ISC Class 12 Chemical Kinetics — practice questions — page 2

To understand why a reaction occurs, we use **Collision Theory**. According to this theory, for a chemical reaction to take place, reactant particles must collide with each other. However, not every collision results in a reaction. For a collision to be 'effective' (leading to product formation), two criteria must be met: 1. **Activation Energy ($E_a$):** The colliding molecules must possess a minimum amount of kinetic energy, known as the activation energy. This energy is required to break the existing chemical bonds in the reactants. If the energy is less than $E_a$, the molecules simply bounce off each other. 2. **Correct Orientation:** Even if molecules have enough energy, they must strike each other at the correct angle or spatial orientation so that the specific atoms involved in the reaction can interact and form new bonds. Therefore, the correct requirement is **Energy $\ge$ activation energy + correct orientation**.

Each group is one summary sentence. Source: syllabus Chemistry Part I, Ch 3 Chemical Kinetics.

1. syllabus introduces the rate law: $\text{rate} = k[A]^a[B]^b$, where $k$ is the rate constant and $a, b$ are reaction orders. 2. The rate constant $k$ captures temperature-dependent intrinsic kinetics: it depends on $T$ via the Arrhenius relation $k = A e^{-E_a/RT}$, and a catalyst can lower $E_a$ (and thus raise $k$). 3. Crucially, $k$ does NOT depend on the concentrations of reactants — those appear separately in the rate law. 4. The rate ITSELF depends on both $k$ and current concentrations so option A is wrong (rate ≠ $k$). Option C confuses rate constant with thermodynamic quantities. Option D is wrong — $k$ and half-life are related but not equal. _Source: syllabus Chemistry Part 1, Ch 3, §3.2 (Rate Constant and Order of Reaction), p. 6–7._