Chemical Bonding

The One Thing

Chemical bonds form when atoms achieve a more stable, lower-energy arrangement, usually by completing their outermost electron shell. The bonding type—electrovalent, covalent or coordinate—depends on whether electrons are transferred, shared mutually or donated by one atom.

Reactions, Processes and Experiments

Key Terms

• Valence electrons: Electrons in the outermost shell that participate in chemical bonding.
• Octet rule: The tendency of atoms to gain, lose or share electrons to obtain eight outer-shell electrons.
• Duplet rule: The tendency of atoms with one shell, such as hydrogen, to obtain two outer-shell electrons.
• Electrovalent bond: A bond formed by complete transfer of one or more electrons, producing oppositely charged ions.
• Cation: A positively charged ion formed when an atom loses electrons.
• Anion: A negatively charged ion formed when an atom gains electrons.
• Electrovalency: The number of electrons lost or gained by an atom to form an ion.
• Ionic compound: A compound of oppositely charged ions held together by electrostatic attraction.
• Covalent bond: A bond formed when two atoms share one or more electron pairs.
• Shared pair: A pair of electrons jointly used by two bonded atoms.
• Single covalent bond: A bond formed by sharing one electron pair, represented by one line, such as H-H.
• Double covalent bond: A bond formed by sharing two electron pairs, represented by two lines, such as O=O.
• Triple covalent bond: A bond formed by sharing three electron pairs, represented by three lines, such as N≡N.
• Covalency: The number of electron pairs shared by an atom in covalent bonding.
• Coordinate bond: A covalent bond in which both electrons in the shared pair are supplied by one atom.
• Donor atom: The atom that provides both electrons for a coordinate bond.
• Acceptor atom: The atom that accepts the shared pair and provides an empty orbital or space for bonding.
• Lone pair: A pair of valence electrons not involved in bonding.
• Dative bond: Another name for a coordinate covalent bond, represented by an arrow from donor to acceptor.
• Lewis dot structure: A diagram showing valence electrons as dots and shared electron pairs as bonds.
• Electronegativity: The tendency of an atom to attract the shared electron pair towards itself.
• Polar covalent bond: A covalent bond in which the shared electrons are attracted unequally by the bonded atoms.
• Non-polar covalent bond: A covalent bond in which the shared electrons are distributed equally or nearly equally.

Easily Confused

• Electrovalent and covalent bonds: Electrovalent bonds involve electron transfer, whereas covalent bonds involve mutual sharing of electrons.
• Covalent and coordinate bonds: Both involve a shared electron pair, but in a coordinate bond both electrons originate from one atom.
• Cation and anion: A cation is formed by electron loss and is positive; an anion is formed by electron gain and is negative.
• Electrovalency and covalency: Electrovalency is the number of electrons lost or gained, whereas covalency is the number of electron pairs shared.
• Single, double and triple bonds: These contain one, two and three shared electron pairs respectively.
• Polar and non-polar covalent bonds: Polar bonds have unequal electron attraction, whereas non-polar bonds have equal or nearly equal electron distribution.
• Solid and molten ionic compounds: Solid ionic compounds do not conduct electricity because their ions are fixed, whereas molten or aqueous ionic compounds conduct because their ions can move.
• Duplet and octet rules: Hydrogen and helium follow the duplet rule, while most other atoms tend towards an octet; the octet rule also has incomplete- and expanded-shell exceptions.

What Gets Asked

• Identify the bonding type from electron behaviour: State whether electrons are transferred, mutually shared or donated as a complete pair. Confusing coordinate donation with ordinary mutual sharing loses the distinction.
• Write ionic formation equations: For sodium chloride, include all three stages: Na -> Na+ + e-, Cl + e- -> Cl- and Na+ + Cl- -> NaCl. Omitting ion formation or the final electrostatic combination loses marks.
• Represent covalent bonds correctly: Use H-H for a single bond, O=O for a double bond and N≡N for a triple bond. The number of lines must match the number of shared electron pairs.
• Explain ammonium and hydronium formation: Use NH3 + H+ -> NH4+ and H2O + H+ -> H3O+, identifying nitrogen or oxygen as the donor of the lone pair. Do not state that the hydrogen ion supplies the shared pair.
• Explain ionic properties: Relate high melting and boiling points, hardness and brittleness to strong ionic attraction, and electrical conductivity in molten or aqueous states to mobile ions. Solid ionic compounds generally do not conduct.
• Compare ionic and covalent compounds: Ionic compounds generally have high melting and boiling points and conduct when ions are mobile; covalent compounds generally have lower melting and boiling points and do not conduct because they lack free ions or mobile electrons.