Rates of most chemical reactions increases with temperature because the rate constant 'k' increases with temperature. The temperature dependence of the rate constant is given by the Arrhenius equation, where

Svante Arrhenius, a Swedish chemist proposed this equation in the year 1889. This is an empirical equation, which was developed based on observations made on a large number of experiments.

Arrhenius equation predicts an exponential dependence on temperature of the rate constant k. In the equation, A is the pre exponential factor, which is almost independent of temperature (temperature dependence is slight and can be conveniently ignored). E_{a}is known as the energy of activation.

The logarithmic form of the Arrhenius equation represents a linear relation, that is,

fig 6.10 - A plot of logarithm of rate constant versus (1/T)

##### log A is determined from the intercept, while the energy of activation E

_{a}is determined from the slope of the plot of ln k or log k versus (1/T).

When the equation was published in 1889, in the paper, Arrhenius had suggested that molecules must be given enough energy so that they become 'activated' before they could react. Collision theory and transition state, which were proposed much later, enlarged the concept of 'activation'.

As mentioned before, the rates of reactions increase with temperature; in most cases, the rate doubles for every 10^{o}rise in temperature. This effect can be represented by the ratio of the rate constants measured at two temperatures, that is T K and (T+10)K. This ratio is called the 'temperature coefficient'. The two temperatures usually selected are 298 K and 308 K. Hence,

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