Chemical Reactions, Energy, Equilibrium, and States of Matter

Vocabulary practice flashcards covering chemical equations, reaction types, oxidation-reduction, stoichiometry, enthalpy and energy changes, equilibrium, gas laws, intermolecular forces, and phase changes.

Chemical equation

A representation that uses 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 that are written on the left side of a chemical equation.

Products

The substances formed in a chemical reaction that are written on the right side of a chemical equation.

Coefficient

A whole number written in front of a formula in a chemical equation showing the number of molecules of a given element or compound that react or are formed.

Combination reaction

A type of reaction characterized by 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—such as atoms or ions—to form two new compounds.

Oxidation

The loss of electrons from an atom, often characterized by the gain of oxygen atoms or the loss of hydrogen atoms.

Reduction

The gain of electrons by an atom, often characterized by the loss of oxygen atoms or the gain of hydrogen atoms.

Redox reaction

A chemical reaction involving the transfer of electrons from one element to another, always containing both an oxidized component and a reduced component.

Oxidizing agent

A compound that causes another compound to be oxidized by gaining electrons and thereby becoming reduced itself.

Percent yield

The measure of the efficiency of a reaction, calculated as $\frac{\text{actual yield}}{\text{theoretical yield}} \times 100\%$.

Potential energy

Stored energy, such as that stored within chemical bonds, where lower potential energy corresponds to higher stability.

Kinetic energy

The energy of motion.

Law of conservation of energy

A fundamental physical principle stating that the total energy in a system does not change and that energy cannot be created or destroyed.

Heat of reaction ($\Delta H$)

The energy absorbed or released in a chemical reaction, also known as enthalpy change.

Endothermic reaction

A chemical reaction where heat is absorbed, $\Delta H$ is positive, the bonds broken in the reactants are stronger than the bonds formed in the products, and products are higher in energy than reactants.

Exothermic reaction

A chemical reaction where heat is released, $\Delta H$ is negative, the bonds formed in the products are stronger than the bonds broken in the reactants, and products are lower in energy than reactants.

Bond dissociation energy

The enthalpy change ($\Delta H$) required to break a covalent bond by equally dividing the electrons between the two atoms, which is always a positive value.

Energy of activation ($E_a$)

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

Catalyst

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

Reversible reaction

A chemical reaction that can occur in either direction, proceeding to the right as a forward reaction and to the left as a reverse reaction.

Chemical equilibrium

The state reached when the rate of the forward reaction equals the rate of the reverse reaction, resulting in constant net concentrations of reactants and products.

Equilibrium constant ($K$)

The ratio relating the concentration of products to the concentration of reactants at equilibrium, given by the expression $K = \frac{[\text{products}]}{[\text{reactants}]}$.

Le Châtelier's principle

The principle stating that if a chemical system at equilibrium is disturbed by changes in concentration, temperature, or pressure, the system reacts in a direction that counteracts the disturbance.

Boyle's law

A gas law stating that pressure ($P$) and volume ($V$) are inversely related for a constant temperature ($T$) and amount ($n$), expressed as $P_1V_1 = P_2V_2$.

Charles's law

A gas law stating that volume ($V$) and temperature ($T$) are directly proportional for a constant pressure ($P$) and amount ($n$), expressed as $\frac{V_1}{T_1} = \frac{V_2}{T_2}$.

Gay-Lussac's law

A gas law stating that pressure ($P$) and temperature ($T$) are directly proportional for a constant volume ($V$) and amount ($n$), expressed as $\frac{P_1}{T_1} = \frac{P_2}{T_2}$.

Combined gas law

A gas law relating pressure, volume, and temperature for a constant number of moles ($n$), expressed as $\frac{P_1V_1}{T_1} = \frac{P_2V_2}{T_2}$.

London dispersion forces

A weak intermolecular attraction exhibited by all molecules.

Dipole-dipole forces

An intermolecular attraction of moderate relative strength exhibited by molecules having a net dipole.

Hydrogen bonding

A strong intermolecular force exhibited by molecules containing an O-H, N-H, or H-F covalent bond.

Sublimation

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

Deposition

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