IMAT Chemistry · Free notes

Equilibrium, Kinetics and Energetics

Equilibrium, Kinetics and Energetics has 12 past questions since 2011, and the ministry papers from 2023 on have asked 5 of them. More than half are about equilibrium: which way it shifts when something is added or the temperature changes, and what the size of Kc means. The ministry papers have also asked short definitions, such as what the activation energy is, what exothermic means for the energy of the products, and how a rate law is found. The difficulty sits in keeping separate ideas apart: rate against yield, collision frequency against collision energy, and the one factor, temperature, that changes K.

The handout has every subtopic, worked example, formula and trap with answers shown. The formula sheet is the chapter’s formulas and traps alone, on one page.

IMAT home All formulas on one page

Subtopic notes

Formula & revision sheet

8 formulas · 9 reference tables · 20 gotchas across all subtopics — the exam-eve cheat-sheet

Exothermic and Endothermic Reactions, Bond Enthalpies and Hess's Law

Formulas (3)

Energy profiles and the activation energy · Energy profile relations $\Delta H = H_{\text{products}} - H_{\text{reactants}} \qquad E_{a,\text{reverse}} = E_{a,\text{forward}} - \Delta H$
Estimating ΔH from bond enthalpies · ΔH from bond enthalpies $\Delta H \approx \sum E(\text{bonds broken}) - \sum E(\text{bonds formed})$
Hess's law and enthalpies of formation · ΔH from enthalpies of formation $\Delta H = \sum \Delta H_f^\circ(\text{products}) - \sum \Delta H_f^\circ(\text{reactants})$

Reference tables (1)

Exothermic and endothermic reactions and the sign of ΔH5 rows

Feature	Exothermic	Endothermic
Heat flow	Released to the surroundings	Absorbed from the surroundings
Sign of ΔH	Negative	Positive
Energy of products	Lower than the reactants	Higher than the reactants
Temperature of surroundings	Rises	Falls
Examples	Combustion, respiration, neutralisation, freezing, condensation	Photosynthesis, thermal decomposition of $\mathrm{CaCO_3}$ , melting, dissolving ammonium nitrate

Watch out for (5)

Exothermic does not mean fast
→ Exothermic and endothermic reactions and the sign of ΔH
Breaking bonds never releases energy
→ Exothermic and endothermic reactions and the sign of ΔH
Activation energy is not the energy difference between products and reactants
→ Energy profiles and the activation energy
Bonds broken minus bonds formed, not the other way round
→ Estimating ΔH from bond enthalpies
Formation enthalpy is zero only for the standard form of an element
→ Hess's law and enthalpies of formation

Entropy, Gibbs Free Energy and Spontaneous Change

Formulas (1)

Gibbs free energy and whether a reaction is spontaneous · Gibbs free energy $\Delta G = \Delta H - T\Delta S$

Reference tables (1)

Entropy: a measure of the spreading of matter and energy5 rows

Change	Entropy of the system	Why
Ice melting	Increases	Particles become free to move
Steam condensing	Decreases	Gas becomes liquid
$\mathrm{CaCO_3(s) \rightarrow CaO(s) + CO_2(g)}$	Increases	A gas is produced from a solid
$\mathrm{N_2(g) + 3H_2(g) \rightarrow 2NH_3(g)}$	Decreases	4 mol of gas become 2 mol
Salt dissolving in water	Usually increases	Ions spread through the solution

Watch out for (3)

The entropy of the system can fall in a spontaneous change
→ Entropy: a measure of the spreading of matter and energy
Endothermic reactions can be spontaneous
→ Gibbs free energy and whether a reaction is spontaneous
Spontaneous does not mean fast
→ Gibbs free energy and whether a reaction is spontaneous

Reaction Rate, Collision Theory, Catalysts and Rate Laws

Formulas (2)

Measuring the rate of a reaction · Average rate $\text{rate} = \frac{\Delta[\text{reactant or product}]}{\Delta t}$
Rate laws and orders of reaction are found by experiment · Rate law $\text{rate} = k\,[\mathrm{A}]^m\,[\mathrm{B}]^n$

Reference tables (2)

Collision theory and the factors that change the rate5 rows

Factor increased	Effect on rate	Why
Concentration of a solution	Faster	More particles per volume, so more frequent collisions
Pressure of a gas	Faster	Particles closer together, so more frequent collisions
Surface area of a solid	Faster	More particles exposed, so more frequent collisions
Temperature	Faster	Mainly a larger fraction of collisions with energy at least $E_a$
Catalyst added	Faster	Lower activation energy, so more collisions succeed

Catalysts: a lower-energy route, not more product6 rows

Feature	What a catalyst does
Activation energy	Lowers it, for both the forward and the reverse reaction
Rate	Increases the rates of both directions equally
ΔH	No change
Equilibrium position, K and yield	No change; equilibrium is reached sooner
Amount at the end	Same as at the start
Examples	$\mathrm{MnO_2}$ for $\mathrm{H_2O_2}$ decomposition, iron in ammonia synthesis, enzymes

Watch out for (4)

A faster reaction does not make more product
→ Measuring the rate of a reaction
Pressure and concentration change how often particles collide, not how hard
→ Collision theory and the factors that change the rate
A catalyst never increases the yield
→ Catalysts: a lower-energy route, not more product
Orders are not the coefficients of the equation
→ Rate laws and orders of reaction are found by experiment

Dynamic Equilibrium and the Equilibrium Constant Kc

Formulas (2)

Writing the Kc expression · Equilibrium constant $K_c = \frac{[\mathrm{C}]^c\,[\mathrm{D}]^d}{[\mathrm{A}]^a\,[\mathrm{B}]^b}$
Calculating Kc and what its size means · Concentration from amount $c = \frac{n}{V}$

Reference tables (1)

Dynamic equilibrium in a closed system5 rows

Feature	At dynamic equilibrium
Forward and reverse rates	Equal, and both greater than zero
Concentrations	Constant, but usually not equal
System	Closed: no matter enters or leaves
Starting point	The same equilibrium is reached from either side
Evidence that it is dynamic	A labelled isotope added to one side soon appears on the other side

Watch out for (3)

Equilibrium does not mean equal amounts
→ Dynamic equilibrium in a closed system
Solids and pure liquids are not in Kc
→ Writing the Kc expression
A small Kc means few products, not fewer moles on one side
→ Calculating Kc and what its size means

Le Chatelier's Principle: Concentration, Pressure, Temperature and Catalysts

Reference tables (4)

Changing a concentration shifts the equilibrium to oppose it5 rows

Change	Shift	Reason
Add a reactant	To the right	The forward reaction uses up the extra reactant
Remove a product	To the right	The forward reaction replaces it
Add a product	To the left	The reverse reaction uses up the extra product
Add a substance that reacts with a product	To the right	That product is being removed
Add more of a pure solid	No shift	A solid is not in Kc

Changing the pressure: count the moles of gas4 rows

Equilibrium	Moles of gas, left : right	Higher pressure shifts
$\mathrm{N_2 + 3H_2 \rightleftharpoons 2NH_3}$	4 : 2	To the right
$\mathrm{N_2O_4 \rightleftharpoons 2NO_2}$	1 : 2	To the left
$\mathrm{H_2 + I_2 \rightleftharpoons 2HI}$	2 : 2	No shift
$\mathrm{CaCO_3(s) \rightleftharpoons CaO(s) + CO_2(g)}$	0 : 1	To the left

Changing the temperature: the only change that alters K3 rows

Forward reaction	Raise the temperature	Lower the temperature
Exothermic (ΔH negative)	Shifts left; K decreases	Shifts right; K increases
Endothermic (ΔH positive)	Shifts right; K increases	Shifts left; K decreases
Self-ionisation of water (endothermic)	Kw increases; pH falls below 7; still neutral	Kw decreases; pH rises above 7; still neutral

Putting the factors together: the compromise in ammonia synthesis4 rows

Condition	Effect on equilibrium yield	Effect on rate	Choice in the Haber process
Higher pressure	Increases	Increases	About 200 atm
Lower temperature	Increases	Decreases	A compromise near 450 °C
Catalyst	No effect	Increases	Iron
Removing ammonia	Increases the conversion	Keeps the forward reaction going	Ammonia condensed out

Watch out for (5)

Adding a solid does not shift an equilibrium
→ Changing a concentration shifts the equilibrium to oppose it
A reagent that removes a product shifts the equilibrium right
→ Changing a concentration shifts the equilibrium to oppose it
Count gas moles only
→ Changing the pressure: count the moles of gas
Heating an exothermic equilibrium lowers K
→ Changing the temperature: the only change that alters K
Pure water is neutral at every temperature
→ Changing the temperature: the only change that alters K

PYQ weightage by concept

17 concepts · 0 PYQs — where the marks actually sit, so you know what to drill first

Exothermic and Endothermic Reactions, Bond Enthalpies and Hess's Law0 PYQs · 0%

Concept	PYQs	Share
Exothermic and endothermic reactions and the sign of ΔHfoundation	—	—
Energy profiles and the activation energyfoundation	—	—
Estimating ΔH from bond enthalpiesfoundation	—	—
Hess's law and enthalpies of formationfoundation	—	—

Entropy, Gibbs Free Energy and Spontaneous Change0 PYQs · 0%

Concept	PYQs	Share
Entropy: a measure of the spreading of matter and energyfoundation	—	—
Gibbs free energy and whether a reaction is spontaneousfoundation	—	—

Reaction Rate, Collision Theory, Catalysts and Rate Laws0 PYQs · 0%

Concept	PYQs	Share
Measuring the rate of a reactionfoundation	—	—
Collision theory and the factors that change the ratefoundation	—	—
Catalysts: a lower-energy route, not more productfoundation	—	—
Rate laws and orders of reaction are found by experimentfoundation	—	—

Dynamic Equilibrium and the Equilibrium Constant Kc0 PYQs · 0%

Concept	PYQs	Share
Dynamic equilibrium in a closed systemfoundation	—	—
Writing the Kc expressionfoundation	—	—
Calculating Kc and what its size meansfoundation	—	—

Le Chatelier's Principle: Concentration, Pressure, Temperature and Catalysts0 PYQs · 0%

Concept	PYQs	Share
Changing a concentration shifts the equilibrium to oppose itfoundation	—	—
Changing the pressure: count the moles of gasfoundation	—	—
Changing the temperature: the only change that alters Kfoundation	—	—
Putting the factors together: the compromise in ammonia synthesisfoundation	—	—