Thermodynamics
Enthalpy of Reaction
Heat absorbed or released — quantifies energy change of chemical reactions
Enthalpy of reaction (ΔH_rxn) is the heat absorbed (positive) or released (negative) during a chemical reaction at constant pressure. ΔH_rxn = Σ ΔH°f(products) - Σ ΔH°f(reactants). Calculated using Hess's law, standard formation enthalpies, or bond energies. Examples: combustion of methane = -890 kJ/mol (exothermic). Photosynthesis: +2870 kJ/mol (endothermic; energy from sunlight). Determines whether reaction releases or absorbs heat — critical for industrial design, safety, biology.
- ΔH_rxn formulaΣ ΔH°f(prod) - Σ ΔH°f(react)
- Combustion CH₄-890 kJ/mol (exothermic)
- Combustion C-394 kJ/mol (CO₂ formation)
- H₂O formation-286 kJ/mol (liquid)
- Standard conditions1 atm, 25°C
- Photosynthesis+2870 kJ/mol (very endothermic)
Interactive visualization
Press play, or step through manually. The visualization is yours to drive — try it before reading on.
Watch the 60-second explainer
A condensed visual walkthrough — narrated, captioned, under a minute.
Why ΔH matters
- Energy production. Fuel selection.
- Industrial. Reactor design, cooling.
- Biology. Metabolism energetics.
- Refrigeration. Phase change endothermic.
- Safety. Runaway reactions.
- Pharmacology. Drug stability.
- Materials. Heat treatment.
Common misconceptions
- Enthalpy = energy. Energy at constant P; specific.
- Bond energies always exact. Approximations.
- Larger |ΔH| means faster. Different concept (rate).
- ΔH = ΔG. Different — entropy term affects spontaneity.
- ΔH is constant with T. Slight variation.
- State changes don't matter. Important for accurate calculations.
Frequently asked questions
How is reaction enthalpy calculated?
Three methods. (1) Hess's law: combine known reaction ΔH values. (2) Standard formation enthalpies: ΔH_rxn = Σ ΔH°f(products) - Σ ΔH°f(reactants). (3) Bond energies: ΔH ≈ sum(bonds broken) - sum(bonds formed). Method choice depends on data available. Standard formation tables most common.
How are formation enthalpies measured?
Calorimetry. Burn or react sample in calorimeter; measure temperature change. Heat released calculates Q. For element-to-compound formation: this is ΔH°f. Example: burning C in O₂ to give CO₂: -394 kJ/mol. Tables of ΔH°f compiled for thousands of compounds.
Why is combustion exothermic?
Bonds in products (CO₂, H₂O) are stronger than reactants (fuel + O₂). Specifically: O=O double bond (498 kJ/mol) is weaker than C=O bonds in CO₂ (799 kJ/mol each, two per CO₂). Net energy release. Nature of fuels.
What's combustion of methane in detail?
CH₄ + 2O₂ → CO₂ + 2H₂O. ΔH = -890 kJ/mol. Bonds broken: 4×C-H (412×4=1648) + 2×O=O (498×2=996) = 2644 kJ. Bonds formed: 2×C=O (799×2=1598) + 4×H-O (463×4=1852) = 3450 kJ. Net: -806 kJ (close to actual -890, difference from approximation).
How do you use bond energies?
Approximate ΔH from bonds. (1) List all bonds broken in reactants. (2) List all bonds formed in products. (3) ΔH ≈ Σ(broken) - Σ(formed). Each bond energy from table (average). Useful when ΔH°f not available. Error ~5% from environment effects on bond strength.
What about state changes?
ΔH for state change important for combustion, melting, etc. ΔH_vap (water): +44 kJ/mol (boiling). ΔH_fus (water): +6 kJ/mol (melting). ΔH_sub (CO₂): +25 kJ/mol (sublimation). Same magnitude regardless of direction; sign reverses (vapor → liquid: -44).
Are reaction enthalpies always constant?
No — slight T dependence. ΔH at different T calculated via Kirchhoff's law: ΔH(T) = ΔH(298 K) + ∫ΔCp dT. Where ΔCp = sum(Cp products) - sum(Cp reactants). Usually small effect; ΔH constant ±5% over 100 K range. Tables give values at 298 K standard state.