Reference: Standard Enthalpies of Formation (kJ/mol, 298 K)
| Compound | ΔH°f | Compound | ΔH°f | Compound | ΔH°f |
|---|
| H₂O(l) | -285.8 | CO₂(g) | -393.5 | CH₄(g) | -74.8 |
| H₂O(g) | -241.8 | CO(g) | -110.5 | C₂H₆(g) | -84.7 |
| NH₃(g) | -45.9 | NO(g) | +90.3 | C₂H₄(g) | +52.3 |
| HCl(g) | -92.3 | NO₂(g) | +33.2 | C₂H₂(g) | +226.7 |
| NaCl(s) | -411.2 | SO₂(g) | -296.8 | C₆H₆(l) | +49.0 |
| CaO(s) | -635.1 | SO₃(g) | -395.7 | CH₃OH(l) | -238.7 |
| CaCO₃(s) | -1206.9 | H₂SO₄(l) | -814.0 | C₂H₅OH(l) | -277.7 |
Bond Energies (kJ/mol): H-H 436, C-H 413, C-C 348, C=C 614, C≡C 839, O-H 463, C-O 358, C=O 799, Cl-Cl 243, H-Cl 431, N≡N 941, O=O 498, N-H 391.
What is thermochemistry?
Thermochemistry is energy bookkeeping for chemical reactions: where the energy goes, in what quantities, and on what schedule. Its core currency is heat (q, joules), measured by letting a reaction proceed inside a calorimeter and watching a known mass of water change temperature. The master rule is the First Law — energy is conserved, never created or destroyed, only moved between chemical bonds and the surroundings — and everything on this page (specific heat, Hess's law, formation enthalpies, Gibbs free energy) is that one law wearing different hats.
The one rule to remember: enthalpy is a state function — it depends only on start and end states, never on the path between them. That single property is what makes Hess's law legal and the whole reference table below usable.
Heat versus temperature
Temperature measures the average jiggle of molecules; heat is the total energy transferred when that jiggle changes. The Specific Heat panel uses q = mcΔT to convert between them: the sample heats 100 g of water (c = 4.184 J/g·°C) from 20 °C to 80 °C, requiring q = 100 × 4.184 × 60 = 25104 J = 25.104 kJ, absorbed endothermically. Water's unusually large c is why it is the coolant of choice and the climate's flywheel: the same joules that would swing a metal pot's temperature wildly barely nudge a kettle. Water's outsized specific heat is hydrogen bonding at work: every joule must partly unwind the network before the molecules can speed up.
Breaking bonds costs, forming bonds pays
A reaction's enthalpy change is simply the accounting difference: ΔH ≈ ΣBE(broken) − ΣBE(formed). Breaking the H–H (436) and Cl–Cl (243) bonds costs 679 kJ/mol; forming two H–Cl bonds (431 each) repays 862 kJ/mol; the Bond Energy panel's sample nets −183 kJ/mol, exothermic — H2 + Cl2 → 2HCl releases heat because the products hold their atoms more tightly than the reactants did. The sign convention is universal: negative ΔH means the system releases heat (exothermic, like combustion), positive means it absorbs (endothermic, like ice melting).
Hess's law and formation enthalpies
Because enthalpy is a state function, you may build any reaction out of any convenient staircase of known steps and add the enthalpies — the sum is path-independent. The bookkeeping is standardized through standard enthalpies of formation, ΔH°f: the enthalpy to build one mole of compound from its elements in their standard states, which makes every element's own value zero by definition. The Formation panel's sample is the archetype: C(s) + O2(g) → CO2(g) is (0 + 0) on the left and −393.5 on the right, so ΔH°rxn = −393.50 kJ/mol — the reference table above is stocked for exactly this kind of sum, products minus reactants, each multiplied by its stoichiometric coefficient.
Entropy: counting arrangements
Entropy (S) counts how many microscopic arrangements a state has — gas beats liquid beats crystal, mixtures beat pure substances. The Second Law demands that the universe's entropy increase in any spontaneous process, which splits neatly: ΔSuniv = ΔSsys + ΔSsurr, where the surroundings' share is ΔSsurr = −ΔH/T — exothermic reactions disperse heat into the surroundings and raise their entropy. The Entropy panel's sample: an exothermic −100 kJ/mol at 298 K pours +335.570 J/mol·K into the surroundings; with ΔSsys = 150, the universe gains 485.570 J/mol·K and the process is spontaneous.
Gibbs free energy: the verdict
J. Willard Gibbs folded both books into one number: ΔG = ΔH − TΔS. The Gibbs panel's sample runs a favorable-but-messy reaction — ΔH = −100 kJ/mol, ΔS = −50 J/mol·K at 298 K — and returns ΔG = −100 − 298(−0.05) = −85.10 kJ/mol, spontaneous. The four sign combinations are worth memorizing: exothermic + entropy up is always spontaneous; endothermic + entropy down never; the two mixed cases depend on temperature, which is exactly why ice melts only above 273 K (TΔS overtakes ΔH). At equilibrium ΔG = 0, and the standard free energy connects to the equilibrium constant through ΔG° = −RT ln K.
Common misconceptions
- Exothermic means spontaneous. No: the entropy book must balance too. Dissolving many salts is endothermic yet spontaneous, because the disorder gain pays the bill.
- Heat and temperature are the same thing. A bathtub at 30 °C holds far more heat than a cup at 90 °C; q = mcΔT says mass and heat capacity get a vote.
- Forming bonds requires energy. The reverse: bond formation always releases energy — that released energy is precisely what exothermic reactions cash.
- Bond energies predict ΔH exactly. They are averages over many molecules; the ≈ sign in the panel's formula is honest. Formation enthalpies are the exact tool.
Related tools: Chemical Equilibrium (where ΔG° meets K), Gas Laws (the PΔV work heat exchanges with), and Kinetics (spontaneous is not the same as fast).
