A Level H2 Chemistry Kinetics Equilibrium Quiz

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A-Level Chemistry H2 Quiz - Kinetics Equilibrium

Name: ___________________________
Class: ___________________________
Date: ___________________________
Score: ___________ / 40

Duration: 60 minutes
Total Marks: 40

Instructions:

• Answer all 20 questions.
• Show all working clearly where calculations are required.
• Use the Data Booklet if needed.
• Section A: Short Answer (1–7). Section B: Structured Response (8–14). Section C: Data Interpretation & Calculation (15–20).

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Section A: Short Answer (1–7)

1. State the rate equation for a reaction where the rate is directly proportional to the concentration of reactant X and independent of reactant Y. [1]

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2. Define the term activation energy, $E_a$, in the context of reaction kinetics. [1]

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3. Write the expression for the equilibrium constant $K_c$ for the reaction:
$\text{N}_2(g) + 3\text{H}_2(g) \rightleftharpoons 2\text{NH}_3(g)$ [1]

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4. According to Le Chatelier’s principle, what is the effect on the equilibrium position of increasing the pressure for a gaseous reaction where there are fewer moles of gas on the product side? [1]

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5. A catalyst increases the rate of a reaction. State whether it changes the value of the equilibrium constant $K_c$. [1]

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6. The rate constant $k$ for a first-order reaction has units of s$^{-1}$. State the overall order of the reaction. [1]

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7. For the exothermic reaction $A + B \rightleftharpoons C + D$, state the effect of increasing temperature on the equilibrium yield of C. [1]

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Section B: Structured Response (8–14)

8. The reaction between iodide ions and peroxodisulfate(VI) ions is:
$2\text{I}^-(aq) + \text{S}_2\text{O}_8^{2-}(aq) \rightarrow \text{I}_2(aq) + 2\text{SO}_4^{2-}(aq)$
(a) Describe how the initial rate of this reaction can be measured experimentally using thiosulfate and starch. [2]
(b) State one variable that must be kept constant. [1]

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9. For the hydrolysis of bromoethane:
$\text{CH}_3\text{CH}_2\text{Br} + \text{OH}^- \rightarrow \text{CH}_3\text{CH}_2\text{OH} + \text{Br}^-$
(a) Suggest a suitable method to follow the rate of this reaction. [1]
(b) Explain how you would determine the order with respect to $\text{OH}^-$ from concentration–time data. [2]

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10. Explain, using the Maxwell–Boltzmann distribution, why increasing temperature increases the rate of a reaction. [3]

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11. The following data were obtained for the reaction $P + Q \rightarrow R$:

Experiment	[P] / mol dm$^{-3}$	[Q] / mol dm$^{-3}$	Initial rate / mol dm$^{-3}$ s$^{-1}$
1	0.10	0.10	$4.0 \times 10^{-4}$
2	0.20	0.10	$8.0 \times 10^{-4}$
3	0.10	0.20	$1.6 \times 10^{-3}$

(a) Determine the order with respect to P. [1]
(b) Determine the order with respect to Q. [1]
(c) Write the rate equation. [1]

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12. State and explain the effect of adding a catalyst on the rate of attainment of equilibrium for the reaction:
$\text{H}_2(g) + \text{I}_2(g) \rightleftharpoons 2\text{HI}(g)$ [2]

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13. For the equilibrium: $\text{Fe}^{3+}(aq) + \text{SCN}^-(aq) \rightleftharpoons \text{FeSCN}^{2+}(aq)$ (blood red)
(a) State the colour change observed when excess $\text{Fe}^{3+}$ is added. [1]
(b) Explain this observation using Le Chatelier’s principle. [2]

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14. The equilibrium constant $K_p$ for $\text{N}_2\text{O}_4(g) \rightleftharpoons 2\text{NO}_2(g)$ is $0.14$ atm at $298$ K.
(a) Write the $K_p$ expression. [1]
(b) State the units of $K_p$ for this reaction. [1]

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Section C: Data Interpretation & Calculation (15–20)

15. The decomposition of $\text{N}_2\text{O}_5$ is first order:
$2\text{N}_2\text{O}_5(g) \rightarrow 4\text{NO}_2(g) + \text{O}_2(g)$
At $320$ K, $k = 1.5 \times 10^{-3}$ s$^{-1}$. Calculate the time required for the concentration to fall from $0.040$ to $0.010$ mol dm$^{-3}$. [3]

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16. For the reaction $\text{CO}(g) + 2\text{H}_2(g) \rightleftharpoons \text{CH}_3\text{OH}(g)$, $K_c = 4.8$ at $500$ K.
A mixture initially contains $[\text{CO}] = 0.20$, $[\text{H}_2] = 0.40$, $[\text{CH}_3\text{OH}] = 0$ mol dm$^{-3}$.
Let equilibrium $[\text{CH}_3\text{OH}] = x$.
(a) Write expressions for equilibrium concentrations of CO and H$_2$ in terms of $x$. [1]
(b) Set up the $K_c$ equation and solve for $x$. [3]

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17. The rate of a reaction doubles when temperature increases from $300$ K to $310$ K. Using the Arrhenius equation concept, explain what happens to the fraction of molecules with energy $> E_a$. [2]

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18.

Image pending generation: graph for Q18.

Using the graph:
(a) State the equilibrium concentration of C. [1]
(b) Calculate $K_c$ for the reaction (no units required). [2]

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19. For the gas equilibrium $2\text{SO}_2(g) + \text{O}_2(g) \rightleftharpoons 2\text{SO}_3(g)$, $\Delta H = -198$ kJ mol$^{-1}$.
(a) State the effect of increasing temperature on $K_p$. [1]
(b) State the effect of increasing pressure on the equilibrium yield of $\text{SO}_3$. [1]
(c) Explain both effects using Le Chatelier’s principle. [2]

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20. A student proposes the mechanism for $2\text{NO} + \text{O}_2 \rightarrow 2\text{NO}_2$:
Step 1: $\text{NO} + \text{NO} \rightleftharpoons \text{N}_2\text{O}_2$ (fast)
Step 2: $\text{N}_2\text{O}_2 + \text{O}_2 \rightarrow 2\text{NO}_2$ (slow)
(a) Identify the rate-determining step. [1]
(b) Deduce the rate equation consistent with this mechanism. [2]

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A-Level Chemistry H2 Quiz - Kinetics Equilibrium: Answer Key

Total Marks: 40
Topic: Kinetics Equilibrium

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Section A: Short Answer

1. [1 mark]
rate = $k[\text{X}]$
Teaching note: Rate proportional to [X] and independent of [Y] means order w.r.t. X = 1, w.r.t. Y = 0. Rate equation shows only X.

2. [1 mark]
The minimum energy required for colliding particles to react / for a successful collision.
Teaching note: Activation energy is the energy barrier above the average kinetic energy that reactants must overcome.

3. [1 mark]
$K_c = \dfrac{[\text{NH}_3]^2}{[\text{N}_2][\text{H}_2]^3}$
Teaching note: Products over reactants, each concentration raised to its stoichiometric coefficient.

4. [1 mark]
Equilibrium shifts to the side with fewer moles of gas (product side).
Teaching note: Higher pressure favours the direction reducing gas moles.

5. [1 mark]
No, it does not change $K_c$.
Teaching note: Catalyst speeds up forward and reverse equally; only temperature changes $K_c$.

6. [1 mark]
First order (overall order = 1).
Teaching note: Units s$^{-1}$ are characteristic of first-order rate constant.

7. [1 mark]
Yield of C decreases.
Teaching note: Exothermic forward reaction is opposed by temperature increase; equilibrium shifts left.

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Section B: Structured Response

8. [3 marks total]
(a) [2] Mix known volumes of I⁻, S₂O₈²⁻, thiosulfate and starch. Reaction with thiosulfate keeps I₂ low; when thiosulfate used up, blue-black colour appears. Time to colour change gives rate via [S₂O₈²⁻] consumed.
(b) [1] Temperature / volume / concentration of thiosulfate.
Marking: 2 for correct method, 1 for valid constant variable.

9. [3 marks total]
(a) [1] Monitor conductivity (Br⁻ replaces OH⁻) or pH (OH⁻ decreases).
(b) [2] Plot [OH⁻] vs time; if linear, zero order; if half-life constant, first order; determine gradient/rate at different [OH⁻].
Marking: 1 + 2.

10. [3 marks]
At higher T, more molecules have energy > $E_a$ (Maxwell–Boltzmann curve shifts right, area beyond $E_a$ increases). More successful collisions per unit time → higher rate.
Marking: 1 for distribution shift, 1 for greater fraction >$E_a$, 1 for rate increase.

11. [3 marks total]
(a) [1] Order w.r.t. P = 1 (doubling [P] doubles rate: exp1→2).
(b) [1] Order w.r.t. Q = 2 (doubling [Q] quadruples rate: exp1→3).
(c) [1] rate = $k[\text{P}][\text{Q}]^2$.
Teaching note: Compare experiments holding one variable constant.

12. [2 marks]
Catalyst provides alternate pathway with lower $E_a$ → both forward and reverse rates increase equally → equilibrium reached faster, position unchanged.
Marking: 1 for faster attainment, 1 for no position change.

13. [3 marks total]
(a) [1] Solution becomes deeper red.
(b) [2] Adding Fe³⁺ increases reactant concentration; equilibrium shifts right to consume Fe³⁺, forming more FeSCN²⁺ (red).
Marking: 1 + 2.

14. [2 marks total]
(a) [1] $K_p = \dfrac{P_{\text{NO}_2}^2}{P_{\text{N}_2\text{O}_4}}$
(b) [1] Units = atm (or pressure units); since 2 – 1 = 1.
Teaching note: $K_p$ units = (pressure)$^{Δn}$, Δn = 1.

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Section C: Data Interpretation & Calculation

15. [3 marks]
First-order: $t = \dfrac{1}{k} \ln\dfrac{[A]_0}{[A]_t}$
$= \dfrac{1}{1.5\times10^{-3}} \ln\dfrac{0.040}{0.010} = \dfrac{1}{1.5\times10^{-3}} \ln 4$
$= 666.7 \times 1.386 = 924$ s (3 s.f.)
Marking: 1 formula, 1 substitution, 1 answer.

16. [4 marks total]
(a) [1] $[\text{CO}] = 0.20 - x$, $[\text{H}_2] = 0.40 - 2x$
(b) [3] $K_c = \dfrac{x}{(0.20-x)(0.40-2x)^2} = 4.8$
Assume $x$ small: $x \approx 4.8 \times 0.20 \times 0.16 = 0.154$ (too large; solve quadratic).
Exact: $(0.20-x)(0.40-2x)^2 = x/4.8$ → using solver $x \approx 0.095$ mol dm⁻³.
Marking: 1 expression, 2 solving, 1 value.

17. [2 marks]
Fraction of molecules with $E > E_a$ increases exponentially with T (Arrhenius). Doubling rate implies significantly larger fraction above barrier at 310 K.
Marking: 1 for increase, 1 for link to rate.

18. [3 marks total]
(a) [1] $[\text{C}]_{eq} = 0.30$ mol dm⁻³
(b) [2] $K_c = \dfrac{0.30}{0.20 \times 0.20} = 7.5$
From image: curves plateau at given values; use equilibrium concentrations.

19. [4 marks total]
(a) [1] $K_p$ decreases.
(b) [1] Yield of SO₃ increases.
(c) [2] Exothermic: T↑ shifts left, Kp↓. Fewer moles product (3→2): P↑ shifts right, yield↑.
Marking: 1+1+2.

20. [3 marks total]
(a) [1] Step 2 (slow).
(b) [2] rate = $k[\text{N}_2\text{O}_2][\text{O}_2]$; from fast eq $[\text{N}_2\text{O}_2] \propto [\text{NO}]^2$ → rate = $k'[\text{NO}]^2[\text{O}_2]$.
Marking: 1 + 2.