2.2.2 Bonding and structure

For ionic bonding, focus on what holds the ions together after electron transfer has happened: the strong electrostatic attraction between oppositely charged ions.

In a dot-and-cross diagram, show the outer electrons clearly, put each ion in brackets and include its charge. I want the diagram to show where the transferred electron came from, but the bond itself is the attraction between the ions, not a shared pair of electrons.

A giant ionic lattice is a three-dimensional repeating arrangement of positive and negative ions. Each ion is attracted electrostatically to oppositely charged ions around it, so the strong attractions act throughout the lattice in all directions.

That “giant” part matters. You are not dealing with separate NaCl molecules; NaCl is a formula ratio within an extended lattice.

Use the lattice to explain the properties. Melting or boiling requires you to overcome strong electrostatic attractions between oppositely charged ions, so ionic compounds usually have high melting and boiling points.

A solid ionic compound does not conduct because its ions are fixed in position. When molten, or when dissolved in water if the compound is soluble, the ions can move and carry charge. For solubility, think about whether attractions between the ions and the solvent can compensate for separating the ionic lattice; many ionic compounds dissolve well in polar water, but not all do.

For a covalent bond, picture the shared pair of electrons between two nuclei. Both positively charged nuclei attract that negative electron pair, and that electrostatic attraction holds the atoms together.

I’d use the words “shared pair of electrons” and “attraction to the nuclei”. Saying only that atoms “share electrons” describes what is present but misses the force that makes it a bond.

Use dots and crosses to keep track of which atom supplied each outer electron. Show every shared pair needed for single, double or triple bonds and include lone pairs where they matter.

For a dative covalent bond, both electrons in the shared pair come from the same atom or ion. If you use an arrow, it points from the lone-pair donor towards the electron-pair acceptor.

Average bond enthalpy gives you a useful measure of covalent bond strength: the larger the bond enthalpy, the more energy you need to break that bond in gaseous molecules, so the bond is stronger.

The word “average” is there because the exact strength of, say, a C–H bond depends slightly on the rest of the molecule. Tabulated values average over different molecular environments.

To work out shape, count the regions of electron density around the central atom, including lone pairs. Those electron pairs repel one another and arrange themselves as far apart as possible.

Then allow for lone pairs: lone pair–lone pair repulsion is stronger than lone pair–bond pair, which is stronger than bond pair–bond pair. That is why lone pairs squeeze bond angles below the ideal value. Count the bonded pairs and lone pairs before you try to name the shape.

• LP–LP repulsion > LP–BP repulsion > BP–BP repulsion

Once you know the number of bonding pairs and lone pairs, match it to the shape and angle. Learn a few anchor examples and then reason from electron-pair repulsion when lone pairs are present.

For example, CO₂ is linear at 180°, BF₃ is trigonal planar at 120°, CH₄ is tetrahedral at 109.5°, NH₃ is pyramidal at about 107°, H₂O is non-linear at about 104.5°, and SF₆ is octahedral with 90° angles.

Electronegativity tells you how strongly an atom attracts the bonding pair of electrons in a covalent bond. If the two bonded atoms have different electronegativities, the electron pair is pulled towards the more electronegative atom and you get δ− and δ+ ends to the bond.

Across a period, electronegativity generally increases; down a group it generally decreases. A larger Pauling electronegativity difference usually means a more polar bond and greater ionic character, but I wouldn’t invent a hard numerical boundary unless the question gives you one.

Do this in two stages. First decide which bonds are polar; then look at the three-dimensional shape and ask whether the bond dipoles cancel.

CO₂ has two polar C=O bonds, but the molecule is linear and symmetrical, so the dipoles cancel and the molecule is non-polar. H₂O is non-linear, so its O–H bond dipoles do not cancel and the molecule is polar. The shape is often the deciding step.

When you compare intermolecular forces, start with something that applies to every atom and molecule: the electron cloud is always fluctuating. That can create a temporary dipole, which then induces a dipole in a neighbouring particle and gives a London dispersion force.

If the molecule is polar, you also have permanent dipole–dipole attractions. Keep the two ideas separate in your explanation: polar molecules have permanent dipoles, but London forces are still present as well; they do not switch off.

Hydrogen bonding is a particularly strong intermolecular attraction. You need H covalently bonded to N, O or F, which makes that H strongly δ+, and a lone pair on N, O or F in a neighbouring molecule.

The attraction is between the δ+ hydrogen on one molecule and the lone pair on the electronegative atom of another. When you draw it, show the covalent N–H, O–H or H–F bond normally and the hydrogen bond as a dotted line between molecules.

Water forms an extensive network of hydrogen bonds. You need extra energy to overcome those intermolecular attractions, which helps explain why water has unusually high melting and boiling points for such a small molecule.

Ice gives you the other useful anomaly. Hydrogen bonding holds the molecules in an open lattice with more empty space than in liquid water, so the molecules are further apart on average and ice is less dense. That is why ice floats.

A simple molecular substance contains separate molecules. Strong covalent bonds hold the atoms together inside each molecule, while intermolecular forces hold neighbouring molecules together in the solid or liquid.

That distinction is crucial when you explain physical changes. Melting a simple molecular solid overcomes intermolecular forces between molecules; it does not break the covalent bonds inside the molecules.

Simple molecular substances often have relatively low melting and boiling points because you only need to overcome intermolecular forces between molecules when they change state; the covalent bonds inside the molecules remain intact.

They usually do not conduct electricity because they have no mobile ions or delocalised electrons. For solubility, compare the intermolecular attractions: polar molecules tend to interact better with polar solvents, while non-polar molecules tend to dissolve better in non-polar solvents. Use the actual molecules in the question rather than treating this as an absolute rule.