General Chemistry: Periodic Properties, Bonding, Lewis Structures, and VSEPR Theory

A comprehensive vocabulary practice set covering periodic trends, chemical bonding, polyatomic ion formulas and charges, Lewis theory, resonance, formal charges, and VSEPR electron/molecular geometries.

Metal

An element that generally tends to lose electrons during chemical reactions to form cations; typically lustrous, malleable, ductile, and a good conductor of heat and electricity.

Nonmetal

An element that generally tends to gain or share electrons during chemical reactions, forming anions or covalent bonds, and typically lacks metallic physical properties.

Metalloid

An element that possesses physical and chemical properties intermediate between those of metals and nonmetals.

Alkali Metal

A reactive element in Group 1 (1A) of the periodic table containing 1 valence electron, which readily loses one electron to form a $+1$ cation.

Alkaline Earth Metal

A reactive metallic element in Group 2 (2A) of the periodic table containing 2 valence electrons, which readily loses two electrons to form a $+2$ cation.

Halogen

A nonmetallic element belonging to Group 17 (7A) of the periodic table containing 7 valence electrons, which readily gains one electron to form a $-1$ anion.

Noble Gas

An element belonging to Group 18 (8A) of the periodic table with an exceptionally stable, complete valence electron shell and very low chemical reactivity.

Atomic Radius

A measure of the size of an atom, generally defined as the distance from the nucleus to the outer boundary of its surrounding electron cloud; it decreases across a period and increases down a column.

Effective Nuclear Charge

The net positive charge experienced by an electron in an atom, equal to the total nuclear charge minus the shielding effect caused by core and other intervening electrons.

Shielding

The electrostatic repulsion exerted by core (inner) electrons that weakens the net attractive pull of the positive nucleus on the outer valence electrons.

Isoelectronic

Describing chemical species (atoms, cations, or anions) that possess the exact same number of electrons and identical electron configurations.

Ionization Energy (IE)

The minimum energy required to remove an electron from an isolated, gaseous neutral atom or ion in its ground state.

First Ionization Energy

The energy required to remove the first and outermost valence electron from a neutral gaseous atom, producing a $+1$ cation.

Second Ionization Energy

The energy required to remove a second electron from an already positively charged gaseous ion ($+1$ cation).

Electron Affinity (EA)

The energy change associated with adding an electron to an isolated, gaseous atom to form an anion.

Paramagnetic

The magnetic property of an atom, ion, or molecule that contains one or more unpaired electrons, causing it to be attracted into an external magnetic field.

Diamagnetic

The magnetic property of an atom, ion, or molecule in which all electrons are paired, causing it to be weakly repelled by an external magnetic field.

Chemical Bond

An attractive electrostatic force between atoms or ions that lowers the potential energy of the interacting species and holds them together.

Ionic Bond

A chemical bond formed via electrostatic attraction between oppositely charged ions, typically resulting from the transfer of electrons from a metal to a nonmetal.

Ionic Compound

A neutral compound composed of cations and anions bound together in a continuous three-dimensional crystal lattice by ionic bonds.

Covalent Bond

A chemical bond formed between nonmetal atoms through the mutual sharing of one or more pairs of valence electrons.

Molecular Compound

A compound composed of discrete, neutral molecules held together internally by covalent bonds.

Chemical Formula

A concise symbolic representation that designates the elements present in a compound and the relative or absolute number of atoms of each element.

Binary Compound

A chemical compound composed of exactly two different elements.

Polyatomic Ion

A group of covalently bonded atoms that carries an overall net electrical charge.

Hydroxide

A polyatomic anion with the chemical formula $\text{OH}^-$ and a charge of $-1$.

Cyanide

A polyatomic anion with the chemical formula $\text{CN}^-$ and a charge of $-1$.

Phosphate

A polyatomic anion with the chemical formula $\text{PO}_4^{3-}$ and a charge of $-3$.

Nitrate

A polyatomic anion with the chemical formula $\text{NO}_3^-$ and a charge of $-1$.

Carbonate

A polyatomic anion with the chemical formula $\text{CO}_3^{2-}$ and a charge of $-2$.

Bicarbonate

A polyatomic anion with the chemical formula $\text{HCO}_3^-$ and a charge of $-1$.

Sulfate

A polyatomic anion with the chemical formula $\text{SO}_4^{2-}$ and a charge of $-2$.

Ammonium

A polyatomic cation with the chemical formula $\text{NH}_4^+$ and a charge of $+1$.

Acetate

A polyatomic anion with the chemical formula $\text{C}_2\text{H}_3\text{O}_2^-$ and a charge of $-1$.

Electronegativity

The relative ability of an atom bonded within a molecule to attract shared bonding electrons toward itself.

Polar Covalent Bond

A covalent bond formed between atoms with different electronegativities, leading to unequal sharing of electrons and a permanent bond dipole.

Lewis Structure

A two-dimensional structural drawing that shows the distribution of valence electrons as bonding pairs (lines) and lone pairs (dots) around bonded atoms.

Octet Rule

The general principle stating that main-group atoms tend to gain, lose, or share electrons until they are surrounded by eight valence electrons.

Bonding Pair

A pair of valence electrons shared between two nuclei to form a covalent bond.

Lone Pair

A pair of nonbonding valence electrons that is localized on a single atom and not shared in a covalent bond.

Single Bond

A covalent bond in which two atoms share exactly one pair of electrons (a total of 2 electrons).

Double Bond

A covalent bond in which two atoms share two pairs of electrons (a total of 4 electrons).

Triple Bond

A covalent bond in which two atoms share three pairs of electrons (a total of 6 electrons).

Resonance Structure

One of two or more valid Lewis structures for a single molecule or ion having the same arrangement of atomic nuclei but different allocations of electrons.

Resonance Hybrid

The actual physical electronic structure of a resonance-stabilized molecule, representing a weighted intermediate blend of all valid resonance structures.

Formal Charge

A hypothetical charge assigned to an atom in a Lewis formula, computed as: $\text{Valence } e^- - \text{Nonbonding } e^- - \frac{1}{2}(\text{Bonding } e^-)$.

Free Radical

An odd-electron molecule or ion that possesses an unpaired valence electron in its Lewis structure.

Expanded Octet

An exception to the octet rule in which a central atom from Period 3 or higher holds more than 8 valence electrons, accommodated by available $d$ orbitals.

Incomplete Octet

An exception to the octet rule where a stable compound has a central atom (such as beryllium or boron) surrounded by fewer than eight valence electrons.

Valence Shell Electron Pair Repulsion (VSEPR) Theory

A model that predicts the 3D shapes of molecules based on minimizing electrostatic repulsions between electron groups surrounding the central atom.

Electron Groups

Collections of valence electrons around a central atom that repel each other, including single bonds, double bonds, triple bonds, lone pairs, or single electrons.

Electron Geometry

The three-dimensional arrangement formed by all electron groups (both bonding pairs and lone pairs) surrounding the central atom.

Molecular Geometry

The three-dimensional spatial arrangement of the bonded nuclei only around the central atom, excluding the nonbonding electron pairs.

Linear Geometry

The arrangement of two electron groups directed at $180^\circ$ from each other around a central atom.

Trigonal Planar Geometry

The arrangement of three electron groups in a single two-dimensional plane around a central atom, with ideal bond angles of $120^\circ$.

Tetrahedral Geometry

The three-dimensional arrangement of four electron groups directed toward the vertices of a regular tetrahedron, with ideal bond angles of $109.5^\circ$.

Trigonal Bipyramidal Geometry

The arrangement of five electron groups around a central atom consisting of three equatorial positions ($120^\circ$ apart) and two axial positions ($90^\circ$ to the equatorial plane).

Octahedral Geometry

The three-dimensional arrangement of six electron groups distributed symmetrically around a central atom with all adjacent bond angles equal to $90^\circ$.

Trigonal Pyramidal Geometry

The molecular geometry resulting from four electron groups consisting of three bonding groups and one lone pair, characterized by bond angles slightly less than $109.5^\circ$.

Bent Geometry

The molecular geometry observed when a central atom has either three electron groups (two bonding, one lone pair) with bond angles $< 120^\circ$, or four electron groups (two bonding, two lone pairs) with bond angles $< 109.5^\circ$.

Seesaw Geometry

The molecular geometry resulting from five electron groups composed of four bonding groups and one equatorial lone pair around a central atom.

T-shaped Geometry

The molecular geometry resulting from five electron groups composed of three bonding groups and two equatorial lone pairs around a central atom.

Square Pyramidal Geometry

The molecular geometry resulting from six electron groups consisting of five bonding groups and one lone pair around a central atom.

Square Planar Geometry

The molecular geometry resulting from six electron groups consisting of four bonding groups in a plane and two opposing lone pairs on opposite sides of the central atom.