Chemistry Final Exam Review

Comprehensive vocabulary flashcards covering phases of matter, gas laws, atomic structure, nomenclature, stoichiometry, and thermochemistry from the chemistry final exam review.

Solid

A phase of matter with a definite shape, close packed particles, and low Kinetic Energy (KE).

Liquid

A phase of matter that takes the form of its container, has particles that are less close together, and has higher Kinetic Energy than a solid.

Gas

A phase of matter that expands to fill its container, has particles very far apart, and the highest Kinetic Energy (KE).

Evaporation

The transition from a liquid to a gas.

Sublimation

When a solid becomes a gas without passing through the liquid phase.

Triple point

The specific temperature and pressure at which solid, liquid, and gas phases all exist at equilibrium.

Ideal Gas

Gases assumed to be small hard spheres with no volume and no forces of attraction or repulsion between particles (perfectly elastic collisions) in constant random motion.

Real Gas

Gases whose particles do have volume and forces of attraction or repulsion; they behave as ideal except under very high pressure or very low temperatures.

Avogadro’s Hypothesis

Gases of the same volume at the same temperature and pressure ($T$ and $P$) must have the same number of particles.

Boyle’s Law

The gas law used to calculate pressure and volume changes when temperature is constant, defined by the equation $P_1V_1 = P_2V_2$.

Charles’ Law

The gas law for volume and temperature changes, defined as $\frac{V_1}{T_1} = \frac{V_2}{T_2}$.

Gay-Lussac's Law

The gas law for pressure and temperature changes, defined as $\frac{P_1}{T_1} = \frac{P_2}{T_2}$.

Element

A pure substance consisting of only one type of atom, such as titanium (Ti) or sulfur (S).

Compound

A substance consisting of more than one element chemically bonded together, such as carbon dioxide, water, or methane ($CH_4$).

Homogeneous mixture

A mixture of multiple elements or compounds in the same state of matter, such as steel (an alloy), salt water, or air.

Heterogeneous mixture

A mixture of multiple elements or compounds in different states of matter, such as chunky vegetable soup.

Isotopes

Atoms of the same element that contain the same number of protons and electrons but different numbers of neutrons, making them chemically alike.

Atomic radius (Period Trend)

The distance from the nucleus to the outer electrons, which decreases from left to right across a row because electrons are drawn to a more positive nucleus.

Ionization energy

The energy required to remove an electron from an atom.

Electronegativity

A measure of the tendency of an atom to attract a bonding pair of electrons; values are higher for nonmetals than for metals.

Aufbau method

A method used to determine the electron configuration orbital diagram of an atom.

Ionic compound

A compound composed of metals and non-metals held together by attraction between oppositely charged particles; they have high melting points and dissociate in water.

Molecular compound

A compound composed of two or more non-metals held together by sharing electrons (covalent bonds) with low melting points.

Formula unit

The smallest representative particle of an ionic compound, consisting of an anion-cation pair with a net charge of zero.

Polar covalent bonds

Intramolecular bonds between atoms within a molecule where electrons are unevenly shared.

Hydrogen bonds

Weak intermolecular forces of attraction between molecules due to dipole-dipole electrostatic attractions.

Double Replacement reaction

A chemical reaction where two compounds exchange ions, such as precipitation or acid-base neutralization reactions.

Limiting reagent

The reactant in a chemical reaction that is completely consumed first and determines the maximum amount of product that can be formed.

Theoretical yield

The maximum amount of product that can be produced from a given amount of reactant according to stoichiometric calculations.

Exothermic process

A process or reaction that releases heat to the surroundings, resulting in a negative enthalpy change ($\Delta H < 0$).

Endothermic process

A process or reaction that absorbs heat from the surroundings, resulting in a positive enthalpy change ($\Delta H > 0$).

Specific heat capacity

The amount of heat required to raise the temperature of one gram of a substance by one degree Celsius (e.g., $4.18\,J/g^\circ C$ for water).