Chemistry Fundamentals and College Admission Review

A comprehensive set of vocabulary flashcards covering physical and chemical properties, atomic structure, periodic trends, bonding, stoichiometry, thermodynamics, and gas laws based on college admission review notes.

Chemical Change

The change of matter that occurs with altering the chemical composition, often irreversible and accompanied by chemical reactions and the formation of new products.

Plasma

A state of matter with no definite shape or volume that contains free electrons and ions, making the particles electrically charged and highly conductive.

Bose-Einstein Condensate (BEC)

A state of matter formed at temperatures very close to absolute zero where bosons occupy the same quantum state and act as a single quantum entity.

Mixture

A physical blend of two or more substances that preserve their identities and are not chemically linked.

Homogeneous Mixture (Solution)

A mixture that has a uniform composition and appearance throughout, where different components are not distinguishable.

Heterogeneous Mixture

A mixture with a non-uniform composition and distinct regions or phases where components are distinguishable.

Colloids

A heterogeneous mixture where microscopically dispersed insoluble particles of one substance are suspended in another substance.

Suspension

A heterogeneous mixture containing solid particles large enough for sedimentation, meaning the solid eventually sinks.

Tyndall Effect

The effect of light scattering in a colloidal suspension while showing no light in a true solution.

Elements

Pure substances consisting of only one type of atom that cannot be broken down into simpler substances by chemical means.

Compounds

Pure substances composed of two or more elements chemically combined in fixed proportions.

Extensive Properties

Properties that depend on the amount of matter present in a system, such as mass, volume, and length.

Intensive Properties

Properties that are independent of the amount of substance present, such as color, boiling point, and density.

Dalton’s Atomic Theory

A theory proposed by John Dalton stating that all matter is made of atoms, which are indivisible and combine in simple whole-number ratios.

Dalton’s Billiard Ball Model

A model representing the atom as a ball-like structure before the discovery of the nucleus and electrons.

Nucleus

The central part of an atom containing positively charged protons and neutral neutrons.

Electron Cloud

The region surrounding the nucleus where negatively charged electrons orbit in various energy levels.

Mass Number (A)

The total number of protons and neutrons in the nucleus of an atom.

Atomic Number (Z)

The number of protons in the nucleus of an atom, identifying the element.

Isotopes

Atoms of the same element that have the same number of protons but different numbers of neutrons, resulting in different mass numbers.

Isobars

Atoms of different elements that have the same mass number ($A$) but different atomic numbers ($Z$).

Atomic Size

The measurement of the atomic radius, defined as the length from the nucleus to the valence electrons.

Ionization Energy

The amount of energy required to remove an electron from an isolated atom or molecule.

Electronegativity

The tendency of an atom or functional group to attract electrons toward itself.

Electron Affinity

The amount of energy released when an electron is added to the outermost shell of an atom.

Quantum Numbers

A set of four numerical values (principal, angular momentum, magnetic, and spin) that describe the properties of atomic orbitals and electrons.

Principal Quantum Number ($n$)

Represents the main energy level or shell of an electron, determining its distance from the nucleus.

Angular Momentum Quantum Number ($l$)

Describes the shape of the orbital where the electron is found, with values from $0$ to $n-1$.

Pauli Exclusion Principle

A principle stating that no two electrons in an atom can have the same set of four quantum numbers.

Aufbau Principle

The rule stating that electrons fill the lowest energy orbitals first before moving to higher energy levels.

Hund’s Rule

The rule that electrons will occupy degenerate orbitals singly before pairing up to minimize repulsion.

Ionic Bonding

A type of intramolecular bonding involving the transfer of electrons from a metal to a nonmetal, held by electrostatic attraction.

Covalent Bonding

A type of intramolecular bonding where two nonmetals share pairs of electrons to achieve a full valence shell.

Metallic Bonding

Bonding characterized by a "sea of electrons" where valence electrons are delocalized among metal cations.

VSEPR Theory

Valence Shell Electron Pair Repulsion theory, which predicts molecular shape based on the repulsion between electron pairs around a central atom.

Hybridization Theory

A theory explaining the mixing of atomic orbitals to form new hybrid orbitals suitable for covalent bonding.

Unsaturated Solution

A solution containing less than the maximum amount of solute that can be dissolved at a given temperature.

Saturated Solution

A solution that has reached its maximum capacity of dissolved solute at a specific temperature and pressure.

Supersaturated Solution

A solution that contains more dissolved solute than is normally possible, often achieved by heating and then cooling gently.

Mole

A chemical unit based on Carbon-12, where $1 \text{ mole} = 6.022 \times 10^{23}$ particles.

Empirical Formula

The simplest whole-number ratio of elements in a compound, such as $HO$ for hydrogen peroxide.

Molecular Formula

The formula representing the actual number of atoms of each element in a molecule, such as $H_2O_2$.

Alkanes

Saturated hydrocarbons containing only single bonds between carbon atoms with the general formula $C_n H_{2n+2}$.

Alkenes

Unsaturated hydrocarbons containing at least one carbon-carbon double bond with the general formula $C_n H_{2n}$.

Alkynes

Unsaturated hydrocarbons containing at least one carbon-carbon triple bond with the general formula $C_n H_{2n-2}$.

Aromatic Hydrocarbons

Hydrocarbons containing one or more benzene rings with delocalized $\pi$-electrons.

Half-Life

The time required for half of the radioactive nuclei in a sample to decay.

Kinetic Molecular Theory (KMT)

A theory stating that gas particles are in constant random motion, have negligible volume, and undergo elastic collisions.

Boyle's Law

The law stating that at constant temperature, gas pressure is inversely proportional to volume: $P_1 V_1 = P_2 V_2$.

Charles's Law

The law stating that at constant pressure, gas volume is directly proportional to absolute temperature ($K$): $\frac{V_1}{T_1} = \frac{V_2}{T_2}$.

Zeroth Law of Thermodynamics

States that if two systems are in thermal equilibrium with a third system, they are in thermal equilibrium with each other.

First Law of Thermodynamics

The Law of Energy Conservation, stating energy cannot be created or destroyed, only transferred or converted.

Second Law of Thermodynamics

States that the entropy of an isolated system always increases over time, introducing the concept of irreversibility.

Third Law of Thermodynamics

States that as temperature approaches absolute zero ($0 \text{ Kelvin}$), the entropy of a perfect crystal approaches zero.

Specific Heat Capacity

The amount of heat required to raise the temperature of one kilogram of a substance by one degree Celsius.

Le Chatelier’s Principle

The principle that a system at equilibrium will adjust itself to counteract any change in concentration, pressure, or temperature.

Oxidation

The loss of electrons by a substance, resulting in an increase in oxidation state.

Reduction

The gain of electrons by a substance, resulting in a decrease in oxidation state.

Arrhenius Theory

A theory defining acids as substances that generate $H^+$ ions and bases as substances that produce $OH^-$ ions in solution.

Bronsted-Lowry Theory

A theory defining an acid as a proton donor and a base as a proton acceptor.

Lewis Theory

A theory defining acids as electron-pair acceptors and bases as electron-pair donors.