Chemistry & Physics Review – Thermodynamics, Equilibrium, and General Chemistry

A comprehensive set of vocabulary flashcards covering thermodynamics, equilibrium, oxidation states, metric prefixes, electromagnetic spectrum, acid–base theory, nuclear decay, and periodic trends.

Enthalpy (ΔH)

Heat content of a system; negative ΔH = exothermic, positive ΔH = endothermic.

Entropy (ΔS)

Measure of disorder or randomness; greater disorder (higher ΔS) is generally favorable.

Gibbs Free Energy (ΔG)

Thermodynamic potential given by ΔG = ΔH – TΔS; predicts spontaneity.

Spontaneous Reaction

Process with ΔG < 0 under given conditions.

Non-spontaneous Reaction

Process with ΔG > 0 under given conditions.

Equilibrium Condition

ΔG = 0; forward and reverse reaction rates are equal.

ΔG° and Equilibrium Constant

ΔG° = –RT ln K; negative value implies K > 1 and reaction favors products.

Exothermic Reaction

Releases heat (ΔH < 0); products lower in energy than reactants.

Endothermic Reaction

Absorbs heat (ΔH > 0); products higher in energy than reactants.

Law of Mass Action

Relates concentrations/pressures of species to the equilibrium constant, K.

Equilibrium Constant (K) Expression

Includes aqueous and gaseous species; omits pure solids and liquids.

Le Châtelier’s Principle

A system at equilibrium shifts to minimize a disturbance in concentration, pressure, or temperature.

Adding Reactant

Shifts equilibrium to the right (toward products).

Decreasing Volume (Gas)

Raises pressure; equilibrium shifts to side with fewer gas moles.

Temperature Increase (Exothermic)

Heat acts as product; equilibrium shifts left (toward reactants).

Temperature Increase (Endothermic)

Heat acts as reactant; equilibrium shifts right (toward products).

Haber Process Enthalpy

Shift left with added heat indicates the synthesis of NH₃ is exothermic (ΔH < 0).

Group 1 Metal Oxidation State

+1 in compounds (e.g., Na⁺).

Group 2 Metal Oxidation State

+2 in compounds (e.g., Mg²⁺).

Hydrogen Oxidation States

+1 with non-metals, –1 with metals.

Oxygen Oxidation States

Usually –2; –1 in peroxides.

Halogen Oxidation States

Generally –1 unless bonded to more electronegative atoms.

Polyatomic Ion Charge Rule

Sum of oxidation numbers equals the ion’s net charge.

STP (Standard Temperature & Pressure)

0 °C (273 K) and 1 atm (760 mm Hg).

Molar Volume at STP

One mole of an ideal gas occupies 22.4 L.

Nano- Prefix (n)

10⁻⁹.

Micro- Prefix (μ)

10⁻⁶.

Milli- Prefix (m)

10⁻³.

Centi- Prefix (c)

10⁻².

Kilo- Prefix (k)

10³.

Mega- Prefix (M)

10⁶.

Giga- Prefix (G)

10⁹.

Electromagnetic Spectrum Order

Radio < Microwaves < Infrared < Visible < Ultraviolet < X-ray < Gamma (increasing energy).

Wavelength–Energy Relationship

Shorter wavelength → higher frequency and higher energy.

Visible Light Range

~400 nm (violet) to 700 nm (red).

Lewis Acid

Species that accepts an electron pair.

Lewis Base

Species that donates an electron pair.

Brønsted–Lowry Acid

Proton (H⁺) donor.

Brønsted–Lowry Base

Proton (H⁺) acceptor.

Alpha (α) Decay

Emission of ⁴₂He nucleus; mass –4, atomic number –2.

Beta Minus (β⁻) Decay

Emission of electron; atomic number +1, mass unchanged.

Beta Plus (β⁺) / Positron Emission

Emission of positron; atomic number –1, mass unchanged.

Gamma (γ) Decay

Emission of high-energy photon; no change in mass or atomic number.

Electronegativity Trend

Increases up and to the right; fluorine is highest.

Ionization Energy

Energy needed to remove an electron; increases up and to the right.

Electron Affinity

Energy released when atom gains an electron; generally becomes more negative up and to the right (largest for halogens).

Atomic Radius Trend

Increases down a group and decreases across a period (left to right).