Enthalpy and Enthalpy Changes Explained Simply

Enthalpy (H) is the total heat content of a system at constant pressure, defined as H = U + PV (internal energy plus pressure×volume). We can never measure absolute H, but we can measure its change: ΔH is simply the heat absorbed or released by a process happening at constant pressure.

Enthalpy matters because it tells us whether a reaction gives out heat (exothermic, ΔH negative — like burning fuel) or needs heat (endothermic, ΔH positive — like cooking an egg). Engineers use ΔH to design heaters, engines, refrigerators and chemical plants.

Key Concepts

• Enthalpy H: SI unit joule (J); a state function — ΔH depends only on initial and final states.
• Exothermic: ΔH < 0, heat released to surroundings (combustion, neutralisation). Products have lower enthalpy than reactants.
• Endothermic: ΔH > 0, heat absorbed from surroundings (melting ice, photosynthesis). Products have higher enthalpy than reactants.
• Standard enthalpy change (ΔH°): measured with all substances in standard states (1 bar, usually 298 K) — the ° just means "under standard conditions".
• ΔH = qp: at constant pressure with only expansion work, the enthalpy change equals the heat exchanged. This is why chemists love enthalpy.
• Enthalpy vs internal energy: ΔH = ΔU + ΔngRT for reactions (Δng = change in moles of gas). For solids/liquids only, ΔH ≈ ΔU.

Formulas

• Definition: H = U + PV. Units: J (U in J, P in Pa, V in m³).
• Constant-pressure heat: ΔH = qp (only PV work allowed).
• Relation to ΔU: ΔH = ΔU + ΔngRT, with R = 8.314 J/mol·K. Assumes ideal gases at constant T.
• Calorimetry: q = mcΔT gives the heat; at constant pressure this equals ΔH for the process.
• Hess's law: ΔH(reaction) = ΣΔHf°(products) − ΣΔHf°(reactants). Enthalpy is a state function, so any path between reactants and products gives the same ΔH.

Worked Example

Burn 8.0 g of methane: CH₄ + 2O₂ → CO₂ + 2H₂O, ΔH° = −890 kJ/mol. How much heat is released?

1. Moles CH₄ = 8.0 g / 16.0 g/mol = 0.50 mol.
2. Heat released = 0.50 mol × 890 kJ/mol = 445 kJ. (ΔH = −445 kJ; the negative sign means exothermic.)
3. Sanity check with formation enthalpies: ΔH = [−393.5 + 2(−285.8)] − [−74.8] = −890.3 kJ/mol. Matches.

Common Misconceptions

• "Exothermic reactions always feel hot and fast." Exothermic only means heat is released — rusting iron is exothermic but slow and barely warm. Rate and heat are separate ideas.
• "ΔH positive means the reaction can't happen." Wrong — endothermic reactions happen all the time (ice melts, salt dissolves). Whether a reaction is spontaneous depends on Gibbs energy (ΔG = ΔH − TΔS), not ΔH alone.
• "Enthalpy and heat are the same thing." Heat is energy in transit (a path function); enthalpy is a property of the state. Only at constant pressure does ΔH equal the heat exchanged.

Real-World Applications

• Fuels: petrol, natural gas and coal are valued by their enthalpy of combustion (kJ per kg) — energy content is ΔH.
• Cold packs: dissolving ammonium nitrate in water is strongly endothermic (ΔH > 0), so instant cold packs chill injuries without refrigeration.
• Industrial reactors: highly exothermic reactions (ΔH very negative) need cooling jackets designed from ΔH data to prevent runaways.

In short, enthalpy H = U + PV is the heat content at constant pressure, and ΔH = qp tells us the heat a process releases or absorbs.