Chemical Reactivity — practice questions

State the three conditions needed for a successful collision between reactant particles.

Define activation energy and explain why a reaction with a high activation energy is slow.

A mixture of methane and oxygen can be kept indefinitely at room temperature without reacting, yet the combustion of methane has ΔrH = −890 kJ mol⁻¹ and once lit continues without further heating. Explain these observations fully.

Name the four factors that affect the rate of a reaction and state the effect of increasing each.

Explain, using collision theory, why powdered calcium carbonate reacts faster with acid than a single large lump of the same mass.

A student claims that increasing the concentration and increasing the temperature both work 'by giving the particles more energy'. Evaluate this claim, explain the actual mechanism of each, and explain why a 10 °C rise typically has a much larger effect on rate than doubling the concentration.

Define a catalyst and state how it increases the rate of a reaction.

Explain, in terms of collision theory, why lowering the activation energy increases the rate of a reaction.

An industrial process uses a catalyst and is run at 450 °C rather than at a higher temperature. Explain the role of the catalyst, why a moderate temperature still gives an acceptable rate, and evaluate the claim that 'using a catalyst is always better than raising the temperature'.

State the two features of a system at dynamic equilibrium.

Explain why an equilibrium is described as 'dynamic' rather than 'static'.

Explain why an equilibrium can only be established in a closed system, using the decomposition of calcium carbonate (CaCO₃ ⇌ CaO + CO₂) as an example. Then explain why 'the concentrations are constant' does not mean 'the concentrations are equal'.

State Le Chatelier's principle.

For N₂ + 3H₂ ⇌ 2NH₃ (ΔrH = −92 kJ mol⁻¹), explain the effect of raising the temperature on the yield of ammonia.

The Haber process (N₂ + 3H₂ ⇌ 2NH₃, ΔrH = −92 kJ mol⁻¹) is run at about 450 °C and 200 atmospheres with an iron catalyst. Explain why these particular conditions are used, discussing the effect of each on both yield and rate, and evaluate the compromise involved.

Write the expression for Kc for the equilibrium N₂ + 3H₂ ⇌ 2NH₃.

A reaction has Kc = 1.2 × 10⁻⁵ at 25 °C. Explain what this tells you about the position of equilibrium, and state whether it tells you anything about the rate.

For N₂O₄ ⇌ 2NO₂ (ΔrH = +57 kJ mol⁻¹), a 2.00 L flask at equilibrium contains 0.400 mol N₂O₄ and 0.200 mol NO₂. Calculate Kc with units, then explain what would happen to both the position of equilibrium and the value of Kc if (i) the flask were compressed to 1.00 L and (ii) the temperature were raised.

Define an acid and a base in terms of proton transfer, and give the conjugate base of HNO₃.

For HCl + H₂O → Cl⁻ + H₃O⁺, identify the acid and base and explain your reasoning.

Water reacts with HCl to give Cl⁻ and H₃O⁺, but with NH₃ to give NH₄⁺ and OH⁻. Explain how water can behave as both an acid and a base, identify the conjugate pairs in each reaction, and explain what this implies about how acid and base strength should be described.

Explain the difference between a strong acid and a weak acid, and write an equation for each.

Explain why 'strong' and 'concentrated' do not mean the same thing, using examples.

Equal volumes of 0.10 mol L⁻¹ hydrochloric acid and 0.10 mol L⁻¹ ethanoic acid are each reacted with excess magnesium ribbon. Predict and explain the difference in the initial rate of reaction and in the total volume of hydrogen produced, and explain why these two answers are not the same kind of comparison.

State which of a strong or a weak acid of the same concentration has the lower pH, and why.

Explain why a 0.1 mol L⁻¹ solution of hydrochloric acid conducts electricity better than a 0.1 mol L⁻¹ solution of ethanoic acid.

A student measures the pH of two acids and finds solution A (0.010 mol L⁻¹) has pH 2.0 and solution B (0.10 mol L⁻¹) has pH 2.9. The student concludes that A is the stronger acid. Evaluate this reasoning, and describe what further evidence would confirm the conclusion.

Calculate the pH of a 0.010 mol L⁻¹ solution of hydrochloric acid.

Calculate the pH of a 0.020 mol L⁻¹ solution of sodium hydroxide, showing all steps.

Calculate the pH of 0.0050 mol L⁻¹ sulfuric acid (H₂SO₄, assume both protons are fully donated). Then explain why the pH of a 0.10 mol L⁻¹ solution of ethanoic acid cannot be calculated using the same method, and state what additional information would be needed.