Learning Exercise


The main objectives of this quantitative experiment are to (1) evaluate the rate constant kR and the reaction orders, (2) investigate the manner in which the reaction rate depends on temperature, and (3) evaluate the activation energy, Ea, of the reaction.
Course: Seconf Year Freshman (General) Chemistry

A simulated experiment and animation based on the iodine clock reaction using iodate and iodide in acetate buffers in... see more



The main objectives of this quantitative experiment are to (1) evaluate the rate
constant kR and the reaction orders, (2) investigate the manner in which the
reaction rate depends on temperature, and (3) evaluate the activation energy,
Ea, of the reaction.


This Post Laboratory assignment uses the Internet Resource, by Gary L.Bertrand,
University of Missouri-Rolla, available at:
The students are strongly advised to print this handout document as well as the
online information posted under Document at the target resource page for
completing the assignment.


There are several different iodine clock reaction systems besides the one used
in the wet laboratory (experiment #22 of reference 1). All of them, however,
have a common feature: the completion of any one of them is signaled by the
sudden appearance of the dark color that is characteristic of the interaction of
molecular iodine (I2) with starch. This co lor appears so abruptly that it can
be as startling as the sudden sound of an alarm.
The iodine clock system considered in this Virtual Laboratory uses the coupling
of an oxidation and reduction reaction as shown below:

The reactions involve the oxidation of iodide ion (I-) to dissolved iodine (I2)
or tri-iodide ion (I3-).

6 H+ + IO3- + 8 I- ?? 3 I3- + 3 H2O (1)

In addition to reaction 1, whose kinetics we will study, a measured amount of a
reducing agent (arsenious acid) is included in the mixture to ensure a
reasonably sharp color development. As compared with reaction 1, this reaction
is essentially instantaneous.

H3AsO3 + I3- + H2O ?? HAsO42- + 3 I- + 4 H+ (2)

The I2 produced in reaction 1 reacts with the reducing agent, arsenious acid,
H3AsO3, present in solution, so that until all the arsenious acid is gone from
the system, the I2 produced by (1) reacts very quickly with tri-iodide ion, 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.

The net resulting ionic equation from summing (1) and (2) can be written as
below (equation 3):

IO3- + 3 H3AsO3 ?? I- + 6 H+ + 3 HAsO42 (3)
can be described as the rate of disappearance of iodate ion. In the initial
stages of the reaction, this is also equal to 1/3 the rate of disappearance of
arsenious acid:
rate = - d[IO3-]/dt = - (1/3)d[H3AsO3]/dt (4)
The initial rate is approximated from the initial concentration of arsenious
acid and the time (tC) from mixing to the color change:
initial rate = (1/3) [H3AsO3]o / tC (5)
The rate of this reaction is expected to be mathematically related to the
concentrations of the reactants through a rate constant (kR), which depends only
on the temperature:
rate = kR[IO3-]a[I-]b[H+]c (6)
with the exponents (also called "order") a, b, and c expected to be integers (0,
1, 2, or 3) or half-integers (1/2, 3/2, 5/2). The reaction is said to have an
"overall order" of a + b + c.
The initial rate is then related to the initial concentrations:
initial rate = kR[IO3-]oa [I-]ob [H+]oc = (1/3) [H3AsO3]o / tC (7)


The procedure of the Virtual Experiment is depicted in figure 1. More details
concerning the experimental procedure are described in the simulation online
Document. To run the virtual experiments follow the steps below:
1. Click on "start". The solutions are mixed, the timer is started.

2. Click on "stop timer" at first appearance of blue color.

3. Record time and initial concentrations.
4. "ReSet" repeats the measurements with the same initial concentrations to
make duplicate/and or triplicate of your runs.

Access the web referenced resource,
http://web.umr.edu/~gbert/IClock/IClock.html, for more details about
experimental procedure and guidelines for using the simulation.

1. Basic Laboratory Studies in GENERAL CHEMISTRY with Semimicro Qualitative
Analysis, Grace R. Hered, Tenth Edition, p.191 ?C Houghton Mifflin Co - 1997


EXP22PL: Post Laboratory Assignment Accompanying Experiment 22, Reaction Rate
(This post laboratory assignment MUST be completed and turned in along with the
wet laboratory report of experiment #22)

Student Name_____________________ Section_________ Year___________

1. Data Collection

Perform the following virtual experiment runs and collect your data to complete
the Data Sheet given below:

Run # [IO3-]o [I-]o [H+]o Tri al 1* Trial 2* Trial 3* Average
1 0.005 0.05 2 x 10-5
2 0.010 0.05 2 x 10-5
3 0.005 0.10 2 x 10-5
4 0.005 0.05 4 x 10-5

* time it takes for the blue color complex to form (in sec.). You must perform
three trials for the runs and compute and average to be used in subsequent
calculations. For your experimental data to be meaningful, YOU MUST record the
exact time at which the blue complex appears. Although three trials are required
in this report, you may need to run ablank experiment to get accustomed to the
Virtual laboratory experiment prior to collecting experimental data.

2. Data Analysis: Determination of Reaction Orders a, b, c, and
Rate constant (kR)

a = ____________________

b = ____________________

c = ____________________

kR = ____________________ (best value of the rate constant)

3. Data Analysis: Temperature Effects and Activation Energy
The activation energy, Ea, for a given reaction is given by the Arrhenius
kR= A !? e - Ea / RT (8)
The equation gives the mathematical relationship between kR and absolute
temperature (T, A, and Ea are constants for the reaction, R is the Gas Law
Constant = 8.314 J/K!? mol)
To determine the activation energy for a given reaction, the concentration of
all the reactants are held constant and the only variable is the temperature of
the reaction. A series of measurements are made at different temperatures with
identical initial concentrations of all of the components. Since all the
concentrations remain constant the changes in kR are inversely related to the
changes in the time required for the blue color to appear (tC) (see equation 7)

kR= constant/tC (9)
or we can write in logarithmic form:
ln(kR) = ln(A) - Ea/RT (10)
Combining (9) and (10) above, one can derive (11)

ln(tC) = ln(constant/A) + Ea/RT (11)
A graph of the logarithm of time, ln(tC), vs the reciprocal of the absolute
temperature (1/T) should be a straight line with slope = Ea/R .
Plot your experimental graph to be included in this report and answer the

Slope = Ea/R = ______________________

Activation Energy, Ea = ________________

4 Advance Study Assignment:
Please use the space provided for answering each question listed below.

4a Define Activation Energy

4b A common rule of thumb says that, near room temperature, a 10!?C increase
results in a doubling of rate for many aqueous reactions. Use the Arrhenius
equation, kR = A!? e (-Ea/RT) , to show whether this rule of thumb holds true
for the iodine clock reaction. Use two temperatures that are 10!?C apart and
evaluate k in terms of the constant A for each of the temperatures. (SHOW WORK

4c Using your results, predict how long it would take for the reaction to occur
at 0!?C. Show your work and/or graphical representation.


Influence of Concentration and Temperature on the Rate of a Chemical Reaction

Reaction Order

Determination of Rate Equation

Activation Energy


Reaction Rate

Type of Task

Learning Objectives

(1) evaluate the rate constant kR and the reaction orders, (2) investigate the manner in which the reaction rate depends on temperature, and (3) evaluate the activation energy, Ea, of the reaction.