The temperature dependence of a reaction rate can be represented by the Arrhenius equation

\[K =Ae^{-E_a/RT}\]

The pre-exponential factor \(A\) is called the frequency factor and \(E_a\) is the energy of activation. The unit of \(E_a\) is J/mol or Kcal/mol.

The rate constants at two different temperatures are related as

\[log\frac{K_2}{K_1} = \frac{E_a}{2.303R}\left[\frac{T_2 – T_3}{T_1T_2}\right]\]

Log K versus 1/T gives a linear graph with negative slope. The reactant molecules collide with each other to cross over an energy barrier existing between the reactants and products. If the value of the difference in the internal energies of reactants and product is positive, the reaction is exothermic and if it is negative, the reaction is endothermic. If the temperature is raised the kinetic energy of the molecules increases which causes increase in (i) number of collisions (ii) number of molecules halving higher energy than threshold energy. For every 10°C rise in temperature, the increase in kinetic energy is about 3.3%. So the increase in number of collisions is about \(\sqrt{3.3}\) . , i.e., 1.8%. Hence the rate of reaction must increase only by about 1.8%. For every 10°C rise in temperature, the rate of reaction increases by 100%, i.e., two times If the rate of reaction is doubled for every rise of 10 K temperature, the rate of reaction increased for rise of temperature from 30°C to 80°C is 32 times. The activation energy does not depend on the concentration. The ratio of the rate constants at two different temperatures (preferably 35°C and 25°C) is known as temperature coefficient. If the activation energy is zero, then all the collisions will be fruitful and the reaction is 100% complete.

Which statement is correct?

Increase in the total pressure of a gas phase reaction increase the fraction of collisions effective in producing products

 The correct answer is option 4. Increase in the total pressure of a gas phase reaction increases the fraction of collisions effective in producing products.

In a gas phase reaction, the reaction rate is determined by the frequency of effective collisions between reactant molecules. Effective collisions are those that have sufficient energy and proper orientation for the reactant molecules to undergo the necessary chemical transformations and form products.

By increasing the total pressure of a gas phase reaction, the number of gas molecules in the reaction mixture is increased. This leads to a higher concentration of reactant molecules, resulting in an increased collision frequency. As the frequency of collisions increases, the chances of effective collisions also increase.

Additionally, an increase in pressure leads to a decrease in the volume available for the gas molecules to move. This results in a higher chance of molecular collisions occurring, as the gas molecules are confined to a smaller space.

The increased collision frequency, along with the reduced volume for molecular motion, increases the likelihood of successful collisions with sufficient energy and proper orientation. As a result, the fraction of collisions that are effective in producing products increases.

It is important to note that while an increase in total pressure generally increases the reaction rate, it does not affect the thermodynamics or equilibrium position of the reaction. The effect of pressure on the rate of a reaction is related to the collision frequency and the fraction of effective collisions, rather than the equilibrium concentrations of the reactants and products.

In summary, the statement (4) is correct. Increasing the total pressure of a gas phase reaction increases the fraction of collisions effective in producing products, leading to an increase in the reaction rate.