Kinetics: every key term you need

Thermodynamics Overview

The study of energy transfer and transformation in chemical systems; includes heat, work, and entropy.

First Law of Thermodynamics

Energy is conserved; ΔE = q + w, where q is heat and w is work. The total energy of a system is constant.

Enthalpy (H)

A state function measuring the heat content of a system; ΔH = ΔE + ΔPV. Used to describe heat flow at constant pressure.

Exothermic Reactions

Reactions that release energy to surroundings; ΔH < 0. Examples: combustion, condensation, freezing.

Endothermic Reactions

Reactions that absorb energy from surroundings; ΔH > 0. Examples: melting, evaporation, photosynthesis.

Hess's Law

The enthalpy change of a reaction is the same regardless of the pathway; ΔH_rxn = Σ ΔH_products - Σ ΔH_reactants.

Standard Enthalpy of Formation (ΔH_f°)

The enthalpy change when one mole of a compound is formed from its elements in their standard states.

Calorimetry

Experimental technique measuring heat transfer; q = m × c × ΔT, where m is mass, c is specific heat, ΔT is temperature change.

Specific Heat Capacity

The amount of heat required to raise the temperature of 1 gram of substance by 1°C; units: J/(g·°C).

Heat of Combustion

The enthalpy change when one mole of substance burns completely in oxygen; typically large negative values.

Entropy (S)

A measure of disorder or randomness in a system; increases with higher temperature, more particles, and phase changes.

Second Law of Thermodynamics

In any spontaneous process, the entropy of the universe increases; ΔS_univ = ΔS_sys + ΔS_surr > 0.

Third Law of Thermodynamics

The entropy of a perfect crystal at absolute zero (0 K) is zero; entropy of all substances > 0 above 0 K.

Standard Entropy (S°)

The entropy of one mole of a substance at 25°C and 1 atm; used to calculate ΔS_rxn = Σ S°_products - Σ S°_reactants.

Gibbs Free Energy (G)

ΔG = ΔH - TΔS; determines spontaneity of a reaction. ΔG < 0 means spontaneous, ΔG > 0 means non-spontaneous.

ΔG and Spontaneity

ΔG < 0: spontaneous; ΔG = 0: equilibrium; ΔG > 0: non-spontaneous. Temperature affects the sign of ΔG.

ΔG and Equilibrium

ΔG° = -RT ln(K); relates standard free energy to the equilibrium constant; used to calculate K or predict reaction favorability.

Temperature Dependence of Spontaneity

For reactions with ΔH < 0, ΔS > 0: spontaneous at all temperatures. For ΔH > 0, ΔS < 0: non-spontaneous at all temperatures.

High Temperature, High Entropy Systems

When TΔS dominates (high T), entropy-driven processes favor disorder; useful for understanding dissolution and diffusion.

Low Temperature, High Enthalpy Systems

When T is low, ΔH dominates; enthalpy-driven processes favor strong bonding and low energy states.

Bond Dissociation Energy

The energy required to break a bond; used to calculate ΔH_rxn = Σ BDE_bonds broken - Σ BDE_bonds formed.

Lattice Energy

The energy required to separate one mole of an ionic solid into gaseous ions; high values indicate strong ionic bonding.

Solvation Energy

The enthalpy change when a solute dissolves; ΔH_solv = ΔH_hydration + ΔH_separation (for ionic compounds).

Born-Haber Cycle

A thermochemical cycle used to calculate lattice energies by summing enthalpy changes from ionization, electron affinity, and other steps.

Phase Diagrams

Graphs showing conditions (T, P) where phases (solid, liquid, gas) are stable; includes triple point and critical point.

Triple Point

The unique condition (T, P) where all three phases (solid, liquid, gas) coexist in equilibrium.

Critical Point

The highest temperature and pressure at which a liquid and gas can coexist; above this, the substance is a supercritical fluid.

Heat of Fusion

The enthalpy change when one mole of solid melts to liquid at its melting point; endothermic (ΔH > 0).

Heat of Vaporization

The enthalpy change when one mole of liquid vaporizes to gas at its boiling point; typically larger than ΔH_fus.

Heat of Sublimation

The enthalpy change when solid converts directly to gas; ΔH_sub = ΔH_fus + ΔH_vap.

Clausius-Clapeyron Equation

ln(P₂/P₁) = -ΔH_vap/R × (1/T₂ - 1/T₁); relates vapor pressure to temperature.

Vapor Pressure

The pressure of a gas in equilibrium with its liquid; increases with temperature and is compound-specific.

Sublimation Pressure

The pressure exerted by solid in equilibrium with its vapor; dry ice and naphthalene sublime at standard conditions.

Heating Curves

Graphs showing temperature vs. heat added; flat regions indicate phase changes (ΔT = 0 during melting/boiling).

Cooling Curves

Reverse of heating curves; show temperature decrease as heat is removed; used to identify phase transitions.

Supercritical Fluids

Substances above their critical point; have properties of both liquids (density) and gases (diffusivity); used as solvents.

Thermal Energy Distribution

At higher temperatures, molecules have greater average kinetic energy; more energetic collisions and faster reaction rates.

Activation Energy (Ea)

The minimum energy required for a reaction to proceed; lowers for catalyzed reactions without changing ΔG.

Arrhenius Equation

k = A × e^(-Ea/RT); relates rate constant to temperature and activation energy; exponential temperature dependence.

Temperature Coefficient (Q₁₀)

Approximate rule: reaction rates roughly double for every 10°C increase (applies when Ea is moderate).

Collision Theory

Reaction rate depends on collision frequency, orientation, and energy; only collisions with E ≥ Ea lead to reaction.

Reaction Coordinate Diagram

Energy diagram showing reactants, transition state (highest point), and products; helps visualize activation energy.

Catalysts and Energy

Catalysts lower Ea by providing alternative pathways; speeds up both forward and reverse reactions equally.

Enzyme Catalysis

Biological catalysts (enzymes) form enzyme-substrate complexes, lowering Ea; highly specific and efficient.

Surface Area Effect

Increasing surface area (grinding solids, stirring) increases collision frequency; speeds up reaction rates.

Concentration Effect on Rate

Higher concentrations increase collision frequency; rate often follows rate laws determined experimentally.

Pressure Effect on Gases

Higher pressure (for gases) increases concentration and collision frequency; liquids/solids minimally affected.

Reaction Order

Zero order: rate = k; first order: rate = k[A]; second order: rate = k[A]² or k[A][B]. Determined experimentally.

Half-Life (t₁/₂)

Time for reactant concentration to decrease to half; for first-order reactions, t₁/₂ = 0.693/k (independent of concentration).

Integrated Rate Laws

Mathematical expressions showing concentration vs. time; first-order: ln[A] = ln[A]₀ - kt; second-order: 1/[A] = 1/[A]₀ + kt.

Pre-Exponential Factor (A)

In Arrhenius equation, represents collision frequency and orientation factor; larger A means faster reaction.

Temperature Jump

Rapid temperature increase triggers instant equilibrium shifts; used to study fast reactions kinetically.

Coupled Reactions

Unfavorable reactions (ΔG > 0) can be driven by favorable ones; ΔG_total = ΔG₁ + ΔG₂ must be negative.

Cell Potential and Spontaneity

In electrochemistry, E°_cell > 0 means spontaneous (ΔG < 0); E°_cell < 0 means non-spontaneous (ΔG > 0).

Work and Thermodynamics

Electrical work (wₑ = -nFE) and mechanical work (w = -PΔV) are both forms of energy transfer from chemical reactions.

Gibbs Energy and Electrochemistry

ΔG° = -nFE°; combines thermodynamics and electrochemistry to predict spontaneity of redox reactions.

Period 5 Big Picture

Thermodynamics and kinetics together determine if and how fast reactions occur. Spontaneity (ΔG), energy (ΔH), disorder (ΔS), and activation barriers (Ea) are interconnected in predicting and controlling chemical change.

More AP Chemistry guides

• Atomic Structure & Properties
• Compound Structure & Properties
• Intermolecular Forces & Properties
• Chemical Reactions