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In every chemical transformation, both the starting materials and the resulting compounds possess a characteristic value known as the standard enthalpy of formation, denoted ΔHf. These values are essential for quantifying the energy exchange that occurs during a reaction.

Step 1: Assemble the Enthalpy Equation

The relationship between the overall reaction enthalpy (ΔH) and the individual ΔHf values is:

ΔH = ΣΔHf(products) – ΣΔHf(reactants)

For illustration, consider the combustion of acetylene:

C₂H₂(g) + 5/2 O₂(g) → 2 CO₂(g) + H₂O(g) ΔH = –1 256 kJ mol⁻¹

Given that ΔHf[CO₂] = –394 kJ mol⁻¹, ΔHf[H₂O] = –242 kJ mol⁻¹, and elemental O₂ has ΔHf = 0, we set up:

–1 256 = [2(–394) + (–242)] – ΔHf[C₂H₂]

Rearranging gives:

ΔHf[C₂H₂] = 2(–394) + (–242) + 1 256

Step 2: Solve for ΔHf

Carrying out the arithmetic:

ΔHf[C₂H₂] = (–788) + (–242) + 1 256 = 226 kJ mol⁻¹

Step 3: Verify the Sign Convention

ΔHf values for products are negative, while those for reactants are positive. Since acetylene is a reactant, the positive 226 kJ mol⁻¹ is correct.

TL;DR

ΔHf and ΔH are expressed in kilojoules per mole. By applying ΔH = ΣΔHf(products) – ΣΔHf(reactants) and respecting sign conventions, you can calculate any missing ΔHf value.