Arrhenius lab
Introduction
Reaction rates can often double or triple with a 10◦C rise in temperature. The effect of
temperature on the rate can be explained by collision theory. According to this theory,
two molecules react after colliding only when the energy of collision is greater than the
activation energy and when the molecules are properly oriented.
It is the rapid increase in the fraction of collisions having energy greater than the
activation energy that explains the large temperature dependence of reaction rates.
The Arrhenius equation is a mathematical relationship showing the dependence of a rate
constant on temperature.
Rate constants for most chemical reactions closely follow an equation of the form:
k = Ae-Ea/RT
Here e is the base of natural logarithms, 2.718………; Ea is the activation energy; R is
the gas constant, 8.31 J/(K·mol); and T is the absolute temperature. The symbol A is
assumed to be a constant, and is termed the frequency factor. The frequency factor is
related to the frequency of collisions with proper orientation.
It is useful to recast the Arrhenius equation in logarithmic form. Taking the natural
logarithm of both sides of the Arrhenius equation gives:
ln k = ln A – Ea/(RT)
This may be equated to the equation for a straight line as follows:
ln k = ln A + (–Ea/R)(1/T)
y = b + m
x
This shows that if you plot ln k against 1/T, you should get a straight line. The slope of
this line is –Ea/R, from which you can obtain the activation energy Ea. The intercept is ln
A.
In the previous experiment – “Reaction Kinetics in a Redox Reaction” we investigated
the reaction kinetics for the oxidation of iodine ion by hydrogen peroxide:
H2O2 + 2I- + H+ -> I2 + 2H2O
The rate of this reaction and its dependence on the concentration of iodine ion and
hydrogen peroxide were determined.
The reaction rate is expressed as the rate of change of concentration of the participating
species. In our case reaction rate, in terms of change of hydrogen peroxide concentration,
was determined to be:
rate = d[H2O2]/dt = k[H2O2][I-]
In our experiment we will perform three runs – repeats of run # 1 (from our last
experiment - Reaction Kinetics in redox reaction). Although, in this experiment, we will
analyze the reaction at three different temperatures.
We will use our collected data to:
1) Calculate the rate of our reactions at the three temperatures,
2) Use our rate of reaction and initial reactant concentrations to calculate the rate
constant, k, at the different temperatures,
3) Calculate ln k and 1/T,
4) Use excel to plot ln k vs 1/T, then
5) From our plot, the slope of this line is –Ea/R, from which we can which we can
calculate the activation energy, Ea.
Arrhenius lab
Procedure
**NOTE: Use instructions provide by the instructor – DO NOT USE INSTRUCTIONS
IN CHEMLAB! To remove the instructions on the screen, and free-up more working
area, perform the following operation: click on the OPTIONS tab; then click on LAB
ONLY. The instructions “disappear” and all of the area is now lab space.**
We will assume the following in our experiment:
Run 1: All chemicals are at 22ºC - for run #1 use chemicals labeled #1 in ChemLab
Run 2: All chemicals are at 42ºC - for run #2 use chemicals labeled #2 in ChemLab
Run 3: All chemicals are at 62ºC - for run #3 use chemicals labeled #3 in ChemLab
Step 1: For Run 1 prepare solution of sodium thiosulfate, place 0.25gm (250 mg) of #1
Sodium thiosulfate – Na2O3S2 in100ml beaker, add 20 ml of #1 Water – H2O and mix
until dissolved.
Step 2: Prepare solution of potassium iodine. Add 10 gm of #1 Potassium iodide - KI to a
100 ml beaker and add 20 ml of #1 Water – H2O and mix until dissolved.
Step 3: Obtain a 600 ml beaker, add #1 KI and #1 Na2S2O3 solutions, then add 60ml of
(1.000M) #1 Hydrogen chloride - HCl. Place on stir plate and stir solution.
Step 4: Finally add 50ml of (0.100M) #1 Hydrogen Peroxide - H2O2, as soon as the H2O2
is added start timing the reaction. Stop timing when the Iodine appears (solution changes
color). Record time.
Step 5: Repeat steps 1-4 for runs 2 and 3. Remember to use #2 chemicals for run 2 and
#3 chemicals for run 3.
Arrhenius lab
Name:__________
Section: ________
Observations
Data
Arrhenius lab – Chemical Composition Data Table
Na2SO3
Run #1
Run #2
Run #3
Grams
added
0.25
0.25
0.25
H2O
added
(ml)
20
20
20
I-
H2O2
M (total
volume)
0.01055
0.01055
0.01055
ml
added
50
50
50
M (total
volume)
0.0333
0.0333
0.0333
Grams
(KI)
added
10
10
10
H2O
added
(ml)
20
20
20
H+
M (total
volume)
0.40160
0.40160
0.40160
ml
HCL
added
60
60
60
M (total
volume)
0.4
0.4
0.4
Times (secs):



run 1: __________
run 2: __________
run 3: __________
rate = d[H2O2]/dt = (moles of H2O2 consumed) / (Total volume * Reaction Time)
moles of H2O2 consumed = .5 * moles S2O3 2moles S2O3 2- = .25 gm / GMW (Na2S2O3) = .25/ 158.1 = .00158
moles of H2O2 consumed = .5 *.00158 = .00079
rates (mol/(L x s)):



run 1: __________
run 2: __________
run 3: __________
rate constant, k (L/(mol x s)): Where: k = rate/[H2O2][I-]



run 1: __________
run 2: __________
run 3: __________
ln k, (unitless):



run 1: __________
run 2: __________
run 3: __________
H2O
(ml)
added
0
0
0
Total
Volume
(ml)
150
150
150
1/T (Kelvin):



run 1: __________
run 2: __________
run 3: __________
Use Excel to graph ln k vs 1/T and determine the slope from the equation for a straight
line using linear regression. Print your graph – including the equation for the straight
line.
Calculations
From your graph, the slope of the regressed line is –Ea/R. Therefore, to calculate
activation energy, Ea:
Ea = -(Slope x R)
Activation Energy: _________J/mol
Discussion