Chemical Energetics: Thermochemistry

Learn and apply Chemical Energetics: Thermochemistry in the published Chemistry course sequence.

GCE A-Level H1 Chemistry 8873-2027

Learning goals

explain that most chemical reactions are accompanied by energy changes, principally in the form of heat usually associated with the breaking and forming of chemical bonds; the reaction can be exothermic (∆H negative) or endothermic (∆H positive)
construct and interpret an energy profile diagram, in terms of the enthalpy change of the reaction and of the activation energy (see also Section 7)
explain and use the terms: — enthalpy change of reaction and standard conditions, with particular reference to: formation; combustion; neutralisation
explain and use the terms: — bond energy (∆H positive, i.e. bond breaking) (see also Section 2)
explain and use the terms: — lattice energy (∆H negative, i.e. gaseous ions to solid lattice)
calculate enthalpy changes from appropriate experimental results, including the use of the relationship: heat change = mc∆T
explain, in qualitative terms, the effect of ionic charge and of ionic radius on the numerical magnitude of a lattice energy
apply Hess’ Law to carry out calculations involving given simple energy cycles and relevant energy terms (restricted to enthalpy changes of formation, combustion and neutralisation), with particular reference to: — determining enthalpy changes that cannot be found by direct experiment, e.g. an enthalpy change of formation from enthalpy changes of combustion
apply Hess’ Law to carry out calculations involving given simple energy cycles and relevant energy terms (restricted to enthalpy changes of formation, combustion and neutralisation), with particular reference to: — average bond energies [construction of energy cycles is not required]

H1 Energetics learning outcomes

Energetics connects observable temperature change to energy transfer, chemical bonding and route-independent enthalpy. The strongest solutions define the relevant energy term, preserve signs and units, and make every stoichiometric scale factor visible.

Follow the five lessons in order for first learning. During revision, classify each error as definition, sign, unit, mole basis, cycle direction, bond count or lattice comparison.

What You’ll Learn

Explain exothermic and endothermic reactions through heat transfer, bond breaking and bond forming, using the correct sign of ΔH.
Construct and interpret energy profiles, including forward and reverse activation energies and the effect of a catalyst.
Define the required standard enthalpy changes of reaction, formation, combustion and neutralisation with an explicit mole basis.
Convert calorimetry evidence through q = mcΔT into a signed molar enthalpy change with consistent J–kJ units.
Apply Hess’ Law to supplied simple cycles involving formation, combustion and neutralisation data.
Estimate reaction enthalpy from average bond enthalpies and explain why the result is approximate.
Explain lattice-energy magnitude by comparing ionic charge and radius without constructing a Born–Haber cycle.

Lessons (Recommended Order)

Begin with energy profiles, build the experimental calorimetry chain, then apply Hess’ Law, average bond enthalpies and qualitative lattice reasoning.

Enthalpy Changes and Energy Profiles
Connect enthalpy signs, energy levels, activation energy and catalyst pathways.
Open lesson
Calorimetry and Molar Enthalpy
Turn measured temperature change into heat and then molar reaction enthalpy.
Open lesson
Hess’ Law from Given Simple Cycles
Use supplied simple cycles without importing H2 cycle-construction depth.
Open lesson
Bond Enthalpy Calculations
Estimate reaction enthalpy from bonds broken and formed, with honest limitations.
Open lesson
Lattice Energy and Ionic Trends
Explain lattice-energy magnitude from ionic charge and radius without Born–Haber construction.
Open lesson

Quick Reference

Idea	Exam-ready reminder
Exothermic	the reacting system releases heat to the surroundings; ΔH is negative
Endothermic	the reacting system absorbs heat from the surroundings; ΔH is positive
Energy profile	vertical difference between products and reactants is ΔH; peak relative to reactants is forward Ea
Catalyst	lowers activation energy through an alternative pathway but does not change ΔH
Calorimetry	qsolution = mcΔT; qreaction = −qsolution; divide by reacting amount and convert J to kJ
Formation-data cycle	ΔHreaction = ΣνΔHf(products) − ΣνΔHf(reactants)
Combustion-data cycle	when both sides reach the same products, use reactant combustion totals minus product totals
Bond enthalpies	ΔH ≈ Σ energy to break bonds − Σ energy released forming bonds
Lattice energy	larger ionic charge product and smaller ionic separation give a larger magnitude

Use supplied simple Hess cycles. Do not construct Born–Haber cycles, and do not import entropy or Gibbs free-energy treatment from H2.

Practice and Check Your Understanding

H1 Energetics Topic Practice
Course-specific 8873 questions across profiles, calorimetry, supplied cycles, bond enthalpy and lattice trends.
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H1 Energetics: Check Your Understanding
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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.

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.

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Answer 6 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.

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Chemical Energetics: Thermochemistry

H1 Energetics learning outcomes

What You’ll Learn

Lessons (Recommended Order)

Enthalpy Changes and Energy Profiles

Calorimetry and Molar Enthalpy

Hess’ Law from Given Simple Cycles

Bond Enthalpy Calculations

Lattice Energy and Ionic Trends

Quick Reference

Practice and Check Your Understanding

H1 Energetics Topic Practice

H1 Energetics: Check Your Understanding

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