Creative Education, 2010, 2, 128-129
doi:10.4236/ce.2010.12019 Published Online September 2010 (http://www.SciRP.org/journal/ce)
How Do We Introduce the Arrhenius
Pre-Exponential Factor (A) to Graduate Students?
Vandanapu Jagannadham
Department of Chemistry, Osmania University, Hyderabad, India.
Received May 21st, 2010; revised July 2nd, 2010; accepted July 19th, 2010.
ABSTRACT
A simple and easy understanding of Arrhenius pre-exponential factor (A) is described in this article for a graduate
class-room lecture.
Keywords: Arrhenius Factor, Pre-Exponential Factor, Frequency Factor, Collision Number, Collision Frequency
1. Introduction
In chemical kinetics, the pre-exponential factor or A factor
is the pre-exponential constant in the Arrhenius equation,
an empirical relationship between temperature and rate
coefficient. It is usually designated by A when deter-
mined from experiment, while Z is usually left for colli-
sion frequency. For a first order reaction it has units of s-1,
for that reason it is often called frequency factor.
In short, the Arrhenius equation gives “the dependence
of the rate constant k of chemical reactions on the tem-
perature T (in Kelvin) and activation energy Ea”, as
shown below:
k = Ae-Ea / RT
where A is the pre-exponential factor or simply the
pre-factor and R is the gas constant. The units of the
pre-exponential factor are identical to those of the rate
constant and will vary depending on the order of the re-
action. If the reaction is first order it has the units s-1, and
for that reason it is often called the frequency factor or
attempt frequency of the reaction. When the activation
energy is given in molecular units, instead of molar units,
e.g. joules per molecule instead of joules per mole, the
Boltzmann constant is used instead of the gas constant. It
can be seen that either increasing the temperature or
decreasing the activation energy (for example through
the use of catalysts) will result in an increase in rate of
reaction.
Given the small temperature range in which kinetic
studies are carried, it is reasonable to approximate the
activation energy as being independent of the tempera-
ture. Similarly, under a wide range of practical condi-
tions, the weak temperature dependence of the pre-ex-
ponential factor is negligible compared to the tempera-
ture dependence of the factor; except in the case of “bar-
rier less” or diffusion-limited reactions, in which case the
pre-exponential factor is dominant and is directly ob-
servable.
2. The Interpretation of Pre-Exponential
Factor in the Arrhenius Equation
So far we have talked if not exhaustively but enough
about the pre-exponential factor for a M. Sc. Class-room.
So what is really the pre-exponential factor in the Ar-
rhenius equation? Did anybody think of its value with
change in temperature? Or did anybody guess its value at
infinite temperature? It can be said as “It is nothing but k
at infinite temperature or it is k of a reaction with zero
activation energy (barrier less)”. So in either way k will
be very large, more so it is a rate constant close to diffu-
sion limit (5 × 109 M-1 sec-1). This can be seen as in the
following by simple mathematics. On right hand side of
the Arrhenius equation Ea and T are the only variables
whereas A and R are constants. k will be equal to A when
the exponential factor becomes equal to one.
Therefore exp (–Ea / RT) = 1, which leads to
–Ea / RT = 0
Hence T should become infinite. Or the other possibility
is Ea should be zero or the reaction should be “barrier
less”.
3. Evaluation of A
In the first case: To get A it is necessary to evaluate k
experimentally, in which it is not possible to determine

How Do We Introduce the Arrhenius Pre-exponential Factor (A) to Graduate Students?
129
such a large rate constant at infinite temperature because
one cannot maintain or achieve infinite temperature. In-
stead one can determine the k at different temperatures
and get A from the antilogarithm of the intercept of the
Arrhenius plot (plot of ln k or log k versus 1 / T).
In the second case: One can use fast reaction kinetic
techniques like stopped flow methods, T-jump or P-jump
methods, laser flash photolysis, and pulse radiolysis.
In the third case: Proceed to evaluate by calculation
[1] assuming the atoms or molecules as Hard-Spheres as
pointed out by Max Trautz – 1916 [2] and W. C. McC.
Lewis – 1918 [3]. The formulation given by these people
was purely based on simple version of the kinetic theory
of gases in which the molecules were treated as hard
spheres.
The final equation obtained was
2
1
22
2





m
Tk
NdZ B
AAA

and if the gas contains two types of molecules say A and
B and where ‘’ is reduced mass, the equation obtained
was
2
1
28









 Tk
dNNZ B
BAAB
Lewis [3] applied this equation to the decomposition
of HI.
2HI  H2 + I2
He obtained a value of 5.5 X 1010 mole-1 sec-1 for ‘A’
both by experiment and by calculation. This clearly
shows that the theory put forward by Trautz and Lewis
assuming the molecules as hard spheres and neglecting
the slight variation of ‘A’ with temperature in good
agreement with each other.
In Table 1, log A values are given for some gas phase
reactions [4]. It is clear that the calculated values are all
Table 1. Some gas phase reactions and their observed and
calculated logarithmic values of Arrhenius frequency fac-
tors (A).
Log A, mole-1 sec-1
Reaction Observed Calculated by simple
collision theory
NO + O3  NO2 + O2 11.9 13.7
NO2 + F2  NO2F + F 12.2 13.8
NO2 + CO  NO + CO2 13.1 13.6
2ClO  Cl2 + O2 10.8 13.4
2NOCl  2NO + Cl2 13.0 13.8
NO + Cl2  NOCl + Cl 12.6 14.0
close to 14 which may be due to the fact that the over
simplification of the assumption that the molecules were
assumed to be hard-spheres and to have the same mo-
lecular diameter. The observed values were smaller than
those of the calculated ones. The reason is that the steric
factors, which are the ratios of the observed values to
those calculated ones, vary from 10-1 to 10-3. The low
steric factors are attributed to the losses of rotational
freedom during the formation of the activated complex.
REFERENCES
[1] K. J. Laidler, “Chemical Kinetics,” 2nd Edition, Ta-
ta-McGraw Hill, Inc., New York, 1973, pp 63-64.
[2] M. Trautz and Z. Anorg, “Evaluation of Arrhenius Fre-
quency Factor (A) by Simple Collision Theory,” Chemis-
try, Vol. 96, No. 1, 1916.
[3] W. C. M. Lewis, “Evaluation of Arrhenius Frequency
Factor (A) by Simple Collision Theory Assuming Atoms
and Molecules As Hard Spheres,” Journal of Chemistry
Socialist, Vol. 113, 1918, p. 471.
[4] D. R. Herschbach, H. S. Johnston, K. S. Pitzer and R. E.
Powell, “Comparison of the Arrhenius Frequency Factors
(A) of Some Gas-Phase Reactions,” Journal of Chemistry
Physics, Vol. 25, 1956, p. 73.