3. Chemical bonding

When two bonded atoms have different electronegativities, the bonding electrons are attracted more strongly towards the atom with the higher value. This unequal attraction can make the bond polar; similar values indicate more equal sharing.

Use electronegativity to describe the direction and relative strength of attraction within a bond. Periodic trends and the factors that cause them belong to the neighbouring objectives; numerical differences used to predict ionic or covalent bonding also belong to the later application card.

A larger atomic radius places the bonding electron pair farther from the nucleus, so attraction is weaker and electronegativity is lower. A smaller radius brings the bonding electrons closer and strengthens the attraction.

Inner-shell electrons shield the outer region from some nuclear attraction. Greater shielding reduces the effective pull on bonding electrons; therefore proton number alone cannot determine electronegativity. Use nuclear charge, radius and shielding together.

Down a group, electronegativity generally decreases. Although nuclear charge increases, new occupied shells increase atomic radius and inner-shell shielding, so the bonding electron pair is farther from the nucleus and feels a weaker effective attraction.

Explain a trend by linking direction to the three factors: nuclear charge, distance/atomic radius and shielding. State the overall trend first, then identify which factor strengthens or weakens attraction and why the dominant effect gives the observed direction.

Equal values give equal attraction. As Δχ increases, sharing becomes more unequal; for a sufficiently large difference, the prediction changes to ions held by ionic bonding. Use the supplied data and the chemical context rather than treating the examples as a memorised list.

Electrostatic means the attraction is caused by charge. In an ionic solid, each ion is attracted to surrounding ions of opposite charge throughout a repeating lattice; the bond is not confined to one isolated ion pair.

A cation carries positive charge and an anion carries negative charge. The ionic compound as a whole is neutral because the total positive and negative charges balance.

For every example, verify two things separately: electrons lost equal electrons gained, and ionic charges sum to zero in the formula unit. Subscripts count the ion ratio; they do not alter an individual ion's charge.

The metallic structure is therefore a lattice of positive metal ions surrounded by a mobile sea of delocalised electrons. The positive ions remain in the lattice while the delocalised electrons move through the structure.

Metallic bonding is the strong electrostatic attraction between the positive metal ions and the surrounding delocalised electrons. This attraction holds the ions in place, counteracts repulsion between positive ions and maintains the stability of the metal structure.

An octet is not an absolute limit for the central Period 3 atom in the specified cases: sulfur has an expanded octet in SO₂ and SF₆, while phosphorus has an expanded octet in PCl₅. This exception does not change the covalent-bond definition.

In this reaction, nitrogen donates its lone pair to hydrogen to form the fourth N–H bond in NH₄⁺; Cl⁻ is the counter-ion. In 2AlCl₃ ⇌ Al₂Cl₆, lone pairs on bridging chlorine atoms are donated to electron-deficient aluminium atoms, forming two coordinate bonds in the dimer.

Every single bond is one σ bond. A double bond contains one σ and one π bond; a triple bond contains one σ and two π bonds formed from two perpendicular pairs of p orbitals.

Hybridisation mixes orbitals on the same atom. sp³ leaves no unhybridised p orbital, sp² leaves one for one π bond, and sp leaves two for two π bonds.

A larger bond energy means a stronger bond and more energy is required to break it. For the same pair or a comparable series of atoms, a shorter bond generally reflects stronger attraction between nuclei and bonding electrons and is associated with a stronger bond.

From HCl to HI, the halogen atom becomes larger, the H–X bond becomes longer and attraction across the bond weakens. Less energy is therefore needed to break H–I, so HI is the most reactive of these hydrogen halides in a comparison where breaking the H–X bond controls the reaction.

Lone pairs occupy domains but are not named as atoms in the molecular shape. Their electron density is concentrated closer to the central atom, so repulsion follows lone pair–lone pair > lone pair–bond pair > bond pair–bond pair and lone pairs compress neighbouring bond angles.

CH₄, NH₃ and H₂O all have four electron domains, but replacing bonding pairs with lone pairs increases repulsion and reduces the observed angle from 109.5° to 107° and then 104.5°. By contrast, the two C=O double bonds in CO₂ count as two domains and point 180° apart.

For NH₄⁺, nitrogen supplies five valence electrons, four hydrogens supply four, and the positive charge removes one: 5 + 4 − 1 = 8 electrons. All four pairs are N–H bonding pairs, so four equivalent domains adopt the tetrahedral arrangement and give 109.5°.

When an analogous species contains lone pairs, begin with the ideal electron-domain angle and use stronger lone-pair repulsion to justify compression. Transfer an exact numerical angle only when the species is directly analogous at the level expected by the question; real angles can vary with the atoms and bonding present.

Ammonia and water both hydrogen-bond because their N–H or O–H bonds are strongly polar and the N or O atoms carry lone pairs. In water, each molecule has two donor hydrogens and two acceptor lone pairs, so an extended intermolecular network can form.

A bond dipole records the direction of this charge separation, with the dipole arrow pointing towards the partially negative end. A larger electronegativity difference gives a more polar bond, but bond polarity is a property of the individual bond before the whole molecular arrangement is considered.

To decide whether a molecule has an overall dipole, identify the polarity and direction of every polar bond, then use the molecular shape to see whether the dipoles cancel. In CH₃Cl they do not cancel, so the molecule is polar; in the symmetric CCl₄ arrangement they cancel, so the molecule is non-polar despite its polar C–Cl bonds.

Id–id forces generally strengthen when a species has more electrons because its electron cloud is more polarisable; greater molecular contact can also strengthen the total attraction. Pd–pd forces additionally align opposite permanent partial charges.

When a simple molecular substance melts or boils, intermolecular attractions are overcome while the covalent bonds inside each molecule normally remain intact. This is why a low boiling point does not mean the molecule's covalent bonds are weak.

Choose the bonding model before drawing: show transfer of valence electrons for an ionic compound, shared pairs for a covalent compound, and a shared pair whose two electrons come from one donor for a coordinate bond. In a displayed formula, the coordinate-bond arrow points away from the donor lone pair.

After drawing, count the shown outer electrons, check each ion's charge and confirm that the shared or transferred electrons match the bonding model. The normal noble-gas/octet pattern is a useful check, but it is not an unconditional rule.