CHEM 203 Wilbur Wright College Rate Law of the Iodine Clock Reaction Lab Report

  

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Experiment 1: Iodine Clock Reaction (Rate Law)
Chemistry 203 – General Chemistry II
Dept. of Physical Sciences & Engineering
Wilbur Wright College
Experiment 1: Iodine Clock Reaction (Rate Law)
Introduction
In this experiment, you will study the kinetics of the following reaction
2−
−
S2 O2−
8 + 2I → 2SO4 + I2
persulfate iodide
sulfate
iodine
The rate law will be determined by investigating the effect of the concentration of the reactants on the rate of
the reaction. Given the general reaction
aA + bB → Products
The dependence of the rate of the reaction on the concentration of the reactants may be expressed by a rate
law of the form
rate = k[A]𝑚 [B]𝑛
where, k, is the rate constant; 𝑚 and 𝑛 are the orders of the reaction with respect to the reactants A and B,
respectively; and the sum of 𝑚 + 𝑛 is the overall reaction order. Unlike the stoichiometric coefficients
determined by calculation, the orders of the reaction are determined from the experimental data based on the
kinetics of the reaction. The orders within the experimentally determined rate law are explained by a plausible
reaction mechanism, which is an account of the sequential molecular events known as elementary reactions.
The rate law determined or derived from the slow-step (also known as the rate-determining step) of the reaction
mechanism must match the experimentally determined rate law.
Description of the “Clock Reaction”
Determining the concentration of reactants or products at any instant is often difficult to measure directly. As
such, indirect measurements are often utilized. The “clock reaction” is a reaction known for its dramatic
colorless-to-blue/black color change that occurs when iodine (I2 ) reacts with starch to form an iodine-starch
complex. Unfortunately, as soon as any iodine (I2 ) is produced, it will instantaneously react with starch to form
the blue/black complex. Thus, the starch only indicates that iodine (I2 ) is produced, but it doesn’t reveal the
concentration and time data needed to calculate the reaction rate. However, if S2 O2−
3 (thiosulfate) is also added
to the mixture a side reaction will take place (2). This side reaction will remove the initial I2 produced by the
primary reaction (1) and will allow time measurements to be made before the solution turns blue/black.
The primary reaction is the oxidation of I- by S2 O2−
8 in aqueous solution:
2−
−
S2 O2−
8 + 2I → 2SO4 + I2
persulfate iodide
sulfate
(slow, rate determining)
(1)
iodine
This reaction will be run in the presence of a known amount S2 O2−
3 , which reacts very rapidly with I2 , and starch.
2−
2−
As long as S2 O3 is present, I2 is consumed by S2 O3 as fast as it is formed, which prevents I2 from reacting
with starch. No color change is observed until S2 O2−
3 is completely used, which allows the dark blue iodinestarch complex to form.
2−
−
2S2 O2−
3 + I2 → S4 O6 + 2I
(fast)
(2)
(fast)
(3)
thiosulfate iodine tetrathionate iodide
I2 + (C6 H10 )5 )n ⦁ H2 O → blue complex
iodine
starch
1
Experiment 1: Iodine Clock Reaction (Rate Law)
Chemistry 203 – General Chemistry II
Dept. of Physical Sciences & Engineering
Wilbur Wright College
The color change signals when the reaction (1) produces the specific amount of I2 required to completely react
2−
with all of the S2 O2−
3 via reaction (2). Since the concentration of S2 O3 initially present for reaction can be
determined, the concentration of the I2 produced in reaction (1) can be calculated using the reaction (2)
stoichiometry.
[I2 ]Produced from Rxn 1 = (1/2)[S2 O2−
3 ]i
(4)
Since the I2 has an initial concentration of zero and a final concentration equal to [I2 ]Produced , the change in
the concentration of I2 is equal to:
∆[I2 ] = [I2 ]final − [I2 ]inital = [I2 ]Produced from Rxn 1 − 0
∆[I2 ] = (1/2)[S2 O2−
3 ]i
(5)
Further, since the initial time is zero and the final time (t f ) is the amount of time for the blue/black color to
appear, the change in time is equal to:
∆t = t f − t i = t f − 0
∆t = t f
(6)
Since the rate of reaction can be expressed as the change in concentration of a reactant or product over the
change in time, the rate of reaction (1) can be calculated from the rate of I2 production shown in equation (7).
Reaction Rate =
Reaction Rate =
∆[I2 ]
∆t
(1/2)[S2 O2−
3 ]i
tf
(7)
Initial Rate Method
In the experiment, you will use the initial rate method to determine the rate law for the reaction of persulfate
and iodide.
2−
−
S2 O2−
8 + 2I → 2SO4 + I2
persulfate iodide
sulfate
iodine
Without experimental data, the rate law takes the following form:
𝑚 − 𝑛
rate = k[S2 O2−
8 ] [I ]
To use the initial rate method, you will generate data by using measurements and performing calculations to
complete Table 3 in the lab report. The reaction rate for each experiment will be calculated using equation (7).
The initial concentrations of reactants are calculated at the moment they are mixed. At that moment, the
solutions have mutually diluted each other (by raising the volume of total solution) but have not started reacting.
For each ion in solution, a new molarity (M2 ) must be calculated that takes into consideration the new total
volume of the solution (V2 ). The concentrations and volumes of the reactants are given in Table 1.
2
Experiment 1: Iodine Clock Reaction (Rate Law)
Chemistry 203 – General Chemistry II
Dept. of Physical Sciences & Engineering
Wilbur Wright College
2−
−
The dilution equation (8) is used to calculate the new molarity (M2 ) for each substance (S2 O2−
3 , S2 O8 and I ) in
experiments 1-3 upon mixing.
Note: The [reactant]i in the Table 3 equals the M2 calculated for each reactant.
M1 V1 = M2 V2
(8)
Preliminary Calculations
1. Using the concentration and volume information provided in Table 1 calculate M2 (the new
concentration after mixing, but before reaction) using the dilution equation (8) for the following
substances in experiments 1-3:
a) S2 O2−
3 (Hint: It will be the same for all three experiments.)
b) S2 O2−
8
c) I−
Table 1: Composition of the Reaction Mixtures
Exp.
H2 O
0.005 M Na2 S2 O3 w/starch
V1 for S2 O2−
3
1.0 M KI
V1 for I−
0.1 M (NH4 )2 S2 O8
V1 for S2 O2−
8
Total Volume
(V2 )
1
48 mL
10 mL
12 mL
30 mL
100 mL
2
63 mL
10 mL
12 mL
15 mL
100 mL
3
54 mL
10 mL
6 mL
30 mL
100 mL
2−
−
2. Record the new concentrations (the calculated M2 values) for S2 O2−
3 , S2 O8 and I in Table 3 of the lab
report.
3. Show the calculation of M2 (the new concentration after mixing, but before reaction) using the dilution
equation (8) for each of the three solutions combined in the Exp. 3 mixture. Show your work for these
three calculations (where indicated) in your lab report.
IMPORTANT – READ BEFORE STARTING THE EXPERIMENT
1. Download and save the file to your computer. If this step is skipped, any data that is inputted into the
PDF file will not be saved. Open the saved copy and start the report.
2. Each student will submit individual lab report. Questions 1, 2 and 4 in the lab report should be written
on separate paper/file. Clearly indicate which question is being answered, otherwise no credit will be
given. If done on paper, written calculations can be submitted by either scanning the pages or taking a
clear photo. Instructor recommends downloading a phone app, such as “CamScanner”, that fixes and
converts photos into a single PDF file (similar to computer taking the scanned pages and converting them
into a PDF file).
3. This lab report (PDF format) and any additional files should be submitted to Brightspace > Assignments
> Lab 1, by indicated due date.
3
Experiment 1: Iodine Clock Reaction (Rate Law)
Chemistry 203 – General Chemistry II
Dept. of Physical Sciences & Engineering
Wilbur Wright College
Simulation
1. To view the simulation of this reaction use the following link:

2. The time data for three experiments is provided in Table 2.
Calculations – Reaction Rate (M/s)
1. Use equation (7) to calculate the reaction rate for experiments 1-3.
Reaction Rate =
∆[I2 ] (1/2)[S2 O2−
3 ]i
=
∆t
tf
2. Record the reaction rates in Table 2 and Table 3 for experiments 1-3.
3. Apply the initial rate method to the data collected in Table 3 and find the orders of 𝑚 and 𝑛 in the
following rate law. (Hint: Select two sets of experiments in which only one of the reactant
concentrations change. Then, write a ratio of rate laws placing the experiment with the larger rate in
the numerator.)
2−
−
S2 O2−
8 + 2I → 2SO4 + I2
𝑚 − 𝑛
rate = k[S2 O2−
8 ] [I ]
4. Demonstrate (in sentences or by calculation) how 𝑚 and 𝑛 were determined. Show your work in the
space provided in the lab report.
5. Complete the remaining questions pertaining to the rate law determination in your lab report.
Data and Questions
1. Show the calculation of M2 (the new concentration after mixing, but before reaction) using the dilution
equation (8) for each of the three solutions that were combined in the Exp. 3 mixture.
a) Calculation of M2 for [S2 O2−
3 ]i
b) Calculation of M2 for [S2 O2−
8 ]i
c) Calculation of M2 for [I − ]i
2. Show the calculation of the reaction rate using equation (7) for Exp. 3.
3. Complete Table 2 and Table 3.
Table 2: Reaction Time & Reaction Rate
Exp.
Time (s)
1
18 s
2
36 s
3
36 s
Rxn Rate (M/s)
4
Experiment 1: Iodine Clock Reaction (Rate Law)
Chemistry 203 – General Chemistry II
Dept. of Physical Sciences & Engineering
Wilbur Wright College
Table 3: Initial Rate Method
Exp.
[S2 O2−
3 ]i (M)
[S2 O2−
8 ]i (M)
[I− ]i (M)
Rxn Rate (M/s)
1
2
3
4. Demonstrate (in sentences or by calculation) how 𝑚 and 𝑛 were determined.
a) Determination of 𝑚.
b) Determination of 𝑛.
5. Fill in the blanks to answer the questions related to the determination of the rate law for the following
reaction.
2−
−
S2 O2−
8 + 2I → 2SO4 + I2
𝑚 − 𝑛
rate = k[S2 O2−
8 ] [I ]
a) To determine the order of ⌈S2 O2−
8 ⌉, 𝑚:
i. Which two experiments were compared?
______ and ______
ii. From one experiment to the other, the [S2 O2−
8 ]i is multiplied by _______.
iii. From one experiment to the other, the rate is multiplied by _________.
iv. Thus, the reaction is _________ order with respect to [S2 O2−
8 ]i and 𝑚 = ______.
b) To determine the order of [I − ], 𝑛:
i. Which two experiments were compared?
_______ and _______
ii. From one experiment to the other, the [I− ]i is multiplied by _________.
iii. From one experiment to the other, the rate is multiplied by __________.
iv. Thus, the reaction is __________ order with respect to [I− ] and 𝑛 = _______.
c) Write the experimentally determined rate law for the following reaction:
2−
−
S2 O2−
8 + 2I → 2SO4 + I2
___________________________
Student Name: __________________________
5
Date: __________________

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