A Level H1 Chemistry Kinetics Equilibrium Quiz

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A-Level Chemistry H1 Quiz - Kinetics Equilibrium

Name: ____________________ Class: ____________________ Date: ____________________ Score: / 45

Duration: 60 Minutes
Total Marks: 45
Instructions: Answer all questions. Show all working for calculations. Use the data booklet where necessary.

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Section A: Reaction Kinetics (Questions 1–10)

1. Define the term rate of reaction. [1]
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2. For a reaction $A + 2B \rightarrow C$, the rate equation is $\text{Rate} = k[A]^1[B]^2$. (a) State the overall order of the reaction. [1] \

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   (b) If the concentration of $B$ is doubled while $[A]$ remains constant, by what factor does the rate increase? [1] \

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3. Explain, in terms of collision theory, why an increase in temperature leads to a significant increase in the rate of reaction. [2]
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4. A catalyst increases the rate of a chemical reaction. Describe how a catalyst achieves this in terms of activation energy. [2]
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5. The decomposition of nitrogen dioxide, $2\text{NO}_2(\text{g}) \rightarrow 2\text{NO}(\text{g}) + \text{O}_2(\text{g})$, is a first-order reaction with respect to $\text{NO}_2$. (a) Write the rate equation for this reaction. [1] \

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   (b) What are the units of the rate constant $k$ for this specific reaction? [1] \

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6. Distinguish between a homogeneous catalyst and a heterogeneous catalyst. [2]
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7. For the reaction $2\text{SO}_2(\text{g}) + \text{O}_2(\text{g}) \rightleftharpoons 2\text{SO}_3(\text{g})$, the rate is found to be independent of the concentration of $\text{O}_2$. (a) What is the order of reaction with respect to $\text{O}_2$? [1] \

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   (b) Suggest a reason why the rate might be independent of $[\text{O}_2]$. [1] \

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8. Draw a Maxwell-Boltzmann distribution curve for gas molecules at two different temperatures, $T_1$ and $T_2$ (where $T_2 > T_1$). Label the activation energy $E_a$ and the shaded area representing molecules with energy $\ge E_a$ at $T_2$. [3]
   
   
   
   
   
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9. A reaction is found to be second-order with respect to reactant $X$. If the initial concentration of $X$ is $0.10\text{ mol dm}^{-3}$ and the initial rate is $2.0 \times 10^{-4}\text{ mol dm}^{-3}\text{s}^{-1}$, calculate the value of the rate constant $k$. [3]
   
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10. Explain why the rate of reaction generally decreases as the reaction progresses toward completion. [2]
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Section B: Chemical Equilibrium (Questions 11–20)

11. Define the term dynamic equilibrium. [2]
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12. For the equilibrium $2\text{SO}_2(\text{g}) + \text{O}_2(\text{g}) \rightleftharpoons 2\text{SO}_3(\text{g})$, write the expression for the equilibrium constant $K_c$. [1]
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13. Consider the reaction: $\text{N}_2(\text{g}) + 3\text{H}_2(\text{g}) \rightleftharpoons 2\text{NH}_3(\text{g})$ ($\Delta H = -92\text{ kJ mol}^{-1}$). Predict and explain the effect of the following changes on the position of equilibrium: (a) Increasing the pressure. [2] \

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    (b) Increasing the temperature. [2] \

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14. State the effect of adding a catalyst to a system at equilibrium. [1]
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15. The equilibrium constant $K_c$ for the reaction $\text{PCl}_5(\text{g}) \rightleftharpoons \text{PCl}_3(\text{g}) + \text{Cl}_2(\text{g})$ is $0.040\text{ mol dm}^{-6}$ at $250^\circ\text{C}$. Calculate the equilibrium concentrations of $\text{PCl}_3$ and $\text{Cl}_2$ if the equilibrium concentration of $\text{PCl}_5$ is $0.20\text{ mol dm}^{-3}$. [3]
    
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16. Explain why the value of $K_c$ changes when the temperature of the system is altered. [2]
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17. For the reaction $\text{H}_2(\text{g}) + \text{I}_2(\text{g}) \rightleftharpoons 2\text{HI}(\text{g})$, the $K_c$ value is $50.0$ at a certain temperature. If the concentrations of $\text{H}_2$ and $\text{I}_2$ are both $0.10\text{ mol dm}^{-3}$ at equilibrium, calculate the equilibrium concentration of $\text{HI}$. [3]
    
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18. A reaction is exothermic in the forward direction. If the temperature is increased, what happens to the value of $K_c$? Explain your answer. [2]
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19. In the Haber process, the reaction is carried out at a high pressure but a relatively moderate temperature (approx $450^\circ\text{C}$), even though a low temperature would shift the equilibrium to the right. Explain this compromise. [3]
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20. For the equilibrium $\text{CO}(\text{g}) + \text{H}_2\text{O}(\text{g}) \rightleftharpoons \text{CO}_2(\text{g}) + \text{H}_2(\text{g})$, the $K_c$ is $1.0$. If the concentration of $\text{CO}$ is increased, describe the shift in equilibrium and the final effect on the value of $K_c$. [2]
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Answer Key - Kinetics Equilibrium Quiz

1. Rate of reaction: The change in concentration of a reactant or product per unit time. [1]

2. (a) Overall order = $1 + 2 = 3$. [1] (b) Rate $\propto [B]^2$. If $[B]$ is doubled, rate increases by $2^2 = 4$ times. [1]

3. Collision Theory:

   • Increase in $T$ increases average kinetic energy of molecules. [1]
   • A much larger fraction of molecules now possess energy $\ge E_a$, leading to a higher frequency of successful collisions. [1]

4. Catalyst: Provides an alternative reaction pathway with a lower activation energy ($E_a$). [2]

5. (a) $\text{Rate} = k[\text{NO}_2]^1$ (or $\text{Rate} = k[\text{NO}_2]$). [1] (b) $\text{mol dm}^{-3}\text{s}^{-1}$. [1]

6. Catalysts:

   • Homogeneous: Catalyst is in the same phase/state as the reactants. [1]
   • Heterogeneous: Catalyst is in a different phase/state (usually solid catalyst for gas/liquid reactants). [1]

7. (a) Zero order. [1] (b) The rate-determining step does not involve $\text{O}_2$. [1]

8. Diagram:

   • X-axis: Energy; Y-axis: Number of molecules. [1]
   • $T_2$ curve should be flatter and shifted to the right compared to $T_1$. [1]
   • $E_a$ marked as a vertical line; area to the right of $E_a$ under $T_2$ curve shaded. [1]

9. Calculation: $\text{Rate} = k[X]^2$ $2.0 \times 10^{-4} = k(0.10)^2$ $k = (2.0 \times 10^{-4}) / 0.01 = 0.020\text{ dm}^3\text{mol}^{-1}\text{s}^{-1}$. [3]

10. Rate decrease: As reactants are consumed, their concentration decreases. [1] This leads to a lower frequency of collisions between reactant particles. [1]

11. Dynamic Equilibrium: A state in a closed system where the rate of the forward reaction equals the rate of the reverse reaction [1], and the concentrations of reactants and products remain constant. [1]

12. $K_c = \frac{[\text{SO}_3]^2}{[\text{SO}_2]^2[\text{O}_2]}$. [1]

13. (a) Shift to the right (products). [1] There are 4 moles of gas on the left and 2 on the right; increase in pressure shifts equilibrium to the side with fewer moles of gas. [1] (b) Shift to the left (reactants). [1] The forward reaction is exothermic; increasing $T$ shifts equilibrium in the endothermic direction to absorb heat. [1]

14. No effect on the position of equilibrium; it only increases the rate at which equilibrium is reached. [1]

15. $K_c = \frac{[\text{PCl}_3][\text{Cl}_2]}{[\text{PCl}_5]}$ $0.040 = \frac{x^2}{0.20}$ (since $[\text{PCl}_3] = [\text{Cl}_2] = x$) $x^2 = 0.008 \Rightarrow x = 0.089\text{ mol dm}^{-3}$. [3]

16. $K_c$ depends on temperature because temperature changes the relative rates of the forward and reverse reactions differently based on their respective activation energies. [2]

17. $K_c = \frac{[\text{HI}]^2}{[\text{H}_2][\text{I}_2]}$ $50.0 = \frac{[\text{HI}]^2}{(0.10)(0.10)}$ $[\text{HI}]^2 = 50.0 \times 0.01 = 0.50$ $[\text{HI}] = \sqrt{0.50} = 0.707\text{ mol dm}^{-3}$. [3]

18. $K_c$ decreases. [1] For an exothermic reaction, increasing $T$ shifts the equilibrium to the left (reactants), decreasing the concentration of products and increasing reactants. [1]

19. Compromise:

    • Low $T$ gives higher yield of $\text{NH}_3$ (equilibrium shift). [1]
    • However, low $T$ results in an extremely slow rate of reaction. [1]
    • $450^\circ\text{C}$ is used to ensure a commercially viable rate of production while maintaining an acceptable yield. [1]

20. Shift: Shift to the right (towards products) to oppose the increase in $[\text{CO}]$. [1] $K_c$: The value of $K_c$ remains unchanged as it is only affected by temperature. [1]