Thermodynamics
Thermochemistry
Energy of chemical reactions — exothermic, endothermic, enthalpy, calorimetry
Thermochemistry studies energy changes accompanying chemical reactions. Heat absorbed (endothermic) or released (exothermic). Quantified by enthalpy change ΔH (heat at constant pressure). Sign convention: ΔH < 0 for exothermic, ΔH > 0 for endothermic. Hess's law: ΔH for overall reaction = sum of ΔH for steps. Calorimetry measures heat. Bond energies estimate ΔH (sum of bonds broken minus bonds formed). Critical for: combustion, batteries, food chemistry, refrigeration. Foundation of thermodynamics.
- Enthalpy changeΔH = q (heat) at constant P
- ExothermicΔH < 0; releases heat
- EndothermicΔH > 0; absorbs heat
- Hess's lawΔH_total = sum of ΔH for steps
- Standard ΔH°fΔH for forming compound from elements
- Bond energiesEstimate ΔH = sum bonds broken - bonds formed
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Why thermochemistry matters
- Energy production. Combustion, batteries.
- Industrial. Designing exothermic/endothermic processes.
- Refrigeration. Heat absorption.
- Biology. Metabolism = combustion of food.
- Materials. Reaction calorimetry.
- Climate. Phase changes affect temperature.
- Safety. Predict heat from runaway reactions.
Common misconceptions
- Endothermic = cold to touch. Reaction absorbs heat from surroundings.
- Exothermic always spontaneous. Need ΔG < 0 (entropy too).
- Enthalpy = energy. Energy at constant P; specific definition.
- ΔH depends on path. State function; only initial/final.
- Larger ΔH means more reactive. Different from rate.
- Bond energies exact. Average values; have error.
Frequently asked questions
What is enthalpy?
Heat content at constant pressure. ΔH = ΔU + PΔV. For most reactions at constant P: ΔH ≈ heat absorbed. Reaction exothermic: ΔH < 0 (e.g., combustion: -890 kJ/mol for CH₄). Endothermic: ΔH > 0 (e.g., melting ice: +6 kJ/mol). Standard state: 1 atm, 25°C, 1 M. ΔH° denotes standard enthalpy.
What's Hess's law?
Enthalpy change for reaction is independent of path; depends only on initial and final states. So: complex reaction can be calculated as sum of simpler reactions whose ΔH known. Critical for: calculating ΔH that can't be measured directly (e.g., C → CO from C → CO₂ and CO → CO₂). Hess's law derives from enthalpy being state function.
What's standard enthalpy of formation?
ΔH°f: enthalpy change when 1 mole of compound forms from its elements in standard states. By convention, ΔH°f of element in standard state = 0. Then: ΔH°rxn = sum(ΔH°f products) - sum(ΔH°f reactants). Compiled in tables; used for calculating reaction enthalpies. Examples: ΔH°f(H₂O liquid) = -286 kJ/mol; CO₂ = -394 kJ/mol.
How is enthalpy measured?
Calorimetry. Reaction in well-insulated container (calorimeter); measure temperature change of surroundings (water). q = mcΔT where m=mass water, c=specific heat (4.18 J/g·K for water), ΔT=temperature change. q reaction = -q surroundings (energy conservation). At constant pressure: q = ΔH. Bomb calorimeter (constant V) gives ΔU; ΔH ≈ ΔU for reactions where Δ(PV) small.
How do bond energies relate?
Bond breaking absorbs energy; forming releases energy. ΔH ≈ sum(bond energies broken) - sum(bond energies formed). Approximate (~5% error) because bond energies depend on environment. Useful for: estimating reaction enthalpies when ΔH°f not available. Formal calculation: average values from tables.
What's specific heat?
Energy needed to raise 1 g of substance by 1°C (or K). Water: 4.18 J/g·K (very high — explains water's heat-storing ability). Metals: lower (Al: 0.90, Fe: 0.45). Affects: cooking time, climate moderation, calorimetry. Higher specific heat = absorbs more heat per degree change.
Why does combustion release energy?
Bonds in CO₂ and H₂O are stronger than bonds in fuel + O₂. Net: bond formation releases more than bond breaking absorbs. Specifically: O=O is weaker than C=O and H-O bonds formed in combustion. Heat released = energy difference. Combustion of CH₄: -890 kJ/mol — basis of natural gas as fuel.