Enthalpy Changes and Standard Conditions

Learning Objectives (OCR Spec 3.2.1)

By the end of this lesson you should be able to:

• Explain that chemical reactions are accompanied by enthalpy changes and classify them as exothermic or endothermic
• Define enthalpy change ΔH and state its units (kJ mol⁻¹)
• State the standard conditions used for thermochemical data
• Define and distinguish the standard enthalpy changes of combustion, formation, neutralisation and reaction
• Draw and interpret enthalpy profile diagrams showing activation energy and ΔH

Enthalpy and Enthalpy Change

The enthalpy H of a system is its total heat content at constant pressure. We cannot measure absolute enthalpy, but we can measure the change in enthalpy when a reaction occurs:

ΔH = H(products) − H(reactants)

• If the products have less enthalpy than the reactants, ΔH is negative and the reaction is exothermic (heat released to the surroundings).
• If the products have more enthalpy than the reactants, ΔH is positive and the reaction is endothermic (heat absorbed from the surroundings).

The units of ΔH are kJ mol⁻¹ (kilojoules per mole) because the value depends on the amount of substance reacting.

Exothermic vs Endothermic

Feature	Exothermic	Endothermic
Direction of heat	Released to surroundings	Absorbed from surroundings
Sign of ΔH	Negative	Positive
Temperature of surroundings	Rises	Falls
Examples	Combustion, neutralisation, most oxidations	Thermal decomposition, photosynthesis, dissolving NH4NO3

Enthalpy Profile Diagrams

An enthalpy profile diagram shows how the enthalpy of the chemical system changes as reactants are converted to products. The activation energy Ea is the minimum energy that colliding particles must have in order to react; it corresponds to the height of the barrier.

graph LR
  A[Reactants] -->|Ea| B[Transition state]
  B --> C[Products]

For an exothermic reaction the products sit lower than the reactants; for an endothermic reaction they sit higher. You must label reactants, products, activation energy Ea and ΔH on any sketch. ΔH is the vertical distance between reactants and products, measured with the correct sign.

Standard Conditions

Enthalpy changes depend on temperature, pressure and the physical state of the reactants and products. To allow data to be compared, chemists quote standard enthalpy changes measured under standard conditions:

• Pressure: 100 kPa (1 bar)
• Temperature: usually stated as 298 K (25 °C)
• Concentration of solutions: 1 mol dm⁻³
• All substances in their standard states at those conditions

The standard symbol is a plimsoll line, written ΔH° or ΔH⊖ (OCR uses the plimsoll). A standard state is the physical state of an element or compound under standard conditions: e.g. H2O(l), CO2(g), Na(s), Cl2(g).

Note: Older data books use 101 kPa; OCR A uses 100 kPa. Quote 100 kPa in answers.

The Four Standard Enthalpy Changes

OCR requires precise definitions for four standard enthalpy changes.

1. Standard Enthalpy Change of Reaction ΔrH°

The enthalpy change when the molar quantities of reactants as shown in the chemical equation react together under standard conditions, all reactants and products in their standard states.

Example: N2(g) + 3H2(g) → 2NH3(g), ΔrH° = −92 kJ mol⁻¹ (per mole of equation as written, i.e. for 2 mol NH3).

2. Standard Enthalpy Change of Formation ΔfH°

The enthalpy change when one mole of a compound is formed from its elements in their standard states under standard conditions.

Example: C(s, graphite) + 2H2(g) → CH4(g), ΔfH° = −74.8 kJ mol⁻¹.

Crucially, ΔfH° of an element in its standard state is zero by definition. So ΔfH°(O2(g)) = 0, ΔfH°(C(s, graphite)) = 0, but ΔfH°(O3(g)) ≠ 0 because ozone is not the standard state of oxygen.

3. Standard Enthalpy Change of Combustion ΔcH°

The enthalpy change when one mole of a substance is completely burned in oxygen under standard conditions.

Example: CH4(g) + 2O2(g) → CO2(g) + 2H2O(l), ΔcH° = −890 kJ mol⁻¹.

Combustion is always exothermic, so ΔcH° is always negative. Note that the water produced must be liquid at 298 K.

4. Standard Enthalpy Change of Neutralisation ΔneutH°

The enthalpy change when an acid and a base react under standard conditions to form one mole of water.

Example: HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l), ΔneutH° = −57.1 kJ mol⁻¹.

For strong acid + strong alkali the value is always close to −57 kJ mol⁻¹ because the same ionic reaction is taking place:

H⁺(aq) + OH⁻(aq) → H2O(l)

Worked Example – Writing Equations with Correct Stoichiometry

Q: Write the thermochemical equation for the standard enthalpy change of combustion of ethanol. Why must the equation be written with a coefficient of 1 in front of ethanol?

A: Because ΔcH° is defined per one mole of fuel.

C2H5OH(l) + 3O2(g) → 2CO2(g) + 3H2O(l)

Balancing requires 3 mol of O2 and produces 2 mol CO2 and 3 mol H2O(l). The enthalpy change quoted (−1367 kJ mol⁻¹) is per mole of ethanol, not per mole of the equation.

Worked Example – Identifying the Standard State

Q: For each reaction below, state which standard enthalpy change is being measured.

(a) C(s) + O2(g) → CO2(g), ΔH = −394 kJ mol⁻¹ (b) ½H2(g) + ½Cl2(g) → HCl(g), ΔH = −92 kJ mol⁻¹ (c) H⁺(aq) + OH⁻(aq) → H2O(l), ΔH = −57 kJ mol⁻¹

A: (a) This is both ΔcH° of carbon and ΔfH° of CO2 - a neat coincidence because carbon is an element and CO2 is its only combustion product. (b) ΔfH° of HCl(g) - one mole formed from elements in standard states (hence the half-coefficients). (c) ΔneutH° - one mole of water formed from acid + base.

Exam Tips

• Memorise the definitions word-for-word. OCR mark schemes often require "one mole", "under standard conditions", "in their standard states" and the correct substance (reaction / combustion / formation / neutralisation).
• Quote standard conditions as 100 kPa and 298 K, not 101.3 kPa or 25 °C in isolation.
• Always give ΔH a sign and units of kJ mol⁻¹.
• When writing ΔfH° equations you may need fractional coefficients to keep one mole of product.
• State symbols are essential in thermochemical equations: swap H2O(l) for H2O(g) and ΔcH° changes by ~44 kJ per mole of water.

Common Exam Mistakes

• Forgetting that ΔfH° of an element is zero.
• Writing ΔcH° with more than one mole of fuel (e.g. 2C2H6 + 7O2 → 4CO2 + 6H2O is not the ΔcH° equation).
• Omitting the sign of ΔH - "890 kJ mol⁻¹" scores zero; it must be −890.
• Writing water as H2O(g) in combustion - at 298 K it condenses to H2O(l).
• Claiming ΔneutH° is −57 kJ mol⁻¹ for weak acids - it is only exactly that for strong acid + strong base, because weak acids absorb energy on ionisation.
• Describing exothermic reactions as "losing heat" - they release heat to the surroundings.

Summary

An enthalpy change ΔH is the heat released (exothermic, ΔH < 0) or absorbed (endothermic, ΔH > 0) at constant pressure, with units of kJ mol⁻¹. Standard enthalpy changes are quoted at 100 kPa and 298 K with reactants and products in their standard states. The four key standard enthalpy changes are of reaction (ΔrH°), formation (ΔfH°, zero for elements), combustion (ΔcH°, always negative) and neutralisation (ΔneutH°, ~−57 kJ mol⁻¹ for strong acid/alkali). Knowing these definitions precisely is the foundation for every calculation in this topic.