Topic 3: Bonding and Structure

For a coloured solution containing a blue cation and a yellow anion, movement of the two colours towards opposite electrodes is especially direct migration evidence. A complete circuit and an ion-conducting path are needed for the observation.

Balance both the number of atoms and the total charge. The electron is on the product side for electron loss and on the reactant side for electron gain.

First determine the ions and their simplest whole-number ratio. Draw each ion in square brackets, show its outer shell after transfer, and write the charge outside the bracket. Repeat each ion the number of times required by the formula.

In a two-dimensional section, show alternating positive and negative ions in both row and column directions. A three-dimensional lattice continues the alternation through additional layers; the exact coordination depends on the compound.

The chemical formula gives the lowest whole-number ratio of ions needed for electrical neutrality. For example, NaCl means a 1:1 ratio and CaCl$_2$ means a 1:2 ratio across the giant lattice.

Every ion interacts with many surrounding ions. Opposite charges attract and like charges repel, but at the stable lattice spacing the overall arrangement has a strong net attraction and lower energy than widely separated ions.

The attraction acts in all directions through the giant lattice. This helps explain why separating ions requires substantial energy rather than breaking one isolated link.

BaCl$_2$ has stronger ionic attraction than CsCl because Ba$^{2+}$ has a higher charge and is smaller than Cs$^+$. More energy is therefore needed to separate its ions, consistent with its higher melting temperature.

For ions with the same charge down a group, ionic radius increases. Each step adds an occupied electron shell; the greater distance and shielding outweigh the increased nuclear charge.

All ions from N$^{3-}$ to Al$^{3+}$ are isoelectronic: they have ten electrons and the same occupied-shell pattern. As proton number increases, the same electron cloud experiences stronger nuclear attraction, so radius decreases.

The anion's electrons are drawn towards the cation, so the electron distribution is no longer symmetrical. Greater distortion increases electron density between the nuclei and gives the bonding more covalent character.

The cation has polarising power; the anion is polarisable. These terms describe the ability to cause distortion and the ease of being distorted, respectively.

The greatest distortion is expected when a small, highly charged cation is next to a large, highly charged anion. The bonding then shows more covalent character than a simple ionic model predicts.

F$^-$ is difficult to polarise because it is small and carries only a single negative charge. Al$^{3+}$ has much greater polarising power than Na$^+$ because it is smaller and has a higher charge.

Both nuclei attract the negatively charged shared pair. A stable bond length is reached where the attractions lower the energy while nucleus–nucleus and electron–electron repulsions prevent collapse.

Show outer-shell electrons only, using one symbol for electrons from each starting atom. A single, double or triple bond contains one, two or three shared pairs. Add every lone pair needed to complete the outer-shell accounting.

In a dative covalent bond, both electrons in the shared pair originally come from the same donor atom. Once formed, the pair is attracted to both nuclei like any other covalent bond.

Applications follow from structure: continuous strong bonds resist deformation, delocalised electrons carry charge, and weak attractions between graphite layers allow sliding.

If two bonded atoms have different electronegativities, the more electronegative atom attracts the pair more strongly and becomes partially negative, $\delta-$; the other becomes partially positive, $\delta+$.

Electronegativity is used comparatively. The difference between the two bonded atoms, rather than either value by itself, helps predict how unevenly the pair is shared.

Covalent and ionic bonding are limiting models on a continuum. As the electronegativity difference between bonded atoms increases, electron sharing becomes more unequal: bond polarity and ionic character increase.

Using Ba 0.9, Be 1.5 and Cl 3.0, $\Delta EN$ is 2.1 for Ba–Cl but 1.5 for Be–Cl. BaCl$_2$ therefore has more ionic character, while BeCl$_2$ has more covalent character.

CO$_2$ has two polar C=O bonds but is linear, so equal opposing dipoles cancel and the molecule is non-polar. H$_2$O is bent, so its O–H dipoles do not cancel and the molecule is polar. Symmetrical BF$_3$ and CCl$_4$ are also non-polar despite polar bonds.

Repulsion strength follows lone pair–lone pair > lone pair–bond pair > bond pair–bond pair because a lone pair is held by only one nucleus and occupies more space near the central atom.

Bond angle is the angle between two bonds that meet at the same central atom. The three atoms defining it must be identified; for H–O–H, oxygen is the vertex.

A bond length is an equilibrium value because attraction and repulsion balance at the minimum-energy separation. Bond angles describe the spatial arrangement produced by electron-region repulsions.

For two to six regions with no lone pairs, maximum separation gives linear, trigonal planar, tetrahedral, trigonal bipyramidal and octahedral arrangements. Lone pairs compress bond angles: CH$_4$ 109.5° becomes NH$_3$ 107° and H$_2$O 104.5° as lone pairs are added.

Include the ion's charge when counting electrons. After arranging all electron regions, ignore lone-pair positions when naming the molecular shape.

Outer electrons are no longer attached to one particular atom or one ion–ion pair. They move throughout the lattice, while the total negative charge of the electrons balances the total positive charge of the ions.

In a particle diagram, show closely packed positive ions in repeating layers and many electrons distributed between them. The diagram represents a continuous giant structure, not separate molecules.

Each ion is attracted to the shared electron sea around many ions. The attraction is non-directional and extends throughout the giant lattice.

More delocalised electrons per ion and a higher ion charge can strengthen the attraction; smaller metal ions also bring charge centres closer, when other structural factors are comparable.

Magnesium usually has stronger metallic bonding than sodium because each Mg atom contributes two delocalised electrons and forms Mg$^{2+}$, whereas sodium contributes one and forms Na$^+$; Mg$^{2+}$ is also smaller. The attraction is therefore stronger and more energy is required to overcome it.

Metals conduct as solids and when molten because their charge carriers are electrons, which remain mobile in both states. This contrasts with an ionic solid, whose ions are fixed until it melts or dissolves.