LABORATORY MANUAL FOR CHEMISTRY 102

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LABORATORY MANUAL FOR
CHEMISTRY 102
Prepared by
Department of Chemistry and Physics
Los Angeles Valley College
This Lab Book Belongs To:
Copyright © 2014 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
TABLE OF CONTENTS .............................................................................................................................. 2
LABORATORY SAFETY RULES .................................................................................................................. 3
FACTORS AFFECTING THE RATE OF A REACTION.................................................................................... 6
CHEMICAL KINETICS ............................................................................................................................. 13
LE CHÂTELIER'S PRINCIPLE ................................................................................................................... 25
WEAK ACIDS AND BASES ...................................................................................................................... 35
DETERMINATION OF Ka BY pH TITRATION ............................................................................................ 43
BUFFERS AND pH ................................................................................................................................. 53
ACID-BASE EQUILIBRIUM PROBLEMS .................................................................................................. 61
A SOLUBILITY INVESTIGATION .............................................................................................................. 66
SOLUBILITY AND Ksp DETERMINATION ................................................................................................. 73
DETERMINATION OF Kf BY SPECTROPHOTOMETRIC METHODS........................................................... 81
INTERNAL ENERGY PROBLEMS ......................................................................................................... 89
BOMB CALORIMETRY ........................................................................................................................... 94
Ksp, G, H, AND S OF POTASSIUM NITRATE DISSOLVING IN WATER .......................................... 105
ELECTROCHEMISTRY .......................................................................................................................... 113
ELECTROLYTIC DETERMINATION OF THE MOLAR MASS OF LEAD...................................................... 127
DETERMINATION OF THE HALF-LIFE OF POTASSIUM-40.................................................................... 131
COORDINATION COMPOUNDS AND COMPLEX IONS......................................................................... 137
EQUILIBRIUM BETWEEN TWO COMPLEX IONS OF Co2+ IN SOLUTION .............................................. 140
SYNTHESIS AND ANALYSIS OF A NICKEL COMPLEX ............................................................................ 151
MOLECULAR MODELS OF TRANSITION METAL COMPLEXES ............................................................. 161
APPENDIX A ........................................................................................................................................ 165
APPENDIX B ........................................................................................................................................ 166
APPENDIX C ........................................................................................................................................ 176
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
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.
7
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.
8
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.
9
Report
Name ___________________________
FACTORS AFFECTING THE RATE OF A REACTION Section ___________________________
A.
The effect of mixing.
Time for color to diffuse throughout the tube
stirred tube
undisturbed tube
B.
__________________________
__________________________
The effect of concentration of a reactant.
Time for color change in first trial
_________________________
Time for color change in second trial
_________________________
C.
The effect of temperature.
Time for color change in heated trial
D.
________________________
The effect of a catalyst.
Amount of gas formed in pure H2O2
_______________________
Amount of gas formed in just the catalyst solution
________________________
Time for tube to fill with gas when mixing
H2O2 and Cu(NO3)2
_______________________
Time for tube to fill with gas when mixing
H2O2 and Fe(NO3)3
_______________________
Time for tube to fill with gas when both
catalysts are used with H2O2
_______________________
10
Report
FACTORS AFFECTING THE RATE OF A REACTION
E.
Name __________________________
The effect of surface area
Color intensity of original copper solution
_______________________
Appearance of the nail initially
_______________________
After at least 10 minutes in the copper solution ______________________
Color intensity of solution with nail after 10 minutes
_______________________
Appearance of the steel wool initially
_______________________
After at least 10 minutes in the copper solution _______________________
Color Intensity of solution with steel wool after
after 10 minutes
_______________________
11
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?
3.
b.
Which tube had the higher concentrations of either or both reactants?
c.
How does concentration affect the rate of the reaction?
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?
12
QUESTIONS FOR FACTORS EXP.
4.
NAME______________________________
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.
13
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:
Rate = k [ A] [B] [C ] 
x
y
z
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.)
14
x
y
k I −   S O2 − 
For reaction (1) the reaction rate is equal to   0  2 8  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
y
2−
−
Rate2 k I  2  S2O8  2
=
Rate1 k I −  x  S O2−  y
 1  2 8 1
" Exp 2"
" Exp 1"
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
x
log
log2.00
=
=
x log2.00
Rate1
further reduces to:
x=
log
( )
Rate2
Rate1
log2.00
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 pre-exponential 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 , R is in J mol 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:
-1
-1
 k1  Ea  1 1 
ln=
 
 − 
 k2  R  T2 T1 
15
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 apparati 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.
Experiment 1
0.200-M KI
0.200-M KCl
0.00500-M Na2S2O3
Starch
0.100-M K2S2O8
0.100-M Na2SO4
10.0
10.0
10.0
10 drops
20.0
0
Volumes of solution, mL
Room Temperature
Experiment 2
Experiment 3
In a flask:
20.0
20.0
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
16
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.
17
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.
18
19
REPORT
NAME_______________________
CHEMICAL KINETICS
SECTION _____________________
Initial reactant concentrations:
Experiment
[Iodide]
[Persulfate]
[Thiosulfate]
1
2
3
Temperature of solution, (°C)
4
Temperature of solution, (°C)
Elapsed time, (s)
Trial A
Trial B
Experiment
Average time, (s)
Rate, (M/s)
1
2
3
4 (ice bath)
For room temperature data:
Reaction Order
With correct sig figs
With respect to Iodide ion
With respect to Persulfate ion
Overall
Value for k from Experiment 1
Value for k from Experiment 2
Value for k from Experiment 3
Average value for k (including units)
Room Temperature Rate Law
SAMPLE CALCULATIONS (use separate sheets)
Include how you solved for both exponents in this experiment.
Rounded to whole
numbers
20
REPORT
NAME________________________________
CHEMICAL KINETICS (cont’d)
For ice bath data:
Value for k (including units)
SAMPLE CALCULATIONS (use separate sheets if needed)
Activation Energy, (kJ mol-1)
SAMPLE CALCULATIONS (use separate sheets if needed)
21
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
Exp2
for 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 .
22
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.
23
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 in 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).
24
25
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 :LB  M(:LB)yz+
Once a formation equilibrium is established, a change in temperature; in the concentration of the
metal ion, (M); in the concentration of a ligand, (:LB), (which may or may not carry an overall charge);
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.
26
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)62+
1.
Clean a test tube.
2.
Observe the 6-M ammonium hydroxide, 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 ammonium hydroxide (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 ammonium hydroxide 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)62+. 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 CoCl42-
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! Immediately neutralize and
clean up any spills!!
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 CoCl42-.)
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.)
27
5. Place the test tube from step 4 in the hot water bath and wait a few minutes for a color 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.
28
29
REPORT
NAME ________________________________
LE CHATELIER’S PRINCIPLE
SECTION _____________________
A. Record colors of:
Fe(NO3)3
KSCN
Fe(SCN)2+
NH3
Ni(NH3)62+
HCl
CoCl42-
Compare the color of the solutions in:
test tubes 1 and 4:
test tubes 2 and 4:
test tubes 3 and 4:
B. Record colors of:
Ni(NO3)2
Color after the addition of 6-M HCl:
C. Record colors of:
Co(NO3)2 
Color after addition of H2O:
Color after heating:
Color after cooling:
30
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
31
2.
a. Write a balanced net ionic equation for the equilibrium reaction in Part B, formation of
Ni(NH3)62+
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.
32
QUESTIONS FOR LE CHATELIER’S EXP.
3.
NAME ________________________________
a. Write a balanced net ionic equation for the (equilibrium) formation of the
tetrachlorocobaltate(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)
33
QUESTIONS FOR LE CHATELIER’S EXP.
NAME ________________________________
4. Consider the hypothetical equilibrium: A + B  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 reestablish equilibrium, right
(→), left (←) or NS?
When the new
A
equilibrium
has been established,
B
is the amount of each
substance present
C
greater (), less (), or
unchanged (NC) from what D
it was before the change?
Changes
C is added
A is added
D is
removed
A catalyst
is added
Temperature
is decreased
34
35
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, ammonium hydroxide. Because
pH =
14.00 − pH
− log H3O +  and pOH =
− 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.
monoprotic acid is (x/y) times 100.
The percent ionization (dissociation) for the
The effect of weak bases on pH is due to the ionization (hydrolysis) of water. If A- represents any weak
base and y is the initial concentration of that base, then
Initial
Change
Equilibrium
A- + H2O  HA + OHy
0
~0
-x
+x
+x
y-x
x
x
and the percent of the weak base involved in ionization is (x/y) times 100.
36
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.
37
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
ammonium hydroxide for the initial solution. (Note: ammonium hydroxide 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 ammonium
hydroxide solution.
38
3.
Calculate the Ka for each of the acetic acid solutions, and the average Ka. Calculate the 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 non-standard conditions, your
experimental error may be quite large. Use analogous calculations to calculate percent relative
average deviation and percent error for your Kb for ammonium hydroxide.
39
REPORT
NAME_________________________________
WEAK ACIDS AND BASES
SECTION _______________________________
Acetic acid:
Solution from
step
Molarity of
HC2H3O2
2
1.0-M
pH
[H3O+], M
% Dissociation
Ka
pH
[OH-], M
% Hydrolysis
Kb
3
4
5
Average Ka
Percent Relative Average Deviation
Percent Error
Ammonium hydroxide (NH3):
Solution from
step
Molarity of
NH3
2
1.0-M
3
4
5
Average Kb
Percent Relative Average Deviation
Percent Error
SHOW CALCULATIONS ON SEPARATE PAGES:
40
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?
41
QUESTIONS FOR WEAK ACIDS & BASES EXP.
NAME ______________________________
3.
A 0.0100 M solution of a weak monoprotic acid is found to be 3.5% 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 4.6×10-4. Calculate the percent hydrolysis of the base and the pH of the
solution if the initial concentration of the weak base is 2.5×10-2 M.
42
43
DETERMINATION OF Ka BY pH TITRATION
INTRODUCTION
From previous chemistry labwork 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 which allows
44
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:
pH
= pK a + log
[conjugate base]
[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 yaxis and log(Vb/(Ve-Vb) on the x-axis for volumes from about 20% to 80% of the equivalence point 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 − pKa
In this experiment, an acid-base pH titration curve will be constructed for potassium hydrogen
phthalate (KHP). KHP, KHC8H4O4, is a monoprotic acid having a structural formula of:
45
O
H
H
C
O - K+
C
C
H
C
C
C
C
O
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 ringstand 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
46
KHP determined in step 4 into each of the two beakers. Record the mass of KHP in each beaker
(±0.0001 g).
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. Place titration beaker 1 under the Drop Counter and lower
the tip of the pH probe through the hole in the Drop
Counter into the solution of beaker 1.
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 about 1 or 2 drops per second.
13. Continuously to swirl the beaker to mix well.
14. When pH reaches about 12 and you have added about 2 to
3 mL of solution 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.
15. Refill the buret with the NaOH solution.
16. Repeat Steps 11 through 14 for beaker 2. Before pressing the “Collect” button, click on the file
cabinet icon to add another run to the data collection (Run 2).
17. Attach a USB flash drive to the USB 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 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).
18. 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:
47
19. In Excel, create a two new columns for your data. The first column
should be labelled pH/V. In this column starting with row
containing the first data point enter the formula “=(B9-B8)/(A9-A8)”
(without the quotation marks). Here we used B8, B9 and A8,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
800 data points so be careful). Do this for both runs.
20. Construct two titration graphs, one for each run. Plot pH (vertical
axis) as a function of the volume of NaOH added (horizontal axis).
Be sure to properly title and label your graphs. See Appendix B for a review on graphing. Do this
for both runs.
21. On the same graph as you created for the pH vs. volume graph the pH/V vs. volume (add a
new data series to the original graph). The point on the x-axis (volume) where this line has its
maximum value is the equivalence point. Do this for both runs.
22. The second column should be labelled log(Vb/(Ve-Vb)). Starting with 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 22. 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 (y=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.
48
49
REPORT
NAME_______________________________
DETERMINATION OF Ka BY pH TITRATION
SECTION _____________________________
Molarity of NaOH, M
Mass of KHP, g Trial 1
pKa (trial 1)
Mass of KHP, g Trial 2
pKa (trial 2)
Average pKa
Ka
50
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 1.5-M K2C8H4O2 solution.
51
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.
52
53
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− 
= pK a + log  
pH
[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.
54
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 0.10-M 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.
55
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.
56
57
REPORT
NAME_________________________________
BUFFERS AND pH
SECTION _______________________________
Solution
0.20-M
HC2H3O2
0.20-M
NaC2H3O2
Combined
Solution
pH
Initial Molarities in Combined Solution
acetate
acetic acid
Initial Moles
Vial Number
acetate
1
2
3
4
5
6
7
8
9
10
11
SAMPLE CALCULATIONS: (Use separate sheets if necessary)
acetic acid
acetate : acetic acid ratio
58
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.
59
QUESTIONS FOR BUFFERS AND pH EXP.
NAME ____________________________
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).
60
61
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.
62
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.
63
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?
64
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?
65
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?
66
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-  HSO42- + 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 reestablished.
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
67
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
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.
68
69
REPORT
NAME ______________________________
A SOLUBILITY INVESTIGATION
SECTION ___________________________
A. Effect of lowering the pH on the solubility of an insoluble salt.
Observations after mixing
the two reagents
Observations after adding
HNO3 to the mixture
B.
Effect of raising the pH on the solubility of an insoluble salt
Observations after mixing the
two reagents
Observations after adding
NH3 to the mixture
Observations after adding
HNO3 to the mixture
70
C. Effect of adding a strong acid or base to an amphoteric hydroxide
Observations after mixing the two
reagents
To the first portion
Observation after adding
mL NaOH
To the second portion
Observation after adding
mL HNO3
71
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 Chatelier’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 nitric acid 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.
72
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
73
SOLUBILITY AND Ksp DETERMINATION
INTRODUCTION
Calcium iodate is an ionic compound that is only slightly soluble in water. In aqueous solution, an
equilibrium forms between the solid salt and its ions:
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- + 24O62-
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.
74
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.)
75
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.
76
77
REPORT
NAME _______________________________
SOLUBILITY AND DETERMINATION OF Ksp
SECTION _____________________________
Molarity of Na2S2O3 solution, (M)
Beaker A
Beaker B
Beaker C
Flask A
Flask B
Flask C
0.0100
0.0100
0.0100
Approx. mass of Ca(IO3)2 used, (g)
Initial Na2S2O3 buret reading, (mL)
Final Na2S2O3 buret reading, (mL)
Volume Na2S2O3 used, (mL)
Moles of IO3-in aliquot, (mol)
Volume of aliquot, (L)
Molarity of IO3-in saturated solution, (M)
Average Molarity of IO3-, (M)
Molar Solubility of Ca(IO3)2
Solubility in g Ca(IO3)2/100 mL
Ksp for Ca(IO3)2 (include units)
SAMPLE CALCULATIONS (use separate sheets)
78
QUESTIONS FOR SOLUBILITY AND DETRMINATION OF KSP
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.
79
QUESTIONS FOR SOLUBILITY AND DETRMINATION OF KSP
NAME_______________________
4.
Should a precipitate form when 25.00 mL of 0.0100 M silver nitrate is added to 15.00 mL of
0.0500 M potassium acetate?
5.
If a 6.00 M potassium chloride solution is added dropwise (no significant volume change) to a
solution containing both 0.010 M silver nitrate and 0.010 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?
80
6.
What is the molar solubility of barium fluoride in a solution that contains 1.00 M acetic acid and
1.00 M sodium acetate? Hint: Combine three equilibria reactions to determine the Kc for:
BaF2 + 2 HC2H3O2  Ba2+ + 2 HF + 2 C2H3O2-
and then solve for the molar solubility using the approximation method. Be sure to validate!
81
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 absorbance and the “standard curve.”
82
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
6. Cover each tube with parafilm and mix well.
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
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
83
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.
84
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).
85
REPORT
NAME ____________________________
DETERMINATION OF Kf …
SECTION __________________________
A.
Standard Curve
Tube No.
1
2
3
4
5
6
7
8
9
10
11
12
b. Kf Determination
Tube No.
1
2
3
4
5
6
7
8
9
10
Tube No.
1
2
3
4
5
6
7
8
9
10
Absorbance
Eq [Fe(SCN)2+] (from calculations)
Absorbance
Eq [Fe(SCN)2+] (from graph)
Initial [Fe3+]
Initial [SCN-]
Eq [Fe3+]
Eq [SCN-]
86
REPORT
NAME _____________________________
DETERMINATION OF Kf …
Tube No.
1
2
3
4
5
6
7
8
9
10
Average Kf
Percent Relative Average Deviation
SAMPLE CALCULATIONS (use separate sheets if necessary)
Kf
87
QUESTIONS FOR DETERMINATION OF Kf …
NAME _____________________________
1. In your opinion, what most affects the precision for this experiment?
2. In your opinion, 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.
88
QUESTIONS FOR DETERMINATION OF Kf …
NAME _____________________________
4. The reaction:
Co2+ + 6 NH3  Co(NH3)62+
has an equilibrium constant of 5.0 x 104 M-6. Solutions were mixed so that the initial concentration
of the cobaltous ion was 0.100 M and the ammonia was 1.00 M. What are the equilibrium
concentrations of all three species in the reaction?
89
INTERNAL ENERGY PROBLEMS
INTRODUCTION
Energy can be broken down into two types: kinetic energy and potential energy. Kinetic energy is
the energy of motion and potential energy is the energy of position. Electrons and nuclei which make
up a substance have both potential energy (because of their relative positions), and kinetic energy
(from motions within the system). The sum of the kinetic and potential energies for each of the
particles that make up a substance is known as the internal energy, E, of the substance. The change
in internal energy, E, that occurs as a result of a chemical reaction or during a phase change is
something that can be measured experimentally.
Another way to express an energy change of a system is by the change in enthalpy, H. The enthalpy
change for a chemical reaction is the amount of heat that is transferred during the reaction at
constant pressure. Heat flows from a region at higher temperature to a region at lower temperature.
A more precise definition of enthalpy is:
H= E + PV
where P is the pressure of the system and V is the volume of the system. If we consider a reaction
under constant pressure we can derive a new expression for H:
∆H =∆E + P∆V
which can be rearranged to solve for E:
∆E =∆H − P∆V
This equation is just another way of stating that for the reaction being studied, the change in internal
energy is equal to heat plus pressure-volume work. (At constant pressure, heat is H and ‒PV
is work.)
When a reaction occurs in a closed, expandable container (such as a piston and cylinder or in
a balloon), the change in volume can be measured and the amount of work can be determined.
However, in an open container, the change in volume cannot be measured and an alternate method
of determining work must be used. To do this, we need the Ideal Gas Law. We need to remember
that the temperature of the system remains constant. Therefore, the only value on the right-hand
side of the equation that can change is the number of moles of gas. Thus, work can be expressed as
−P∆V = −∆n ( RT )
which makes the expression for E:
∆E = ∆H − ∆n ( RT )
When using this equation there are several things to remember.
• The values for E and H are for the reaction as written and the units for E and H are
J/mol rxn or kJ/mol rxn.
• n is the change in number of moles of gas in the balanced chemical equation. The units
are moles of gas / mole of reaction. R has the value of 8.314 J/mol K.
• T is the standard thermodynamic temperature of 298.15 K.
90
Example:
Hydrogen reacts with oxygen to produce water according to the following reaction:
H2 + ½ O2  H2O
H = –241.826 kJ/mol rxn
What is the change in internal energy of this reaction?
Solution:
The change in number of moles of gas in this reaction is:
1 mol of gas final – 1.5 moles of gas initial = –0.5 mol gas /mol rxn
We can now calculate the change in internal energy of the reaction.
∆E = ∆H − ∆n ( RT )
=
=
1 kJ 
−241.826 kJ  0.5 mol gas   8.314 J 
−  −
(298.15 K )  3  


mol rxn
mol rxn   mol gas ⋅ K 
 10 J  

−240.587 kJ
mol rxn
91
PROBLEMS (1 Latm = 101.325 J, 1 Lbar = 100.000 J)
1. In an exothermic process, the volume of a system expanded from 186 mL to 1997 mL against
a constant pressure of 745 torr. During the process, 18.6 calories of heat energy were given
off. What was the internal energy change for the system in joules?
2. Calculate the change in internal energy for the thermal decomposition of 1.000 g of potassium
chlorate at a constant external pressure of 943.2 mmHg. The decomposition reaction is
2 KClO3
∆

→ 2KCl + 3 O2
MnO2
Potassium chlorate’s heat of formation is ‒391.20 kJ/mol, potassium chloride’s heat of
formation is ‒435.87 kJ/mol, and oxygen has a density of 1.308 g/L at the reaction temperature.
92
3. (a) A gas at expands from 2.0 L to 6.0 L at a constant pressure of 912 mmHg. If qp was zero in
the process, what would be the change in internal energy?
(b) What would be the change in internal energy for the process described in 3 (a) if the expansion
occurred when the external pressure was zero?
4. The oxidation of nitric oxide
2 NO + O2  2 NO2

∆Hrxn
=
−113.1 kJ
is a key step in the production of photochemical smog. Calculate the change in internal energy
(in kJ) that occurs when 15.4 g of NO reacts with excess Oxygen at 35.0°C.
93
5. A gaseous mixture was enclosed in a piston and cylinder system having a volume of 1545
mL. When a chemical reaction occurred, the volume decreased to 375 mL and 321 calories of
heat was absorbed. The external pressure was 753 mmHg. Calculate the change in enthalpy,
work, and the change in internal energy for this system in kJ.
6. When 3.55 g of methane (CH4) gas, was burned in excess oxygen at 45C, the internal energy
change was ‒196.3 kJ. Calculate the enthalpy change for the combustion of one mole of
methane at 45°C.
94
BOMB CALORIMETRY
INTRODUCTION
The enthalpy of reaction involving gases can be conveniently determined using an apparatus called a
"bomb" calorimeter (see the diagram above). A bomb calorimeter, a device in which volume, not
pressure, is constant, is used to measure the change in internal energy (E) that occurs during a
physical or chemical change. The enthalpy of the process can then be calculated. In a bomb
calorimeter, the reaction takes place in a sealed, rigid, container (called the bomb) which is enclosed
in an insulated water jacket. A spark from an electrical circuit is used to start the combustion reaction
inside the bomb. The resulting temperature change of the water is measured and used to calculate
the energy change.
The bomb calorimeter must be calibrated to determine its heat capacity. A weighed sample of pure
benzoic acid having a known enthalpy of combustion is burned and the water’s temperature change
is then used to calculate the heat capacity of the calorimeter. The calorimeter is then set up in an
identical manner for additional experiments involving various fuels.
The video experiment demonstrates the use of a bomb calorimeter. You will use data obtained from
the video to calculate the molar enthalpy of combustion for each of the compounds listed on the
following page.
95
Straight Chain Alcohols
Formula
M. W.
1-Propanol (Propan-1-ol)
CH3(CH2)2OH
60.11
1-Butanol (Butan-1-ol)
1-Pentanol (Pentan-1-ol)
CH3(CH2)3OH
74.12
CH3(CH2)4OH
Six Carbon Cyclic Compounds
Formula
88.15
M. W.
H2
C
H2C
CH2
H2C
CH2
84.16
Cyclohexane
C
H2
H2
C
H2C
CH
82.15
Cyclohexene
H2C
CH
C
H2
H2
C
HC
CH
80.14
1,4-Cyclohexadiene
HC
CH
C
H2
H
C
1,3,5-Cyclohexatriene (Benzene)
HC
CH
HC
CH
C
H
78.12
96
For these combustion reactions, a fuse wire and cotton ignition thread are used. It has been found
that the length of fuse wire and ignition thread used to initiate the combustion does not contribute
enough heat to make a significant difference in the total heat evolved when sample sizes are as large
as those used in the video.
The products for the combustion of hydrocarbons are carbon dioxide gas and liquid water (when the
reactions occur under the conditions employed in the video). To make sure that the water generated
by the reaction condenses to the liquid form, a few drops of water are added to the reaction chamber
during assembly. This added water avoids supersaturation of water in the gas phase.
PROCEDURE
Note: Some of the data needed for this experiment are already recorded on your data sheet.
A. Determination of the Heat Capacity of the Bomb Calorimeter
1. Watch the video. Calculate the mass of benzoic acid (molecular formula: C6H5COOH; molecular
weight: 122.13 g/mole) used.
2. Write a balanced equation for the combustion of one mole of benzoic acid. (Remember, the other
reactant is oxygen gas and the products will be carbon dioxide gas and liquid water.)
3. The standard enthalpy of combustion (H) for benzoic acid is -3227 kJ/mol and standard
temperature is 298.15 K.
a. Determine the standard change in internal energy (E) for the combustion of one mole of
benzoic acid.
∆E = ∆H − ∆nRT
Where n = moles gaseous products - moles gaseous reactants from the balanced equation for
the combustion of one mole of benzoic acid.
b. Calculate q for the moles of benzoic acid actually burned in this experiment (this is qexp used
in step 4c).
c. Remember that
qcalorimeter = −qexp
and because in this experiment volume (not pressure) was constant
qcalorimeter =
nrxn ∆Ecalorimeter =
C calorimeter ∆T
97
98
4. Use the information/equations outlined in step 3 and the experimental data from the video, to
calculate the heat capacity of the calorimeter.
B. Relationship between Enthalpy of Combustion and chain length in Straight Chain Alcohols
1. The masses of the sample plus crucible and the empty crucible are recorded on your data sheet.
Calculate the mass of each compound used. Using the molecular weight given in the introduction,
calculate the moles of each compound used.
2. Observe the video. The initial and final temperatures for the combustion of each compound are
given in the data tables of the report pages. Calculate T for each.
3. Using the formulae shown in the introduction, write a balanced equation for the combustion of
one mole of each compound. Calculate n for each reaction.
4. Calculate the standard molar enthalpy of combustion (enthalpy per mole of substance burned) for
each compound (use the heat capacity for the calorimeter determined in Part A).
Note that for all of these experiments the masses of reactants used are similar (± 0.1g). Because the
chemicals involved in the reaction are all part of the calorimeter assembly that absorbs the heat given
off during combustion, similar masses of reactants must be used in order for the heat capacity
determined in Part A to be valid for all the experiments. Also note that the same equations presented
in Part A apply to Parts B and C.
5. Answer the questions concerning the relationship between chain length of straight-chain alcohols
and their enthalpies of combustion.
C. Relationship between Enthalpy of Combustion and number of double bonds in Cyclic Six-Carbon
Compounds
1. Follow steps 1-4 for Part B above.
2. Answer the question about Benzene concerning the relationship between carbon-carbon double
bonds and enthalpy of combustion.
99
REPORT
NAME ______________________________
BOMB CALORIMETRY EXP.
SECTION _____________________________
A.
Determination of Heat Capacity
Mass of benzoic acid + thread, (g)
0.7934
Mass of thread, (g)
0.0047
Mass of benzoic acid, (g)
Final temperature, (°C)
17.107
Initial temperature, (°C)
15.070
T, (°C)
Write the balanced equation for the combustion of one mole of Benzoic acid:
n=________
Heat capacity of the calorimeter _____________________________
SAMPLE CALCULATIONS (use separate sheets if necessary)
100
REPORT FOR BOMB CALORIMETERY EXP. (cont.) NAME ________________________________
B.
Straight Chain Alcohols
Propanol
Butanol
Pentanol
Mass of sample + crucible, (g)
11.6465
11.7158
11.7357
Mass of crucible, (g)
10.8902
10.8925
10.8997
Final temperature, (°C)
22.085
22.091
20.030
Initial temperature, (°C)
19.623
19.230
16.972
Mass of sample, (g)
T, (°C)
Molar enthalpy of combustion
Write a balanced equation for the combustion of one mole of each compound.
1.
n=__________
2.
n=__________
3.
n=__________
SAMPLE CALCULATIONS (use separate sheets if necessary)
101
REPORT FOR BOMB CALORIMETERY EXP. (cont.) NAME _________________________________
C.
Cyclic compounds
Cyclohexane
Cyclohexene
1,4-Cyclohexadiene
Benzene
Mass of sample + crucible, (g)
11.7403
11.7017
11.5666
11.6982
Mass of crucible, (g)
10.8978
10.8977
10.8990
10.8993
Final temperature, (°C)
20.987
21.369
23.778
20.879
Initial temperature, (°C)
17.203
17.825
20.891
17.661
Mass of sample, (g)
T, (°C)
Molar enthalpy of combustion
Write a balanced equation for the combustion of one mole of each compound:
1.
n=________
2.
n=________
3.
n=________
4.
n=________
SAMPLE CALCULATIONS (use separate sheets if necessary)
102
QUESTIONS FOR BOMB CALORIMETRY EXP.
NAME ________________________________
1. What kind of changes in enthalpies of combustion did you observe as the chain length of the
alcohols increased? What was the approximate change for each CH2 unit added?
2. Benzene does not follow the trend established by the enthalpy values obtained for the other
compounds in Part C. What value would be predicted for the molar enthalpy of combustion for
benzene from the observed trend obtained from cyclohexane, cyclohexene, and 1,4cyclohexadiene? Examine the cyclic fuels’ structures, and review resonance structures, delocalized
bonding, and bond energies in Chapters 9 and 10 of the Tro text. Use the information obtained to
explain why the molar enthalpy of combustion for benzene does not follow the trend.
103
QUESTIONS FOR BOMB CALORIMETRY EXP.
NAME ________________________________
3. Hexanol (C6H14O) is a liquid alcohol similar to the ones used in the experiment. In another bomb
calorimetry experiment 0.8278 g of hexanol is burned and the temperature of the calorimeter
increased from 16.834C to 19.203C. The heat capacity of the calorimeter is 13.52 kJ C-1.
Calculate the enthalpy of formation of hexanol in kJ/mol C6H14O. Compare this to the accepted
value of -377.5 kJ mol-1.
104
105
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 10-10, 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)
106
where R is the 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 hotwater 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).
107
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 (K-1).
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.
∆S
22. Calculate S° from the y-intercept from your graph ( b =
average S° from step 21.

R
) and compare this value to the
108
109
REPORT
NAME ___________________________
Ksp, G, H, and S of KNO3
SECTION _________________________
Mass of KNO3 used, (g)
Moles of KNO3 used, (mol)
Temperature at
which crystals
appear, (°C)
Volume of
Solution, (mL)
Temperature at
which crystals
appear, (K)
1
T , (K-1)
Solution 1
Solution 2
Solution 3
Solution 4
Solution 5
Solution 6
Molar
solubility of
KNO3
Ksp
ln Ksp
Solution 1
Solution 2
Solution 3
Solution 4
Solution 5
Solution 6
Show Calculations (use separate sheets if necessary)
G°
(J mol-1)
H
(J mol-1)
S°
(J mol-1 K-1)
110
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.
111
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.
112
113
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 1M, all partial pressures of reactant and product gases are 1 atmosphere. 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 nonstandard 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.
114
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.
115
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.
116
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.
117
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.
118
119
REPORT
NAME___________________________
ELECTROCHEMISTRY
SECTION _________________________
A. Voltaic Cells
Cu/Cu2+; Fe/Fe2+
Experimental
Cell Voltage
Cathode half-reaction

Ered
Anode half-reaction

−Ered

Ecell
Overall chemical reaction
Experimental Cell Notation
Cu/Cu2+; Fe2+/Fe3+
Experimental
Cell Voltage
Cathode half-reaction

Ered
Anode half-reaction

−Ered

Ecell
Overall chemical reaction
Experimental Cell Notation
Cu/Cu2+; Zn/Zn2+
Experimental
Cell Voltage
Cathode half-reaction

Ered
Anode half-reaction

−Ered

Ecell
Overall chemical reaction
Experimental Cell Notation
120
REPORT
NAME ___________________________
ELECTROCHEMISTRY (cont.)
Fe/Fe2+; Fe2+/Fe3+
Experimental
Cell Voltage
Cathode half-reaction

Ered
Anode half-reaction

−Ered

Ecell
Overall chemical reaction
Experimental Cell Notation
Fe/Fe2+; Zn/Zn2+
Experimental
Cell Voltage
Cathode half-reaction

Ered
Anode half-reaction

−Ered

Ecell
Overall chemical reaction
Experimental Cell Notation
Fe3+/Fe2+; Zn/Zn2+
Experimental
Cell Voltage
Cathode half-reaction

Ered
Anode half-reaction

−Ered

Ecell
Overall chemical reaction
Experimental Cell Notation
121
REPORT
NAME ___________________________
ELECTROCHEMISTRY (cont.)
B. Concentration Cell
Temperature of copper(II) nitrate solution
Molarity of the diluted copper(II) nitrate solution
Initial Voltage
Voltage after addition of 10 drops of 0.10-M copper(II) nitrate
C. Electrolysis of Aqueous Potassium Iodide
Volume of solution
Initial pH
Initial [OH-]
Elapsed time of electrolysis
Final pH
Final [OH-]
Change in [OH-]
Observations:
Before electrolysis
During electrolysis
After electrolysis
122
REPORT ELECTROCHEMISTRY EXP. (cont.)
NAME _____________________________
D. Determination of the solubility product constant of copper(II) carbonate.
Standard Cell Potential of a copper/zinc cell, V
Cell Potential of experimental cell, V
Number of moles of electrons transferred in the
chemical equation, n
Concentration of copper(II) ions in the saturated
copper(II) carbonate half-cell, M
Temperature of solution in beaker, C
Experimental Ksp
SAMPLE CALCULATIONS (Use additional paper as needed)
123
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?
124
QUESTIONS FOR ELECTROCHEMISTRY EXP.
NAME___________________________
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).
125
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?
126
127
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 half-reaction:
Pb2+ + 2 e-  Pb
represented by the half-reaction:
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
Pb  Pb2+ + 2 eThe 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.
128
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. The initial masses of the cathode and anode lead strips will be the final
masses from step 8.
129
REPORT
NAME ____________________________
ELECTROLYTIC DETERMINATION …
SECTION __________________________
Before Electrolysis
Trial 1
After Electrolysis
Mass of Anode, g
Mass of Anode, g
Mass of Cathode, g
Mass of Cathode, g
Initial Current, A
Mass of lead
transferred, g
Current @ 5 min, A
Time Elapsed, s
Current @ 10 min, A
Average Current, A
Current @ 15 min, A
Current @ 20 min, A
Current @ 25 min, A
Moles of electrons
transferred, mol
Moles of lead
transferred, mol
Molar mass of lead, g
mol-1
Current @ 30 min, A
Before Electrolysis
Trial 2
After Electrolysis
Mass of Anode, g
Mass of Anode, g
Mass of Cathode, g
Mass of Cathode, g
Initial Current, A
Mass of lead
transferred, g
Current @ 5 min, A
Time Elapsed, s
Current @ 10 min, A
Average Current, A
Current @ 15 min, A
Current @ 20 min, A
Current @ 25 min, A
Current @ 30 min, A
Moles of electrons
transferred, mol
Moles of lead
transferred, mol
Molar mass of lead,
g mol-1
Average molar
mass of trials 1 & 2,
SHOW CALCULATIONS (attach separate sheets if necessary)
130
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?
131
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
132
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 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 selfabsorption 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
133
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 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.
134
135
REPORT
NAME___________________________
DETERMINATION OF THE HALF-LIFE OF K-40
SECTION_________________________
30 min
background
count
Initial
mass, g
Background
CPM
Added
mass, g
Total
mass, g
Counts in
10 min
CPM
Extrapolated ln(CPM/g)
Extrapolated CPM/g
Adjusted CPM/g
Number of K-40 nuclei
in 1.00 g
Rate constant (1/min)
Half-life (min)
Half-life (yr)
SAMPLE CALCULATIONS: (Use separate sheets if necessary)
Net
CPM
Net CPM
g KCl
 Net CPM 
ln 

 g KCl 
136
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?
137
COORDINATION COMPOUNDS AND COMPLEX IONS
Nomenclature System
The textbook provides an account of how to write names of complex ions and coordination
compounds. However, the method in the textbook is not entirely correct according to the
International Union of Pure and Applied Chemistry (IUPAC). IUPAC is the governing body that sets the
rules for naming compounds and approving the names of newly created elements. For inorganic
compounds, compounds that are not comprised mainly of carbon and hydrogen, the document is
“NOMENCLATURE OF INORGANIC CHEMISTRY IUPAC Recommendations 2005” also known as the Red
Book. For organic chemistry, the relevant document is “A Guide to IUPAC Nomenclature of Organic
Compounds Recommendations 1993,” also known as the Blue Book. The rules outlined here are from
the Red Book.
The following general rules are used when naming coordination compounds:
(i)
ligand names are listed before the name of the central atom,
(ii)
no spaces are left between parts of the name that refer to the same coordination entity,
(iii)
ligand names are listed in alphabetical order (multiplicative prefixes indicating
quantities are not considered in determining that order),
(iv)
the use of abbreviations in names is discouraged.
The general process for naming a complex ion or coordination compound is:
• identify the central atom
• identify the ligands
• name the ligands
• specify the mode of coordination (which atom provides the lone pair of electrons that
forms the coordinate covalent bond to the central atom, if known)
• order the ligands – this will be alphabetical without regards to any multiplicative
modifiers
The central atom (in this class) will be a transition metal. This can become more complicated in higher
level chemistry classes.
The ligands can be small molecules or monatomic ions or polyatomic ions or more complicated organic
molecules.
In order to indicate the number of a particular ligand we use the prefixes that are used in the
nomenclature of covalently bonded binary compounds.
2
di3
tri4
tetraIf the ligand is a complex ligand (for example a larger molecule, usually an organic compound) we use
the following prefixes with parentheses around the ligand name:
2
bis3
tris4
tetrakisFor example, we would write diammine for (NH3)2 but we would write bis(methylamine) for (CH3NH2)2
to avoid confusion with dimethylamine ((CH3)2NH). Use of bis-, tris-, etc. does not require the ligand
138
name to contain a prefix as stated in the textbook.
Consecutive vowels are retained in the names (tetraammine or diaqua, for example).
The names of ligands which are anions are altered to end it “o.” –ide becomes –ido (not –o as in the
textbook), –ate becomes –ato, and –ite becomes –ito.
The names of neutral and cationic ligands (which are rare) are used in an unaltered form:
(NH2)2CO
urea
CH3CONH2
acetamide (not acetamido)
NH2CH2CH2NH2 ethylenediamine
N2H4
hydrazine
Exceptions to this are:
H2O
aqua
CO
carbonyl
SO2
sulfonyl
NH3
NO
ammine
nitrosyl
Enclosing marks (parentheses) are required for ligand names in order to avoid ambiguity. The ligands
in the list above generally do not require enclosing marks unless there is an ambiguity.
Which atom connects to the central atom is denoted with a κ (Greek letter kappa) followed by the
atom bonded to the central atom in italics. This is not required for single atom ligands or ligands where
it is obvious which atoms is forming the bond (i.e., water or ammonia). If there is more than one atom
in a ligand that bonds to the central atom in the ligand, then we use κ2 followed by the atoms forming
the bonds in alphabetical order (-κ2N,O, use a ‘ if the atoms are the same kind, i.e., - κ2N,N’). The
difference in bonding of the cyanide ion, the nitrite ion, the cyanate ion and thiocyanate ion is
indicated in this way. They are NOT nitro and nitrito for NO2-, they are nitrito-κN or nitrito-κO. For
cyanide (CN-) it’s either cyanide (the rules specify that the –κ is not required when CO or CN- bonds
via the carbon atom, via the oxygen atom in carboxylate (like acetate), or via the nitrogen atom in NO)
or cyanide-κN. For cyanate (OCN-) it’s either cyanato-κN or cyanato-κO. For thiocyanate (SCN-) it’s
either thiocyanato-κN or thiocyanato-κS. The textbook is incorrect about this. There is no iso- prefix
according to IUPAC. However, sometimes the κ may be omitted such as cyanato-O or nitrito-N.
The name of the central atom is modified with the Latin form of the element’s name and an –ate
ending when the complex ion is an anion. If the complex is neutral or a cation, the IUPAC name of the
central atom is unaltered.
Element Name
Iron
copper
gold
Tin
silver
Lead
Latin Form
ferrum
cuprum
aurum
stannum
argentum
plumbum
Anionic Complex Ion Form
ferrate
cuprate
aurate
stannate
argentate
plumbate
A couple of exceptions to this are mercury which will be mercurate instead of hydrargyrate (from the
Greek hydrargyrum –“water + silver”) and tungsten which will be tungstate instead of wolframate
(from the German wolfram –“wolf’s rahm” or “wolf’s cream”).
139
The last part of the name is either the oxidation state of the central atom, which is represented in
Roman numerals in parentheses, or the overall charge of the complex, which is represented with
Arabic numerals with the sign of the charge after the number in parentheses.
Examples:
[CoCl(NH3)5]2+
[Fe(CH3COO)3(H2O)3]
[AuBr6]5[Fe(NH2CH2CH2NH2)3]3+
pentaamminechloridocobalt(III) or pentaamminechloridocobalt(2+)
triaquatris(acetato)iron(III) or triaquatris(acetato)iron(0)
hexabromidoaurate(I) or hexabromidoaurate(5-)
tris(ethylenediamine-κ2N,N’)iron(III)
or tris(ethylenediamine-κ2N,N’)iron(3+)
This is indicating that the bonds are formed by both of the nitrogen
atoms in the ligand.
Writing Formulas of Complex Ions
The central atom (metal ion) is listed first.
Polyatomic ligands are enclosed in parentheses. Usually, the element that is forming the bond to the
central atoms is listed first in the ligand formula (even for water, for example (OH2) instead of (H2O)).
The ligand symbols are then listed alphabetically without regard to charge.
The entire complex is then enclosed in square brackets regardless if it is charged.
If there is a charge and no counter-ion is listed, the charge is a following superscript with the number
preceding the sign of the charge.
Abbreviations of complicated organic ligands may be used in formulas even though the same is
discouraged in writing names.
Examples:
tetraamminechloridonitrito-κN-cobalt(III)
amminebromidochloridonitrito-κO-platinate(1-)
bis(ethylenediamine-κ2N,N’)diiodidonitrito-κO-chromate(III)
(ONO)
[CoCl(NH3)4(NO2)]+
[PtBrCl(NH3)(NO2)]or [PtBrCl(NH3)(ONO)][CrI2(NH2CH2CH2NH2)2(NO2)]
or
or [Cr(en)2I2(NO2)] or (ONO)
(because e comes before i)
or [Cr(C2H8N2)2I2(NO2)] or (ONO)
140
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
Density
Solvent molecule:
Co2+ coordination
solution color
Solvent
-3
(g cm )
group attach to –OH
geometry
Water
1.000
Hlight pink
octahedral
methanol
0.7914
CH3light pink
octahedral
ethanol
0.7893
CH3-CH2dark blue
tetrahedral
1-propanol
0.7796
CH3-CH2-CH2dark blue
tetrahedral
2-propanol
0.7851
(CH3)2-CHdark blue
tetrahedral
When we dissolve CoCl2 in a mixture of methanol (CH3OH) and 2-propanol (CH3CH(OH)CH3) the
solution will contain two types of Co2+ complex ions: octahedral ones (with methanol) and tetrahedral
ones (with 2-propanol). 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 2-propanol and methanol molecules, respectively. Because this equilibrium
is between tetrahedral and octahedral complex ions, we can also represent it as
141
[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
Co( oct )  P

K eq = 
Co(tet )  M 3


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 2-propanol 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-tored 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 2propanol is 657 nm, so we use 657 nm as the analytical wavelength for these solutions. The
absorbance at 657 nm for a CoCl2–2-propanol 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 2-propanol, we can write Beer’s law as
A = ε b Co(tet ) 
142
We can determine the value of ε as follows. First, we measure the absorbances for various CoCl2–2propanol 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 CoCl2–2-propanol 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 2-propanol. To
do so, you will prepare several solutions, each containing a known concentration of CoCl2 in 2propanol. 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 2-propanol and methanol. To do so, you will
prepare solutions of CoCl2 in various mixtures of 2-propanol 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 2-propanol. Rinse again with
another 1 mL of 2-propanol. Also rinse a clean, dry 13×100-mm test tube twice, using about 1 mL
of 2-propanol 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.
143
4. Use your molecular model kit to construct models of the following Co2+ coordination complex ions:
(a) [CoClP3]+: one Cl- and three 2-propanols tetrahedrally arranged around a central Co2+
(b) [CoClP5]+: one Cl- and five 2-propanols octahedrally arranged around a central Co2+
Remember that the oxygen atom in the 2-propanol 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 2propanols 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.
spectrophotometer to stabilize while you do Steps 7–14.
Allow the
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 2-propanol. Collect the rinses in the “Discarded Solutions” container.
8. Fill the buret to the 0.00-mL mark with 2-propanol. 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 2-propanol 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.
14. Prepare various mixtures of your CoCl2 stock solution and 2-propanol 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 2propanol. Record the colors of the four solutions.
144
Cup
1
2
3
4
Volume from Buret Co, mL
0.50
1.00
1.50
2.00
Volume from Buret P, mL
9.50
9.00
8.50
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 2-propanol from buret P each
time. Pour the rinses into your “Discarded Solutions” container. Fill the cuvet with 2- propanol
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 2-propanol 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 2-propanol.
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.
145
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.
Volume from Buret M,
Volume from Buret P,
mL
mL
1.00
A
1.00
8.00
B
1.00
1.50
7.50
C
1.00
2.00
7.00
D
1.00
2.50
6.50
E
1.00
3.00
6.00
F
1.00
2.00
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.
Cup
Volume from Buret Co, mL
24. Prepare an ice-water bath by adding about 25 mL of water to about 100 mL of ice in a 250-mL
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.
146
147
REPORT
NAME _________________________
EQUILIBRIUM BETWEEN TWO COMPLEX IONS …
SECTION _______________________
Color
Co2+ in 2-propanol
Co2+ in methanol
Unstable Coordination Complex ions of Co2+
Too much crowding in 2-propanol structures
Too much room in methanol structures
Mass of CoCl2  6H2O in 50.00 mL 2-propanol, g
Cup
Volume of stock
CoCl2 solution, mL
Volume of 2propanol, mL
Solution color Absorbance
1
2
3
4
Cup
A
B
C
D
E
F
Volume of
CoCl2, mL
Volume of M,
Volume of P, mL
mL
Solution color
Abs
T, °C
148
REPORT (Cont’d)
NAME ________________________________
EQUILIBRIUM BETWEEN TWO COMPLEX IONS OF Co2+ IN SOLUTION
Solvent
Molecular formula of complex
Geometry of complex
Color
2-propanol
methanol
Co2+ concentration of stock CoCl2 solution, M
Cup
[Co(tet)], M
Abs
1
2
3
4
Slope = ε, M-1
Cup
[Co(total)]
Abs
[Co(tet)]
[Co(oct)]
[P]
[M]
Keq
A
B
C
D
E
F
Average Keq in cups A-E
Average temperature in cups A-E, °C
T
149
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 2-propanol and methanol.
[CoClP3]+ + 3 M  [CoClP2M3]+ + P
(a) For the five solutions of CoCl2, 2-propanol, 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.
150
QUESTIONS FOR EQUILIBRIUM BETWEEN TWO…
NAME ___________________________
(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.
2. An alternative equilibrium that could be used to explain the conversion of tetrahedrally
coordinated Co2+ in 2-propanol 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 2-propanol 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.
151
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(lI) 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 [NiCI2(NH3)n], with n representing a small whole number.
Synthesizing [NiCI2(NH3)n]
You will synthesize [NiCI2(NH3)n] by reacting nickel(II) chloride hexahydrate (NiCI2 6H2O) and NH3. This
reaction is shown below
Ni2+ (green) + 2 CI- + n NH3  [NiCI2(NH3)n] (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(lI) 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 [NiCI2(NH3)n] is also
somewhat soluble in water, you must keep the volume of water used in the synthesis to an absolute
minimum. Nickel(lI) 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 NiCI2. By
maintaining the reaction temperature at 60°C, you will maximize the [NiCI2(NH3)n] yield. Once the
reaction is complete, you will cool the reaction mixture to 0°C in an ice-water bath. Because
[NiCI2(NH3)n] 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 [NiCI2(NH3)n] is insoluble in ethanol. You will filter the [NiCI2(NH3)n]
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 [NiCI2(NH3)n], as shown in Equation 4.
Ni(OH)2 (green) + n NH3 + 2 Cl-  [NiCI2(NH3)n] (bluish purple) + 2 OH- (Eq.4)
152
Finally, you will dry and weigh the crystals to determine the actual yield of your synthesis.
Analyzing [NiCI2(NH3)n]
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 [NiCI2(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 = ε bc
(Eq. 8)
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 [NiCI2(NH3)n] 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 [NiCI2(NH3)n]
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.
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On a tared piece of weighing paper, weigh 8.0 g of NiCI2  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 NiCI2  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 icewater bath.
Assemble a second ice-water bath in a third 600-mL beaker. After the reaction in the warm-water 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.
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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 [NiCI2(NH3)n] 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 [NiCI2(NH3)n] 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 [NiCI2(NH3)n]. Transfer the sample to the "unk" beaker. Record
the mass of [NiCI2(NH3)n] 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 the
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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 NiCI2/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 NiCI2/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, [NiCI2(NH3)n],
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.
156
157
REPORT
NAME _________________________________
SYNTHESIS & ANALYSIS OF A Ni2+ COMPOUND
SECTION _______________________________
DAY ONE – SYNTHESIZING [NiCI2(NH3)n]
Mass of NiCl2  6H2O, g
Molar mass of NiCl2  6H2O, g mol-1
Moles of NiCl2  6H2O used in the synthesis, mol
Mass of capped shell vial and synthesized [NiCI2(NH3)n], g
Mass of capped shell vial, g
Mass of synthesized [NiCI2(NH3)n], g
237.69
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DAY TWO – SPECTROPHOTOMETRIC DETERMINATION OF MASS PERCENT OF Ni2+ ION IN [NiCI2(NH3)n]
Initial color and appearance of NiSO4 solution
Initial color and appearance of [NiCI2(NH3)n] Solution
Color after adding NH3 to NiSO4 solution
Color after adding NH3 to [NiCI2(NH3)n] solution
Absorbance for [Ni(NH3)n]2+ ion in solution using NiSO4  6H2O
Mass of NiSO4  6H2O used, g
Molar mass of NiSO4  6H2O, g mol-1
262.85
Concentration of standard [Ni(NH3)n]2+, M
Wavelength, nm
Abs
620
610
600
590
580
570
560
550
Approximate λmax of the [Ni(NH3)n]2+ ion, nm
Approximate λmax - 5 nm, nm
Approximate λmax + 5 nm, nm
Chosen λmax of the [Ni(NH3)n]2+ ion, nm
Abs of the standard [Ni(NH3)n]2+ solution at λmax
Absorbance of the unknown [Ni(NH3)n]2+ ion solution
Chosen Wavelength, nm
Abs of the [Ni(NH3)n]2+ ion in the synthsized solution at the chosen wavelength
Mass of [NiCI2(NH3)n] used, g
Concentration of [NiCI2(NH3)n] synthesized, M (from calculation)
APPROXIMATE molar mass of [NiCI2(NH3)n], g mol-1
REPORT
NAME _________________________________
159
DETERMINATION OF THE EMPIRICAL FORMULA AND PERCENT YIELD OF [NiCI2(NH3)n]
Empirical formula for [NiCI2(NH3)n]
*Mass of [NiCI2(NH3)n] synthesized, g
Molar mass of [NiCI2(NH3)n], g mol-1 (calculated
from empirical formula)
Moles of [NiCI2(NH3)n] synthesized, mol
*Moles of NiCl2  6H2O used in the synthesis, mol
Percent yield of [NiCI2(NH3)n], %
* these are the values from Day 1.
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QUESTIONS FOR SYNTHESIS & ANALYSIS …
NAME _________________________________
1. Based on your determined empirical formula for [NiCI2(NH3)n], 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 [NiCI2(NH3)n] complex to be paramagnetic or diamagnetic? Use the answers to
questions 1 and 2 to support your answer.
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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 polyhedron (with a hole in each
face) representing the transition metal as the central atom. 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
1. Construct four ammonia molecules:
“extra” bond position on the nitrogen.
H
These will attach to the metal atom through the
2. Construct a model of cis-diamminedichloroplatinum(II).
3. Construct a model of trans-diamminedichloroplatinum (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 ammineaquabromochloroiron(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.
5. Replace the aqua ligand with another ammine ligand in both of the models.
162
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.
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REPORT
NAME ________________________________
MOLECULAR MODELS OF TRANSITION…
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 ammineaquabromochloriron(II) match up exactly
(superimpose) with one another?
Did the mirror images of the models of diamminebromochloroiron(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 cis-dichlorobis(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-diamminedichloro
(ethylenediamine)cobalt(II) match up exactly (superimpose) with one another?
Which of these compounds exist as enantiomers if any?
164
165
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.
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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.
167
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 “straight-line” 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 y-intercept, 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
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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
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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 (yaxis) 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.
170
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 y-axis 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 chart title, where it currently says “y data,” you can
171
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 doubleclicking 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.
172
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:
173
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:
174
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
175
“Axis Options.” The first section here is “Bounds.” 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
Wavelength (nm)
At this point, you can click on a blank area of the graph and press Ctrl-P you can print it.
880.0
176
APPENDIX C
BALANCING REDOX REACTIONS USING THE HALF-REACTION METHOD
There are many methods that can be used when balancing chemical reactions that involve oxidationreduction. 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 half-reaction.
7. Multiply each half-reaction by the factors needed to make the electrons in each half-reaction
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.
177
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:
N2H4 + Pu2O3 → N2O + Pu(OH)2
1.
2
3.
4.
(in base)
N2H4 + Pu2O3 → N2O + Pu(OH)2 (in OH- and H2O)
-2 +1
+3 -2
+1 -2
+2 -2 +1
-2 +1
N2H4 → N2O
Nitrogen is balanced
N2H4 → N2O

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
+1 -2
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 side reactant side and 2 OH- to the product 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 OHNote: 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 OH7.
8.
Multiply equation above by 1
Multiply the equation above by 3
6 OH + H2O + N2H4 → N2O + 6 H2O + 6e6 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.
178
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.
179
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
+6 -2
3 e- + 4 H2O + KMnO4  MnO2 + K+ + 2 H2O +4 OH8OH- + 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.
180
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.

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