HSC Chemistry Module 5: Everything You Need to Know About Equilibrium

If you breezed through Modules 1 to 4 by memorising definitions and plugging numbers into formulas, Module 5 is going to feel like a different subject. Equilibrium is the first topic in HSC Chemistry that genuinely demands conceptual thinking. You cannot just learn "what happens"; you need to understand why it happens. That is what catches students off guard, and it is also what makes this module one of the most heavily examined in the HSC.

The good news? Once the core logic clicks, every question in this module becomes a variation of the same few ideas. Let's break them down.

What Is Chemical Equilibrium?

Chemical equilibrium is not about reactions stopping. That is the single biggest misconception in this module. At equilibrium, both the forward and reverse reactions are still occurring. They are just happening at the same rate, so the overall concentrations of reactants and products remain constant. This is why we call it dynamic equilibrium, not static equilibrium.

For equilibrium to be established, you need two conditions: the reaction must be reversible, and it must take place in a closed system (nothing enters or leaves). In an open beaker where gas can escape, the system can never reach equilibrium because products are being removed.

Think of it like two escalators running in opposite directions. People are still moving (the reactions are still going), but the number of people on each floor stays the same. That is dynamic equilibrium.

Le Chatelier's Principle: The Golden Rule

If there is one idea that carries you through almost every Module 5 question, it is Le Chatelier's Principle. It states that if a system at equilibrium is subjected to a change, the system will shift in the direction that partially opposes that change. It does not fully reverse it. It partially counteracts it.

There are three types of changes you need to know:

• Concentration: If you increase the concentration of a reactant, the system shifts forward to use up some of the excess. If you remove a product, the system also shifts forward to replace what was lost. The classic example is the N₂O₄/NO₂ equilibrium. Add more N₂O₄, and you will observe the mixture becoming browner as more NO₂ is produced.
• Temperature: This one depends on whether the reaction is exothermic or endothermic. Increasing temperature favours the endothermic direction, because the system shifts to absorb the added heat. This is the only change that actually alters the value of K_eq.
• Pressure (for gaseous systems): Increasing pressure causes the system to shift toward the side with fewer moles of gas. The system reduces pressure by reducing the total number of gas molecules.

A common mistake is saying the system "restores" the original conditions. It does not. It partially counteracts the change. After the shift, the new equilibrium position is different from the original one.

K_eq and ICE Tables

The equilibrium constant, K_eq, is a ratio of the concentrations of products to reactants at equilibrium, each raised to the power of their coefficients. For a generic reaction aA + bB ⇌ cC + dD, the expression is:

K_eq = [C]^c[D]^d / [A]^a[B]^b

What K_eq tells you is straightforward. A large K_eq (much greater than 1) means products are favoured at equilibrium. A small K_eq (much less than 1) means reactants are favoured. A K_eq close to 1 means neither side strongly dominates.

To solve quantitative equilibrium problems, you will almost always use an ICE table (Initial, Change, Equilibrium). The process is: write the balanced equation, fill in the initial concentrations, express the changes in terms of a variable (usually x), calculate the equilibrium row, and substitute into the K_eq expression to solve.

You should also understand the reaction quotient Q. Q uses the same expression as K_eq, but with current concentrations rather than equilibrium concentrations. If Q < K_eq, the reaction shifts forward. If Q > K_eq, it shifts in reverse. If Q = K_eq, the system is at equilibrium.

Solubility Equilibrium

This is where Module 5 extends into slightly different territory. When a sparingly soluble salt dissolves in water, the dissolved ions reach equilibrium with the undissolved solid. The equilibrium constant for this process is called the solubility product, K_sp.

For example, for AgCl(s) ⇌ Ag⁺(aq) + Cl^-(aq), the expression is K_sp = [Ag⁺][Cl^-]. The solid does not appear in the expression because its concentration is constant.

The common ion effect is an important application here. If you add a solution containing Cl^- ions to a saturated AgCl solution, the increased [Cl^-] means the ion product exceeds K_sp, so the equilibrium shifts to the left and more AgCl precipitates out. This is Le Chatelier's Principle applied to solubility.

Exam Tips for Module 5

The HSC examiners are looking for specific things in your equilibrium responses. Knowing the content is only half the battle. You also need to communicate it in the way they expect.

1. Explain Le Chatelier's shifts using collision theory. Do not just say "the equilibrium shifts forward." Explain why. For a concentration increase, say that the higher concentration of reactant molecules increases the frequency of effective collisions in the forward direction, so the forward rate temporarily exceeds the reverse rate until a new equilibrium is established.
2. Always state conditions clearly. When discussing temperature effects, state whether the forward reaction is exothermic or endothermic. When discussing pressure, note the number of moles of gas on each side.
3. Show all working in calculations. Write the balanced equation, the K_eq expression, the ICE table, and every substitution step. Marks are awarded for method, not just the final answer.
4. Do not confuse "rate" with "position." A catalyst increases the rate of both forward and reverse reactions equally. It does not change the equilibrium position or K_eq.

Module 5 rewards students who think carefully about cause and effect. If you can explain the "why" behind every shift and every calculation, you are in a strong position for the exam.