Part 1: Ionic, covalent and coordinate bonds

Explain bonding as electrostatic attraction and connect the particles in a structure to its properties.

Ionic bonding extends throughout a lattice

Ionic bonding is the strong electrostatic attraction between oppositely charged ions. Electron transfer explains ion formation; the bond itself is the attraction, not the transfer. A giant ionic lattice contains many ions arranged regularly in three dimensions, not individual salt molecules.

Mg loses two electrons and oxygen gains two, producing Mg²⁺ and O²⁻. In a dot-and-cross diagram, put each ion in brackets with its charge and show the appropriate outer electrons. One Mg²⁺ balances one O²⁻; MgCl₂ instead needs two chloride ions.

Melting requires much energy to overcome lattice attractions. Solid ionic compounds do not conduct because ions are fixed; molten compounds conduct because ions can move. Aqueous solutions conduct if enough ions dissolve. Solubility varies: do not claim every ionic solid is soluble.

Explain a melting-point comparison with particles and forces

MgO generally melts at a much higher temperature than NaCl. Both have giant ionic structures, so merely writing “MgO is ionic” does not explain the difference. Mg²⁺ and O²⁻ have larger charges than Na⁺ and Cl⁻; they also have smaller ionic radii in these comparisons. These factors produce stronger electrostatic attractions and more energy is needed to overcome them.

An ionic bond acts throughout a lattice, not only between a single electron donor and receiver. In a solid, ions vibrate about fixed positions but cannot migrate through the structure. Melting allows ions to move; it does not create delocalised metallic electrons. In an aqueous solution, mobile solvated ions carry charge.

Conductivity therefore needs both charge and mobility. A sodium chloride crystal already contains charged particles, but they are not mobile. A liquid molecular substance has mobile molecules, but ordinarily lacks enough mobile charged particles. “It can move” or “it has ions” alone is only half the explanation.

A shared pair attracts both nuclei

A covalent bond is the electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms. A single, double or triple bond contains one, two or three shared pairs. A larger average bond enthalpy indicates a stronger covalent bond.

In H₂O, oxygen makes two O–H bonds and keeps two lone pairs. In N₂, three shared pairs form a triple bond and each nitrogen retains one lone pair. Count the total valence electrons before drawing; hydrogen needs a duet, and some third-period atoms can have more than eight electrons around them.

In a coordinate (dative covalent) bond, both shared electrons come from the same atom. NH₃ donates its nitrogen lone pair to H⁺ to form NH₄⁺. The arrow indicating donation points N → H. After formation, the bonds in ammonium behave as equivalent covalent bonds.

For BF₃, show three B–F shared pairs and three lone pairs on each F; boron has six surrounding electrons. For SF₆, show six S–F shared pairs and three lone pairs on each F, with no central lone pair: sulfur has twelve surrounding electrons. These diagrams illustrate why an octet rule is not a universal instruction for every central atom.

Draw a covalent structure from a valence-electron inventory

Water has 6 + 1 + 1 = 8 valence electrons. Two O–H bonds use four, leaving four electrons as two oxygen lone pairs. Nitrogen, N₂, has 5 + 5 = 10; three shared pairs use six, leaving a lone pair on each nitrogen. A double bond is two shared pairs, not one exceptionally large pair.

For NH₄⁺ count 5 + 4(1) − 1 = 8 valence electrons: subtract one because the ion is positive. Four N–H bonds use all eight. To show formation from NH₃ and H⁺, one pair must be shown with both electrons originating from nitrogen. The final nitrogen has no lone pair, so ammonium is tetrahedral rather than pyramidal.

For H₃O⁺ count 6 + 3(1) − 1 = 8. Three bonds use six, leaving one lone pair on oxygen. Oxygen donates one original water lone pair to H⁺. The resulting ion has three bonds and one lone pair, giving a pyramidal atom arrangement. Donation moves electrons into a shared bond; it does not add new electrons to the total.

Molecules have bonds within and forces between

A simple molecular substance has strong covalent bonds inside each molecule but intermolecular forces between molecules. Melting or boiling separates molecules; it normally does not break their internal covalent bonds. This explains low melting/boiling points compared with many giant structures.

Iodine crystals contain I₂ molecules held by London forces. Ice contains H₂O molecules held in an extended hydrogen-bonded arrangement. Most simple molecular substances do not conduct because they lack mobile charged particles. Dissolving and reacting can create ions, so distinguish a pure molecule from its aqueous chemistry.

Infer a structure from several properties together

A substance has a high melting point, is non-conducting as a solid and conducts when molten. A giant ionic lattice is consistent: strong ion attractions explain melting, fixed ions explain the solid, and mobile ions explain the melt. No single observation alone proves the structure.

A low-boiling liquid that does not conduct is consistent with simple molecules. The molecules can separate without breaking the covalent bonds inside them. A solid conducting before it melts might instead be a metal or a form of carbon with mobile electrons; combine conductivity with other supplied evidence.

The metallic and giant covalent structures are developed in Periodicity. Use that cross-topic knowledge when comparing diamond, graphite, silicon and metals. Do not call a diamond a “large molecule” if the intended model is a continuous giant covalent lattice.

Quick checks

Original Finesse questions. Reveal the indicative worked solutions after attempting each question; these are not official OCR A mark allocations.

Sources