Question

Each of the statements given below is false. Explain why.

(a) The activation energy of a reaction depends on the overall energy change $\mathbf{∆}\mathit{E}$ for the reaction.

(b) The rate law for a reaction can be deduced from the examination of the overall balanced equation for the reaction.

(c) Most reactions occur by one-step mechanisms.

Short answer

(a) The energy of activation is the energy required in order for a reaction to get past the ‘saddle point’ in the diagram while overall energy change is simply

$∆E=E_{Products}-E_{Reac\tan ts}$

(b) Reaction orders, with respect to all the reactants, can but don't necessarily have to be the same as coefficients in a balanced equation.

(c) Most reactions occur by multiple-step mechanisms.

Step-by-step solution

Understanding the non-dependence of activation energy of a reaction on total energy change.

Overall energy change and energy of activation are two completely different things. When observing an energy diagram of a reaction, energy of activation is the energy required in order for a reaction to get past the ‘saddle point’ in the diagram while overall energy change is simply:

$∆E=E_{Products}-E_{Reac\tan ts}$

Understanding the reaction orders relation with reaction coefficient.

Reaction orders, with respect to all the reactants, can but don't necessarily have to be the same as coefficients in a balanced equation which is why we can't simply deduce the rate law from a balanced equation.

Understanding the concept behind multiple step mechanism of reactions.

Most reactions occur by multiple-step mechanism which is also one of the reasons why the rate law cannot be directly deduced from the stoichiometry of a balanced chemical equation.