Lesson Notes By Weeks and Term v5 - Grade 12

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Performance objectives

• Define chemical equilibrium and its dynamic nature.
• State and apply Le Chatelier's Principle to predict shifts in equilibrium.
• Write the equilibrium constant expression (Kc) and calculate its value.
• Explain the effect of a catalyst on equilibrium.
• Use ICE tables to calculate equilibrium concentrations.

Lesson summary

Chemical equilibrium is a fundamental concept in chemistry that describes the state where the rate of the forward reaction equals the rate of the reverse reaction. This doesn't mean the reaction has stopped; rather, the concentrations of reactants and products remain constant over time. Understanding chemical equilibrium is crucial because it allows us to predict and control the outcomes of chemical reactions, which has wide-ranging implications in various fields. Consider the Haber process used to produce ammonia (NH₃) for fertilizers, vital for South African agriculture.

Lesson notes

2.1 Definition of Chemical Equilibrium Chemical equilibrium is a dynamic state in a reversible reaction where the rate of the forward reaction is equal to the rate of the reverse reaction. This means that reactants are being converted into products at the same rate that products are being converted back into reactants. At equilibrium, the net change in concentrations of reactants and products is zero.

Key Features: Dynamic: The reaction is still occurring in both directions.

Reversible Reaction: A reaction that can proceed in both the forward and reverse directions.

Closed System: A system where no reactants or products are added or removed.

Constant Macroscopic Properties: Measurable properties like concentration, pressure, and temperature remain constant at equilibrium. 2.2 Le Chatelier's Principle Le Chatelier's Principle states that if a change of condition (stress) is applied to a system in equilibrium, the system will shift in a direction that relieves the stress.

These stresses can be: Change in Concentration: Adding reactants will shift the equilibrium to favour product formation. Removing reactants will shift the equilibrium to favour reactant reformation. Adding products will shift the equilibrium to favour reactant reformation. Removing products will shift the equilibrium to favour product formation.

Change in Pressure: This only affects reactions involving gases. Increasing the pressure will shift the equilibrium towards the side with fewer moles of gas. Decreasing the pressure will shift the equilibrium towards the side with more moles of gas.

Change in Temperature: Increasing the temperature will favour the endothermic reaction (heat absorbed). Decreasing the temperature will favour the exothermic reaction (heat released). Important

Note: Le Chatelier's principle only predicts the direction of the shift, not the extent of the shift. 2.3 Equilibrium Constant (Kc) The equilibrium constant (Kc) is a numerical value that indicates the relative amounts of reactants and products at equilibrium. It is calculated using the following expression for a general reversible reaction: aA + bB ⇌ cC + dD Kc = ([C]^c [D]^d) / ([A]^a [B]^b)

Where: [A], [B], [C], and [D] are the equilibrium concentrations of reactants and products, respectively. a, b, c, and d are the stoichiometric coefficients from the balanced chemical equation.

Important Notes: Kc is temperature-dependent. Kc has no units. Pure solids and liquids are not included in the Kc expression because their concentrations effectively remain constant. This is crucial for heterogeneous equilibria.

Interpreting the Value of Kc: Kc > 1: The equilibrium lies to the right, favouring the formation of products. Kc 0 What will be the effect on the equilibrium position if: (a) The temperature is decreased? (b) The pressure is increased? (c) More N₂ is added? (d) A catalyst is added?

For the equilibrium: H₂(g) + I₂(g) ⇌ 2HI(g) At 700 K, Kc =

5

7. If 1.0 mol of H₂, 1.0 mol of I₂, and 1.0 mol of HI are mixed in a 1.0 L container, what are the concentrations of each substance at equilibrium?

Consider the following reaction: 2CO(g) + O₂(g) ⇌ 2CO₂(g) Initially, a container is filled with 0.500 M CO and 1.00 M O₂. At equilibrium, the concentration of CO₂ is found to be 0.35

M. Calculate the value of Kc.

For the reaction: N₂(g) + 3H₂(g) ⇌ 2NH₃(g) Kc = 0.060 at 500 K. If a flask contains 0.200 M N₂, 0.200 M H₂, and 0.100 M NH₃ at equilibrium, is the system truly at equilibrium, or will it shift to the left or the right to reach equilibrium? (Hint: calculate the reaction quotient, Q, and compare it to Kc). Write the equilibrium constant expression for each of the following reactions: (a) Fe(s) + 2HCl(aq) ⇌ FeCl₂(aq) + H₂(g) (b) 2H₂O(l) ⇌ 2H₂(g) + O₂(g) (c) AgCl(s) ⇌ Ag⁺(aq) + Cl⁻(aq) Consider the following reaction at equilibrium in a closed container: C(s) + H₂O(g) ⇌ CO(g) + H₂(g) ΔH > 0 How will each of the following changes affect the amount of CO(g) at equilibrium? (a) Adding more C(s) (b) Adding more H₂O(g) (c) Increasing the temperature (d) Decreasing the volume of the container The equilibrium constant (Kc) for the reaction: 2SO₂(g) + O₂(g) ⇌ 2SO₃(g) is 4.0 x 10⁻³ at 375 °C. If the initial concentrations of SO₂ and O₂ are both 0.50 M, what is the equilibrium concentration of SO₃?