions, (show the
voltages in justifying your answers):
(a) What will happen to an iron nail in cupric nitrate solution?
(b) What will happen when a piece of copper metal is added to a solution of zinc nitrate?
2. Use the Nernst equation to calculate the expected voltage of the concentration cell before
the addition of the 10 extra drops of 0.10 M copper(II) nitrate solution. Use your
experimental temperature.
How did the addition of 10 drops of 0.10 M cupric nitrate affect the concentration cell's
voltage? Why?
3. How do your experimental voltages obtained for the voltaic cells compare with the
standard potentials for those cells? Give some reasons to explain why the experimental
voltages are probably different than the standard voltages.
4. (a) Assuming that there was no overvoltage, what half reaction occurred at the anode of
the electrolytic cell?
(b) At the cathode?
(c) Write a balanced net ionic equation for the overall reaction that occurred in the cell.
(d) Calculate the standard potential for the reaction.
5. Calculate the average current that flowed through the electrolytic cell (using the change in
hydroxide ion concentration, the volume of the solution, and the total elapsed time of the
electrolysis).
74
QUESTIONS FOR ELECTROCHEMISTRY EXP. NAME________________________
6. An experimental cell is set up such that one half cell contains a solid silver electrode
dipping into a saturated solution of silver oxalate. The oxalate ion comes from added
sodium oxalate so that equilibrium concentration of oxalate ions is 0.50 M. This half-cell
is connected to a standard hydrogen electrode (SHE) half-cell. The measured potential
for this cell is 0.476 V at 25°C. What is the experimental value of the solubility product
constant for silver oxalate?
75
76
ELECTROLYTIC DETERMINATION OF THE MOLAR MASS OF
LEAD
INTRODUCTION
If an electric current is allowed to pass through a solution containing ionic species, the ions
experience a force that causes the positive ions to move in one direction while the negative
ions move in the opposite direction. This movement of ions allows the current to pass through
the solution. In order to maintain the current, oxidation-reduction chemical reactions must
occur in the solution.
The amount of electric current that passes through the solution and the amount of chemical
reaction that occurs are related by Faraday's Laws. The transfer of Avogadro's number of
electrons corresponds to one faraday of charge. One faraday is equal to 96,485.3399±0.0024
coulombs of electrical charge. The number of electrons transferred, and hence the number of
faradays, can be found by multiplying the current by the time during which the current flowed.
In this experiment, two lead strips are placed in a lead(ll) nitrate solution and a wire is
attached to each of these strips. The wires are then connected to the Constant Current System.
The Constant Current System is also connected to the Vernier™ LabQuest to monitor current
flowing in the cell. Electrons flow between the current source and the lead strips. One of the
strips receives electrons while the other strip loses electrons. In order to have electron flow,
electrons must be used up at the strip that
gains electrons and must be released at
the strip that loses electrons.
The
positively charged ions will move to the
electron-rich strip and accept electrons;
thus, a reduction occurs at this strip
which can be represented by the halfreaction:
Pb2+ + 2 e−  Pb
The strip where reduction occurs is called
the cathode. Negatively charged ions
migrate to the other strip and electrons
are released; thus, an oxidation must
occur which can be represented by the half-reaction:
Pb  Pb2+ + 2 e−
The strip where oxidation takes place is called the anode.
The mass of lead that dissolves at the anode must be the same as the mass of lead deposited
at the cathode.
77
PROCEDURE
1. Obtain two lead strips from the laboratory instructor. Sandpaper each strip to remove
surface oxides and dirt. Use a file to mark each strip. One scratch mark will designate the
strip to be used as the anode and two scratch marks will identify the strip to be used as the
cathode. Weigh each strip to the nearest ±0.0001 g. Record the mass of the anode and
cathode strips on the Data Sheet.
2. Assemble the apparatus by adding about 100 mL of 0.50 M lead(ll) nitrate solution to a
250-mL beaker containing the two lead strips from step 1.
3. Gently turn the dial of the Constant Current System™ counterclockwise to confirm that it
is in the minimum current position.
4. Place the lead strips into the solution in the beaker. Be sure to keep them as far apart as
possible. You may find it easier if you bend the strips over the edge as shown in the figure
on the previous page.
5. Connect the lead strip you marked as the anode (two scratches) to the positive (red) clip
of the Constant Current System. Connect the other Lead strip (the cathode, one scratch
mark) to the negative (black) clip.
6. Plug the Constant Current System™ into a powered electrical outlet. Connect the sensor
cable to the LabQuest™ and choose New from the File menu.
7. Start data collection by pressing the start button or the green arrow on the screen and now
adjust the current to the 0.1–0.2 A range. Data collection will run for 30 minutes.
8. When data collection is complete, disconnect the DC power source and carefully remove
the lead strips from the solution. Gently rinse with distilled water. Allow the strips to dry
in air at room temperature. Weigh each strip to the nearest ±0.0001 gram, adding the
mass of any solid that may have fallen off to the cathode and record on the Data Sheet.
9. Perform a second determination following the same procedure as before. Use the same
equipment as before. Swap the cathode and the anode from the first trial. The initial
masses of the cathode and anode lead strips will be the final masses of the anode and
cathode from step 8.
78
QUESTIONS FOR ELECTROLYTIC DETERMINATION … NAME__________________
1. A electrolytic cell is set up as was done in this experiment but with a different metal. An
average current of 135.4 mA is delivered for 15 minutes and 23 seconds. The cathode gains
0.0728 g in mass. If there are two moles of electrons transferred per mole of the metal,
what is the molar mass of the metal?
2. What mass of sodium metal can be obtained from the electrolysis of molten sodium
chloride if a current of 10.0 amps is allowed to pass though the cell for 45 minutes?
3. How many hours are required to obtain 150.0 g of chromium metal from a solution of
chromium(III) nitrate if a current of 2.50 amps is passing through the cell?
79
FACTORS AFFECTING THE RATE OF A REACTION
INTRODUCTION
There are several factors that affect the rate of a reaction. Some of these factors are:
•
•
•
•
•
Mixing
Concentration of a reactant
Temperature
The presence of a catalyst
Surface area in a heterogeneous reaction
In this experiment we will examine these factors. This experiment is an introduction to the
more commonly encountered factors that affect the rate of reactions. There are other factors
that do influence the rate of a reaction such as light, molecular geometry and the type of
solvent used; however, we do not have the time to explore all facets of all factors that affect
reaction rate.
PROCEDURE
A.
The effect of mixing.
1. Fill two small test tubes ¼ full with water.
2. Into each tube add a small crystal of solid potassium permanganate.
3. Let one tube sit undisturbed. Swirl the other tube to dissolve the potassium permanganate.
Note the amount of time it takes the swirled sample to dissolve (form a solution).
4. Continue with the experiment (Parts B through E) and observe the undisturbed tube every
few minutes. Note approximately how long it takes for the potassium permanganate to
dissolve and diffuse throughout this tube.
To complete Parts B and C each group will need a timer, a total of 7.0 mL of 3%(m/m)
hydrogen peroxide solution (H2O2) and a total of 15.0 mL of solution A (a mixture of starch
(as an indicator), acetic acid, potassium iodide, and sodium thiosulfate). Do not waste
reagents by taking more than you need for the experiment!
B.
The effect of concentration of a reactant.
1. In your smallest beaker place 5.0 mL of solution A and add 5.0 mL of the hydrogen
peroxide solution. Start your timer as soon as the solutions are mixed in the beaker.
2. Record the number of seconds that elapse until the solution turns blue/black.
3. Repeat steps 1 and 2 using 5.0 mL of solution A and 4.0 mL of deionized water which has
been added to 1.0 mL of the hydrogen peroxide solution.
80
C.
The effect of temperature.
1. Place a test tube containing 5.0 mL of solution A and another test tube containing 4.0 mL
of deionized water combined with 1.0 mL of the 3% hydrogen peroxide solution into a
warm water bath for about 5 minutes.
2. Measure the temperature of the water bath.
3. Mix the two solutions into a small beaker and start your timer.
4. Record the number of seconds required for the solution to turn blue/black.
5. Compare this number with the elapsed time from the second (diluted) mixture in part B.
D.
The effect of a catalyst.
1. Fill a large (400 mL or larger) beaker about 2/3 full with water.
2. Fill a small test tube all the way with water. Place your finger over the opening and invert
it into the beaker. Do not allow any gas to enter the tube as you remove your finger.
3. Obtain a gas collection apparatus. Place the gas evolution tube under the inverted test tube
in the beaker.
4. Do you notice gas formation in the 3% H2O2 reagent bottle?
5. Place 20 drops of 3 M copper(II) nitrate solution in the flask. Swirl the contents of the
flask. Is any gas formed in the catalyst solution alone?
6. Quickly add about 20 mL of 3% hydrogen peroxide solution to the flask and quickly put
the stoppered end of the tube into the flask. Continuously swirl the flask’s contents.
7. Record the number of seconds required for the tube to fill with the gas produced.
8. Empty the contents of the flask and beaker, clean them and set up the experiment for the
next trial.
9. Place about 5 mL of 3% hydrogen peroxide solution in the flask. Remove the stopper just
long enough to add 2 drops of 3 M iron(III) nitrate solution and quickly put the stoppered
end of the tube into the flask. Continuously swirl the flask’s contents.
10. Record the number of seconds required for the tube to fill with the gas produced.
11. Set up the experiment again with 5 drops of 3 M copper(II) nitrate, 5 drops of 3 M iron(III)
nitrate and 20 mL of 3% hydrogen peroxide solution in that order. Record the number of
seconds required for the tube to fill with the gas produced.
81
E.
The effect of surface area in a heterogeneous reaction.
1. Place a small iron nail that has been cleaned into a test tube. In a second test tube place a
small ball of steel wool.
2. Into each test tube add 5 mL of 1 M Cu(NO3)2 solution and place each tube into a warm
water bath. Occasionally stir the tubes and let them heat for at least 10 minutes.
3. Observe each tube closely and note any color change in the solutions. The intensity of the
color change is an indication of the progress of the reaction. Which tube has a quicker
color change? Record the color intensity (lighter or darker) for each tube.
4. After 10 minutes, decant the solution from each tube and note the appearance of the nail
and of the steel wool.
82
83
QUESTIONS FOR FACTORS EXP.
NAME _____________________
1.
Looking at part A of this experiment, how would you expect the rate of the reaction to
change if you were to stir a reaction mixture instead of just letting it sit?
2.
Recall that the rate of the reaction is inversely proportional to the time measured.
a.
Based on this, in part B, which tube had a higher rate of reaction?
b.
Which tube had the higher concentrations of either or both reactants?
c.
How does concentration affect the rate of the reaction?
3.
In part C of this experiment the temperature was raised above that used in part B.
Comparing the rate of the reaction in the second part of part B and the rate of the
reaction in part C, how does temperature affect the rate of the reaction?
4.
a.
In part D of this experiment how did the presence of the copper(II) nitrate
affect the rate of the reaction?
b.
How did the presence of iron(III) nitrate affect the rate of the reaction?
c.
Which compound is a better catalyst?
d.
Does having both catalysts present increase the rate of the reaction more than
either one alone?
5.
If you have a choice of either to grind up a reactant or leave it in a large lump, which
would you choose so that the reaction rate is increased? Refer to your results from part
E.
84
CHEMICAL KINETICS
INTRODUCTION
It is not possible for one to predict a reaction rate or rate law from a balanced, overall equation.
Information about the reaction mechanism (pathway) must be known to make such
predictions.
Through numerous laboratory studies, experimental rate laws have been found to obey the
general expression:
x
y
z
Rate = k  A  B  C  
where [A], [B], [C], ... represent molarities of all chemical species that affect the rate, and x, y,
z, ... are the experimentally determined exponents for each species. (The overall order of the
reaction is equal to the sum of x + y + z +... .) The term k is known as the rate constant for the
reaction.
Usually, when a reaction is initiated, the rate (known as the initial rate) is found to be at its
maximum value. As the reaction progresses, reactants are consumed (lowering their
concentrations) and the rate slows. One can avoid difficult concentration measurements by
monitoring the initial rate. The concentrations at the time of the initial rate are simply the
initial concentrations after taking dilutions into account. If the initial concentration of one
reactant is varied while all others are held constant, then the resulting change in initial
reaction rate yields the order with respect to that one reactant. This is the initial rate method
used to determine reaction order.
In this experiment you will be measuring the initial rate for the iodide ion and persulfate ion,
S2O82-, reaction:
2 I− + S2O82−  I2 + 2 SO42−
(1)
To detect the extent to which reaction (1) has proceeded, an additional, simultaneous process
must also occur:
I2 + 2 S2O32−  2 I− + S4O62−
(2)
In reaction (2) the thiosulfate ion, S2O32−, instantly reduces iodine molecules back to iodide
ions. Only when the thiosulfate ions have been completely consumed can the iodine formed
in reaction (1) be available to react with the indicator to form the blue-black starch complex.
Therefore, the rate for reaction (1) is equal to the change in thiosulfate concentration per time.
What will be measured in this experiment is the time required to use all of the thiosulfate (the
time required to change the thiosulfate concentration to zero). The elapsed time depends
upon the rate of reaction (1) as well as the amount of thiosulfate added to the reaction mixture.
(Thus, the amount of thiosulfate must be carefully controlled so that the only variables are the
concentrations of iodide and persulfate ions.)
85
For reaction (1) the reaction rate is equal to
x
y
k I −  S2O82− 
0
0
(where the subscripts of zero indicate
initial concentrations for the molarity terms). Thus,
x
y
Rate = k I −  S2O82− 
0
0
Because of the experimental conditions employed, applying the initial rate method to the
above expression yields
x
−
Rate2 k I  2
=
Rate1 k I −  x
 1
y
S2O82 − 
2
" Exp 2"
" Exp 1"
y
2−
8
1
S2O 
Because the initial concentration of iodide ion in experiment 2 equals twice its initial
concentration in experiment 1, after cancelling constant terms we obtain:
Rate2
= 2.00 x
Rate1
Solving for x is simplified by taking the logarithm of both sides. Thus,
Rate2
=
=
log
log
2.00 x x log 2.00
Rate1
further reduces to:
x=
log
(
Rate2
Rate1
)
log 2.00
Calculations similar to those presented above may be derived for y.
As we increase the temperature of a reaction, its rate typically increases. We find the
temperature dependence of the rate of a reaction is a result of the temperature dependence of
the rate constant. In class, we examine two possible explanations of the temperature
dependence of the rate constant. Here, we will restrict ourselves to Arrhenius theory.
Arrhenius conjectured that the rate constant would be a product of two factors; A, the preexponential factor, and an exponential factor that depends on the activation energy and the
temperature:
k = Ae −
Ea
RT
Typically, the activation energy is in J mol-1, R is in J mol-1 K-1, and the temperature is in K.
This equation in itself cannot help us to determine the temperature dependence of k or the
activation energy. To do this we need to use the two-point form of this equation. If we take
the natural logarithm of both sides and subtract the equation at one temperature from the
equation for another temperature, we get:
 k1  Ea  1 1 
ln=
 
 − 
 k2  R  T2 T1 
86
Using this equation, we can calculate the rate constant at different temperatures if we know
the activation energy or calculate the activation energy if we know the rate constant at two
different temperatures. In this experiment we will determine the rate constant at two different
temperatures and then calculate the activation energy of the reaction.
PROCEDURE
1.
Students may work in small groups (2 to 3 students/team). Each team needs a timer.
2.
Each of the 4 experiments will be performed in duplicate. For each experiment, every
team will need two clean flasks and two clean beakers – these can be wet, but should be
well drained.
3.
All reagent bottles have been fitted with Dispensette III bottletop dispensers. These are
designed to deliver an exact volume when used properly. To dispense, turn the red cap
counter- clockwise to remove, position the container you are using under the spout, pull
the piston gently all the way up, then push gently all the way down dispensing into the
container. Finally, replace the cap turning clockwise. Note: all of the Dispensette
apparatus are set to dispense 10.0 mL so for those experiments that require 20.0 mL you
will need to dispense two times. Potassium chloride and sodium sulfate solutions are
used to maintain a constant ionic strength while diluting reactants in this experiment.
Read labels carefully.
4.
Obtain reagents and perform one experiment at a time. For experiment 1, prepare 2
flasks, each containing the volume of potassium iodide, potassium chloride, sodium
thiosulfate, and starch (indicator) solutions provided in the table below.
5.
Also for experiment 1, prepare 2 beakers each with the volume of potassium persulfate
shown in the table below.
0.200-M KI
0.200-M KCl
0.00500-M
Na2S2O3
Starch
0.100-M K2S2O8
0.100-M Na2SO4
Volumes of solution, mL
Room Temperature
Experiment 1
Experiment 2
Experiment 3
In a flask:
10.0
20.0
20.0
10.0
0
0
10.0
10 drops
20.0
0
10.0
10.0
10 drops
10 drops
In a beaker:
20.0
10.0
0
10.0
Ice bath
Experiment 4
20.0
0
10.0
10 drops
20.0
0
87
6.
Start the timer as the contents of one beaker are added to the contents of one flask. Mix
the reagents by quickly pouring the contents of the flask into the beaker and then
returning the solution to the flask. Allow the flask to sit undisturbed while observing its
contents constantly. Stop the timer when the blue-black color appears. (Constant
observation of the flask is necessary because of the sudden appearance of the blue-black
color.) Record the elapsed time in seconds. Repeat step 6 using the second beaker and
flask. Elapsed times for duplicate sets of experimental conditions should agree within
about 2-3 seconds. (If not, do a third trial for that experiment.)
7.
Clean your beakers and flasks and allow them to drain well.
8.
Repeat steps 4, 5, 6, and 7 for experiment 2 and then experiment 3 with the same mixing
procedure for each experiment. Note: In experiments 2, 3, and 4, no KCl is required in
the flask, but sodium sulfate as well as potassium persulfate is needed for experiment 3.
(Read steps 10 through 12 before doing experiment 4.)
9.
Record the temperature of one of the reaction solutions. (Because all of the solutions
have been sitting at room temperature, you can assume this is the temperature for all the
solutions.)
10. For experiment 4, place the volumes of the solutions indicated into clean flasks and
beakers. Fill four of your largest beakers with ice and set the reagent flasks and beakers
onto the ice. Allow the flasks and beakers to remain on ice for at least 5 minutes.
11. Record the temperature of one of the solutions in the ice bath.
12. Repeat step 6. Return the solutions to the ice bath immediately after mixing. When the
reaction flask starts to show a color change, swirl the flask and determine the elapsed
time.
13. Perform dilution calculations to determine the molarities at the instant of mixing for the
iodide, persulfate, and thiosulfate ions used in each experiment. You can assume that the
10 drops of starch is 0.5 mL and that the total volume in each experiment is 50.5 mL.
14. Calculate the average elapsed time for each experiment (in seconds).
15. Calculate the Rate for each reaction. The Rate is equal to:
 thiosulfate  f − thiosulfate i
−

elapsed time





Remember that the final concentration of thiosulfate for each experiment is zero!
16. Using the average rates from experiments 1, 2, and 3 (but not 4), calculate the order of
reaction (1) with respect to the iodide ion and the persulfate ion. Record the orders using
the number of significant figures appropriate for your data.
88
17. Round off the orders you have determined for iodide and persulfate to the nearest whole
number and then calculate the overall order for the reaction (1). Use the Rate and the
rounded orders to calculate the rate constant, k, for each of experiments 1, 2, and 3.
Report your average value of k (include units).
18. Write the complete rate law for the reaction for the reaction that occurred at room
temperature.
19. Determine the rate constant, k, (include units) for the reaction in the ice bath. Using the
average rate constant from the room temperature experiments and the rate constant from
the ice bath, determine the activation energy of the reaction.
89
QUESTIONS FOR CHEMICAL KINETICS EXP.
NAME _______________________
1. The reaction:
2 NO + Cl2  2 NOCl
has been studied and found to be second order with respect to nitrogen monoxide and first
order with respect to chlorine.
a. What is the overall order for the reaction?
b. How does the reaction rate change when the nitrogen monoxide concentration is
halved and the chlorine concentration is doubled? Define terms (e.g., [NO]1 for initial
concentration in experiment 1, [NO]2 for initial concentration in experiment 2, [NO]2 = ½
[NO]1), set up the rate law ratios and show cancellations for Exp2 .
Exp1
2. The initial rate of a reaction is found to increase by a factor of sixteen when the
concentration of one reagent is doubled while all other reagent concentrations are held
constant. What is the order of the reaction with respect to that one reagent? Define
terms, set up the rate law ratios and show cancellations for Exp2 .
Exp1
3.
At 593K a particular decomposition’s rate constant had a value of 2.88×10−4 and at 673K
the same reaction’s rate constant was 1.94×10−3. It was noticed that when the reactant’s
initial concentration was 0.1250 M (with a 593K reaction temperature), the initial
reaction rate was identical to the initial rate when the decomposition was run at 673K
with an initial reactant concentration of 0.04816 M. Recall that rate laws have the form
rate = k [A]x and, showing work, determine the order of the decomposition reaction.
90
4. The following data was obtained for a reaction in which a chemical, X, decomposed.
Concentration of X (in Molarity)
5.00
3.52
2.48
1.75
1.23
Time (in seconds)
0
5.00  102
10.00  102
15.00  102
20.00  102
Chem 102 students are expected to prepare proper graphs (or lose points!) Appendix B of this
document is a reprint of the Chem 101 lab manual’s graphing exercise which includes
instructions for proper graph construction by hand and using Excel™.
a. Prepare plots of concentration versus time using the provided data in a manner
appropriate for zero, first and second order processes. You must include all three graphs
with your report.
b. Based on your graphs, is this reaction zero, first, or second order for X?
c. Determine the slope for the straight line graph. Show how you arrived at the value for
the slope of the line. Calculate the rate constant for this reaction from the slope.
d. Write the complete rate law for the reaction including the value of k (with units).
91
DETERMINATION OF THE HALF-LIFE OF POTASSIUM-40
INTRODUCTION
All radioactive isotopic decay follows first order kinetics. We explored kinetics in the first
week of this class and found that first order kinetics obeys the following equation:
−
∆N
kN0
=
∆t
where k is the rate constant with units of reciprocal time and N0 is the initial number of nuclei.
We can get a measurement of the rate (the left hand side of the equation above) by using a
Geiger-Müller detector. The detector we are using is the Digital Radiation Monitor made by
Vernier. Once we have a rate measurement, we can calculate the number of radioactive nuclei
in the sample from the mass. The rate and the number of nuclei then give us the rate constant.
The half-life, t1/2, can be found from the following equation:
t 12 =
ln2
k
In this experiment we will make adjustments to the count to take into account the beta
counting efficiency of the detector (not all beta particles emitted get measured by the detector)
and to adjust to a (hypothetically) infinitely thin disk of KCl (the KCl will start to absorb some
of the beta particles as the sample gets thicker).
PROCEDURE
1. Obtain a Digital Radiation Monitor (DRM). Turn it on using the bottom switch on the
front of the device. In order to avoid annoying your instructor and classmates do not put
the switch in the “Audio” position.
2. The first task is to obtain a background count. The background count is more accurate the
longer it is measured. We are going to do a 30-minute count for the background. To set
the timer on the DRM for 30 minutes we need to put the top switch into the “Total/Timer”
position. The display should show a time measure and the word “SET” in the upper right.
Use the “+” and “−” buttons on the top left of the DRM to set it to 30 minutes (display will
show “0:30”).
3. Press the “Set” button (between the “+” and “−” buttons). The DRM will start totaling the
counts it measures. At the end of the counting period the DRM will beep three times.
Record the number in the display as the Background Count on your Report Sheet. Divide
this number by 30 and record it as the “Background Counts Per Minute (CPM)” on your
Report Sheet. After the background count has been completed, reset the timer on the DRM
by placing the switch in the CPM/CPS position and then moving it back to the
“Timer/Total” position. Use the “+” and “−” buttons to adjust the timer to 10 minutes
(0:10 in the display) which is the time used for the remainder of the experiment.
4. Measure approximately 0.7 gram (±0.0001 g) of potassium chloride on a piece of tared
weighing paper. Record the mass in your Report Sheet. Place the potassium chloride into
a clean dry shell vial. Tightly cover the shell vial with a piece of Parafilm™. This is your
sample holder. We will place the shell vial upside-down in a buret clamp to hold it in place
over the counter window of the DRM (located on the top right side of the device). We will
92
also clamp the DRM in place with a buret clamp to make sure that it doesn’t move out of
position (See picture below).
5. Press the “Set” button to start the timer and the count. When
the DRM beeps, record the total count in your Report Sheet.
Divide the count by 10 to get the CPM and record this
number as the CPM in your data sheet.
6. Measure approximately 0.3 gram (±0.0001 g) of potassium
chloride and add it to the potassium chloride already in the
shell vial. Record the amount measured in your data sheet.
Cover the shell vial tightly with another piece of Parafilm™.
Reset the timer as before. Set everything up as before and
press the “Set” button to start the count again for 10 minutes.
Record the total and the CPM as before.
7. Measure approximately 1.0 gram (±0.0001 g) of potassium
chloride and add it to the potassium chloride already in the
shell vial. Record the amount measured in your data sheet.
Cover the shell vial tightly with another piece of Parafilm™.
Reset the timer as before. Set everything up as before and
press the “Set” button to start the count again for 10 minutes.
Record the total and the CPM as before.
8. Measure approximately 1.0 gram (±0.0001 g) of potassium chloride and add it to the
potassium chloride already in the shell vial. Record the amount measured in your data
sheet. Cover the shell vial tightly with another piece of Parafilm™. Reset the timer as
before. Set everything up as before and press the “Set” button to start the count again for
10 minutes. Record the total and the CPM as before.
9. Return the potassium chloride to the container. Turn off the DRM and return it to the
same place from which you got it.
10. For each of the samples, calculate the Net CPM by subtracting the background CPM from
the measured CPM. Then calculate CPM/g KCl and ln(CPM/g KCl) for each sample and
record these values on your Report Sheet.
11. Construct a graph of ln(CPM/g KCl) vs g KCl and extrapolate back to zero grams KCl. The
intercept corresponds to an infinitely thin disk of KCl. This process eliminates the effect
of the self-absorption of beta particles by the potassium chloride. Calculate the
extrapolated CPM/g KCl from the intercept and record this on your Report Sheet.
12. Next we need to make an adjustment for the efficiency of the detector. Prior
measurements of samples with known activities have determined the efficiency of the
detector to be about 1.5%. We can then calculate the Adjusted CPM/g KCl from
=
Adjusted CPM/g KCl
Extrapolated CPM/g KCl Extrapolated CPM/g KCl
=
efficiency
0.015
Record the Adjusted CPM/g KCl on your report sheet.
13. The Adjusted CPM/g KCl is our experimental activity. We need to calculate the number
of potassium-40 nuclei in a 1.000 g sample (because we adjusted everything to be per
93
gram). To accomplish this, you will need the isotopic abundance of potassium-40 which
is 0.0118%. Record the number of K-40 nuclei/g KCl in your Report Sheet.
14. Calculate the rate constant using
k=
Adjusted CPM/g KCl
number of K-40 nuclei/g KCl
Record this value on your Report Sheet.
15. Calculate the half-life, t1/2, by taking
t 12 =
ln2
k
This value will have units of minutes. Convert your answer to years. Record this on your
Report Sheet.
94
95
QUESTIONS FOR HALF-LIFE OF K-40 EXP.
NAME ______________________
1. Calculate the activity of a 15.0 g sample of natural potassium. Express your answer both
in Curies (Ci) and in Becquerel (Bq).
1 Bq = 1 nuclei s−1 1 Ci = 3.700×1010 nuclei s−1
2. Technetium-99m is a metastable form of technetium-99 (isotopic mass = 98.9062547
amu) and has a half-life of 6.0058 hours. How many grams of Tc-99m are required to
have an activity of 1.00 µCi?
3. Cesium-137 is a radioactive isotope. 10.0 g of pure Cs-137 (isotopic mass = 136.9070835)
has an activity of 871.8 Ci. What is the half-life of Cs-137 in years?
96
EQUILIBRIUM BETWEEN TWO COMPLEX IONS OF Co2+ IN
SOLUTION
INTRODUCTION
Co2+ in solution can be surrounded by either four or six species in tetrahedral or octahedral
geometries, respectively. Such structures, called complex ions, are stable because the central,
positively charged Co2+ attracts the negatively charged, or electron-rich, portions of the
coordinating species. The number of species surrounding the Co2+ depends on the charges
and the sizes of the ligands. The complex’s structure determines its resulting color:
tetrahedral Co2+ complex ions are deep blue, while octahedral ones are light pink.
When we dissolve cobalt(II) chloride (CoCl2) in water (H2O), the Co2+ retains one chloride ion
and attracts the electronegative, electron rich, oxygen end of water molecules. The resulting
complex ion consists of one Co2+ ion, one chloride ion and five water molecules in an
octahedral configuration with a light pink color.
The size of the ligands is one of the factors that determine the structure of a Co2+ complex ion.
Table 1 shows the geometries and colors of the complex ions formed when CoCl2 is dissolved
in a variety of solvents. Alcohols are structurally similar to water in that they all have –OH
groups with the other hydrogen replaced by an organic group. We can see in Table 1 that the
geometry of all Co2+ complex ions is either tetrahedral or octahedral and not anything else.
Table 1 Co2+ coordination complex ion color and structure for CoCl2 dissolved
in various alcohols
Co2+ coordinatio
Density
Solvent molecule:
Solvent
solution color
-3
n
(g cm )
group attach to –OH
Water
1.000
Hlight pink
octahedral
methanol
0.7914
CH3light pink
octahedral
Ethanol
0.7893
CH3-CH2dark blue
tetrahedral
0.7796
CH3-CH2-CH2dark blue
tetrahedral
Propan-1-ol
0.7851
(CH3)2-CHdark blue
tetrahedral
Propan-2-ol
When we dissolve CoCl2 in a mixture of methanol (CH3OH) and propan-2-ol
(CH3CH(OH)CH3) the solution will contain two types of Co2+ complex ions: octahedral ones
(with methanol) and tetrahedral ones (with propan-2-ol). An equilibrium will be established
between the two complex ions.
[CoCl(CH3CH(OH)CH3)3]+ + 3 CH3OH  [CoCl(CH3CH(OH)CH3)2(CH3OH)3]+ +
CH3CHOHCH3
This equation can be simplified as
[CoClP3]+ + 3 M  [CoClP2M3]+ + P
where P and M represent propan-2-ol and methanol molecules, respectively. Because this
equilibrium is between tetrahedral and octahedral complex ions, we can also represent it as
97
[Co(tet)] + 3 M  [Co(oct)] + P
where the subscripts indicate the geometries of the complex ions. We can then use this last
equation to write an equilibrium constant expression for this system
K eq
Co  P
( oct ) 
= 
Co  M3
 (tet ) 
This equilibrium constant, like all equilibrium constants, depends only on the temperature of
the system.
The color of the system is determined by the relative proportions of the two complex ions. If
we add methanol to the equilibrium mixture, the equilibrium position will shift in accordance
with Le Châtelier’s principle. Thus, the dark blue solution of Co2+ in pure propan-2-ol
becomes a lighter blue upon addition of methanol, because of production of the octahedral
complex ion at the expense of the tetrahedral complex ion.
As previously stated, the equilibrium constant for a system will only change if the system
temperature changes. Octahedral complex ions of Co2+ are reported to be favored over
tetrahedral complex ions by 31 kJ/mol. Given this information, we can conclude that the
reaction to form an octahedral Co2+ complex ion from a tetrahedral Co2+ complex ion is
exothermic, with H equal to −31 kJ/mol.
Color results from the absorption or transmittance by matter of certain wavelengths of light
within the visible spectrum. Tetrahedral Co2+ complex ions have a strong absorption of light
in the yellow-to-red wavelengths that results in a dark blue solution color. Octahedral Co2+
complex ions have a weak absorption of light in the blue-to-green wavelengths that results in
a pale pink solution color.
Initially, we use molecular models to demonstrate the relationship between the geometry and
color of a Co2+ coordination complex ion. You will construct models of different Co2+
coordination complex ions in order to illustrate how the geometry of a complex ion depends
on the sizes of the coordinating molecules. Then you will identify the color associated with
each of the geometries.
In this experiment you will take advantage of the intense absorbance of the tetrahedrally
coordinated Co2+ in solution. The wavelength of maximum absorbance for solutions of CoCl2
dissolved in propan-2-ol is 657 nm, so we use 657 nm as the analytical wavelength for these
solutions. The absorbance at 657 nm for a CoCl2–propan-2-ol solution is proportional to the
concentration of tetrahedrally coordinated Co2+ present. The general mathematical
relationship between absorbance (A) and concentration (c) of the absorbing species is known
as Beer’s law. One form of Beer’s law is
A = ε bc
where ε is the proportionality constant, relating absorbance to concentration for solutions
measured in cuvets of a constant size, b. For the absorbance at 657 nm for solutions of
tetrahedrally coordinated Co2+ formed by dissolving CoCl2 in propan-2-ol, we can write Beer’s
law as
98
A = ε b Co(tet ) 


We can determine the value of ε as follows. First, we measure the absorbances for various
CoCl2–propan-2-ol solutions with known concentrations of tetrahedrally coordinated Co2+.
We plot each absorbance as a function of the corresponding Co2+ concentration, and then draw
the best straight line through the data points and the origin (because a solution with zero
concentration has zero absorbance). The slope of this straight line is the value of ε.
Consequently, we can determine [Co(tet)] for any CoCl2–propan-2-ol solution from its
absorbance at 657 nm and this value of ε.
You will follow the above method to determine the value of ε for solutions of CoCl2 in 2propanol. To do so, you will prepare several solutions, each containing a known concentration
of CoCl2 in propan-2-ol. All of the Co2+ in these solutions is in a tetrahedral geometry, and the
intensity of the blue solution color is proportional to [Co(tet)].
You will use a
spectrophotometer to measure the absorbances of the solutions at 657 nm. Based on these
data, we determine the value of ε.
You will then determine the equilibrium constant for the conversion of tetrahedrally
coordinated Co2+ to octahedrally coordinated Co2+ in solutions of CoCl2 in propan-2-ol and
methanol. To do so, you will prepare solutions of CoCl2 in various mixtures of propan-2-ol
and methanol, measure the absorbance of each solution at 657 nm, and relate the absorbance
to the corresponding concentration of tetrahedrally coordinated Co2+. The equilibrium
constant for this system defines the relative stability of tetrahedrally coordinated Co2+ versus
that of octahedrally coordinated Co2+.
PROCEDURE
1. Rinse a clean, dry 10-mL graduated cylinder with about 1 mL of propan-2-ol. Rinse again
with another 1 mL of propan-2-ol. Also rinse a clean, dry 13×100-mm test tube twice, using
about 1 mL of propan-2-ol each time. Pour all rinses into your “Discarded Solutions’’
container. Use a spatula to place about 0.01 g of cobalt chloride hexahydrate, CoCl2 ∙
6H2O, in the test tube. Using the graduated cylinder, add 5 mL of 2-propanol to the test
tube. Use a clean glass stirring rod and stir to dissolve the solid. Record the color of the
resulting solution.
2. Rinse the 10-mL graduated cylinder twice, using about 1 mL of methanol each time. Rinse
a second clean, dry 13×100-mm test tube twice, using about 1 mL of methanol each time.
Pour all rinses into your ‘‘Discarded Solutions’’ container. Use a spatula to place about
0.01 g of CoCl2∙6H2O in the test tube. Using the graduated cylinder, add 5 mL of methanol
to the test tube. Use a second clean, dry glass stirring rod to dissolve the solid. Record the
color of the resulting solution.
3. Pour the contents of the two test tubes into your “Discarded Solutions” container.
99
4. Use your molecular model kit to construct models of the following Co2+ coordination
complex ions:
(a) [CoClP3]+: one Cl- and three propan-2-ols tetrahedrally arranged around a central
Co2+
(b) [CoClP5]+: one Cl- and five propan-2-ols octahedrally arranged around a central
Co2+
Remember that the oxygen atom in the propan-2-ol is what links the alcohol to the central
Co2+. Compare the two structures. Identify and record which structure is too crowded for
all of the propan-2-ols to easily fit around the Co2+, and therefore will be unstable.
5. Using your molecular model kit, construct models of the following Co2+ coordination
complex ions:
(a) [CoClM3]+: one Cl- and three methanols tetrahedrally arranged around a central
Co2+
(b) [CoClM5]+: one Cl- and five methanols octahedrally arranged around a central Co2+
Compare the two structures. Identify and record which structure has too much open space
between the methanols to be stable.
6. Turn on the spectrophotometer, and adjust the wavelength control to 657 nm. Allow the
spectrophotometer to stabilize while you do Steps 7–14.
7. Attach a clean, dry 50-mL buret to a support stand using a buret clamp. Rinse the buret
with three 5-mL portions of propan-2-ol. Collect the rinses in the “Discarded Solutions”
container.
8. Fill the buret to the 0.00-mL mark with propan-2-ol. Use tape to label the base of the
support stand “P”.
9. Weigh between 0.16 and 0.20 g of CoCl2 ∙ 6H2O. Transfer the CoCl2 ∙ 6H2O to a clean, dry
250-mL Erlenmeyer flask. Record your exact mass of CoCl2 ∙ 6H2O.
10. Dispense exactly 50.00 mL of propan-2-ol from buret P into the Erlenmeyer flask. Swirl
the flask to dissolve the solid. Make sure that the entire solid has dissolved before going
on to the next step. Use tape to label the flask “stock CoCl2 solution”.
11. Attach another clean, dry 50-mL buret to the same support stand (on the other side of the
buret clamp). Rinse this buret with three 5-mL portions of your CoCl2 stock solution. Pour
the rinses into your “Discarded Solutions” container. Pour all of the remaining CoCl2 stock
solution into the buret. The stock solution will only fill the buret to about the 25-mL mark.
Use tape to label the base of the support stand “Co”.
12. Refill buret P with 2-propanol to the 0.00-mL mark.
13. Label four clean, dry 3-oz plastic cups “1”, “2”, “3”, and “4”. The cups must be free of all
traces of water.
100
14. Prepare various mixtures of your CoCl2 stock solution and propan-2-ol in the cups by
dispensing the following volumes from buret Co and buret P. Record the exact volumes
dispensed from each buret. Swirl the mixtures. Each cup now contains a dilution of the
CoCl2 stock solution in propan-2-ol. Record the colors of the four solutions.
Cup
1
2
3
4
Volume from Buret Co, Volume from Buret P, Ml
0.50
9.50
1.00
9.00
1.50
8.50
2.00
8.00
15. Press the A/T/C button until the display shows “A.”
16. Rinse a clean spectrophotometer cuvet twice, using about 1 mL of propan-2-ol from buret
P each time. Pour the rinses into your “Discarded Solutions” container. Fill the cuvet with
propa-2-nol from buret P, wipe the outside of the cuvet with lint-free tissue, and place the
cuvet in the spectrophotometer’s sample holder. Always position the grooved sides of the
cuvet in the spectrophotometer in the same orientation, facing the sides, for this and all
subsequent analyses. Close the sample holder cover.
17. Press the 100%T/0A button to zero the spectrometer. Remove the cuvet from the sample
holder, and pour the propan-2-ol from the cuvet into your “Discarded Solutions”
container.
18. Rinse the cuvet twice, using about 1 mL of the solution from cup 1 each time. Pour the
rinses into the “Discarded Solutions” container. Fill the cuvet with the solution from cup
1. Wipe the outside of the cuvet with lint-free tissue, and place the cuvet in the
spectrophotometer’s sample holder.
Close the sample holder cover. Record the
absorbance of this solution. Remove the cuvet, and pour its contents into your “Discarded
Solutions” container.
19. Repeat Step 18 using the solutions in cups 2, 3, and 4. Remember to rinse the cuvet twice,
using 1-mL portions of the solution each time before filling the cuvet and determining the
solution’s absorbance. Record the absorbance of each solution.
20. Attach a third clean, dry 50-mL buret to a second support stand, using another buret
clamp. Rinse the buret with three 5-mL portions of methanol. Pour the rinses into your
“Discarded Solutions” container. Fill the buret to the 35-mL mark with methanol. Use
tape to label the base of the support stand “M”.
21. Refill buret P to the 0.00-mL mark with propan-2-ol.
22. Label six clean, dry, 3-oz plastic cups as “A”, “B”, “C”, “D”, “E”, and “F”. The cups must be
free of all traces of water.
101
23. Prepare six solutions as prescribed in the following table, using the liquids in your three
burets. Swirl the cups to ensure complete solution mixing.
Cup
A
B
C
D
E
F
Volume from Buret Co,
Ml
1.00
1.00
1.00
1.00
1.00
1.00
Volume from Buret M,
mL
1.00
1.50
2.00
2.50
3.00
2.00
Volume from Buret P,
mL
8.00
7.50
7.00
6.50
6.00
7.00
Record the exact volumes dispensed from each buret. Note that the contents of cup F are
the same as those of cup C.
24. Prepare an ice-water bath by adding about 25 mL of water to about 100 mL of ice in a 250mL beaker. Position cup F in the ice-water bath in a way that will prevent it from tipping
over.
25. Repeat Steps 15–17 to recalibrate the spectrophotometer at 0 %T and 100 %T.
26. Measure the temperatures of the solutions in cups A–E by inserting a thermometer in each
and allowing 1 min for equilibration. Record the temperatures.
27. Rinse the cuvet with two 1-mL portions of the solution in cup A. Pour the rinses into your
“Discarded Solutions” container. Fill the cuvet with the solution in cup A, place the cuvet
in the spectrophotometer’s sample holder, and measure and record the solution’s
absorbance. Also record the solution color. Remove the cuvet, and pour its contents into
your “Discarded Solutions’’ container.
28. Repeat Step 27 for the solutions in cups B–E.
29. Remove cup F from the ice-water bath, insert a thermometer into the solution, and record
the solution temperature. Repeat Step 27 for the solution in cup F.
30. Remove the cuvet from the spectrophotometer. Pour the contents of the cuvet, the three
burets, and the ten cups into your “Discarded Solutions” container. Discard the material
in this container as directed by your laboratory instructor. Wash all glassware with
detergent. Discard the plastic cups and tissue as directed by your laboratory instructor.
Turn off the spectrophotometer.
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QUESTIONS FOR EQUILIBRIUM BETWEEN TWO…
NAME _________________
1. The equation in the introduction is one way to represent the equilibrium between Co2+
complex ions in mixtures of propan-2-ol and methanol.
[CoClP3]+ + 3 M  [CoClP2M3]+ + P
(a) For the five solutions of CoCl2, propan-2-ol, and methanol at approximately the same
temperature (cups A–E):
(i) Did you find that Keq is indeed constant for these solutions? Briefly explain, stating
your criteria for constancy.
(ii) Describe the effect on [Co(tet)] of increasing the proportion of methanol in the
solutions. Is this effect consistent with Le Châtelier’s principle? Briefly explain.
(b) For the two solutions with the same composition (cups C and F):
(i) Describe the effect of increasing temperature on [Co(tet)].
(ii) Describe the effect of increasing temperature on Keq.
(iii) Determine the sign of H for the conversion of tetrahedrally coordinated Co2+ to
the octahedral form. Briefly explain.
(iv) The van’t Hoff equation uses the temperature dependence of the equilibrium
constant to calculate the value of the enthalpy change, H, for a reaction. One form of
this equation is,
 K  ∆H   1 1 
ln  1 
=
 − 
R  T2 T1 
 K2 
where R is the ideal gas constant (8.314 J/mol ∙ K) and T is the temperature, in Kelvin.
Use this equation and your experimental data to calculate H for the conversion of
tetrahedrally coordinated Co2+ to the octahedral form. Compare your answer to the
reported value for H for this conversion in the introduction.
103
2. An alternative equilibrium that could be used to explain the conversion of tetrahedrally
coordinated Co2+ in propan-2-ol to the octahedral form upon addition of methanol is
represented by the following equation:
[CoClP3]+ + 5M  [CoClM5]+ + 3P
In this equilibrium, five methanol molecules replace the three propan-2-ol molecules in
the complex ion.
(a) Write an expression for Keq for this equilibrium.
(b) Use your experimental data and the expression for Keq from (a) to calculate the values
of Keq for the solutions in cups A, C, and E.
(c) Based on your answers to (b), confirm or reject this alternative equilibrium. Briefly
explain.
104
SYNTHESIS AND ANALYSIS OF A NICKEL COMPLEX
Introduction
Two important tasks many chemists perform are the synthesis and analysis of compounds.
Synthesis involves not only preparing a compound, but also maximizing the yield of pure
product. After isolating the product, the chemist must analyze it to ascertain its chemical
composition (formula). Both tasks require good lab technique and close attention to what
might seem minor procedural details. Therefore, a technically skilled chemist with a good
understanding of the purposes of each step in both the synthesis and analysis procedures will
get the most accurate results. This experiment involves the preparation of a coordination
compound containing nickel(II) ion (Ni2+), ammonia (NH3), and chloride ion (Cl−) and the
determination of its empirical formula. Until the formula has been determined, it will be
represented as [NiClx(NH3)n]Cl2-x, with n representing a small whole number and x is an
integer from 0 to 2 and n + x = 6.
Synthesizing [NiCIx(NH3)n]Cl2-x
You will synthesize [NiClx(NH3)n]Clx−2 by reacting nickel(II) chloride hexahydrate (NiCl2
6H2O) and NH3. This reaction is shown below
Ni2+ (green) + 2 Cl− + n NH3  [NiCIx(NH3)n]Cl2−x (bluish purple)
(Eq. 1)
A complication arises because NH3 in aqueous solution is involved in the equilibrium reaction
shown as follows:
NH3 + H2O  NH4+ + OH−
(Eq.2)
Although the equilibrium constant for Equation 2 is small (1.75 x 10-5) some of the Ni2+ ion
will react with hydroxide ion (OH−) to form nickel(II) hydroxide, Ni(OH)2, as shown below:
Ni2+ (green) + 2 OH−  Ni(OH)2 (green)
(Eq. 3)
To the extent that the reaction in Equation 3 occurs, the product formed in Equation 1 will be
impure, and the synthesis reaction yield will therefore decrease.
Water is a convenient
solvent for the synthesis reaction because the reactants are water soluble. However, because
[NiCIx(NH3)n]Cl2-x, which is bluish-purple, is also somewhat soluble in water, you must keep
the volume of water used in the synthesis to an absolute minimum. Nickel(II) chloride
hexahydrate is more soluble in hot water than in cold water, so heating the reactants will
enable you to dissolve more of this compound in a smaller volume of water. Unfortunately,
the water solubility of NH3 is greatly decreased with increasing temperature. In this case, at
temperatures approaching 100°C, NH3 volatilizes before it can react with the NiCl2. By
maintaining the reaction temperature at 60°C, you will maximize the [NiClx(NH3)n]Cl2−x yield.
Once the reaction is complete, you will cool the reaction mixture to 0°C in an ice-water bath.
Because [NiClx(NH3)n]Cl2-x is less soluble in cold water than in hot water, this step decreases
the solubility of the product in Equation 1. You will add cold ethanol to the cold reaction
mixture to further reduce the product’s solubility, because [NiCIx(NH3)n]Cl2-x is insoluble in
ethanol. You will filter the [NiClx(NH3)n]Cl2−x crystals from the cold ethanolic solution and
wash them with cold concentrated NH3. This treatment should help to convert any Ni(OH)2
in the sample to [NiClx(NH3)n]Cl2−x, as shown in Equation 4.
Ni(OH)2 + n NH3 + 2 Cl-  [NiCIx(NH3)n]Cl2-x + 2 OH-
(Eq.4)
105
Finally, you will dry and weigh the crystals to determine the actual yield of your synthesis.
Analyzing [NiClx(NH3)n]Cl2-x
Determining the Molar Mass of the Compound, Mass % Ni2+ Ion
The [Ni(NH3)n]2+ ion absorbs light in the visible region of the spectrum. You will take
advantage of this property in order to determine the molar mass of the [NiCl2(NH3)n].
Solutions containing [Ni(NH3)n]2+ ion are colored. The observed color is produced by those
visible wavelengths that are not being absorbed. You can determine which wavelengths are
absorbed by using a spectrophotometer to measure the absorbance of the solution throughout
the visible region of the spectrum. The wavelength at which the species absorbs the most light
is called the analytical wavelength (λmax) for that species. Absorbance is directly proportional
to the concentration of the absorbing species in solution. This relationship, known as Beer's
law, is represented by Equation 8. A is absorbance, ε is molar absorptivity, b is the length of
the light path through the solution, and c is the molar concentration of the absorbing species.
A = εbc
(Eq. 5)
Molar absorptivity is a proportionality constant relating absorbance and molar concentration
of the absorbing species at the wavelength being measured. The value of ε varies with
wavelength, reaching a maximum at the analytical wavelength.
You will prepare a standard [Ni(NH3)n]2+ ion solution by dissolving a known mass of nickel(II)
sulfate hexahydrate (NiSO4  6H2O) in water and adding excess concentrated NH3. Then you
will dilute the mixture with water to a known volume. The Ni2+ ion in the sample converts to
[Ni(NH3)n]2+ ion. You will use your standard [Ni(NH3)n]2+ ion solution to establish the
analytical wavelength (λmax) for [Ni(NH3)n]2+ ion, the wavelength where the complex has the
maximum absorbance. Then, you will compare the absorbance of your standard [Ni(NH3)n]2+
ion solution with that of a solution you will prepare from a known mass of the
[NiClx(NH3)n]Cl2−x you synthesized. Because the absorbing species, the [Ni(NH3)n]2+ ion, is
identical in both solutions, ε is the same for both solutions. This will allow you to determine
the concentration of the Ni(NH3)n2+ and therefore the molar mass of the synthesized complex.
This will allow you to determine the empirical formula.
PROCEDURE
DAY ONE – SYNTHESIZING [NiClx(NH3)n]Cl2−x
Prepare a warm-water bath. Half fill a 600-mL or larger beaker with tap water and place the
beaker on a hot plate. Monitor the water temperature with a thermometer. Suspend the
thermometer into the beaker making sure that the thermometer extends into the water but
does not touch the side or bottom of the beaker. Heat the beaker and its contents until the
water temperature reaches 50°C. Adjust the hot plate setting so that the water temperature
remains between 50 and 60°C.
106
On a tared piece of weighing paper, weigh 8.0 g of NiCl2  6H2O. Record the mass of the solid
on the Data Sheet for day 1. Transfer the solid to a clean 125-mL Erlenmeyer flask. Add 10 mL
of distilled or deionized water to the solid in the flask. Place the flask in the 60°C water bath
and clamp the flask in position. Stir the mixture in the flask with a clean glass stirring rod until
it has all dissolved.
Loosen the clamp on the ring stand, and while holding the end of the clamp, remove the flask
from the bath. Attach the clamp to another ring stand, and let the flask and contents cool in
air for 1-2 min.
Slowly, with stirring, in the fume hood, add 25 mL of concentrated NH3 solution to the NiCl2
 6H2O solution in the flask.
Cover the top of the flask with a wet paper towel and suspend the flask in the warm-water bath
by clamping it to the ring stand. Make sure the water temperature is between 50°C and 60°C.
Leave the flask in the bath for 15 min. During this time, periodically swirl the mixture.
Prepare an ice-water bath in another 600-mL beaker by adding 150 mL of water and several
pieces of ice to the beaker. Transfer 20 mL of concentrated NH3 solution into a labeled, 18 x
150-mm test tube. Stopper the test tube with a No. 2 solid rubber stopper. Place the test tube
in the ice-water bath. Obtain 60 mL of 95% ethanol in a labeled 100-mL beaker. Place the
beaker and its contents in the ice- water bath.
Assemble a second ice-water bath in a third 600-mL beaker. After the reaction in the warmwater bath has proceeded for 15 min, unclamp the flask from the ring stand. Carefully clamp
the reaction flask on another ring stand so that the flask is suspended in the second ice-water
bath. Remove the damp paper towel covering the mouth of the flask. While holding the flask,
loosen the clamp and swirl the reaction mixture for 5 min while it is cooling. Add 10 mL of
ice-cold 95% ethanol to the flask and stir. Remove the reaction flask from the ice-water bath.
Wipe any water off the bottom of the flask using a paper towel.
Place a Büchner funnel with an adapter in a 250 mL filter flask. Place a circle of filter paper in
the funnel. Wet the paper with a small amount of distilled water and turn on the vacuum,
making sure the paper is firmly seated. Slowly pour the liquid-solid mixture from the flask
into the Büchner funnel, as follows. Pour the mixture smoothly at such a rate that it passes
through the filter quickly, but does not cause a build-up of liquid in the funnel.
Rinse any remaining solid down the inside wall of the reaction flask using 5 mL of ice-cold,
concentrated NH3 solution from your test tube. Swirl the solid and rinse solution mixture,
and quickly pour the mixture into the Büchner funnel.
In the same manner, rinse any remaining solid from the flask using two additional 5-mL
portions of the cold, concentrated NH3 solution.
Dry the solid by drawing air through the solid for 3-5 min. Break up the solid with a spatula,
being careful not to tear the filter paper. Pour 15 mL of cold 95% ethanol over the solid.
Repeat the ethanol washing two more times, using 15 mL of cold 95% ethanol each time. Make
sure to carefully break up the solid before each ethanol wash.
Pour 15 mL of acetone over the solid in the funnel. Break up the solid with a spatula to expose
all its surfaces to the acetone. Draw air through the solid for 10-15 min. Turn off the vacuum.
107
Determine the mass of a labeled beaker and record this mass on the Data Sheet for day 1. After
the solid is completely dry, add it to the beaker. Determine the mass of the beaker and solid
and record this mass on your Data Sheet. Cover your beaker with a watch glass and store it in
your locker.
Transfer the solution in the filter flask into the container provided that is labeled "Discarded
NiCl2/Ethanol/Acetone Solution Mixture." Rinse the filter flask and Büchner funnel once with
20 mL of tap water. Transfer the rinse to the same discard container.
DAY TWO – DETERMINING THE MOLAR MASS OF [NiClx(NH3)n]Cl2−x BY
SPECTROPHOTOMETRY
On the frosted or white circle on a clean, dry 100-mL beaker, write "std" to indicate the NiSO4
 6H2O standard sample solution. Write "unk" on a second clean, dry 100-mL beaker, which
you will use for your [NiClx(NH3)n]Cl2−x sample solution with unknown %Ni2+ ion. Clean two
50-mL volumetric flasks and stoppers. Label one flask "std" and the other flask "unk."
Using an analytical balance, weigh on tared piece of weighing paper approximately 0.30 g
NiSO4  6H2O. Transfer the sample to the "std" beaker. Record the mass of NiSO4  6H2O to
the nearest 0.1 mg on the Data Sheet for day 2. Using an analytical balance, weigh on a tared
piece of weighing paper approximately 0.35 g sample of your [NiClx(NH3)n]Cl2−x. Transfer the
sample to the "unk" beaker. Record the mass of [NiClx(NH3)n]Cl2−x to the nearest 0.1 mg on
the Data Sheet for day 2 (near the bottom of the page).
Using a graduated cylinder, add 20 mL of distilled water to both samples. Note and record the
color and appearance of each solution on the Data Sheet for day 2. Using separate glass stirring
rods, stir each mixture until most of the solid has dissolved. Leave the rods in the beakers to
avoid losing any solution adhering to the rods.
Measure out 10 mL of concentrated NH3 solution in a 10-mL graduated cylinder, which need
not be dry. Add the NH3 solution to the solution in the "std" beaker. Then measure another
10 mL of concentrated NH3 solution, and add it to the solution in the "unk" beaker. Stir each
mixture until no solid remains. Note and record the color and appearance of each solution on
The Data Sheet for day 2.
Using a short-stem funnel, transfer the "std" solution into the "std" 50-mL volumetric flask.
Rinse the beaker and rod with a minimum amount of distilled water from a wash bottle, and
pour the rinses into the "std" volumetric flask. Rinse the beaker two more times, using distilled
water, but do not allow the volume of solution in the volumetric flask to exceed 50 mL.
Add distilled water to the solution in the flask until the solution level coincides with the
junction of the neck and body of the flask. Stopper the flask. Firmly holding the stopper in
place, invert the flask 10 times to thoroughly mix the solution. Then fill the flask exactly to
the etched mark, 50.00 mL, by adding distilled water from a disposable pipet or medicine
dropper. Stopper the flask. Thoroughly mix the solution by inverting the flask at least 25
times, while holding the stopper firmly in place.
Follow the same procedure to transfer your "unk" solution to the "unk" 50-mL volumetric
flask. Dilute the "unk" solution to the 50.00 mL mark, stopper the flask, and thoroughly mix.
Obtain two spectrophotometer cuvets, and place them in a dry beaker or test tube rack. Clean
108
the cuvets and rinse them distilled water. Fill one cuvet about three-quarters full with distilled
water. Set the wavelength on the spectrophotometer to 620 nm. Place the cuvet in the holder
in the spectrophotometer. Press the 100%T/0A button to zero the spectrophotometer. Press
the A/T/C button until the display shows an “A”.
Rinse the other cuvet three times, using a 1-2 mL portion of your "std" solution each time.
Dispose of the rinses into the container provided that is labeled "Discarded NiCl2/NiSO4/NH3
Solutions." Fill this second cuvet about three-quarters full with "std" solution. Place the cuvet
in the holder in the spectrophotometer and read the absorbance. Record this in your Data
Sheet.
Adjust the wavelength to 600 nm. Following the above procedure, first with the reference
(H2O-filled) cuvet and then the sample cuvet obtain and record the absorbance for the
standard solution at 600 nm. Repeat the procedure at wavelength intervals of 20 nm down
to 540 nm. From among your five absorbance readings, determine the approximate λmax for
the [Ni(NH3)n]2+ ion and record it on The Data Sheet for day 2. To more precisely establish
λmax for the [Ni(NH3)n]2+ ion, measure the absorbance of the standard solution at wavelengths
10 nm less and 10 nm greater than the wavelength you estimated as λmax. Record these
additional absorbance measurements on The Data Sheet for day 2. Select the λmax for the
[Ni(NH3)n]2+ ion and record it on the Data Sheet for day 2.
Empty the sample cuvette into the "Discarded NiCl2/NiSO4/NH3 Solutions" container. Rinse
the cuvette with distilled water, and then rinse it three times with your "unk" solution, using
1 mL of solution each time. Transfer all rinses to the discard container. Fill the cuvette about
three-quarters full with "unk" solution. Set the spectrophotometer at the λmax you determined
earlier. Check the 0%T setting with the cuvette compartment empty and its cover closed.
Check the 0 Abs setting using the water-filled reference cuvette. Determine the absorbance of
the unknown solution at λmax. Record this absorbance on The Data Sheet for day 2. Transfer
the solutions in your cuvettes to the discard container. Rinse and wash the cuvettes and add
any rinses and washings to the discard container.
Transfer the solutions in your volumetric flasks into the appropriate discard container. Rinse
the volumetric flasks twice with 10 mL of tap water each time and twice with 10 mL of distilled
water each time. Transfer the rinses into the appropriate discard container. Allow the flasks
to drain. Empty the cuvettes appropriately and allow them to dry.
Calculate the molar mass of the synthesized compound using the mass of the unknown,
[NiClx(NH3)n]Cl2−x, dissolved in the 50.0 mL solution and the solution’s molarity obtained
from its absorbance.
Recall that there must be 1 mole of Ni2+ ions and 2 moles of Cl- ions per mole of the synthesized
compound and that any remaining mass in the compound must be due to the ammonia
ligands. Calculate the number of moles of ammonia equivalent to the “remaining mass” and
round to the nearest whole number to arrive at the empirical formula of the synthesized
compound.
109
QUESTIONS FOR SYNTHESIS & ANALYSIS …
NAME _______________________
1. Based on your determined empirical formula for [NiClx(NH3)n]Cl2−x, what is the
coordination number of the nickel(II) ion?
2. Based on the electron configuration of the nickel(II) ion and the coordination number
stated above, what is the hybridization used by the Ni2+ ion (i.e.,sp3, dsp2, d2sp3 or sp3d2)?
Why? Explain your answer based on the electron configuration coordination number and
what you know about the different types of hybridization.
3. Do you expect the [NiClx(NH3)n]Cl2−x complex to be paramagnetic or diamagnetic? Use the
answers to questions 1 and 2 to support your answer.
110
MOLECULAR MODELS OF TRANSITION METAL
COMPLEXES
INTRODUCTION
The chemical and physical properties of a substance are influenced by the distribution of outer
shell (valence) electrons and the three-dimensional arrangement of its nuclei. A variety of
experimental methods are employed to map out the relative positions of the nuclei in a
molecule or an ion. We will be using molecular models to determine some of the properties
of transition metal complexes.
All transition metal complexes in this lab have a transition metal ion as the central atom. You
will determine the distribution of electrons and bonded atoms about a central atom. In so
doing, you will be able to determine the probable hybridization of the central atom, electron
pair and molecular geometries, and polarity of the species in question. We can also determine
if the species is optically active by looking at the presence or lack of a superimposable mirror
image of the compound or ion.
PROCEDURE
All of the models we will be constructing will use the gray 14 sided polyhedrons (with a hole
in each face) representing the transition metal as the central atom. Only the square sides will
be used. In order to simplify things, we will be omitting the hydrogen atoms on
ethylenediamine. Carbon atoms are the black polyhedra with 4 holes, nitrogen atoms are the
blue polyhedra with 4 holes, oxygen atoms are the red polyhedra with 4 holes and chlorine
atoms are the green polyhedra with 4 holes.
A.
Square Planar Complexes
H
H
N
H
1. Construct four ammonia molecules:
the “extra” bond position on the nitrogen.
These will attach to the metal atom through
2. Construct a model of cis-diamminedichloridoplatinum(II).
3. Construct a model of trans-diamminedichloridoplatinum (II).
B. Tetrahedral Complexes.
1. Construct a water molecule. This will attach to the metal atom through one of the “extra”
bond positions on the oxygen.
2. Construct a model of ammineaquabromidochloridoiron(II).
3. Using the model from step 1 as guide construct the mirror image of the model.
4. Rotate one of the models to try to make it match exactly (superimpose) with the other
model.
111
5. Replace the aqua ligand with another ammine ligand in both of the models.
6. Rotate one of the models to try to make it match exactly (superimpose) with the other
model.
C. Octahedral complexes.
C
N
C
N
1. Construct 6 ethylenediamine molecules:
(we are not showing the hydrogen
atoms here). These will attach to the metal atom through the nitrogen atom’s lone pairs.
2. Construct the model for the tris(ethylenediamine)cobalt(II) ion.
3. Using the model from step 2 as a guide, construct the mirror image of the model in step 2.
4. Rotate the model from step 3 to try to make it match up exactly (superimpose) with the
model from step 2.
5. Replace one of the ethylenediamines in each model with 2 chloride ions. The two chloride
ions should be in adjacent positions around the central metal atom (cis- conformation).
6. Rotate one of the modified models to try to make it match up exactly (superimpose) with
the other one.
7. Replace one more of the ethylenediamines with 2 ammonia molecules in each of the
models.
a. While the chlorides are in the cis- position, examine the structures to see if the models
match up exactly (superimpose).
8. In each model assembled in step 7 swap the positions of two of the appropriate ligands so
that the chloride ions are on opposite sides of the metal atom (trans- conformation).
9. Rotate one of the modified models to try to make it match up exactly (superimpose) with
the other one.
112
REPORT
MOLECULAR MODELS OF TRANSITION…
NAME _______________________
SECTION ____________________
A. Refer to the models that you constructed for this portion of the lab.
Which conformation (cis- or trans-) should be polar?
________________
Which conformation should be soluble in water?
________________
B. Refer to the models that you constructed for this portion of the lab.
Did the mirror images of the models of ammineaquabromidochloridoiron(II) match
up exactly (superimpose) with one another?
Did the mirror images of the models of diamminebromidochloridoiron(II) match up
exactly (superimpose) with one another?
Which of these compounds exists as stereoisomers?
C. Refer to the models that you constructed for this portion of the lab.
Did the mirror images of the models of the tris(ethylenediamine)cobalt(II) ion match
up exactly (superimpose) with one another?
Does this complex ion exist as enantiomers?
Did the mirror images of the models of the cisdichloridobis(ethylenediamine)cobalt(II) match up exactly (superimpose) with one
another?
Does this compound exist as enantiomers?
Did the mirror images of the models of either cis- or trans-diamminedichlorido
(ethylenediamine)cobalt(II) match up exactly (superimpose) with one another?
Which of these compounds exist as enantiomers if any?
113
114
CHECK OUT INSTRUCTIONS
1. Clean!
•
•
•
West Bench – Benchtop, West Fume Hoods, Balance Room
Middle Bench – Benchtop, Rear Balance Area, Rear Fume Hoods
East Bench – Benchtop, East Fume Hoods, Rear Sink Area
2. Return Locks!
•
•
•
•
Obtain a white tag (if you no longer have the original)
Write your combination on the tag
Lock it to your lock
Place the lock on the middle bench in the front
3. Equipment Check
•
•
•
•
•
Ensure your drawer has a complete set of equipment
Remove extra items
CLEAN AND RETURN SHELL VIALS
Obtain missing items
When EVERYONE is ready the stockroom will verify that
your drawer is complete
THANK YOU & ENJOY!
¦|
115
Quantity
2
2
2
2
2
1
1
1
1
1
1
1
3
1
1
1
1
1
2
1
1
2
4
10
10
1
1
1
1
1
Description
Beaker, 20 mL
Beaker, 50 mL
Beaker, 100 mL
Beaker, 150 mL
Beaker, 250 mL
Beaker, 400 mL
Beaker, 600 mL
Bottle, 500 mL, Screw Cap
Bulb, Pipet
Clamp, Buret
Cylinder, Graduated, 10 mL
Cylinder, Graduated, 50 mL
Flask, Erlenmeyer, 250 mL
Forceps
Funnel, Small, 45 mm
Holder, Test Tube
Microspatula
Pipet, Graduated 1.0 mL
Pipet, Graduated 5.0 mL or 10.0 mL
Pipet, Volumetric 10 mL
Pipet, Volumetric 25 mL
Rack, Test tube
Shell Vials
Test Tube, 10 mm x 75 mm
Test Tube, 13 mm x 100 mm
Test Tube Brush, 12 mm x 62 mm
Thermometer, -20°C to 110°C
Tongs
Wash Bottle, Polyethylene, 250 mL
Watch Glass, 75 mm
116
APPENDIX A
Calculations Involving Precision and Accuracy
Precision
Precision is a measure of how well multiple (repeated) measurements agree with each other.
It is an indication of consistency. One method of evaluating the precision of a set of data is to
determine the percent relative average deviation. The procedure for this calculation is as
follows:
1. Determine the average value for at least three experimental trials.
2. Subtract each individual value from the average value to get the deviation for each trial.
3. Add together the absolute values of the deviations and divide by the number of trials and
the average value to get the relative average deviation.
4. Multiply by 100 to get percent relative average deviation.
n
Percent Relative Average Deviation =
average deviation
× 100 =
average value
∑x
i =1
i
−x
nx
× 100
Accuracy
Accuracy is a measure of how close an experimental value (usually an average value) is to the
accepted value (also called the "true" value). One method of evaluating the accuracy of an
experimental result is to determine the percent error as follows:
Percent Error =
experimental value - accepted value
× 100
accepted value
Do not use absolute values when calculating accuracy. The sign simply indicates that the
experimental value is higher than the accepted value when the percent error is a positive
number, lower if negative.
117
APPENDIX B
GRAPHS
INTRODUCTION
Relationships between experimental quantities are often represented in the form of graphs.
Straight line graphs are easier to construct and to interpret than curved ones. Data that
initially result in a curve when graphed are sometimes mathematically rearranged to result in
a straight line relationship. This can often be accomplished by taking the logarithms of the
values for one or both of the quantities that were being plotted and then graphing these new
log values. When data that has been graphed forms a straight line plot, the mathematical
relationship between the quantities can be determined from the equation for a line:
=
y mx + b
PROCEDURE
A. Construction of a graph
A number of rules must be followed when constructing graphs. Your score for this exercise
will depend upon how well you follow these rules.
1. Select a good quality graph paper that is easy to use with the metric scale. Graph paper
that has divisions marked in blocks with different shades of lines is easier to use (less
counting) than paper that has uniform shading. Choose paper that is divided into five by
five or ten by ten small squares within a larger grid. Avoid paper in which the large squares
are divided into four by four or eight by eight blocks (this type of graph paper is for drafting
classes that use English system units).
2. It is customary to plot the quantity that is varied (the independent variable) on the x
(horizontal) axis and the quantity that is measured (the dependent variable) on the y
(vertical) axis. In mathematical terms, the quantity on the y-axis is a function of the
quantity on the x-axis.
3. Use a scale for each axis that will spread the data points to be plotted over the full page (or
over the space assigned). Do not crowd the data into one corner. However, your scale
should result in convenient units (such as 10, 20, 30, etc. or 2, 4, 6, 8, etc.) for each major
division on the graph. A compromise may be necessary.
4. Use a constant scale (the same number of divisions/unit) along each axis. However,
because different quantities are plotted on each axis you would not necessarily expect the
scale on the x and y axes to be the same.
5. It is only necessary to mark (and label) the intervals at 4 to 6 places along an axis (more
than that gets cluttered). For example, if you had mass readings ranging from 7 to 68 g,
you might mark and label the axis at 0, 20, 40, 60, and 80 g. Do NOT mark your axes at
the data points. The coordinates for the data plotted on the graph should be presented in
a table on an unused section of the graph paper (away from the data points) or on a
separate piece of paper.
6. The precision in the labels for the axes intervals should reflect the precision in the data
being plotted. For example, if masses were determined to one place after the decimal
(such as 9.1 g, 15.4 g, etc.) the intervals on the graph should be labeled 0.0, 20.0, 40.0 and
so on.
118
Note: The precision for measurements plotted on the y-axis may differ from those for the x
axis.
7. If you do not have any data close to a zero value, you need not place “zero” in the lower
left-hand corner of the graph. The graph origin can begin at any convenient value
(provided it is labeled). However, if the graph is to be assessed to determine a “straightline” relationship between data, and you wish to read the y-intercept directly from the
graph, then you must use intervals and plot the data so that the y-intercept is NOT off the
graph.
8. Label each axis with the appropriate label.
9. Title each graph. The title should reflect what quantities are being plotted. The title might
simply be an equation that has been provided or it might be the description of
experimental quantities.
10. After the data have been plotted, draw either a straight line or a smooth curve that best
represents the data points. Do NOT connect the dots with individual straight lines. When
data being plotted has been experimentally obtained, you should not expect the line to
pass directly through every data point due to experimental errors. Construct a “best-fit”
plot in which the points that do not fall on the line are randomly scattered. The sum of the
distances between the line and the points above it should be the same as the sum of the
distances between the line and the points below it. In addition, the line should be drawn
so that these distances are minimized.
B. Determination of a Mathematical Relationship from a Straight Line Graph.
The straight line relationship between quantities x and y can be represented by:
=
y mx + b
where y (the quantity plotted on the vertical axis) is a function of x (the quantity plotted on
the horizontal axis). The “m” is the slope of the line and “b” is called the y-intercept. Linear
regression analysis and substitution can be used to obtain the exact value for the slope and yintercept, but in this exercise these values will be estimated by reading them directly from the
graph.
1. Graph the data and draw a “best-fit” straight line (see Part A of the Procedure).
2. Determine the slope of the line. Choose two points on the line (not necessarily data points)
that can be read accurately. To maximize precision, these two points should be fairly far
apart. Read the coordinate values for each point. Point number one is the data point
having an x value closest to the origin and the values for point one will be (x1, yl). The other
point will have values of (x2, y2). The slope of the line is:
m=
y 2 − y1
x 2 − x1
119
3. The sign of the slope can be negative (indicating an inverse relationship between the
quantities x and y). Note that the number of significant figures for the slope will be
artificially reduced if the points on the line selected for slope determination are too close
together. Be sure to include units (unit for y/unit for x) with the value for the slope.
4. To determine the y-intercept value from the graph, extrapolate (extend) the line until it
reaches the y-axis (x = 0) and read the value for y at that point (include units).
5. Write the mathematical relationship for the quantities that have been graphed. Into the
equation:
=
y mx + b
substitute (each with its appropriate unit):
for y – the quantity (what is being graphed) on the y axis
for m – the value (number) for the slope
for x – the quantity (what is being graphed) on the x axis
for b – the value (number) for the y-intercept
For example: distance(m) = time(s)5.26 m/s + 6.35 m
120
Constructing Scientific Graphs in Excel™
Excel™ makes constructing scientific graphs easy. You can plot points and find the line of
best fit (linear or otherwise) and the equation for that line. You can also import scientific
data (such as that from the Vernier™ LabQuest) and visually represent the data. Excel’s
“Scatter Plot” chart function allows you to do both.
Input the data
One version of the scatter chart is used when you only have a few data points (5 to 15 or so).
You will enter the independent variable (x-axis) in the first column, “A,” and the dependent
variable (y-axis) in the second column, “B.” You can also add descriptions of the data in the
first row if you like. For example:
Once the data has been entered you can use the cursor to select all of the data, including the
header row. Then you will click on the “Insert” tab.
On the “Insert” tab you want to choose the icon under the “Charts” section that indicates a
scatter chart “Insert Scatter (X, Y) or Bubble Chart” and click on it.
After clicking the icon you will have a choice of what kind of scatter plot you want to make.
Click on the first one which just shows the points.
121
With this data set you will have something that looks like this.
y data
Double-click here to
change the title.
120
100
80
60
40
20
0
0
2
4
6
8
10
12
Presently, this is not that useful. We need to have the graph properly formatted with the
axes labelled, the correct precision, and a descriptive title. We should also have more grid
lines. In Excel 2013 clicking on the Scatter Chart icon will then bring up a two new toolbars
at the top. One for the Design and one for the Format of the chart. We are mainly interested
in the Design toolbar. The first part of the toolbar is labelled “Add Chart Element.” In Excel
2010, three new toolbars are created, “Design,” “Layout,” and “Format.” The Chart
Elements are in the “Layout” toolbar in Excel 2010. With this we can add our axes labels
and the minor gridlines and format them as we need.
Click on “Add Chart Element” then click on “Axis Title” then on “Primary Horizontal.” In
the formula bar just under the toolbar you can then put in the title for that axis. For
instance, time in minutes (“Time (min)”) and press ENTER. We can do the same for the yaxis by clicking through to “Primary Vertical” and put in the label in the box. For instance,
the distance in kilometers (“Distance (km)”) and press ENTER. If you double click on the
122
chart title, where it currently says “y data,” you can change that to a descriptive title such as
“Plot of distance vs. time for a road trip.” The graph should now look like this:
Plot of distance vs. time for a road trip
Distance (km)
120
100
80
60
40
20
0
0
2
4
6
8
10
12
Time (min)
Going back to the “Add Chart Element” box we can add minor gridlines by clicking on
“Gridlines” then either “Primary Minor Horizontal” or “Primary Minor Vertical.” You will
want to change both of them but you can only change on at a time. Adding these with the
default values gives this:
Plot of distance vs. time for a road trip
120
Distance (km)
100
80
60
40
20
0
0
2
4
6
8
10
12
Time (min)
The values on the axes can also be modified to indicate values with appropriate precision by
double-clicking on the numbers on the axis. On the “Format” panel that opens on the right
click on the icon with 3 vertical bars and the last option is “Number.” Clicking this and
changing the option from “General” to “Number” allows you to specify the number of
decimal places the values on that axis have.
123
We can also make the data
points smaller (the default is
too big). On the right side of
the program is the formatting
options. Clicking on the drop
down arrow where it says
“AXIS OPTIONS” and then
clicking on the “Series ‘y data’”
option (the last one in the list),
allows you to change the size,
type and color of the data
point markers. In the new
panel that comes up click on the icon that looks like a paint can and then click on
“MARKER” then on “MARKER OPTIONS.” Change it to “Built In” and reduce the size. The
smallest you can make it is “2” which works well.
At this point we also want to
add a trend line which will be
the best-fit line for the data.
We can do this with the “Add
Chart Element” option under
“Trendline.” I would suggest
using the “More Trendline
Options.” Here on the right
side under the “Format…,”
click on the icon that looks
like three vertical bars. You
can then select which kind of
trend line you want and click on the option to display the equation on the chart. You can
also choose to set the intercept to 0.0 (or any value that you know it should be). You can
also format the line under the icon that looks like a paint can being poured out. Set the
trend line to be a solid line and the thickness to be thin (0.5 pt). Our finished graph then
looks like:
124
Plot of distance vs. time for a road trip
120
Distance (km)
100
y = 8.2517x
80
60
40
20
0
0
2
4
6
8
10
12
Time (min)
Clicking on a blank area of the chart and pressing “Ctrl-P” will allow you to print the chart.
You will usually want to print only one graph on a page.
Smooth Line Scatter Chart
Inputting the data
In this case we will usually input data
from another source so we won’t we
typing it in by hand. The process for
creating the chart will be essentially the
same as above, we’ll select “Insert
Scatter (X, Y) or Bubble Chart” and
choose the option “Scatter with Smooth
Lines.” Here we have spectrum data
created with the Vernier™ SpectroVis
Module from three different discharge
lamps. There are 645 rows of data!
Obviously, there’s far too much data to
enter (or graph) by hand but Excel will
handle it nicely. We need to select the data we want to plot. Here, we will just plot the first
spectrum out of the three. To select the first spectrum we can just click on the column “A”
label and drag over to the column “B” label. This gives us:
125
As before, we then click on the
“Insert” tab and click on the
“Insert Scatter (X, Y) or
Bubble Chart” option. Then
well click on the option that
shows a smooth line. If we
wanted to plot all three spectra
on the same plot we would
then hold down the Ctrl key
and click on the other two intensity columns.
This will create a chart that looks like this:
Intensity
1.2
1
0.8
0.6
0.4
0.2
0
0
200
400
600
800
1000
Again, we need to do some formatting here to make it useful as a scientific graph. First we
need to set the x-axis correctly because we do not need to show the area from 0 to almost
400 nm and from about 900 nm to 1000 nm. To set the x-axis scale correctly we then click
in the “Format” area on the right on the “Horizontal (Value) Axis.” Click on the icon that
looks like three vertical bars and on “Axis Options.” The first section here is “Bounds.”
126
Change the Minimum to 380 and the maximum to 900.
Next we need to add the axis labels. This is done exactly as we did before. Click on “Add
Chart Elements” then “Axis Titles” then “Primary Horizontal.” Type “Wavelength (nm)” in
the box and press ENTER. Do the same for “Primary Vertical” and type “Intensity” in the
box. This has no units so we don’t have to add anything else. We should also change the
graph title to something more descriptive such as “Plot of Intensity vs. Wavelength for the
Hydrogen lamp.” We can also add in the minor gridlines to make the graph easier to read by
clicking “Add Chart Elements” then “Gridlines” then “Primary Minor Vertical.” Then in the
“Format” section on the right click on “Vertical (Value) Axis” and click on the icon that looks
like 3 bars. Click on “Axis Options” and change the value for the minor units to 0.02 (1/10th
of the major unit). Then do the same thing for the Horizontal axis. We can change the
precision of the labels on the axes in the same way we did before with the “Number” option
at the bottom of the “Format Axis” panel (3 decimals for intensity and 1 for the wavelength).
Finally, again the default value for the line is too thick. Click on “Series ‘Intensity’” and then
on the paint bucket icon and change the line thickness to 0.5 pt. We then have a graph that
looks like:
Plot of intensity of light vs. wavelength for the hydrogen lamp
1.200
1.000
Intensity
0.800
0.600
0.400
0.200
0.000
380.0
480.0
580.0
680.0
780.0
880.0
Wavelength (nm)
At this point, you can click on a blank area of the graph and press Ctrl-P you can print it.
127
APPENDIX C
BALANCING REDOX REACTIONS USING THE HALF-REACTION METHOD
There are many methods that can be used when balancing chemical reactions that involve
oxidation-reduction. The following steps are used in a “half-reaction” method:
1. The initial “skeleton” reaction to be balanced for an oxidation-reduction reaction
occurring in aqueous solution often does not include the H2O, H+(in acid), or OH− (in base)
that will be added later as the reaction is balanced. Sometimes spectator ions are not
included either.
Write the skeleton reaction and assign oxidation numbers to each element.
2. Split the reaction into two half-reactions, one containing the oxidation and one containing
the reduction. (Note: In some reactions, more than one element is oxidized or more than
one is reduced. Sometimes the mole to mole relationships between these elements can be
determined from the formulas of the chemicals involved in the reaction. However, in some
cases, experimental data is needed to help determine the correctly balanced equation.)
3. For each half-reaction, balance of all the elements present except oxygen and hydrogen.
4. Balance oxygen by adding H2O to the side of each half-reaction needing oxygen.
5. The method for balancing hydrogen in each half-reaction depends on whether the reaction
is taking place in acidic or basic solution.
a. in acid, add H+ to the side of the reaction needing more hydrogen.
b. in base, count the number of hydrogen atoms that are needed. Add one H2O for every
hydrogen atom needed to the side with insufficient hydrogen and simultaneously add
the same number of OH− ions to the opposite side
Note: # of H needed = # of H2O added to the side with insufficient H = # of OH− added to
opposite side
6. Balance overall charge by adding electrons (e-) to the more positive side of the halfreaction.
7. Multiply each half-reaction by the factors needed to make the electrons in each halfreaction equal.
8. Add the half-reactions (combining any like terms) and cancel species that appear on both
sides of the equation (electrons must cancel).
9. If needed, divide by the largest common factor to reduce the coefficients to the lowest
whole number ratio.
10. CHECK to make certain that the number of atoms of each element and overall charge are
balanced.
128
This method of balancing redox reactions will now be applied to a problem. The numbers
shown to the left of each step in the process correspond to the numbers for the steps in the
instructions given on the previous page.
Balance the following redox reaction:
1.
2
3.
4.
N2H4 + Pu2O3 → N2O + Pu(OH)2
(in base)
N2H4 + Pu2O3 → N2O + Pu(OH)2
(in OH− and H2O)
−2 +1
N2H4 → N2O
Nitrogen is balanced
N2H4 → N2O
+3 −2
+1 −2
−2 +1
+1 −2

Pu2O3 → Pu(OH)2

(Pu needs to be balanced)
Pu2O3 → 2 Pu(OH)2
one oxygen needed on the reactant side,
add one H2O to the reactant side
H2O + N2H4 → N2O
+2 −2 +1
one oxygen needed on the reactant side,
add one H2O to the reactant side

H2O + Pu2O3 → 2 Pu(OH)2
5. (in base)
6 H needed on the product side, add 6 H2O to 2H needed on the reactant side, add 2 H2O
to the product side and 6 OH− to the reactant  reactant side and 2 OH- to the product side
side
6OH− + H2O + N2H4 → N2O + 6 H2O  2 H2O + H2O + Pu2O3 → 2 Pu(OH)2 + 2 OH−
(H2O on the reactant side could be combined)
| 3 H2O + Pu2O3 → 2 Pu(OH)2 + 2 OH−
Note: It does not matter that there is H2O on both sides of the nitrogen equation at this point.
They will be canceled later. Hydrogen and Oxygen are balanced in each half reaction.
6.
add 6 e− to the product side
add 2 e− to the reactant side
6 OH− + H2O + N2H4 → N2O + 6 H2O + 6e−  2 e− + 3 H2O + Pu2O3 → 2 Pu(OH)2 + 2 OH−
7.
8.
Multiply equation above by 1
Multiply the equation above by 3
−
6 OH + H2O + N2H4 → N2O + 6 H2O + 6e−
6 e− + 9 H2O + 3 Pu2O3 → 6 Pu(OH)2 + 6 OH−
(add equations and combine like terms)
6 e− + 6 OH− + 10 H2O + N2H4 + 3 Pu2O3 → N2O + 6 H2O + 6 Pu(OH)2 + 6 OH− + 6e(cancel 6 e−, 6 OH−, and 6 H2O from each side of the reaction)
4 H2O + N2H4 + 3 Pu2O3 → N2O + 6 Pu(OH)2
9. Because the coefficients are in the lowest whole number ratio, the equation is complete.
10. Check to make sure the number of atoms and overall charge are balanced in the completed
equation.
129
Apply the method outlined for the half-reaction method to balance the following redox
reactions.
1. NO3− + Zn → Zn2+ + N2 (in acid)
2. O2 + I− → I2 (in base)
Hint: one of the half-reactions has nothing on the product side.
3. CrO2− + ClO− → Cl− + CrO42− (in base)
4. HNO3 + Bi2S3 → Bi(NO3)3 + NO + S (in acid)
Hint: remember that most metal sulfides are insoluble.
5. S2O32− + I2 → S4O62− + I− (in base)
6. Cr2O72− + Sn2+ → Sn4+ + Cr3+ (in acid)
7. SCN− + H2O2 → NH4+ + HCO3− + HSO4− (in acid)
Hint: same as in number 2.
130
DERIVING CHEMICAL EQUATIONS FROM BALANCED NET IONIC EQUATIONS
For oxidation-reduction reactions, often it is easier to balance the net ionic form of the
equation first and then to derive the chemical equation from the net ionic equation. The
following is a method for this procedure:
1.
Write the skeleton equation from the information given for the reactants and products.
2. Assign oxidation numbers to every element (including the elements of any acids and bases).
3. The elements which are spectator ions do not change oxidation numbers. However, sometimes an
ion can be involved as both a spectator ion and in oxidation-reduction.
4. Balance the net ionic equation following the rules given in the previous section. Remember to add
in the spectator ions on the side needing them when you balance the atoms other than oxygen and
hydrogen.
5. If needed divide to reduce the coefficients to the lowest whole-number ratio. At this point you will
need to add in the counter ion for the acid or base used. Add one counter ion for each H+ (for
sulfuric acid you will add HSO4−) or OH− in the equation to each side. The result of this step is the
ionic equation. Check to make sure that the net charge on each side of the reaction is zero.
6. Combine anions and cations to create the balanced chemical equation. No uncombined ions should
remain. Check to make sure the number of atoms is still balanced and that the coefficients are in
the lowest whole-number ratio.
Potassium permanganate reacts with chromium(III) chloride to produce manganese(IV) oxide and the
chromate ion in potassium hydroxide.
KMnO4 + CrCl3  MnO2 + CrO42−
+1 +7 −2
−
+3 −1
+4 −2
3 e + 4 H2O + KMnO4  MnO2 +
+6 −2
K+
+ 2 H2O +4 OH−
8OH− + 4 H2O + CrCl3  CrO42− + 3 Cl− + 8 H2O + 3 e—
4 OH− + KMnO4 + CrCl3  MnO2 + CrO42− + 3 Cl− + K+ + 2 H2O
+4 K+
+4 K+
4 KOH + KMnO4 + CrCl3  MnO2 + K2CrO4 + 3 KCl + 2 H2O
Balance atom other than H or O.
Balance O by adding water.
Balance H by adding H+.
Balance charge by adding e-.
Add in counter ion to acid or base.
131
MORE PRACTICE REDOX PROBLEMS
1. NaCl + MnO2 → Mn2+ + Cl2
(in H2SO4)
2. K4Fe(CN)6 + CeCl4 → Ce(OH)3 + Fe(OH)3 + CO32− + NO
(in KOH)
3. NaNO2 + Al → NH3 + AlO2−
(in NaOH)
4. NaIO3 + NaI → NaI3
(in HI)
5. Fe + HCl → HFeCl4 + H2
6. Fe(OH)2 + H2O2 → Fe(OH)3
(in KOH)
7. Na2S2O8 + CrCl3 → Cr2O72− + SO42−
(in HCl)
8. KCN + KMnO4 → CNO− + MnO2
(in KOH)
9. CrI3 + Cl2 → CrO42− + IO4− + Cl−
(in NaOH)
10. Potassium permanganate and nitrous acid react in sulfuric acid. Two of the products of this
reaction are manganese(II) bisulfate and nitric acid.