Chemical Bonding

Vocabulary flashcards covering chemical bonds, Lewis symbols, octet rule, ions, transition metal behavior, covalent vs ionic properties, electronegativity, and dipole moments.

Chemical bond

Attractive force holding two or more atoms together.

Covalent bond

Bond formed by sharing electrons between two nonmetals; usually found between nonmetals.

Ionic bond

Bond formed by transfer of electrons from a metal to a nonmetal.

Metallic bond

Attractive force holding together pure metals.

Lewis symbol

A representation of an element's valence electrons as dots around its symbol.

Unpaired electrons

Dots in a Lewis symbol that indicate electrons available for bonding.

Electron dot placement

Electrons are placed around the element symbol, often on four sides in a square pattern.

Octet rule

Atoms tend to gain, lose, or share electrons to achieve eight valence electrons (four pairs); many exceptions exist.

Noble gas configuration

A stable valence shell configuration (often s2p6) that many atoms imitate.

Exceptions to the octet rule

There are many cases where atoms do not achieve an octet like some expanded or deficient octets.

Electron transfer in ionic bonding

In NaCl formation, Na loses an electron to become Na+, Cl gains an electron to become Cl−.

Octet around ions

Ions like Na+ and Cl− achieve a complete octet around the ion, analogous to Ne and Ar configurations.

Mg2+ configuration

Mg2+ has a noble-gas like configuration, [Ne].

Mg+ configuration

Mg+ has [Ne]3s1 and is not considered a stable ion in this context.

Cl− configuration

Cl− achieves a filled shell equivalent to [Ar].

Polyatomic ion

An ion composed of two or more atoms bonded covalently with an overall charge.

Transition metal ionization order

For many transition metals, electrons are removed first from 4s before 3d orbitals.

Single bond

One shared pair of electrons between two atoms.

Double bond

Two shared pairs of electrons between two atoms.

Triple bond

Three shared pairs of electrons between two atoms.

Lewis structures

Covalent bonds shown with lines representing shared electron pairs; lone pairs may be shown as dots.

Dipole moment

Magnitude of molecular polarity, measured in Debye (D); HF is a classic polar molecule.

Polar bond

A covalent bond with unequal sharing of electrons, leading to partial charges.

Electronegativity

Ability of an atom in a molecule to attract electrons to itself; Pauli (Pauling) scale ranges roughly from 0.7 to 4.0.

Electronegativity trend (as given in notes)

Increases across a period and down a group (as stated in the notes; note that standard chemistry states increases across a period and up a group).

Electronegativity scale (Pauling)

Scale from about 0.7 (very low) to 4.0 (very high) used to compare atoms’ pull on electrons.

Bond polarity vs electronegativity difference

Polarity arises when electronegativity differences cause uneven electron sharing in a bond.

Bond distance trend

Bond lengths typically decrease from single to double to triple bonds.

Ionic compound properties (as per notes)

Usually solids with high melting points; conduct electricity when dissolved or molten; soluble in water; insoluble in organic solvents.

Covalent compound properties (as per notes)

Can be solids, liquids, or gases; lower melting points; do not conduct electricity; soluble in organic solvents; often insoluble in water.

Lewis structures

A schematic showing how atoms share electrons to form bonds and how lone pairs are arranged.

Bond polarity and polarity indicators

Polar bonds show partial charges (δ+ and δ−) due to unequal electron sharing.