General Chemistry: Chemical Reactions, Energy, and States of Matter

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

Chemical equation

A symbolic representation of a chemical reaction using chemical formulas and other symbols to identify the starting materials (reactants) and the substances formed (products).

Reactants

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

Products

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

Coefficients

Numbers placed before chemical formulas in a balanced chemical equation that indicate the number of molecules or moles of a given element or compound that react or are formed.

Combination reaction

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

Decomposition reaction

A chemical reaction in which a single compound breaks down into two or more simpler substances.

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 contexts, it corresponds to the gain of oxygen atoms or the loss of hydrogen atoms.

Reduction

The gain of electrons by an atom; in organic contexts, it corresponds to 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 coupled oxidation and reduction components occurring together.

Oxidizing agent

A substance that causes another compound to be oxidized by accepting electrons, undergoing reduction itself in the process.

Avogadro’s number

The quantity $6.02 \times 10^{23}$, representing the number of atoms, molecules, or formula units in $1\,\text{mol}$ of any substance.

Mole–mole conversion factor

A conversion ratio derived from the stoichiometric coefficients of a balanced chemical equation used to relate the moles of one substance to the moles of another.

Theoretical yield

The maximum calculated quantity of product that can be formed in a chemical reaction from a given amount of starting reactant.

Actual yield

The measured amount of product that is experimentally obtained and isolated from a chemical reaction.

Percent yield

The ratio comparing actual yield to theoretical yield, expressed as $\text{Percent yield} = \frac{\text{actual yield}}{\text{theoretical yield}} \times 100\%$.

Law of conservation of energy

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

Potential energy

Stored energy, such as the chemical energy held within the bonds of molecules.

Kinetic energy

The energy associated with the motion of an object or particle.

Heat of reaction ($\Delta H$)

The enthalpy change representing the energy absorbed or released during a chemical reaction, calculated as the difference in potential energy between products and reactants.

Endothermic reaction

A reaction that absorbs heat from its surroundings, characterized by a positive $\Delta H$, products that are higher in energy than reactants, and broken reactant bonds that are stronger than formed product bonds.

Exothermic reaction

A reaction that releases heat to its surroundings, characterized by a negative $\Delta H$, products that are lower in energy than reactants, and formed product bonds that are stronger than broken reactant bonds.

Bond dissociation energy

The enthalpy change ($\Delta H$) required to cleave a covalent bond by equally dividing the shared electrons between the two atoms; it is always a positive value ($\Delta H > 0$).

Energy of activation ($E_a$)

The energy barrier that reactants must overcome to convert to products, measured as the energy difference between the reactants and the transition state.

Transition state

The unstable, high-energy arrangement of atoms at the maximum energy point along a reaction coordinate diagram.

Catalyst

A substance that accelerates the rate of a chemical reaction by lowering the energy of activation ($E_a$) without changing the overall heat of reaction ($\Delta H$).

Reversible reaction

A chemical reaction that can simultaneously proceed in the forward direction (towards the right) and in the reverse direction (towards the left).

Chemical equilibrium

The state in a reversible reaction where the rate of the forward reaction equals the rate of the reverse reaction, keeping the net concentrations of reactants and products unchanged over time.

Equilibrium constant ($K$)

The mathematical relationship relating the molar concentrations of products to reactants at equilibrium, given by $K = \frac{[\text{products}]}{[\text{reactants}]}$.

Le Châtelier’s principle

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

Boyle’s law

A gas law stating that gas volume is inversely proportional to pressure at constant temperature and amount of moles: $P_1V_1 = P_2V_2$.

Charles’s law

A gas law stating that gas volume is directly proportional to absolute temperature at constant pressure and amount of moles: $\frac{V_1}{T_1} = \frac{V_2}{T_2}$.

Gay-Lussac’s law

A gas law stating that gas pressure is directly proportional to absolute temperature at constant volume and amount of moles: $\frac{P_1}{T_1} = \frac{P_2}{T_2}$.

Combined gas law

A gas law showing the mathematical relationship among pressure, volume, and absolute temperature when the number of moles ($n$) is held constant: $\frac{P_1V_1}{T_1} = \frac{P_2V_2}{T_2}$.

London dispersion forces

Weak intermolecular attractive forces exhibited by all molecules due to temporary, fluctuating electron dipoles.

Dipole–dipole interactions

Moderate electrostatic intermolecular forces that occur between molecules possessing permanent net dipoles.

Hydrogen bonding

A strong intermolecular force occurring when a hydrogen atom covalently bonded to an electronegative atom ($\text{O-H}$, $\text{N-H}$, or $\text{H-F}$) interacts with a lone pair on another electronegative atom.

Sublimation

An endothermic phase change in which a substance converts directly from a solid to a gas, absorbing energy.

Deposition

An exothermic phase change in which a substance converts directly from a gas to a solid, releasing energy.

Melting

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

Freezing

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

Vaporization

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

Condensation

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