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Chemistry Unit 8: Thermochemistry and electrochemistry

Reactions absorb or release energy. Thermodynamics predicts whether a reaction happens on its own, and electrochemistry turns redox reactions into electricity.

Covers: Endothermic vs exothermic, enthalpy, specific heat q = mcΔT, Hess's law, entropy and Gibbs free energy, galvanic cells

Study guide

Thermochemistry studies heat in reactions. Exothermic reactions release heat to the surroundings (ΔH negative, surroundings get warmer), and endothermic reactions absorb heat (ΔH positive, surroundings get cooler). Enthalpy change ΔH is the heat exchanged at constant pressure. It can come from bond energies (energy to break bonds minus energy released forming bonds) or from standard enthalpies of formation: ΔH = Σ ΔHf(products) − Σ ΔHf(reactants).

Specific heat is the energy needed to raise 1 gram of a substance by 1°C. The heat equation is q = mcΔT, where water's specific heat is 4.18 J/(g·°C). In calorimetry, the heat lost by one substance equals the heat gained by another, so you can measure ΔH experimentally.

Hess's law says the total enthalpy change is the same no matter which steps a reaction takes. You can add known reactions to get a target reaction: reverse a reaction and its ΔH changes sign, multiply a reaction and its ΔH multiplies too.

Entropy (S) measures disorder or the spread of energy; it increases from solid to liquid to gas and when more gas molecules form. Gibbs free energy combines enthalpy and entropy: ΔG = ΔH − TΔS. A reaction is spontaneous (thermodynamically favorable) when ΔG < 0. Temperature can flip spontaneity when ΔH and ΔS have the same sign.

Electrochemistry connects redox reactions to electricity. In a galvanic (voltaic) cell, a spontaneous redox reaction produces current. Oxidation happens at the anode and reduction at the cathode, electrons flow through the wire from anode to cathode, and a salt bridge balances charge. Cell voltage comes from the difference in reduction potentials, and electrolytic cells use outside electricity to drive nonspontaneous reactions.

The reaction's heat is opposite in sign to the heat gained by the solution: q_rxn = −q_solution. Divide by the moles that reacted to get ΔH in kJ/mol. Because ΔH per mole is intensive, using less reactant releases less total heat but gives the same ΔH per mole. Charge is q = I·t, in coulombs. Dividing by Faraday's constant (96,485 C per mole of electrons) gives moles of electrons, and the half-reaction converts those to moles of metal. For Cr³⁺ + 3e⁻ → Cr, for example, every 3 mol of electrons deposits 1 mol of chromium.

ΔG° = −nFE°, so a positive cell potential means a negative ΔG° and a spontaneous reaction. ΔG° also equals −RT ln K, so a reaction with ΔG° 1. In a galvanic cell, the anode metal is oxidized into ions and loses mass, and metal ions are reduced onto the cathode, which gains mass. To get the largest cell voltage from a table, pair the half-reaction with the most positive reduction potential as the cathode with the one with the most negative potential as the anode.

Key ideas

• Exothermic: releases heat, ΔH 0.
• q = mcΔT measures heat gained or lost.
• Hess's law: ΔH for a reaction is the sum of ΔH for its steps.
• ΔG = ΔH − TΔS; ΔG < 0 means the reaction is spontaneous.
• In a galvanic cell, oxidation happens at the anode and reduction at the cathode.

Formulas and rules

Common mistakes

• Mixing up the signs: exothermic is negative ΔH.
• Thinking spontaneous means fast. It only means thermodynamically favored.

Words to know

Exothermic reaction

Releases heat; ΔH is negative.

Specific heat equation

q = mcΔT: heat = mass × specific heat × temperature change.

Spontaneous reaction

ΔG < 0.

Anode

The electrode where oxidation happens.

Specific heat of water

4.18 J/(g·°C)

Endothermic reaction

Absorbs heat; ΔH is positive.

Enthalpy (ΔH)

Heat exchanged at constant pressure.

Calorimetry

Measuring heat: heat lost by one substance = heat gained by another.

Hess's law

ΔH is the same whatever path the reaction takes.

Enthalpy of formation

ΔH to form 1 mol of a compound from its elements.

Entropy (S)

Disorder, or the spread of energy.

Gibbs free energy

ΔG = ΔH − TΔS

Cathode

The electrode where reduction happens.

Salt bridge

Balances charge between the halves of a galvanic cell.

Electrolytic cell

Uses outside electricity to drive a nonspontaneous reaction.

Sample practice questions

15 more questions like these, from level 1 to 5, are in the practice set.