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The results given in the below table were obtained during kinetic studies of the following reaction: 2A + B → C + D Experiment | [A]/mol L^-1 | [B]/mol L^-1 | Initial rate/mol L^-1 min^-1 I | 0.1 | 0.1 | 6.00 × 10^-3 II | 0.1 | 0.2 | 2.40 × 10^-2 III | 0.2 | 0.1 | 1.20 × 10^-2 IV | X | 0.2 | 7.20 × 10^-2 V | 0.3 | Y | 2.88 × 10^-1 X and Y in the given table are respectively:
- 0.3, 0.4
- 0.4, 0.3
- 0.4, 0.4
- 0.3, 0.3
Correct answer: 0.4, 0.3
Solution
The correct values for X and Y can be determined by analyzing the initial rates in relation to the concentrations of A and B. In Experiment II, doubling the concentration of B while keeping A constant results in a fourfold increase in the rate, indicating that the reaction is second order with respect to B. Similarly, in Experiment IV, the rate is consistent with the expected increase when A is increased to 0.4 mol L^-1 and B is held at 0.2 mol L^-1, confirming that X is 0.4 and Y is 0.3.
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