Effect of Temperature on rate of Reaction

The
rate of reaction is a function of many
things, including surface area, concentration and temperature. For
a reaction to take place, the molecuels must collide
with enough energy (called the Activation Energy, E_{A})
Consider the
reactant molecules to be at a particular temperature. By definition,
the
molecules have an average kinetic energy, some with more than the
average, some
with less. If we increase the temperature of a system, the average
kinetic
energy will increase and the number of molecules with the 'minimum'
energy to
react, thus increasing the frequency (rate) of reaction.

Svante Arrhenius has shown that for simple reactions, the equation

k
= Ae^{-Ea/RT} →
ln k = -E_{A}/RT
+
ln A

where k = rate constant, A is the
collision frequency factor
(function of the nature of reaction) R is the gas constant, T is the
absolute
temperature, E_{A} is the activation energy, and e
is the base of the
natural logarithms. We want to know the relationship between rate of
reactions
and two different temperatures. After some creative algebra, we can
transform
the Arrehenius equation into something managable:

k is the rate constant at temperature
T, and k' is the rate
at T'.

In this
experiment, the reaction to be studied is the oxidation of HI to I_{2}
by the action of H_{2}O_{2}
according to the equation

1
H_{2}O_{2}
+ 2 HI → 2 H_{2}O + 1 I_{2}

This reaction has been shown to be a
second order reaction,
first with respect to [HI], first with respect to [H_{2}O_{2}].
The experimentally determined rate law is of the form:

rate
= k[HI][H_{2}O_{2}]

The total volume of the solution containing the components of the reaction is kept constant during the course of this experiment; as another simplification, conditions are adjusted so that the reaction follows pseudo first order kinetics. This is done by maintaining the concentration of [HI], so the rate law becomes

rate
= k'[H_{2}O_{2}] where k' = k[HI]

This reaction will thus behave is a pseudo first order kinetics under these conditions.

To
calculate the activation energy E_{a}, we simply
need to determine the ratio of
the rate constants k_{1}/k_{2}
at two different temperatures T_{2}
and T_{1}. There exists many different ways of
determining rate
constants. Here, we will measure the time taken to produce identical
amounts of
I_{2} at two temperatures, T and T'. We will have a
series of data that
is amount of Na_{2}S_{2}O_{3}
and time. If we graph this (amount of
Na_{2}S_{2}O_{3}
(x) vs. time(y) using a computer and
getting the equation and R^{2}) the slope will be
how much time is
required per mole of Na_{2}S_{2}O_{3}.
Since time and
rate are inversely proportional we can obtain the rate constant ratios
by:

One experiment will be performed at
temperature T_{2},
and another at different temperature T_{1}, using
the same volumes and
concentrations of reactants in both experiments. The time required to
produce
the same amount of I_{2} at each of the two
temperatures will be
measured and used to calculate k_{1}/k_{2}.

In this
experiment, HI is produced by dissolving KI in a solution containing an
excess
of H_{2}SO_{4}:

1
H_{2}SO_{4}
+ 1 KI → 1 HI + KHSO_{4}

Starch is
added to this solution as the indicator of free iodine. We start the
stopwatch
at the moment we add our second reactant, H_{2}O_{2}.
Of course
the production of I_{2} begins immediately, and the
blue color will
instantly appear. A measured volume of 0.1 M Na_{2}S_{2}O_{3}
is then added, causing the blue color to disappear due to the reduction
of the
I_{2} by thiosulfate. The solution will stay
colorless until enough
additional I_{2} is produced by H_{2}O_{2}-HI
oxidation
to react with the remaining Na_{2}S_{2}O_{3}
in the
solution. At this point the blue color will reappear. The time is
noted. This
represents the time required to produce I_{2}
equivalent to the Na_{2}S_{2}O_{3}
in the volume of 0.1 M Na_{2}S_{2}O_{3}
which was
added.

The entire reaction will be repeated with the same concentrations, but at a different temperature.

11 min screencast of the 'concept' of the lab

Chemicals needed- KI, 0.100 M Na_{2}S_{2}O_{3},
3% H_{2}O_{2}, starch solution,
6M H_{2}SO_{4},
ice

2. Place a 400 mL beaker on the lab balance, and add 1.00 g
of KI to it. It is important to get as close to 1.00 g as possible as
you want
the masses to be the same for both runs.

3. Dissolve the KI in about 250 mL of water. Add 15 mL of 6M
H_{2}SO_{4} (caution!) and
carefully place the beaker in the
1000 mL water (ice?) bath. Place a stir bar in the solution and place
the
entire system on a stir plate and turn it on (no heat!!)

4. Make 50 mL of a 0.6% H_{2}O_{2}
solution
(10 mL 3% H_{2}O_{2} and 40 mL
DI H_{2}O) exact volumes
are best.

5. Carefully measure out 5.0 mL of this solution and place
it in a test tube, which is sitting in the 250 mL water bath beaker.
Put a
stopper on the test tube.

6. Rinse your buret twice and then fill it with 0.100 M Na_{2}S_{2}O_{3}.

7. You want to wait until everything is at the desired
temperature before you go any further. Once you start the reaction, you
will
want to keep the system stirring to keep everything at the same (water
bath)
temperature.

8. Position the 400 mL beaker (plus its water bath) under
the buret. Record the temperature of the solution.

9. Add 4-6 drops of starch indicator to the 400 mL beaker.

10. Add the H_{2}O_{2 }to the
solution and
begin timing. The solution should immediately turn blue.
That’s a good thing.
The stopwatch does not get stopped until you have your (~12-20)
measurements.

11. One student should be reading the buret and recording
data (known as buret person), whilst the other watch the clock and the
solution
for the color change (clock person)

12. Buret person measures the volume in the buret. Then they
add about 1 mL of solution and write down the new volume (how much was
just
added?) The solution should turn clear. At the instant the solution
turns blue
again, (clock person calls out the time, since they are watching the
reaction)
record the time (buret person).

13. Buret person adds 1 more mL, measures and records the
new volume. The solution turns clear. When it turns blue again, measure
and
record the time.

14. Continue this procedure until you have added more than
20 mL of Na_{2}S_{2}O_{3}
or until the blue color just
stops coming back.

**When you are determining the slope, you
should use only data that gives a linear shape. If your data has a
logarithmic
curve to it, use only the first 6-8 data points to determine the slope.**

_{2}
by the action of H_{2}O_{2}. You must submit computer generated graphs of
VNa_{2}S_{2}O_{3}
vs. time. (where should they be?) There are no prelab questions, however if you want
to
start the lab, you will need to show the instructor your premade data
sheet
(recycled paper, of course) before you step foot in the lab. You also
need to
write up a flow-chart of the procedure. No data sheet or flow chart, no
lab for
you!