Collision Theory Of Chemical Reactions

Step 1: Activation energy depends on enthalpy change of reaction.
Step 2: For exothermic reactions, reverse activation energy is higher and vice versa.

Step 1: For a first order reaction, ((V_∞)/(V_∞ - Vt)) t
Step 2: The given data satisfy first-order kinetics.

Step 1: Comparing runs. Rate changes linearly with both [A] and [B].
Step 2: Order determination. Reaction is first order with respect to both A and B.

Step 1: Use rate equation.
From the rate law, the rate of the reaction is related to the concentration of : Substituting the given values for rate and rate constant, we find that the concentration of is 0.04 M.
Thus, the correct answer is (3).

Step 1: Use the Arrhenius equation.
The rate constant depends on temperature as .
Step 2: Calculate the rate constant at 300 K.
Using the Arrhenius equation, the rate constant at 300 K is . Final Answer:

Step 1: Stoichiometric ratio:
Step 2:

Step 1: Pressure increase in .
Step 2:
Step 3: From stoichiometry, disappearance rate of equals pressure increase rate.

Step 1: For second order reaction:
Step 2: Units:

Step 1: First order rate law:
Step 2: Concentration decreases exponentially with time.

Step 1: Given time relation indicates second-order dependence on P.
Step 2: Linear decrease of Q with time indicates zero order in Q.
Step 3: Overall order: 2+0=2

Step 1: Rate law depends on reactants in slow step.
Step 2: Mechanism A has slow step involving Cl₂ and H₂S directly.
Step 3: Mechanism B involves HS⁻ in slow step, which contradicts given rate law.

Step 1: Molecularity is number of reacting molecules in the equation = 2.
Step 2: Order of reaction = 1 + (1)/(2) = (3)/(2)

2ⁿ=(1)/(4) ⟹ n=-2

Step 1: Arrhenius equation:

Step 2: Substitute .
Step 3: Rearranging gives option (B).

Step 1: At the point of intersection, [A]=[B].
Step 2: For reaction A→ B, this occurs when half of A has reacted.
Step 3: Thus the time corresponds to half-life t₁/2.

For a second order reaction, Rate = k[A]² Hence, k = fracmol L⁻1s⁻1(mol L⁻1)² = L mol⁻1s⁻1

For first order reactions, [A] = [A]₀ e⁻kt Thus concentration decreases exponentially with time.

Mechanism A: The slow (rate-determining) step is: Cl₂ + H₂S → ⋯ Hence, rate = k[Cl₂][H₂S] This matches the given rate law. Mechanism B: From the fast equilibrium: [HS⁻] = K[H₂S] Rate-determining step: rate = k[Cl₂][HS⁻] = k'[Cl₂][H₂S] However, since HS⁻ is an intermediate explicitly appearing, the mechanism does not directly justify the observed rate law without assumptions. Thus, only mechanism A is consistent.

Step 1: Let the rate law be: Rate ∝ [B]ⁿ Step 2: When concentration of B is doubled: fracNew rateOld rate = (2)ⁿ Step 3: Given that rate decreases by a factor of 4: (2)ⁿ = (1)/(4) = 2⁻2 ⟹ n = -1

Step 1: Order of reaction: Order = 1 + (1)/(2) = (3)/(2) Step 2: Overall reaction involves two reacting molecules: Molecularity = 2 Step 3: Hence, molecularity =2 and order =3/2.

Step 1: Understanding the Concept
Different experimental methods exist to determine the order of a reaction and the rate constant. The choice depends on the type of data provided.

Step 2: Detailed Explanation
1. Initial Rate Method: This method involves running the reaction multiple times with different starting (initial) concentrations of reactants and measuring the initial rate for each. By comparing how the rate changes as concentration changes, the order and then the rate constant ( ) can be determined. 2. Integration Method: Used when concentration data is given as a function of time ( ) for a single run. 3. Half-life Method: Used when the time taken for concentration to reduce by half is measured at different initial concentrations. 4. Differential Method: Relates the rate of reaction to the concentration at any point in time during the reaction.

Step 3: Final Answer
The method used when initial concentrations and corresponding rates are given is the Initial rate method.

Step 1: Understanding the Concept
For a first-order reaction, the time required for a certain percentage of completion depends on the rate constant , which remains constant throughout the reaction.

Step 2: Key Formula or Approach
The integrated rate law for a first-order reaction is: where is the initial concentration and is the amount reacted.

Step 3: Detailed Explanation
1. Find from 25% completion: - , , so . - 2. Calculate for 75% completion: - , so . - 3. Substitute : - - Using and : - . - The closest standard option provided in such exam sets is 150 min.

Step 4: Final Answer
The reaction will be 75% complete in approximately 150 minutes.

Step 1: Understanding the Concept:
For a zero-order reaction, the rate is independent of the concentration of the reactant. The half-life ( ) is directly proportional to the initial concentration ( ).

Step 2: Key Formula or Approach:
1. Zero order half-life:
2. Integrated rate law:

Step 3: Detailed Explanation:
1. Find the rate constant ( ): Using and : 2. Calculate the time ( ) for concentration change: Initial concentration ( ) = Final concentration ( ) = 3. Convert to minutes:

Step 4: Final Answer
The time required is 15 min.