General Chemistry: Reactions, Energy, Equilibrium, Gases, and Phase Changes

Vocabulary practice flashcards covering chemical equations, reaction types, redox processes, stoichiometry, energy changes, reaction rates, chemical equilibrium, gas laws, intermolecular forces, and phase transitions.

Chemical Equation

A representation using chemical formulas and other symbols showing what reactants are the starting materials in a reaction and what products are formed.

Reactants

The starting materials in a chemical reaction, which are written on the left side of a chemical equation.

Products

The substances formed during a chemical reaction, which are written on the right side of a chemical equation.

Coefficients

The numbers placed before formulas in a chemical equation that show the number of molecules of a given element or compound that react or are formed.

Combination Reaction

A type of chemical reaction involving the joining of two or more reactants to form a single product.

Single Replacement Reaction

A reaction in which one element replaces another element in a compound to form a different compound and element as products.

Double Replacement Reaction

A reaction in which two compounds exchange parts—atoms or ions—to form two new compounds.

Oxidation

The loss of electrons from an atom; in organic and redox reactions, it also manifests as the gain of oxygen atoms or the loss of hydrogen atoms.

Reduction

The gain of electrons by an atom; in organic and redox reactions, it also manifests as the loss of oxygen atoms or the gain of hydrogen atoms.

Redox Reaction

An oxidation-reduction reaction involving the transfer of electrons from one element to another, composed of one component that is oxidized and one that is reduced.

Oxidizing Agent

A compound that causes another compound to be oxidized by gaining electrons, while itself being reduced.

Avogadro's Number

The number of items, atoms, or molecules contained in one mole of a substance, equal to $6.02 \times 10^{23}$.

Percent Yield

A measure of reaction efficiency defined by the formula: $\text{Percent yield} = \frac{\text{actual yield}}{\text{theoretical yield}} \times 100\%$.

Potential Energy

Stored energy, such as the chemical energy stored in the chemical bonds of a compound.

Kinetic Energy

The energy possessed by an object or particles due to motion.

Law of Conservation of Energy

A law stating that the total energy in a system does not change; energy cannot be created or destroyed.

Bond Dissociation Energy

The enthalpy change ($\Delta H$) required for breaking a covalent bond by equally dividing the electrons between the two atoms; it is always a positive value because bond breaking is endothermic.

Endothermic Reaction

A reaction in which heat is absorbed and $\Delta H$ is positive ($\Delta H > 0$); bonds broken in reactants are stronger than bonds formed in products, making products higher in energy than reactants.

Exothermic Reaction

A reaction in which heat is released and $\Delta H$ is negative ($\Delta H < 0$); bonds formed in products are stronger than bonds broken in reactants, making products lower in energy than reactants.

Energy of Activation ($E_a$)

The difference in energy between the reactants and the transition state on a reaction energy diagram.

Catalyst

A substance that speeds up a chemical reaction by lowering the energy of activation ($E_a$) without altering the $\Delta H$ of the reaction.

Chemical Equilibrium

The state of a reversible reaction where the rate of the forward reaction equals the rate of the reverse reaction, resulting in constant net concentrations of reactants and products.

Equilibrium Constant ($K$)

A numerical ratio expressing the relationship between the concentration of products and reactants at equilibrium: $K = \frac{[\text{products}]}{[\text{reactants}]}$.

Le Châtelier's Principle

The principle stating that if a chemical system at equilibrium is disturbed or stressed, the system will react in a direction that counteracts the disturbance or relieves the stress.

Boyle's Law

A gas law stating that pressure and volume are inversely related for constant temperature ($T$) and moles ($n$): $P_1V_1 = P_2V_2$.

Charles's Law

A gas law stating that volume and temperature are directly proportional for constant pressure ($P$) and moles ($n$): $\frac{V_1}{T_1} = \frac{V_2}{T_2}$.

Gay-Lussac's Law

A gas law stating that pressure and temperature are directly proportional for constant volume ($V$) and moles ($n$): $\frac{P_1}{T_1} = \frac{P_2}{T_2}$.

Combined Gas Law

A gas law that shows the relationship between $P$, $V$, and $T$ when two quantities are changed while moles ($n$) remain constant: $\frac{P_1V_1}{T_1} = \frac{P_2V_2}{T_2}$.

London Dispersion Forces

Weak intermolecular forces that are exhibited by all molecules.

Dipole-Dipole Forces

Moderate intermolecular attractive forces exhibited by molecules that possess a net dipole.

Hydrogen Bonding

Strong intermolecular forces exhibited by molecules containing an $O-H$, $N-H$, or $H-F$ bond.

Sublimation

An endothermic phase transition in which energy is absorbed as a substance changes directly from a solid to a gas.

Deposition

An exothermic phase transition in which energy is released as a substance changes directly from a gas to a solid.