Chemical Bonds

Chemical Bond

Attractive force holding atoms together to form compounds

Ionic Bond

Transfer of electrons from one atom to another (metal + nonmetal)

Covalent Bond

Sharing of electrons between two nonmetals

Polar Covalent Bond

Unequal sharing of electrons causing partial (+) and (–) charges

Nonpolar Covalent Bond

Equal sharing of electrons; no charge difference across molecule

Metallic Bond

Attraction between metal cations and a sea of delocalized electrons

Electronegativity

Measure of an atom’s attraction for shared electrons in a bond

Electronegativity Difference

Determines bond type: 0 = nonpolar, 0.4–1.7 = polar, >1.7 = ionic

Ion

Atom that has gained or lost electrons to become charged

Cation

Positively charged ion formed when an atom loses electrons (usually a metal)

Anion

Negatively charged ion formed when an atom gains electrons (usually a nonmetal)

Crystal Lattice

Repeating 3-D pattern of positive and negative ions in an ionic compound

Bond Energy

Energy required to break one mole of bonds in a substance

Single Bond

One shared pair of electrons between atoms

Double Bond

Two shared pairs of electrons between atoms

Triple Bond

Three shared pairs of electrons between atoms

Lewis Structure

Diagram showing bonding and lone-pair electrons in a molecule

Lone Pair

Pair of valence electrons not involved in bonding

Valence Electrons

Electrons in the outermost energy level; involved in bonding

Octet Rule

Atoms bond to achieve eight electrons in their outer shell for stability

VSEPR Theory

Predicts molecular shapes based on electron-pair repulsion

Molecular Geometry

Three-dimensional shape of a molecule (e.g., linear, bent, tetrahedral)

Polar Molecule

Molecule with uneven charge distribution; has a dipole moment

Nonpolar Molecule

Molecule with even charge distribution; no dipole moment

Dipole

Separation of charge in a polar bond or molecule (partial + and – ends)

Intermolecular Forces (IMFs)

Forces between molecules that affect melting, boiling, and solubility

Hydrogen Bond

Strong IMF between H and N, O, or F atoms

Dipole–Dipole Force

Attraction between positive and negative ends of polar molecules

London Dispersion Force

Weak IMF caused by temporary electron shifts in all molecules

Network Covalent Bonding

Covalent bonding forming a continuous 3-D network (e.g., diamond, SiO₂)

Metallic Properties

Conductivity, malleability, and luster due to delocalized electrons in metals

Bond Length

Distance between nuclei of two bonded atoms; shorter = stronger bond

Bond Polarity

Degree to which electrons are shared unequally between atoms

Resonance Structure

When multiple Lewis structures represent the same molecule’s bonding