Iodine Clock Reaction Lab Report

Hi, I need help with my chemistry assignment. Please follow the lab format and organize all work in order. Please make sure you attach the graph in the lab as well. Read all of the lab instructions pdf and follow it carefully. I will also send you more links that might be helpful to solve.

A) Background: The Iodine Clock Reaction
A little background about how we got the rate
Use this website to run the simulation to collect the data in the form of time that the
reaction occurs
https://web.mst.edu/~gbert/IClock/Clock.html
FORMAT:
Orgnize it in order
1. A title page showing your name and the experiment title
2. Data table showing experimental conditions and experimental observations (based
on Table 2)
3. Results table showing experimental conditions for each run and calculated values of
initial rates. Show at least one example of calculation work. (Section F step 1)
4. Organized calculations showing how you determined the reaction order for IO3 – ,
Iand H+ (section F step 2)
5. Organized calculations showing how you determined the average value of the rate
constant, k. (Section F step 3)
6. Results table showing experimental conditions where temperature is varied and
calculations for average initial rate, value of the rate constant k, ln(k), temperature T in
C and K, and 1/T (K-1 ) (as described in Section F step 4)
7. Graph of ln(k) vs (1/T). Fully annotated as described in Section F step 5)
8. Answers to the post-lab questions shown on the next page.
Fill in the table and show 1 work not all
Runs 1 through 4 = 1+2+3+4
Runs 1&5 through 8 = 1+5+6+7+8
***
Make sure you do the graph and attach it
Determination of the Rate Law and Activation Energy for the Iodine Clock Reaction.
A) Background: The Iodine Clock Reaction
“Iodine Clock” refers to a group of reactions which involve the mixing of two colorless solutions
to produce a solution which remains colorless for a precise amount of time, then suddenly changes
to a deep purple-blue color. The time is controlled by the temperature and/or the concentrations of
the reactants.
The reactions involve the oxidation of iodide ion (I-) to dissolved iodine (I2) or tri-iodide ion (I3-).
Either of these combine with starch indicator to produce the characteristic purple-blue color. This
color is visible to the eye when the concentration of iodine or tri-iodide ion exceeds 10-5
moles/liter.
A typical reaction is: 6 H+ + IO3- + 8 I-  3 I3- + 3 H2O
(Eq. 1)
The rate of this reaction depends on the temperature, and on the concentrations of iodate (IO3-),
iodide (I-), and hydrogen ions (H+). This reaction alone does not give very impressive delays and
color changes. The time delay until the appearance of the blue color is inversely related to the rate
of the reaction (the faster the reaction, the shorter the delay) but the color development is directly
related to the rate (a sharp change in the color requires a moderately fast reaction). Consequently,
if the time delay is more than a second, the color development appears relatively slow.
In order to obtain a time delay of a few seconds to a few minutes with a reasonably sharp color
development, a measured amount of a reducing agent (arsenous acid, H3AsO3, or thiosulfate ion,
HSO3-) is included in the mixture. These react very quickly with tri-iodide ion (I3-), and very slowly
with iodate ion, removing the tri-iodide ion as quickly as it is produced, so that the concentration
does not reach the visible level until all of the reducing agent is consumed:
H3AsO3 + I3- + H2O  HAsO42- + 3 I- + 4 H+
(Eq. 2)
While the reducing agent is present, the net reaction is:
IO3- + 3 H3AsO3  I- + 6 H+ + 3 HAsO42-
(Eq. 3)
This reaction has been used extensively for laboratory experiments involving investigation of the
rate expressions in Chemical Kinetics. The concentration of hydrogen ion is maintained by using
Acetic Acid/Sodium Acetate buffers in the pH range 4 -5. The initial concentration of arsenous
acid is less than 1/3 of the initial concentrations of iodate and iodide ions, so that these
concentrations change by less than 10% before the reducing agent is consumed and the sharp color
change to blue is observed.
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B) Protocol for Data Collection
Solutions are prepared with the iodate ion, starch, arsenous acid, and buffer solution in one flask
and a potassium iodide solution in another flask. These solutions are equilibrated at a known
temperature in a constant temperature bath. A timer is started as the solutions are mixed, and
stopped when the blue color appears. Times are measured for a series of mixtures with the same
initial concentration of arsenous acid while varying the initial concentrations of iodate, iodide, and
hydrogen ion and possibly the temperature. Typically, one of the initial concentrations will be
changed by a factor (2, 3, or 4) while the other two initial concentrations are held constant. The
process is then repeated for each of the other two components.
A typical experiment may consist of a set of “runs” with different initial concentrations of reactants
as follows ([A]o indicates the initial concentration of component A at the start of the reaction). The
different conditions are chosen so that pairs of runs are created where one reactant concentration
changes (typically by a factor of 2 or 3 or 4) while the other reactant concentrations remain
constant. This variation is created for each reactant so that the reaction order for each reactant can
be determined by the initial rates method.
Run # [IO3-]o [I-]o [H+]o
time to blue color, tc (sec)
1
0.005 0.05 2 × 10-5 t1
2
0.010 0.05 2 × 10-5 t2
3
0.005 0.10 2 × 10-5 t3
4
0.005 0.05 4 × 10-5 t4
Table 1: Reactant Concentrations in Experimental Runs
Initial Rate
(1/3) [H3AsO3]o / t1
(1/3) [H3AsO3]o / t2
(1/3) [H3AsO3]o / t3
(1/3) [H3AsO3]o / t4
C) Determining Rate Law by Initial Rates Method
The reaction being studied is 6 H+ + IO3- + 8 I- -> 3 I3- + 3 H2O
The initial rate of can be calculated as the change in concentration of triiodide (I3-) over time:
rate = (1/3) [I3-]/t
Triiodide reacts with starch to produce a dark blue solution. However, arsenous acid reacts quickly
and completely with triiodide in the reaction (Eq. 2), which removes triiodide until all the arsenous
acid has reacted. At that point, triiodide reacts with the starch and the solution turns blue. Since
the stoichiometric ratio of triiodide to arsenous acid in Eq. 2 is 1:1, we know that when the solution
turns blue, the number of moles of triiodide that has been formed is the same as the number of
moles of arsenous acid that were added to the reaction. Therefore:
rate = (1/3) [I3-]/t = – (1/3)[H3AsO3]/t
or rate = -(1/3) ([H3AsO3]c – [H3AsO3]o)/(tc – to)
where [H3AsO3]c and [H3AsO3]o are the concentration of arsenous acid at time tc and to
respectively. to is the time the solutions are mixed (0 sec) and tc is the time the solution turns blue.
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Since [H3AsO3]c = 0 (all the arsenous acid has reacted when the solution turns blue) then:
rate = -(1/3) (0 – [H3AsO3]o)/(tc – 0) = (1/3) [H3AsO3]o / tc
Therefore, the initial rate is approximated from the initial concentration of arsenous acid,
[H3AsO3]o, and the time taken from mixing to the solution turns blue (tc):
initial rate = (1/3) [H3AsO3]o / tc
This initial rate can be therefore be calculated for each “run” as shown in the Table 1 above based
on the observed time from mixing to the color change. For precision purposes and to minimize
experimental errors, each “run” may be repeated several times and the average value for the initial
rate is calculated.
The initial rate of this reaction is related to the initial concentrations of the reactants through a
general rate law expression involving a rate constant (k) which depends only on the temperature:
rate = k[IO3-]x[I-]y[H+]z
The goal is therefore to determine the reaction order for iodate ion (x), for iodide ion (y) and for
hydrogen ion (z) and the to determine the value of the rate constant, k. This will then allow the
rate law for the reaction to be stated explicitly.
Each reaction order can be determined mathematically by ratio of different experimental runs. For
example, comparison of the rate law expressions for run 2 and run 1 from the table above leads to:
Initial rate run 2 = k[IO3-]x[I-]y[H+]z
Initial rate run 1 k[IO3-]x[I-]y[H+]z
Substituting the appropriate initial reactant concentrations for run 1 and run 2 yields:
Initial rate run 2 = k[0.010]x[0.05]y[2.0×10-5]z
Initial rate run 1 k[0.005]x[0.05]y[2.0×10-5]z
from which the value of “x”, the reaction order for IO3- can be isolated and calculated.
Similar analysis of appropriate runs in the table allow the value of “y”, the reaction order for I-,
and “z”, the reaction order for H+ to be calculated.
Reaction orders are usually whole numbers (0, 1 or 2 are common) but fractions are possible. The
calculations for reaction order are therefore interpreted so that whole number or appropriate values
for reaction order are chosen. For example, if the reaction order is calculated to be 1.9 it would be
rounded to 2. If it were calculated to 0.48, it would be rounded to 0.5.
Once the reaction orders have been determined, the experimental data can be used to calculate the
value of the rate constant, k, for each experimental condition.
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For example, for run 1 in the table above: Initial rate run 1 = k[0.005]x[0.05]y[2.0×10-5]z
Therefore, k = ___initial rate run 1____
[0.005]x[0.05]y[2.0×10-5]z
Similar calculations for k can be performed for each reaction run, substituting the appropriate
values for concentrations and x, y and z. The average value of k can then be calculated.
At this point, all the information necessary is available to explicitly state the rate law for the
reaction in the form:
rate = k[IO3-]x[I-]y[H+]z where k = ___ unit at ___ C
In this expression, x, y and z are replaced by the determined order numbers, and the value of k,
with appropriate unit, and the temperature of the experiment are written in the spaces.
D) Temperature Effects: Determining the Activation Energy of the Reaction
To study the effect of temperature on this reaction, a series of measurements are made at different
temperatures with identical initial concentrations of all of the components. For example, run 1
shown in the table above, where [IO3-] = 0.005 M, [I-] = 0.05 M, [H+] = 2.02.0×10-5 M, may be
repeated at 5 C, 15 C, 25 C, 35 C and 45 C and the time from mixing to the color change (tc)
is recorded for each. As indicated previously, the initial rate of the reaction at each temperature
can be calculated: initial rate = (1/3) [H3AsO3]o / tc
Again, for precision purposes, each temperature measurement may be repeated several times and
the average initial rate calculated for each.
Since the rate law for the reaction is known, the value for k for at each temperature can be
calculated from:
k = ______Initial rate_____
[0.005]x[0.05]y[2.0×10-5]z using the previously determined values for x, y and z.
The Arrhenius equation relates the rate constant, k to the Activation energy (Ea), the Arrhenius
constant (A), absolute temperature (T, in K), and the universal gas constant (R = 8.314 J/mol.K):
k = A e (- Ea / RT)
or in logarithmic form:
ln(k) = ln(A) – (Ea / RT)
A graph of ln(k), vs the reciprocal of the absolute temperature (1/T) should be a straight line with
slope = – (Ea / R) and y-intercept = ln(A). Thus, the temperature data can be transformed to
calculate ln(k) and 1/T for each run (remembering to convert T from C to K) and can then be
plotted as ln(k) on the y-axis vs (1/T) on the x-axis.
A best fit line can be fit to the data, from which the slope and intercept can be used to calculate the
Arrhenius constant A, and the value for the activation energy Ea (usually in kJ/mol).
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E) Online Lab Experiment Instructions: Collecting Experimental Data
1. To collect experimental data for analysis go to the following link in a web browser:
http://web.mst.edu/~gbert/IClock/IClock.html
This site called “The Iodine Clock Reaction: A Simulated Experiment” contains an online
simulation of the iodine clock reaction. It was created by Prof. Gary Bertrand of the Department
of Chemistry at Missouri University of Science and Technology (https://chem.mst.edu/) and we
thank him for our use of this simulation.
You will see an animation on this page showing what happens when the solutions for the iodine
clock reaction are mixed as described earlier.
2. Click on “Experiment” above the animation to go to the simulation set up. This brings you to a
page where you can prepare the solutions for mixing in each experiment.
On the left is a panel where you can select appropriate stock solutions of each component in
solution A and solution B. Solution A contains KIO3, a buffer that controls the H+ concentration,
the arsenous acid and starch. Solution B contains KI. The temperature for the experiment can
also be selected.
On the right is a panel that will show you the initial concentration of each component in the
reaction. Notice that 50 mL of solution A and 50 mL of solution B are mixed to initiate the
reaction. So the initial concentration of each component in the reaction is calculated from M1V1
= M2V2. For example, if a stock concentration of [KIO3] = 0.010 M is selected, the initial
concentration of IO3- is 50 mL × 0.010 M / 100 mL = 0.005 M.
3. Use the dropdown menus in Solution A for [KIO3] and buffer pH to select the appropriate
initial concentrations of IO3- and H+. To change the concentration of H+ you need to change the
buffer pH to an appropriate value. You can check if your selection is appropriate by looking in
the panel on the right that shows the initial concentrations.
4. Use the dropdown menu in Solution B for [KI] to select the appropriate initial concentrations
of I-. Again, check your selection by looking at the panel on the right.
5. Once you have selected the appropriate settings to prepare the solutions, click the “Mix the
Solutions” link near the top left of the page. This brings you to a video that will show what will
happen when the solutions are mixed. When you are ready, click “start” above the image.
You will see the solutions mixed and placed on a stirrer (a magnetic stirrer is in the reaction
vessel to continuously mix the reaction as it proceeds). The timer is automatically started at the
moment that solution A and B are first mixed.
Watch the reaction and click “stop timer” as soon as you see any appearance of blue color.
Record the time (tc, sec). Also make any observations of what you saw.
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If you mess up you can repeat any reaction condition by clicking the “ReSet” above the image.
This will create two new solution A and B with the same composition so that you can repeat the
same experimental condition.
6. When you are done with a particular reaction condition, click “prepare solutions” to go back to
the page to select new reaction conditions.
Your goal is to collect data for each of the following conditions (“runs”):
Run # [IO3-]o [I-]o [H+]o
[H3AsO3]o Temperature (C) time, tc (sec)
1
0.005 0.05 2 × 10-5 0.0015
25
-5
2
0.010 0.05 2 × 10 0.0015
25
-5
3
0.005 0.10 2 × 10 0.0015
25
4
0.005 0.05 4 × 10-5 0.0015
25
-5
5
0.005 0.05 2 × 10 0.0015
5
6
0.005 0.05 2 × 10-5 0.0015
15
-5
7
0.005 0.05 2 × 10 0.0015
35
-5
8
0.005 0.05 2 × 10 0.0015
45
Table 2: Data Collection Conditions for Experiment
Observation
Following the instructions given above, set up each run, start the reaction and record the time taken
for the blue color to appear and any observations. Use the ReSet button to repeat each run at least
3 times (get at least 3 good measurements where the time taken appears consistent) so that you can
eventually calculate an average initial rate for each condition. That means you’ll repeat the
experiment at least 8 runs x 3 repeats/run = 24 times.
To record your data, draw up a data table like Table 2 shown above but give yourself room to
record the times and observations of each reaction (in ink!). You can draw it by hand (neatly, use
a ruler for straight lines) on a piece of paper or you can create a table in an electronic document
(such as MS-Word or MS-Excel)
F) Data Analysis
1. For each run calculate the average time taken from your three (or more) measurements for the
run. Following the instructions in section C, use the average time taken to calculate the average
initial rate for each run:
Average initial rate = (1/3) [H3AsO3]o / tc (where tc is the average time for blue color to appear)
2. Organize the results of these calculations into an appropriate table showing initial reactant
concentrations and the average initial rate for each run. Remember to include units!
3. Use the data for run 1 through run 4 to determine the rate law for the iodine clock reaction as
described in section C above. First determine the order for each reactant (IO3-, I- and H+) and
then determine the average value of the rate constant, k.
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4. Use the data for run 1 and run 5 through run 8 to determine the activation energy and Arrhenius
constant for the reaction. Create a data table with columns showing values for each run as follows
(include appropriate units):





average initial rate,
value of the rate constant k
ln(k)
temperature T in C and K
1/T (K-1)
5. Create a plot of ln(k) vs (1/T) from this table. ln(k) should be on the y-xis, (1/T) on the x-axis
Use a graphing program like Microsoft Excel to do this. Plot your data as an x, y scatter plot. Add
a best fit straight line and show the equation and R-value of the best fit line on your graph.
Remember to include a graph title, and include appropriate axis labels (for x and y axes).
If you need help creating a graph in Excel this link has good step by step instructions:
https://dataprivacylab.org/courses/popd/lab2/ExcelLine.pdf
You can also refer to the MS-Excel file from Heat of Vaporization of Water experiment for an
example.
6. Use your graph to calculate the activation energy (Ea) and determine the value of the
Arrhenius constant (A) as described in Section D above.
G) Preparation of Lab Report
When you have completed the experiment you will be able to state the rate law for iodine clock
reaction and present experimental values for the Arrhenius constant and activation energy.
To present and submit your work you must prepare a lab report that contains the following items,
in this order:
1. A title page showing your name and the experiment title
2. Data table showing experimental conditions and experimental observations (based on Table 2)
3. Results table showing experimental conditions for each run and calculated values of initial
rates. Show at least one example of calculation work. (Section F step 1)
4. Organized calculations showing how you determined the reaction order for IO3-, I- and H+
(section F step 2)
5. Organized calculations showing how you determined the average value of the rate constant, k.
(Section F step 3)
6. Results table showing experimental conditions where temperature is varied and calculations
for average initial rate, value of the rate constant k, ln(k), temperature T in C and K, and 1/T
(K-1) (as described in Section F step 4)
7. Graph of ln(k) vs (1/T). Fully annotated as described in Section F step 5)
8. Answers to the post-lab questions shown on the next page.
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Post Lab Questions:
1. Based on your experimental observations and data analysis, what is the rate law for the iodine
clock reaction?
2. Consider running the iodine clock reaction with the following initial reactant concentrations:
[IO3-] = 0.020 M, [[I-] = 0.15 M, [H+] = 0.00001 M at a temperature = 5 C.
a) Calculate the initial rate of the reaction for this set of conditions (show your work)
b) Based on your answer for 2a), and your knowledge of the experimental set up and
calculations, calculate the predicted time it would take (in seconds) for the blue color to
appear for the given initial reactant concentrations. (show your work)
c) Use the iodine clock simulation to run the reaction at the given conditions. What is the
“observed value” for tc (the time taken for blue color to appear)?
d) Calculate the % error between your observed time and the predicted time
% error = (observed time – predicted time) × 100
predicted time
3. What value did you calculate for the activation energy for the iodine clock reaction?
4. What value did you calculate for the Arrhenius constant for the iodine clock reaction?
5. What do the activation energy and Arrhenius constant tell us about a chemical reaction?
6. Several mechanisms have been proposed for the iodine clock reaction. Five of them are shown
below along the predicted experimental rate law. Which of these mechanism(s) is/are plausible
based on your observations? Explain your reasoning.
Mechanism 1:
Mechanism 2:
IO3- + I- + 2 H+ ⇄ H2I2O3 (fast)
H2I2O3  HIO + HIO2 (slow)
Followed by fast reactions
rate = k[IO3-][I-][H+]2
IO3- + H+ ⇄ IO2+ + OH- (fast)
H+ + OH- ⇄ H2O (fast)
IO2+ + I- ⇄ IOIO (fast)
IOIO + I-  I+ + 2 IO- (slow)
Followed by fast reactons
rate = k[IO3-][I-]2[H+]
Mechanism 3:
Mechanism 4:
IO3- + H+ ⇄ IO2+ + OH- (fast)
H+ + OH- ⇄ H2O (fast)
IO2+ + I- ⇄ IOIO (fast)
IOIO  I+ + IO+ + IO- (slow)
Followed by fast reactons
rate = k[IO3-][I-][H+]
IO3- + I- + 2 H+ ⇄ H2I2O3 (fast)
H2I2O3 ⇄ I2O2 + H2O (fast)
I2O2 + I-  I3O2- (slow)
Followed by fast reactions
rate = k[IO3-][I-]2[H+]2
Mechanism 5:
IO3- + 2I- + 2 H+  2 HIO + IO- (slow)
Followed by fast reactions
rate = k[IO3-][I-]2[H+]2
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H) Lab Report Submission
Create a single document containing your lab report with the work organized in the order given in
section G. You can include segments written by hand and segments generated digitally (e.g.
graphs). Write clearly and legibly and show examples of your work as indicated in the instructions.
Make sure that each section of the report is organized in a clear and logical fashion (use headings
and subheadings as appropriate).
Use whatever technology you have available to convert your work into a single digital document
(scan/take pictures, convert to PDF) that is in PDF format. You can use the following free online
tools to help you convert and merge files:
https://smallpdf.com/pdf-converter
https://smallpdf.com/merge-pdf
Name your PDF file based on this convention. lastname_initial_kinetics.PDF
Submit this document by uploading into the assignment on Canvas. This must be done by
deadline indicated in the assignment on Canvas.
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