Idea Transcript
CHEM 122L General Chemistry Laboratory Revision 2.0
Determination of the Rate Constant for an Iodine Clock Reaction
To learn about Integrated Rate Laws. To learn how to measure a Rate Constant. To learn about Clock Reactions.
In this laboratory exercise, we will measure the Rate Constant k for the oxidation of Iodide (I-) by Peroxysulfate (S2O82-). 2 I-(aq) + S2O82-(aq)
I2(aq) + 2 SO42-(aq)
(Eq. 1)
In this experiment, the reaction is Clocked by a secondary Clocking reaction, which consumes the product I2 as soon as it is produced and which triggers a color change due to the persistent presence of I2 when the Clocking Reagent itself is completely consumed. We will use Thiosulfate (S2O32-) as the Clocking Reagent because it can reduce the I2 back to I-: 2 S2O32-(aq) + I2(aq)
S4O62-(aq) + 2 I-(aq)
(Eq. 2)
As long as any Thiosulfate Ion is present, none of the Iodine produced in (Eq. 1) remains; it is consumed as quickly as it is produced. (This, of course, requires that (Eq. 2) proceed very rapidly; which is in fact the case.) As soon as the Thiosulfate is used up, Iodine will begin to appear in the solution. The presence of Iodine is then dramatically detected because of the formation of a Starch-Iodine molecular complex which is deep blue in color. (Actually, the starch molecules complex with I3- which is formed by a reaction between I2 and I-.) When the deep blue color appears the reaction is said to “Clock”. We can set the Clock Point by simply adding differing amounts of Thiosulfate Ion to the reacting solution. The Rate at which a chemical reaction proceeds is typically influenced by the concentration of each reactant present and by the temperature of the reaction vessel. And, typically, the relationship between the Reaction Rate and Reagent Concentration takes a simple form known as the Rate Law: Rate = k [A]n [B]m
(Eq. 3)
Here A and B are generic reacting Species, k is a reaction specific proportionality constant known as the Rate Constant, and n and m are the Reaction Order. The Rate Law parameters k, n and m must be determined experimentally. In our case, (Eq. 1), it has been determined experimentally that the reaction kinetics is MixedSecond Order:
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Rate = k [I-] [S2O82-]
(Eq. 4)
For reasons that will become apparent, we will use a vast excess of Iodide during the reaction. Under these conditions, the Iodide concentration [I-] will barely change over the course of the reaction and we can say: [I-] ~ [I-]o
(Eq. 5)
which means we can re-write our Rate Law as:
Rate = k [I-]o [S2O82-]
(Eq. 6)
This allows us to define a new Rate Constant k’ as: k’ = k [I-]o
(Eq. 7)
And, under the experimental constraint that [I-] is very large, the Rate Law for our reaction becomes First Order:
Rate = k’ [S2O82-]
(Eq. 8)
(Eq. 8) is said to be a Pseudo-First Order Rate Law with a Pseudo-First Order Rate Constant k’,
because the large [I-] constraint means the reaction is only behaving as though it is first order. Keep in mind, it is k and not k’ we are after. However, if we can measure k’ using a known Iconcentration, then we can use (Eq. 7) to determine k. We can measure k’ by obtaining [S2O82-] vs. Time data and plotting this data according to the Integrated form of the First Order Rate Law: ln [S2O82-] = ln [S2O82-]o - k’ t
(Eq. 9)
The slope of this plot will yield k': slope = - k'
(Eq. 10)
We can obtain [S2O82-] vs. Time data fairly easily by varying the amount of Thiosulfate used in the Clocking Reaction, (Eq. 2). The amount of S2O82- remaining at the Clocking Point is related to the amount of Thiosulfate used by the reaction stoichiometry of (Eqs. 1 and 2):
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At the Clocking Point, the number of moles of Peroxysulfate (S2O82-) which has reacted is exactly 1/2 the number of moles of Thiosulfate (S2O32-) orginally added. As an example, suppose we mix the following: 5mL 7mL 5mL 1mL ____ 18mL
of of of of
1M KI 0.2M Na2S2O3 0.15M K2S2O8 Starch Sol'n
total
Then, at the Clocking Point, when the blue Iodine-Starch complex appears: # mmoles S2O32- added
= ( 0.2M ) x ( 0.007 L ) x ( 1000 mmole / 1 mole ) = 1.40 mmole
# mmoles S2O82- consumed
= 1.4 mmole S2O32- x ( 1 mmole I2 / 2 mmole S2O32- ) x ( 1 mmole S2O82- / 1 mmole I2 ) = 0.70 mmole
# mmole S2O82- remaining
= ( 0.15M ) x ( 0.005 L ) x ( 1000 mmole / 1 mole ) - 0.7mmole = 0.75 mmole - 0.70 mmole = 0.05 mmole
[S2O82-] remaining
= 0.05 mmole / 18 mL = 0.00278
ln ( [S2O82-] ) remaining
= ln ( 0.00278 ) = -5.886
Thus, by arranging for the reaction to run under Pseudo-First Order conditions, and by varying the Clocking Reagent (Thiosulfate Ion) concentration, we can generate [S2O82-] vs time data and thus determine first k' and then k for (Eq. 1).
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Pre-Lab Safety Questions 1.
Why are Strong Oxidizing Agents dangerous? Examine (Eq. 1) and (Eq. 2) and identify the Oxidizing Agents. Are these Oxidizing Agents considered "Strong Oxidizing Agents"? (You may wish to consult an SDS.)
2.
When purchased at the pharmacy, Iodine comes in the form of a tincture. What is the Molar concentration of a typical pharmaceutical tincture? How does that compare to Iodine concentrations we will be using in this laboratory exercise? (A rough approximation is all that is needed here.) Should you use gloves during this exercise? Explain.
3.
What about goggles? Explain.
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Procedure 1.
Obtain 4 small beakers and label them: Label KI Na2S2O3 K2S2O8 Starch Water
Amount ~40 mL ~40 mL ~40 mL ~10 mL ~40 mL
Obtain the indicated amount of each solution. 2.
Place six 15 cm test tubes in a test tube rack. Label them 1 - 6. Add, by means of a Mohr pipet, the amounts of the reagents listed:
Tube #
mL 1M KI
1 2 3 4 5 6
5 5 5 5 5 5
mL of 0.2M Na2S2O3 7.0 6.5 6.0 5.0 4.0 2.0
mL Water
mL Starch
0 0.5 1.0 2.0 3.0 5.0
1 1 1 1 1 1
3.
To tube #1, add 5.0 mL of 0.15 M Potassium Peroxysulfate (S2O82-) by pipet. Record the time to the nearest second when addition is begun. As soon as delivery is complete, stopper the test tube and INVERT the test tube at least 5 times. (If the system is not mixed well, only part of the solution will “clock” and you will have to repeat that run.) Place this in a large beaker filled with Room Temperature water.
4.
Repeat this for test tubes #2 - #6. Start each reaction no more than one minute after the preceding one.
5.
Record the exact time when the blue color appears; i.e., the Time at which the reaction Clocks.
6.
Record the temperature of the Water bath.
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Data Analysis 1.
Prepare a Table of the following form and fill-in the needed Concentration values.
Tube #
mL Na2S2O3
mmoles S2O32added
mmoles S2O82consumed
mmoles S2O82remaining
[S2O82-] remaining
ln [S2O82-] remaining
elapsed time
1 2 3 4 5 6
2.
Prepare a graph of ln [S2O82-] vs. Time using a software package such as Excel. Use the software package’s Linear Least Squares analysis routine to determine the Slope of the line.
3.
Determine the Pseudo-First Order Rate Constant k’ for the Peroxysulfate oxidation of Iodide.
4.
Calculate the Rate Constant k for the Peroxysulfate oxidation of Iodide; see (Eq. 7).
Note: You must provide appropriate Units of Measurement for k’ and k.
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Post Lab Questions 1.
Why is it important that our test tubes be placed in a Water bath? Be thorough in your answer.
2.
In theory, we could perform a single run for this experiment and use the equation: ln([S2O82-]) = ln([S2O82-]o) - k' t to determine k', and hence k. Why have we gone to the trouble of performing this experiment several times, preparing a graph of the results, and determining the slope of the graph in order to determine k’. Be specific.
3.
The chemical reaction of (Eq. 1) is written in Net Ionic form. If Potassium Ion (K+) is the only Spectator Ion present, write the Total Molecular Equation for this reaction.
4.
Write the Lewis Structure for the Thiosulfate Ion (S2O32-) used in the Clocking reaction (Eq. 2) . Hint: Thiosulfate is the “thiol” (Sulfur based) analog of Sulfate (SO42-). So, start by working out the Lewis Structure for the Sulfate Ion.
5.
The Clocking reaction (Eq. 2) results in the Reduction of I2 to I-. Show the Sulfur is oxidized by explicitly assigning Oxidation Sates to the S in each relevant species.