Lesson Notes By Weeks and Term v5 - Grade 12

Lesson Video

This page supports the lesson note with a companion video and a short classroom-ready summary.

For class groups and homework, share this lesson page so learners also get the summary, objectives, and full lesson context.

Performance objectives

• Define chemical equilibrium as a dynamic state.
• State and apply Le Chatelier's principle to predict shifts in equilibrium.
• Write equilibrium constant (Kc) expressions and calculate Kc.
• Use Kc to predict the extent of a reaction.
• Explain the effect of a catalyst on equilibrium.

Lesson summary

Chemical equilibrium is a fundamental concept in chemistry that governs the extent to which reactions proceed. Understanding equilibrium is crucial for predicting the yield of products in chemical reactions, optimizing industrial processes, and understanding natural phenomena. In South Africa, many industries, such as fertilizer production (essential for agriculture) and mining (where chemical processes are used in ore extraction and refining), heavily rely on principles of chemical equilibrium. For example, the Haber process, used to produce ammonia for fertilizers, relies heavily on shifting the equilibrium to maximise ammonia production.

Lesson notes

2.1 Chemical Equilibrium: A Dynamic State Chemical equilibrium is a dynamic process where the rate of the forward reaction is equal to the rate of the reverse reaction. This means that reactants are constantly being converted into products, and products are constantly being converted back into reactants, but the net concentrations of reactants and products remain constant at equilibrium. This does NOT mean that the amounts of reactants and products are equal, only that their concentrations are no longer changing. For example, consider the reversible reaction: N 2 (g) + 3H 2 (g) ⇌ 2NH 3 (g) Initially, if we start with only N 2 and H 2 , the forward reaction (formation of NH 3 ) will be faster than the reverse reaction (decomposition of NH 3 ). As the concentration of NH 3 increases, the rate of the reverse reaction increases. Eventually, the rates of the forward and reverse reactions become equal, and the system reaches equilibrium.

Conditions for Equilibrium: The reaction must occur in a closed system (no exchange of matter with the surroundings). The temperature must be constant. 2.2 Le Chatelier's Principle Le Chatelier's principle states that if a change of condition is applied to a system in equilibrium, the system will shift in a direction that relieves the stress. The "stress" refers to changes in: Concentration: Adding reactants will shift the equilibrium towards the products. Adding products will shift the equilibrium towards the reactants. Removing reactants will shift the equilibrium towards the reactants. Removing products will shift the equilibrium towards the products.

Pressure: (Important for gaseous reactions) 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. If the number of moles of gas is the same on both sides, pressure has no effect.

Temperature: Increasing the temperature will shift the equilibrium towards the endothermic reaction (heat is absorbed). Decreasing the temperature will shift the equilibrium towards the exothermic reaction (heat is released).

Example of Le Chatelier's Principle: Consider the Haber process: N 2 (g) + 3H 2 (g) ⇌ 2NH 3 (g) ΔH = -92 kJ/mol (exothermic) Increasing the concentration of N 2 or H 2 : Shifts the equilibrium to the right, favoring the formation of NH 3 .

Increasing the pressure: Shifts the equilibrium to the right because there are 4 moles of gas on the reactant side (1 N 2 + 3 H 2 ) and only 2 moles of gas on the product side (2 NH 3 ).

Decreasing the temperature: Shifts the equilibrium to the right because the forward reaction is exothermic (releases heat). To favour product formation, a lower temperature is preferred, but a very low temperature decreases the rate of reaction drastically. Thus, in the real industrial setting, a compromise temperature is used. 2.3 Equilibrium Constant (K c ) The equilibrium constant, K c , is a quantitative measure of the extent to which a reaction proceeds to completion at a given temperature.

For the general reversible reaction: aA + bB ⇌ cC + dD The equilibrium constant expression is: K c = ([C] c [D] d ) / ([A] a [B] b ) where [A], [B], [C], and [D] are the equilibrium concentrations of the reactants and products, and a, b, c, and d are their respective stoichiometric coefficients. A large K c value (K c >> 1) indicates that the equilibrium lies to the right, favoring the formation of products. A small K c value (K c c ≈ 1 suggests that the equilibrium mixture contains significant amounts of both reactants and products. Important

Note: K c is temperature dependent. Changing the temperature will change the value of K c . Also, solids and liquids do not appear in the K c expression, only gases and aqueous solutions. 2.4 Calculating K c To calculate K c , you need to know the equilibrium concentrations of all reactants and products. You can obtain these values from experimental data, typically presented in an ICE (Initial, Change, Equilibrium) table.