Effects of temperature and concentration: Practice Questions

For a given reversible reaction at equilibrium in a closed system, which of the following changes will result in a change in the value of the equilibrium constant (\(K_c\))?

The gaseous reaction below is exothermic and has reached equilibrium in a closed container:
\(2\text{NO(g)} + \text{O}_2\text{(g)} \rightleftharpoons 2\text{NO}_2\text{(g)}\)
Which of the following changes would decrease the rate of the forward reaction but shift the equilibrium position to the right?

Consider the reaction \( X + Y \rightarrow Z \) with an activation energy \( E_{a} \). If the temperature of the system is raised from \( T_{1} \) to \( T_{2} \), which of the following statements regarding the Maxwell-Boltzmann distribution curve and the reaction rate is CORRECT?

Which of the following statements best explains why the rate of a chemical reaction generally increases as the concentration of the reactants increases at a constant temperature?

Consider the following gaseous equilibrium system in a closed container:
\(A(g) + 2B(g) \rightleftharpoons C(g) \quad \Delta H < 0\)
Which of the following combinations of changes will always result in the equilibrium position shifting to the left?

Using collision theory, state the effect of increasing the concentration of reactants on the number of collisions between reactant particles per unit time.

Consider the reaction between marble chips and hydrochloric acid. If the concentration of the acid is doubled while maintaining the same temperature and total volume, describe the change in the number of reactant particles and how this affects the frequency of effective collisions.

A student compares the initial rates of reaction for two mixtures: (A) 0.5 M HCl at 25 °C and (B) 0.1 M HCl at 60 °C reacting with excess zinc. Even though mixture B has a lower concentration, its initial rate is higher. Suggest two reasons, based on kinetic theory, why the temperature factor outweighs the concentration factor in this specific scenario.

A student conducts an experiment to investigate the reaction between zinc granules and dilute hydrochloric acid:

\(\text{Zn(s)} + 2\text{HCl(aq)} \to \text{ZnCl}_2\text{(aq)} + \text{H}_2\text{(g)}\)

(a) Predict the change in the initial rate of reaction if the concentration of the hydrochloric acid is increased while keeping the temperature and mass of zinc constant. (1 mark)

(b) Use the collision theory to explain why the reaction rate increases when the temperature of the reaction mixture is raised. (2 marks)

A student investigates the reaction between magnesium ribbon and dilute hydrochloric acid at different conditions. The reaction is represented by the following equation:
\(\text{Mg(s)} + 2\text{HCl(aq)} \to \text{MgCl}_2\text{(aq)} + \text{H}_2\text{(g)}\)

(a) In Experiment 1, \(0.1\text{ g}\) of magnesium ribbon reacts with excess \(1.0\text{ M HCl}\) at \(25^\circ\text{C}\). In Experiment 2, the same mass of magnesium and volume of acid are used, but the concentration of \(\text{HCl}\) is increased to \(2.0\text{ M}\). Explain, in terms of the collision theory, why the initial rate of reaction in Experiment 2 is higher than in Experiment 1. (2 marks)

(b) The student then conducts Experiment 3 at \(45^\circ\text{C}\) using \(1.0\text{ M HCl}\). It is observed that the reaction rate increases significantly compared to Experiment 1. Using the Maxwell-Boltzmann distribution curve, explain why this increase in temperature has a more pronounced effect on the rate than doubling the concentration. (3 marks)