General Chemistry: Matter, Atomic Structure, and Periodic Trends

Comprehensive vocabulary flashcards covering basic chemical classification of matter, measurements, subatomic structure, historical atomic models, fundamental chemical laws, and periodic trends.

Matter

Anything that has mass and occupies space, existing in solid, liquid, or gas states.

Element

A pure substance consisting of only one type of atom, such as Oxygen ($O_2$) or Gold ($Au$).

Compound

A pure substance composed of two or more types of atoms chemically bonded in fixed ratios, such as Water ($H_2O$) or Sodium Chloride ($NaCl$).

Mixture

A physical combination of two or more substances in which each substance retains its own individual identity and properties.

Physical Property

A characteristic of a substance that can be observed or measured without changing its chemical identity, such as color, odor, density, melting point, and hardness.

Chemical Property

A characteristic that describes a substance's ability to undergo chemical changes and form new substances, such as flammability, toxicity, and acidity.

Supercritical Fluids

Substances at temperature and pressure conditions above their critical point that exhibit physical properties intermediate between liquids and gases.

Accuracy and Precision

Accuracy refers to how close a measurement is to the true value, whereas precision refers to how closely repeated measurements agree with each other.

Potential Energy

The energy stored within an object due to its position or chemical composition.

Temperature

A measure of the average kinetic energy of the particles within a substance.

Molecular Formula

A formula showing the actual number of atoms of each element present in a single molecule of a compound (e.g., $C_6H_{12}O_6$ for glucose).

Empirical Formula

A formula representing the simplest whole-number ratio of the atoms of each element in a compound (e.g., $CH_2O$ for glucose).

Nucleons

Subatomic particles located within the atomic nucleus, specifically protons (positive charge, mass of $1\,amu$) and neutrons (neutral charge, mass of $1\,amu$).

Atomic Number ($Z$)

The total number of protons in the nucleus of an atom, which defines the identity of the element.

Mass Number ($A$)

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

Isotopes

Atoms of the same element that have the same atomic number ($Z$) but different mass numbers ($A$) due to varying numbers of neutrons.

Atomic Weight

The average mass of an element calculated by accounting for the mass and natural abundance of all its naturally occurring isotopes.

Periodic Table

An arrangement of chemical elements organized by increasing atomic number into horizontal rows (periods) and vertical columns (groups) of elements with shared characteristics.

Periodic Trends

Systematic patterns in element properties across periods and down groups, including atomic radius, ionization energy, electron affinity, and electronegativity.

Atomic Radius

The size of an atom, defined as half the distance between bonded nuclei; it decreases across a period due to increasing $Z_{eff}$ and increases down a group as principal energy levels are added.

Ionic Radius

The size of an ion relative to its parent atom; cations are smaller than parent atoms due to reduced electron-electron repulsion, whereas anions are larger due to increased repulsion.

Isoelectronic Series

A series of atoms or ions that share identical electron configurations, where ionic size decreases as the nuclear charge ($Z$) increases.

Ionization Energy

The minimum amount of energy required to remove an electron from a gaseous atom or ion in its ground state.

Electron Affinity

The energy change associated with adding an electron to a gaseous atom to form a negative ion.

Electronegativity

A measure of the ability of a nucleus to pull shared electrons toward itself within a chemical bond.

Polarizability

A measure of how easily an atom's or ion's electron cloud can be distorted by an external electrical force.

Quantization

The concept that electrons exist only at specific, discrete energy levels rather than across a continuous spectrum.

Fluorescence

A short-lived emission of light occurring immediately after a molecule absorbs higher-energy (lower-wavelength) light.

Law of Constant Composition

The law stating that the elemental composition of a pure substance is fixed and never varies regardless of its source.

Law of Conservation of Mass

The law stating that total mass remains constant before and after a chemical reaction.

Law of Multiple Proportions

The law stating that if two elements form more than one compound, the masses of one element that combine with a fixed mass of the other are in ratios of small whole numbers.

Thomson's Cathode Ray Experiment

The experiment conducted by J.J. Thomson that discovered the electron and calculated its charge-to-mass ratio ($e/m = 1.76 \times 10^8\,C/g$).

Millikan's Oil Drop Experiment

The experiment by Robert Millikan that determined the electron charge ($1.60 \times 10^{-19}\,C$) and enabled the calculation of electron mass ($9.11 \times 10^{-28}\,g$).

Rutherford's Gold Foil Experiment

The experiment that established the nuclear atomic model by proving atoms contain a small, dense, positively charged nucleus.

Alkali Metals (Group 1 / Group 1A)

Soft, highly reactive metals with low densities and melting points that exhibit a $+1$ oxidation state and react vigorously with water.

Alkaline Earth Metals (Group 2 / Group 2A)

Reactive metallic elements exhibiting a $+2$ oxidation state that are harder and denser than alkali metals.

Halogens (Group 17 / Group 7A)

Highly electronegative nonmetals with a $-1$ oxidation state that form salts upon reacting with metals.

Noble Gases (Group 18 / Group 8A)

Extremely unreactive, inert nonmetal gases characterized by full valence octets.

Transition Metals (Groups 3–12)

Metals that exhibit variable oxidation states and can act as complex ions to form coordination complexes.