Energetics & Thermodynamics

Learn and apply Energetics & Thermodynamics in the published Chemistry course sequence.

GCE A-Level H2 Chemistry 9476-2027

Learning goals

Enthalpy Changes and Energy Profiles
Calorimetry (q = mcΔT)
Hess’ Law and Cycles
Bond Enthalpy Calculations
Lattice Energy and Born–Haber Cycles
Entropy and Gibbs Free Energy

H2 Energetics and Thermodynamics learning outcomes

This hub focuses on the exam-ready cycles and decision logic, not memorising isolated definitions.

Before you start

Beginner path: follow the Lessons (Recommended Order) from top to bottom.

Revision: Jump to: Quick Reference · What You Must Memorise · Common Exam Traps · Hub Quiz.

Prerequisites: skim the Prerequisites list first if you’re rusty.

Useful links: A Level portal · Exam Skills

What You’ll Learn

Use sign conventions and energy profile diagrams correctly.
Do calorimetry, Hess’ law, and Born–Haber calculations with clean working and units.
Use Δ G = Δ H - TΔ S to decide feasibility under given conditions.

Lessons (Recommended Order)

Enthalpy Changes and Energy Profiles
Signs, Ea, catalyst diagrams.
Open lesson
Calorimetry (q = mcΔT)
q to ΔH with correct signs and units.
Open lesson
Hess’ Law and Cycles
Build cycles and avoid sign/scale traps.
Open lesson
Bond Enthalpy Calculations
Estimate ΔH from bond data.
Open lesson
Lattice Energy and Born–Haber Cycles
Lattice definitions plus full Born–Haber sums.
Open lesson
Entropy and Gibbs Free Energy
ΔS and ΔG feasibility decisions.
Open lesson

Quick Reference

Task	Use	Unit check
calorimetry heat	q = mcΔ T	m in g, c in J g⁻¹ K⁻¹, Δ T in K (a °C change is numerically the same)
convert to enthalpy	Δ H = -qₛₒₗᵤₜᵢₒₙ/n	convert J → kJ; n in mol
Hess (formation)	Δ H_reaction⦵ = ∑ ν Δ H_f⦵(products) - ∑ ν Δ H_f⦵(reactants)	include coefficients
bond enthalpy estimate	Δ H ≈ ∑ E(broken) - ∑ E(formed)	estimate only
Gibbs feasibility	Δ G = Δ H - TΔ S	T in K; match kJ vs J
standard conditions	T = 298 K; p = 100 kPa; Δ S usually in J mol⁻¹ K⁻¹ (convert before Δ G)	quote conditions when using ⦵ data
Δ G temperature logic	feasible if Δ G < 0; increasing T favours feasibility when Δ S > 0 (since -TΔ S becomes more negative)	always use T in K
Born–Haber checklist	atomisation → IE₁/IE₂… → EA₁/EA₂… → lattice → formation	include signs and coefficients

What You Must Memorise

Exothermic vs endothermic: exothermic Δ H < 0; endothermic Δ H > 0.
Standard conditions (A Level): typically 298 K and 100 kPa; solutions at 1.00 mol dm⁻³ when relevant.
Calorimetry chain: qₛₒₗᵤₜᵢₒₙ = mcΔ T; q_reaction = -qₛₒₗᵤₜᵢₒₙ; Δ H = q_reaction/n (report in kJ mol⁻¹).
Hess rules: reverse equation → change sign; multiply equation by k → multiply Δ H by k.
Bond enthalpy estimate: Δ H ≈ ∑ E(broken) - ∑ E(formed) (mean values → estimate).
Lattice enthalpy conventions: dissociation is the reverse of formation, so signs are opposite.
Gibbs unit check: use T in K and make the energy units match before using Δ G = Δ H - TΔ S; feasibility requires Δ G < 0 under the stated conditions.

Hub Quiz and Check Your Understanding

Use the quiz without notes first. Work out whether the difficult step was the sign convention, energy-term definition, stoichiometric scaling, cycle direction, unit conversion or feasibility conclusion.

H2 Energetics and Thermodynamics Quiz
Check profiles, calorimetry, Hess and Born–Haber cycles, bond enthalpy, entropy and Gibbs reasoning.
Start hub quiz
H2 Energetics Knowledge Check
Find the reasoning step that needs more work, use the feedback, then try a fresh question independently.
Practice onlyCheck understanding

Practise

Work through questions with marking and feedback as you learn.

About 10 minutes

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A reaction has ΔH = +58 kJ mol⁻¹ and a forward activation energy of +75 kJ mol⁻¹. Which description of its energy profile is correct?

The products are 58 kJ mol⁻¹ above the reactants, the barrier is 75 kJ mol⁻¹ above the reactants, and the reverse activation energy is 17 kJ mol⁻¹.
The products are 58 kJ mol⁻¹ below the reactants, and the barrier is 75 kJ mol⁻¹ above the reactants.
The products are 58 kJ mol⁻¹ above the reactants, and the reverse activation energy is 133 kJ mol⁻¹.
The products are 58 kJ mol⁻¹ above the reactants, and the barrier is 75 kJ mol⁻¹ above the products.

Which statement defines the standard enthalpy change of neutralisation?

The enthalpy change when an acid and a base react in dilute aqueous solution to form 1 mol of water under standard conditions
The enthalpy change when 1 mol of acid reacts completely with a base under standard conditions
The enthalpy change when an acid and a base react to form 1 mol of water at 273 K and 1 atm
The heat absorbed when 1 mol of salt is formed from an acid and a base

Burning 0.460 g of ethanol (Mᵣ 46.0) heats 200 g of water by 12.0 K (c = 4.18 J g⁻¹ K⁻¹). Assuming no heat loss, what is the enthalpy change of combustion of ethanol?

−1.00 × 10³ kJ mol⁻¹
−2.31 kJ mol⁻¹
−10.0 kJ mol⁻¹
+1.00 × 10³ kJ mol⁻¹

Standard enthalpy changes of formation: C₂H₄(g) +52; C₂H₆(g) −85 kJ mol⁻¹. What is ΔH for C₂H₄(g) + H₂(g) → C₂H₆(g)?

−137 kJ mol⁻¹
+137 kJ mol⁻¹
−33 kJ mol⁻¹
−85 kJ mol⁻¹

Average bond energies (kJ mol⁻¹): N≡N 994, H–H 436, N–H 390. Estimate ΔH for N₂(g) + 3H₂(g) → 2NH₃(g).

−38 kJ mol⁻¹
+38 kJ mol⁻¹
+1132 kJ mol⁻¹
−910 kJ mol⁻¹

Which statement about lattice energy is correct?

It is the enthalpy change when 1 mol of an ionic solid forms from its gaseous ions; LiF has a more exothermic value than CsI because its ions are smaller and closer together.
It is the enthalpy change when 1 mol of an ionic solid forms from its elements in their standard states.
It is the energy absorbed when gaseous ions form 1 mol of an ionic solid.
It is the enthalpy change when 1 mol of an ionic solid forms from its gaseous ions; CsI has a more exothermic value than LiF because its ions have more electrons.

Data (kJ mol⁻¹): ΔH_f MgO(s) −602; ΔH_at Mg +148; first and second ionisation energies of Mg +738 and +1451; ΔH_at O (½O₂ → O) +249; first electron affinity O −141; second electron affinity O +798. What is the lattice energy of MgO?

−3845 kJ mol⁻¹
−2394 kJ mol⁻¹
−2249 kJ mol⁻¹
−4094 kJ mol⁻¹

Which reaction has the most positive entropy change, ΔS?

2NaHCO₃(s) → Na₂CO₃(s) + H₂O(g) + CO₂(g)
N₂(g) + 3H₂(g) → 2NH₃(g)
H₂O(l) → H₂O(s)
Ag⁺(aq) + Cl⁻(aq) → AgCl(s)

Practise after feedback

After a check, practise the skills it showed you need to work on.

About 10 minutes

Questions are picked at random each time you start. You'll see the answer after each question. It's for practice only and doesn't count towards mastery.

Recent attempts

History is stored only in this browser.

No completed attempts are saved yet.

This check needs JavaScript to record answers. The reviewed questions are listed below for study.

A reaction has ΔH = +58 kJ mol⁻¹ and a forward activation energy of +75 kJ mol⁻¹. Which description of its energy profile is correct?

The products are 58 kJ mol⁻¹ above the reactants, the barrier is 75 kJ mol⁻¹ above the reactants, and the reverse activation energy is 17 kJ mol⁻¹.
The products are 58 kJ mol⁻¹ below the reactants, and the barrier is 75 kJ mol⁻¹ above the reactants.
The products are 58 kJ mol⁻¹ above the reactants, and the reverse activation energy is 133 kJ mol⁻¹.
The products are 58 kJ mol⁻¹ above the reactants, and the barrier is 75 kJ mol⁻¹ above the products.

Which statement defines the standard enthalpy change of neutralisation?

The enthalpy change when an acid and a base react in dilute aqueous solution to form 1 mol of water under standard conditions
The enthalpy change when 1 mol of acid reacts completely with a base under standard conditions
The enthalpy change when an acid and a base react to form 1 mol of water at 273 K and 1 atm
The heat absorbed when 1 mol of salt is formed from an acid and a base

Burning 0.460 g of ethanol (Mᵣ 46.0) heats 200 g of water by 12.0 K (c = 4.18 J g⁻¹ K⁻¹). Assuming no heat loss, what is the enthalpy change of combustion of ethanol?

−1.00 × 10³ kJ mol⁻¹
−2.31 kJ mol⁻¹
−10.0 kJ mol⁻¹
+1.00 × 10³ kJ mol⁻¹

Standard enthalpy changes of formation: C₂H₄(g) +52; C₂H₆(g) −85 kJ mol⁻¹. What is ΔH for C₂H₄(g) + H₂(g) → C₂H₆(g)?

−137 kJ mol⁻¹
+137 kJ mol⁻¹
−33 kJ mol⁻¹
−85 kJ mol⁻¹

Average bond energies (kJ mol⁻¹): N≡N 994, H–H 436, N–H 390. Estimate ΔH for N₂(g) + 3H₂(g) → 2NH₃(g).

−38 kJ mol⁻¹
+38 kJ mol⁻¹
+1132 kJ mol⁻¹
−910 kJ mol⁻¹

Which statement about lattice energy is correct?

It is the enthalpy change when 1 mol of an ionic solid forms from its gaseous ions; LiF has a more exothermic value than CsI because its ions are smaller and closer together.
It is the enthalpy change when 1 mol of an ionic solid forms from its elements in their standard states.
It is the energy absorbed when gaseous ions form 1 mol of an ionic solid.
It is the enthalpy change when 1 mol of an ionic solid forms from its gaseous ions; CsI has a more exothermic value than LiF because its ions have more electrons.

−3845 kJ mol⁻¹
−2394 kJ mol⁻¹
−2249 kJ mol⁻¹
−4094 kJ mol⁻¹

Which reaction has the most positive entropy change, ΔS?

2NaHCO₃(s) → Na₂CO₃(s) + H₂O(g) + CO₂(g)
N₂(g) + 3H₂(g) → 2NH₃(g)
H₂O(l) → H₂O(s)
Ag⁺(aq) + Cl⁻(aq) → AgCl(s)

Check what I know

Start here to see which parts you already know.

About 8 minutes

Answer 8 short questions. It shows what to work on next and doesn't count towards mastery.

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This check needs JavaScript to record answers. The reviewed questions are listed below for study.

An energy profile diagram shows the reactants at 0 kJ mol⁻¹, the top of the energy barrier at +180 kJ mol⁻¹ and the products at −90 kJ mol⁻¹. Which statement is correct?

ΔH = −90 kJ mol⁻¹ and Eₐ = +180 kJ mol⁻¹; the reaction is exothermic because more energy is released forming bonds than is absorbed breaking bonds.
ΔH = −90 kJ mol⁻¹ and Eₐ = +270 kJ mol⁻¹; the reaction is exothermic.
ΔH = +90 kJ mol⁻¹ and Eₐ = +180 kJ mol⁻¹; the reaction is endothermic.
ΔH = −90 kJ mol⁻¹ and Eₐ = +180 kJ mol⁻¹; the reaction is exothermic because breaking bonds releases energy.

Which equation has an enthalpy change equal to the standard enthalpy change of formation of ethanol, C₂H₅OH(l)?

2C(s, graphite) + 3H₂(g) + ½O₂(g) → C₂H₅OH(l)
C₂H₅OH(l) + 3O₂(g) → 2CO₂(g) + 3H₂O(l)
2C(g) + 6H(g) + O(g) → C₂H₅OH(l)
4C(s, graphite) + 6H₂(g) + O₂(g) → 2C₂H₅OH(l)

50.0 cm³ of 1.00 mol dm⁻³ HCl is mixed with 50.0 cm³ of 1.00 mol dm⁻³ NaOH, and the temperature rises by 6.8 °C. Assume the solution has density 1.00 g cm⁻³ and c = 4.18 J g⁻¹ K⁻¹. What is the enthalpy change of neutralisation?

−56.8 kJ mol⁻¹
−28.4 kJ mol⁻¹
+56.8 kJ mol⁻¹
−5.68 × 10⁴ kJ mol⁻¹

Standard enthalpy changes of combustion (kJ mol⁻¹): C(graphite) −394; H₂(g) −286; C₃H₈(g) −2222. What is the standard enthalpy change of formation of propane, 3C(graphite) + 4H₂(g) → C₃H₈(g)?

−104 kJ mol⁻¹
+104 kJ mol⁻¹
−4548 kJ mol⁻¹
+1542 kJ mol⁻¹

Average bond energies (kJ mol⁻¹): C–H 410, Cl–Cl 244, C–Cl 340, H–Cl 431. Use them to estimate ΔH for CH₄(g) + Cl₂(g) → CH₃Cl(g) + HCl(g).

−117 kJ mol⁻¹
+117 kJ mol⁻¹
−771 kJ mol⁻¹
+1425 kJ mol⁻¹

Which compound has the most exothermic lattice energy, and why? NaCl, KCl, NaF or MgO

MgO: its ions have charges of 2+ and 2− and small radii, so the most energy is released when its gaseous ions form the solid lattice.
KCl: its larger ions release more energy when they come together.
NaF: fluorine is the most electronegative element.
NaCl: it needs the most energy to break its lattice into gaseous ions.

Data (kJ mol⁻¹): ΔH_f NaCl(s) −411; ΔH_at Na +107; first ionisation energy Na +496; ΔH_at Cl (½Cl₂ → Cl) +122; first electron affinity Cl −349. What is the lattice energy of NaCl?

−787 kJ mol⁻¹
−1485 kJ mol⁻¹
−909 kJ mol⁻¹
+787 kJ mol⁻¹

For CaCO₃(s) → CaO(s) + CO₂(g), ΔH = +178 kJ mol⁻¹ and ΔS = +161 J K⁻¹ mol⁻¹. Assuming these do not change with temperature, when is the reaction feasible?

Above about 1106 K
Above about 1.11 K
At all temperatures, because ΔS is positive
At no temperature, because ΔH is positive

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Energetics & Thermodynamics

H2 Energetics and Thermodynamics learning outcomes

What You’ll Learn

Lessons (Recommended Order)

Enthalpy Changes and Energy Profiles

Calorimetry (q = mcΔT)

Hess’ Law and Cycles

Bond Enthalpy Calculations

Lattice Energy and Born–Haber Cycles

Entropy and Gibbs Free Energy

Quick Reference

What You Must Memorise

Hub Quiz and Check Your Understanding

H2 Energetics and Thermodynamics Quiz

H2 Energetics Knowledge Check

Practise

Practice complete

Your answers

Practise after feedback

Practice complete

Your answers

Check what I know

Starting check complete

Your answers

Check my progress

Progress check complete

Your answers

Check again

Progress check complete

Your answers

Review

Review complete

Your answers