Energetics, kinetics and equilibria are the quantitative heart of H2 Chemistry's Physical Chemistry pillar — three topics that together explain whether a reaction happens, how fast, and how far. They are calculation-heavy, and the students who do well keep the three questions — feasibility, rate and extent — clearly separated. This guide is from Ancourage Academy, whose JC H2 Chemistry tuition teaches these topics in small groups of 3–6 at Bishan and Woodlands.
This is a single-topic deep-dive — a sibling to our H2 Chemistry organic mastery and inorganic chemistry guides, and part of our wider H2 Chemistry overview.
If energetics or equilibria calculations are where marks slip, Ancourage Academy's JC1 H2 Chemistry programme drills the quantitative methods directly — book a trial class (usually $18) for a diagnostic assessment.
What Does Physical Chemistry Cover in H2 Chemistry?
In H2 Chemistry (9476), this part of the Physical Chemistry pillar covers energetics (enthalpy, entropy and Gibbs free energy), reaction kinetics (rate equations, reaction orders, and the qualitative effect of temperature via the Boltzmann distribution), and chemical and ionic equilibria (Kc, Kp, Le Chatelier's principle, pH and buffers). The SEAB Chemistry syllabus (9476) defines what is examinable, and these topics are tested with heavy cross-integration.
How Do You Handle Energetics?
Energetics asks whether a reaction is energetically feasible, using enthalpy (heat), entropy (disorder) and Gibbs free energy to combine them.
- Enthalpy (ΔH): found via Hess's law, energy cycles and Born–Haber cycles for ionic compounds.
- Entropy (ΔS): a measure of disorder; it increases when gases are produced or more particles result.
- Gibbs free energy (ΔG = ΔH − TΔS): the deciding quantity — a reaction is feasible when ΔG is negative.
The key insight students miss is that feasibility is about ΔG, not ΔH alone: an endothermic reaction can still be feasible if the entropy increase is large enough at a high enough temperature. Temperature units must be in kelvin in the ΔG equation.
How Do Reaction Kinetics Work?
Kinetics is about rate — how concentration, temperature and catalysts change the speed of a reaction — and it is governed by the experimentally determined rate equation.
| Concept | Key idea |
|---|---|
| Rate equation | rate = k[A]ᵐ[B]ⁿ, with orders m and n found by experiment |
| Order of reaction | Determined from data, not from the stoichiometric equation |
| Rate constant k | Increases with temperature |
| Boltzmann distribution | Explains qualitatively how raising temperature increases the rate constant k |
| Catalysis | Provides an alternative pathway with lower activation energy |
The most common kinetics error is reading the order of reaction off the balanced equation. Orders are experimental quantities — they must come from rate data, half-life analysis, or initial-rate experiments.
How Far Does a Reaction Go? Chemical and Ionic Equilibria
Equilibrium describes the extent of a reversible reaction, quantified by the equilibrium constant and predicted qualitatively by Le Chatelier's principle.
- Kc and Kp: equilibrium constants in terms of concentrations or partial pressures; their size shows how far a reaction proceeds.
- Le Chatelier's principle: a system at equilibrium shifts to oppose any imposed change in concentration, pressure or temperature.
- Ionic equilibria: acid–base behaviour, pH, the acid dissociation constant Ka, buffers, titration curves, and solubility equilibria — the solubility product Ksp and the common-ion effect.
A crucial distinction: only temperature changes the value of the equilibrium constant. Changing concentration can shift the position of equilibrium, and a pressure change shifts a gaseous equilibrium only when the two sides have different total moles of gas — but neither alters K, a point examiners test repeatedly.
The Most Common Physical Chemistry Mistakes
In our H2 Chemistry classes at Ancourage Academy, a handful of recurring errors cause most avoidable mark loss in these topics.
| Mistake | Why it happens | How to fix it |
|---|---|---|
| Judging feasibility by ΔH | Forgetting entropy's role | Use ΔG = ΔH − TΔS; feasible when ΔG < 0 |
| Order from the equation | Assuming stoichiometry gives order | Orders are experimental — read them from rate data |
| Thinking concentration changes K | Confusing position with constant | Only temperature changes K; concentration shifts position (pressure only for gas equilibria with unequal moles) |
| Celsius in ΔG | Using the wrong temperature unit | Always convert temperature to kelvin |
| Buffer set up wrongly | Mismatching acid and conjugate base | A buffer needs a weak acid with its conjugate base (or vice versa) |
How Does Physical Chemistry Connect to the Rest of H2 Chemistry?
These three topics underpin the whole subject and integrate with the other pillars.
- Inorganic chemistry: energetics explains Group 2 and Group 17 trends. See our inorganic chemistry guide.
- Organic chemistry: kinetics underlies reaction mechanisms. See our organic mastery guide.
- Maths foundation: pH and pKa calculations rely on logarithms, and rate and equilibrium problems use graphs. See our A-Maths logarithms guide.
A Study Plan for Mastering Physical Chemistry
Work these topics in order: energetics, then kinetics, then equilibria.
- Weeks 1–2 — energetics: master Hess's law, Born–Haber cycles, and ΔG = ΔH − TΔS feasibility.
- Weeks 3–4 — kinetics: determine orders from data and explain temperature and catalyst effects using the Boltzmann distribution.
- Weeks 5–6 — chemical equilibria: drill Kc, Kp and Le Chatelier's principle.
- Weeks 7–8 — ionic equilibria: work pH, Ka, buffers and titration curves under timed conditions.
Ancourage Academy's JC1 and JC2 H2 Chemistry programmes work through physical chemistry on this progression in small groups of 3–6. Book a trial class (usually $18) for a diagnostic, or WhatsApp us with any questions.
Common Questions About H2 Chemistry Physical Chemistry
What determines whether a reaction is feasible?
Feasibility is determined by the Gibbs free energy change, ΔG = ΔH − TΔS, not by enthalpy alone. A reaction is energetically feasible when ΔG is negative. This means an endothermic reaction (positive ΔH) can still be feasible if the entropy change ΔS is sufficiently positive and the temperature is high enough. Temperature must be in kelvin in the equation, and overlooking the entropy term is a frequent mistake.
How do you find the order of a reaction?
The order of a reaction with respect to each reactant is found experimentally, not from the balanced equation. Common methods are the initial-rates method (comparing how rate changes when one concentration is varied) and half-life analysis (a constant half-life indicates first order). The overall order is the sum of the individual orders. Reading orders off the stoichiometric coefficients is one of the most common kinetics errors.
What changes the value of the equilibrium constant?
Only a change in temperature changes the value of the equilibrium constant Kc or Kp. Changing the concentration of a species can shift the position of equilibrium, and a pressure change shifts a gaseous equilibrium only when the two sides have different total moles of gas — but K stays the same. A catalyst changes neither K nor the position of equilibrium; it speeds up the forward and reverse reactions equally, so equilibrium is simply reached faster. Confusing a shift in position with a change in the constant is a distinction examiners test repeatedly, so state it carefully.
What makes a good buffer solution?
A buffer resists changes in pH when small amounts of acid or base are added. It is made from a weak acid together with its conjugate base (for example ethanoic acid and ethanoate ions) or a weak base with its conjugate acid. The weak acid neutralises added base and the conjugate base neutralises added acid. Setting up the wrong acid–base pair, or using a strong acid, is a common reason buffer questions go wrong.
Related: H2 Chemistry Overview · Organic Chemistry Mastery · Inorganic Chemistry · A-Maths Logarithms · Electrochemistry (H2 Chem) · Atomic structure & chemical bonding