WACE Chemistry · Unit 3
WACE Chemistry Unit 3: Equilibrium & Acid-Base — Flashcards & Quiz
WACE Chemistry ATAR Unit 3 explores chemical equilibrium systems and acid-base chemistry. These free flashcards and true/false questions cover dynamic equilibrium, Le Chatelier's principle, equilibrium constants (Kc, Kp), ICE tables, Brønsted-Lowry acids and bases, pH and pOH calculations, strong vs weak acids, Ka and Kb, buffer solutions, titration curves and indicator selection. Every card is aligned to the SCSA Chemistry ATAR syllabus so you revise exactly what appears in your Year 12 exams.
Key Terms
- Equilibrium Constant (K)
- A numerical value expressing the ratio of product concentrations to reactant concentrations at equilibrium, each raised to the power of their stoichiometric coefficients. The SCSA WACE Chemistry ATAR Unit 3 course requires students to write K expressions and interpret their magnitude.
- Le Chatelier's Principle
- When a system at equilibrium is subjected to a change in concentration, pressure or temperature, the system shifts to partially counteract that change. SCSA expects WACE ATAR students to predict equilibrium shifts and distinguish between changes that alter K (temperature only) and those that do not.
- Conjugate Acid-Base Pair
- Two species that differ by a single proton (H+), where one acts as a Bronsted-Lowry acid and the other as its conjugate base. The WACE ATAR Unit 3 course assessed by SCSA requires identification of conjugate pairs in acid-base equilibrium reactions.
- Buffer Solution
- A solution that resists changes in pH when small amounts of acid or base are added, typically composed of a weak acid and its conjugate base. SCSA expects Western Australian WACE students to explain buffer action using equilibrium shift principles.
- ICE Table
- A systematic method (Initial-Change-Equilibrium) for organising concentration data when calculating equilibrium constants or unknown concentrations. The SCSA WACE ATAR exam frequently requires students to set up and solve ICE tables in calculation questions.
- Dynamic Equilibrium
- A state in a reversible reaction where the rates of the forward and reverse reactions are equal, resulting in constant macroscopic concentrations despite ongoing molecular-level changes. SCSA assesses this concept as foundational to all equilibrium content in the WACE ATAR Unit 3 course.
Sample Flashcards
Q1: Define dynamic equilibrium in a chemical system.
Dynamic equilibrium occurs in a closed system when the rate of the forward reaction equals the rate of the reverse reaction. Macroscopic properties (concentrations, pressure, colour) remain constant while both reactions continue at the molecular level. It is "dynamic" because reactions have not stopped and "equilibrium" because there is no net change in composition.
Q2: State Le Chatelier's principle and explain its significance.
If a change in conditions (concentration, pressure, temperature) is imposed on a system at equilibrium, the system will shift to partially counteract the change and establish a new equilibrium. It predicts the direction of shift but not the extent. Only temperature changes alter the value of K.
Q3: Write the equilibrium constant expression (Kc) for a general reaction and explain what its magnitude indicates.
For aA + bB ⇌ cC + dD: Kc = [C]^c[D]^d / [A]^a[B]^b. Concentrations are equilibrium values in mol L⁻¹. Pure solids and liquids are excluded. Kc ≫ 1 means products are favoured; Kc ≪ 1 means reactants are favoured; Kc ≈ 1 means significant amounts of both are present.
Q4: What factors change the value of Kc?
Only temperature changes Kc. Increasing temperature increases Kc for endothermic reactions and decreases Kc for exothermic reactions. Changes in concentration, pressure or the addition of a catalyst do NOT change Kc — they may shift the position of equilibrium but the ratio of products to reactants at the new equilibrium still gives the same Kc value.
Q5: Explain how to use an ICE table to calculate equilibrium concentrations.
ICE stands for Initial, Change, Equilibrium. Steps: 1) Write initial concentrations. 2) Define the change using stoichiometric ratios and variable x. 3) Express equilibrium concentrations as Initial ± Change. 4) Substitute into the Kc expression and solve for x. 5) Calculate equilibrium concentrations.
Q6: How does changing pressure affect a gaseous equilibrium?
Increasing pressure (decreasing volume) shifts equilibrium toward the side with fewer moles of gas. Decreasing pressure shifts toward more moles of gas. If both sides have equal gas moles, pressure has no effect. Adding an inert gas at constant volume does NOT shift equilibrium because partial pressures of reactants and products are unchanged.
Q7: How does temperature change affect an exothermic equilibrium?
For an exothermic forward reaction: increasing temperature shifts equilibrium LEFT (toward reactants) and decreases Kc, because the system absorbs excess heat via the endothermic reverse reaction. Decreasing temperature shifts RIGHT and increases Kc. Temperature is the only factor that changes the value of K.
Q8: Define Brønsted-Lowry acids and bases and identify conjugate pairs.
A Brønsted-Lowry acid is a proton (H⁺) donor; a base is a proton acceptor. Every acid-base reaction involves two conjugate pairs. A conjugate acid-base pair differs by one proton. Water is amphoteric — it acts as an acid or base depending on the other reactant.
Sample Quiz Questions
Q1: At dynamic equilibrium, the concentrations of reactants and products are always equal.
Answer: FALSE
At equilibrium the concentrations are CONSTANT but not necessarily equal. The ratio is determined by the value of Kc.
Q2: Dynamic equilibrium can only be established in a closed system.
Answer: TRUE
A closed system prevents matter from escaping. In an open system, products or reactants can leave, preventing equilibrium from being established.
Q3: Adding a catalyst to a system at equilibrium changes the value of Kc.
Answer: FALSE
A catalyst speeds up both forward and reverse reactions equally. It does not change Kc or the equilibrium position — only temperature changes Kc.
Q4: Pure solids and liquids are excluded from the Kc expression.
Answer: TRUE
Pure solids and pure liquids have constant concentrations (activities) and are excluded from equilibrium expressions. Only aqueous and gaseous species are included.
Q5: Increasing pressure always shifts equilibrium toward the products.
Answer: FALSE
Increasing pressure shifts equilibrium toward the side with FEWER moles of gas, which could be reactants or products depending on the reaction.
Why It Matters
Chemical equilibrium is one of the most conceptually demanding topics in WACE Chemistry, but mastering it unlocks your ability to predict and explain reaction behaviour in both exam questions and laboratory contexts. Understanding how reversible reactions reach a dynamic balance, how Le Chatelier's principle predicts responses to disturbances, and how to calculate and interpret equilibrium constants is essential for the external exam. This topic also provides the theoretical foundation for acid-base chemistry and industrial chemistry applications. Examiners frequently test equilibrium through multi-step problems that integrate calculation with conceptual explanation. Equilibrium concepts underpin acid-base and redox chemistry later in the course, so a solid foundation here is essential for sustained success. Exam questions on equilibrium often appear in the calculation section and require you to set up ICE tables correctly, so practise these under timed conditions.
Key Concepts
Dynamic Equilibrium
At equilibrium, forward and reverse reactions occur at equal rates, producing constant concentrations of reactants and products. Understand that equilibrium is dynamic (reactions continue) and can only be established in closed systems. Recognise equilibrium on concentration-time and rate-time graphs.
Le Chatelier's Principle
When a system at equilibrium is disturbed by changes in concentration, temperature, or pressure, it shifts to partially counteract the change. Practise predicting the direction of shift for each type of disturbance and explaining the effect on equilibrium position and the value of the equilibrium constant.
Equilibrium Constant (K)
The equilibrium constant expression relates product and reactant concentrations at equilibrium. Learn to write Keq expressions, calculate unknown concentrations using ICE tables, and interpret the magnitude of K — a large K favours products, a small K favours reactants. Remember that K changes only with temperature.
Industrial Applications of Equilibrium
Industrial processes like the Haber process and Contact process apply equilibrium principles to maximise yield. Study how temperature, pressure, and catalyst choices represent compromises between equilibrium position and reaction rate, and explain why actual operating conditions differ from those predicted by Le Chatelier's principle alone.
Common Mistakes to Avoid
- Stating that the equilibrium constant K changes when concentration or pressure is altered — only temperature changes K; SCSA WACE ATAR marking guides specifically check that students understand this distinction.
- Including pure solids or pure liquids in equilibrium constant expressions — the SCSA WACE course requires students to omit species with constant activity (pure solids and liquids) from K expressions.
- Confusing equilibrium shift with reaching a new equilibrium — WACE examiners expect students to explain that Le Chatelier's principle predicts the direction of shift, and that the system then establishes a new equilibrium with different concentrations but the same K (unless temperature changed).
- Mixing up strong acids with weak acids when applying equilibrium concepts — SCSA WACE exam questions require students to recognise that strong acids dissociate completely and therefore equilibrium calculations apply only to weak acid and weak base systems.
Study Tips
- Master ICE tables by working through at least ten practice problems of increasing difficulty — this structured approach prevents calculation errors under exam pressure.
- Build a flashcard deck for Le Chatelier's principle covering every disturbance type, and use spaced repetition to ensure you can predict shifts instantly.
- Always state whether K changes when analysing a disturbance — only temperature changes K, and examiners specifically check for this understanding.
- Practise writing equilibrium constant expressions for unfamiliar reactions, remembering to exclude pure solids and liquids from the expression.
- For industrial chemistry questions, explicitly discuss the trade-off between thermodynamic yield and kinetic rate to demonstrate deeper understanding.
- Before your exam, work through the practice questions in this set at least twice using spaced repetition. Testing yourself repeatedly is the most effective revision strategy for long-term retention.
Related Topics
Frequently Asked Questions
What does WACE Chemistry Unit 3 Equilibrium cover?
Unit 3 covers dynamic equilibrium, Le Chatelier's principle, equilibrium constants (Kc and Kp), ICE tables, Brønsted-Lowry acid-base theory, pH calculations, strong and weak acids, buffers, titrations and indicator selection.
How many flashcards are in this set?
This free set contains 20 flashcards and 20 true/false quiz questions covering all key equilibrium and acid-base concepts, aligned to the SCSA WACE Chemistry ATAR syllabus.
Are these flashcards aligned to the WA SCSA syllabus?
Yes — every flashcard and quiz question is mapped to SCSA syllabus content for WACE Chemistry ATAR Unit 3.
Last updated: March 2026 · 20 flashcards · 20 quiz questions · Content aligned to the SCSA Curriculum