Covalent Bonding

a shared pair of electrons between two atoms

the shared pair of electrons is attracted to the positive nucleus of each atom; this electrostatic attraction between the nuclei and the shared electrons bonds the atoms together

three shared pairs / a triple bond

each oxygen atom has six outer electrons and needs two more to complete its outer shell; therefore two shared pairs (a double bond) are needed

one lone pair

each oxygen atom shares two pairs of electrons with the carbon atom; this forms two double covalent bonds

four

each carbon atom needs to form four bonds to have a complete outer shell; each carbon bonds to two hydrogen atoms and shares two pairs of electrons with the other carbon, forming a C=C double bond

diagram shows five C–H bonds (one shared pair each); one C–C bond (one shared pair); one C–Cl bond (one shared pair); three lone pairs on Cl

chlorine has a simple molecular structure; the intermolecular forces between Cl₂ molecules are weak; little energy is required to overcome these forces, so the boiling point is very low

intermolecular forces of attraction (between water molecules)

bromine has a greater relative molecular mass than chlorine; the intermolecular forces between Br₂ molecules are stronger; more energy is required to overcome these stronger intermolecular forces

higher; iodine has a greater relative molecular mass so the intermolecular forces between I₂ molecules are stronger, requiring more energy to overcome

silicon dioxide has a giant covalent structure; it contains many strong covalent bonds throughout the structure; a large amount of energy is required to break these bonds

when water boils, weak intermolecular forces between molecules are overcome (covalent bonds within molecules are not broken); when silicon dioxide melts, strong covalent bonds throughout the giant structure are broken

each carbon atom is bonded to four other carbon atoms by strong covalent bonds; these bonds extend in all directions throughout the giant structure, making it very rigid

in graphite, each carbon atom forms only three covalent bonds; this leaves one delocalised electron per carbon atom that is free to move along the layers and carry charge; in diamond, all four outer electrons are used in covalent bonds so there are no free electrons to carry charge

C₆₀ fullerene has a simple molecular structure made of individual C₆₀ molecules; the intermolecular forces between C₆₀ molecules are weak and require little energy to overcome; diamond has a giant covalent structure where many strong covalent bonds must be broken, requiring much more energy

the layers in graphite are held together by weak intermolecular forces; the layers can slide over each other easily

methane is a covalent compound with no free electrons or ions; there are no charged particles free to move and carry charge