Enthalpy

Enthalpy H = E + PV is a state function defined so that at constant pressure, ΔH = q_p — heat measured directly gives enthalpy change.
ΔH < 0 → exothermic (releases heat, reactants higher-energy). ΔH > 0 → endothermic (absorbs heat, reactants lower-energy).
Phase changes: melting/evaporation/sublimation = endothermic; freezing/condensation/deposition = exothermic.
Heating curves: sloped segments use q = mcΔT (single phase); flat segments use q = mL (phase change in progress, T constant).
Hess's Law: ΔH is path-independent — sum ΔH of any steps that add up to the overall reaction, or use ΔH = Σ(bonds broken) − Σ(bonds formed).
Standard enthalpy of formation ΔH°f of an element in its standard state = 0. ΔH°rxn = ΣnΔH°f(products) − ΣnΔH°f(reactants).

Enthalpy (H): State function H = E + PV; at constant pressure, ΔH = q_p.
Exothermic: Process with ΔH < 0 that releases heat; reactants start at higher energy than products.
Endothermic: Process with ΔH > 0 that absorbs heat; reactants start at lower energy than products.
Transition state: The highest-energy point on a reaction energy diagram, between reactants and products.
Activation energy: Height of the transition state above the reactants; lower activation energy means a faster reaction.
Latent heat (L): Heat needed to change the phase of 1 g of substance without changing its temperature (fusion or vaporization).
Triple point: The single temperature/pressure combination where solid, liquid, and gas phases all coexist.
Critical point: Highest T and P where a distinct liquid phase exists; beyond it, liquid/gas merge into a supercritical fluid.
Hess's Law: Total ΔH for a reaction equals the sum of ΔH for any set of steps that add up to it, since enthalpy is a state function.
Bond dissociation energy: Average energy required to break a specific bond type in the gas phase.
Standard enthalpy of formation (ΔH°f): Enthalpy change when 1 mole of a substance forms from its elements in their standard states.

H = E + PV → ΔH = q_p

ΔH = Σ(bonds broken) − Σ(bonds formed)

Breaking bonds is always endothermic (costs energy); forming bonds is always exothermic (releases energy).
Worked example: C(graphite) + O₂ → CO₂ directly (ΔH = −393.5 kJ/mol) matches the sum of two steps via CO intermediate (−110.5 + −283.0 = −393.5 kJ/mol).

ΔH°rxn = ΣnΔH°f(products) − ΣnΔH°f(reactants)

Worked example: CH₄ combustion → ΔH°rxn = −965.1 − (−74.6) = −890.5 kJ/mol (strongly exothermic).

Endothermic phase changes: energy to break the more-ordered phase's attractions EXCEEDS energy released forming the less-ordered phase's weaker attractions.
Heating curve flat segments = phase change happening (use q = mL); sloped segments = temperature changing within one phase (use q = mcΔT).
Phase diagram axes: temperature (x) vs. pressure (y); boundary lines = two phases coexisting in equilibrium.

Don't confuse the reaction energy diagram's peak (transition state, activation energy) with ΔH — activation energy affects RATE, ΔH determines whether the reaction is exo/endothermic overall.
ΔH°f of an element in its standard state is always zero — don't assign it a nonzero value even for a common element like O₂(g) or N₂(g).
Hess's Law works for full reaction steps AND for bond dissociation energies — but the bond-energy sign convention flips (broken − formed), easy to invert by mistake.

Is sublimation endothermic or exothermic?

On a heating curve, what's happening during a flat segment?

What is the ΔH°f of O₂(g) at standard conditions?

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