Bonding and structure

Every chemical bond is an electrostatic attraction between positive and negative charges. What differs is where the charges are: between ions, between nuclei and shared electrons, or between metal ions and a sea of electrons. This topic explains each type of bond, the shapes that molecules take, how bonds become polar, and how the structure of a substance explains its melting temperature, conductivity and solubility.

1. Ionic bonding

When a metal reacts with a non-metal, electrons are transferred. The metal atoms lose their outer electrons to become positive ions, and the non-metal atoms gain electrons to become negative ions. In simple cases each ion ends up with the electron configuration of a noble gas: sodium (2,8,1) becomes Na⁺ (2,8) and chlorine (2,8,7) becomes Cl⁻ (2,8,8). In terms of sub-shells, Mg 1s² 2s² 2p⁶ 3s² becomes Mg²⁺ 1s² 2s² 2p⁶, and O 1s² 2s² 2p⁴ becomes O²⁻ 1s² 2s² 2p⁶.

The attraction acts equally in every direction, so each ion attracts as many oppositely charged ions as can pack around it. The result is a giant ionic lattice, a regular repeating arrangement of a huge number of ions. In sodium chloride every Na⁺ is surrounded by six Cl⁻ ions and every Cl⁻ by six Na⁺ ions. The formula NaCl gives only the ratio of ions; there are no molecules.

What makes ionic bonding stronger

The attraction between two ions is greater when their charges are larger and when their centres are closer together, which happens when the ions are smaller. Stronger attraction means more energy is needed to separate the ions, so the melting temperature is higher.

Ionic radii

• Positive ions are smaller than their atoms. The atom loses its whole outer shell, and the same number of protons now pulls on fewer electrons, so they are held closer.
• Negative ions are larger than their atoms. The extra electrons repel the others, and the same nuclear charge is shared among more electrons, so the attraction per electron is less.
• Down a group, ionic radius increases, because each ion has one more shell of electrons.
• Isoelectronic ions (with the same number of electrons) get smaller as the number of protons increases.

Evidence that ions exist

Electron density maps. When X-rays pass through a crystal they are scattered by the electrons, and the pattern can be turned into a map of electron density, drawn with contour lines like a map of hills. For sodium chloride the map shows separate, almost spherical regions of electron density, arranged regularly, with the density falling almost to zero between neighbouring ions. That is what separate ions would look like. The chloride ions appear larger than the sodium ions. The map does not show a sharp edge to each ion, so ionic radii cannot be read exactly from it.

Migration of ions. If ions exist, a potential difference should move them: positive ions towards the negative electrode, negative ions towards the positive one. Green copper(II) chromate(VI) solution placed under a colourless electrolyte in a U-tube shows this. After some time a blue colour (Cu²⁺ ions) rises towards the negative electrode and a yellow colour (CrO₄²⁻ ions) towards the positive electrode. Similarly, a crystal of potassium manganate(VII) on filter paper moistened with an electrolyte, connected to a low-voltage d.c. supply, gives a purple streak that moves towards the positive electrode, because the coloured particle is the negative MnO₄⁻ ion.

Properties of ionic compounds

• High melting and boiling temperatures: melting means overcoming strong electrostatic attractions between many ions throughout the giant lattice, which needs a lot of energy.
• Conduct electricity when molten or dissolved, but not when solid: the ions carry the charge. In the solid they are held in fixed positions; when the lattice breaks down on melting or dissolving, they are free to move.
• Often soluble in water: polar water molecules are attracted to the ions and pull them out of the lattice (the details come later in the course).
• Hard but brittle: a sharp blow can shift one layer of ions relative to the next. Ions of the same charge then line up next to each other, they repel, and the crystal splits apart.

2. Covalent bonding

Covalent bonds form between non-metal atoms. Electron density maps of molecules such as H₂ show a high concentration of electron density between the two nuclei. That negative charge is attracted to both positive nuclei, holding them together, and these attractions outweigh the repulsion between the two nuclei. Each atom usually gains a share of enough electrons to fill its outer shell.

Atoms can share more than one pair. Oxygen, O=O, has a double bond (two shared pairs), and nitrogen, N≡N, has a triple bond (three shared pairs). The bond is strong: it takes a lot of energy to break, which is why giant covalent structures such as diamond have such high melting temperatures. The forces between separate molecules are a different and much weaker thing.

Not every atom ends up with eight outer electrons. Boron in BF₃ and BCl₃ has only six, which is why these molecules readily accept a lone pair. Elements in period 3 and below can share more than eight: phosphorus in PCl₅ has ten outer electrons and sulfur in SF₆ has twelve.

Bond length and bond strength

The bond length is the distance between the two nuclei. At that distance the attraction of both nuclei for the shared electrons exactly balances the repulsion between the two nuclei (and between the inner electrons). A double or triple bond puts more electron density between the nuclei, so the attraction is greater: the nuclei are pulled closer together and the bond is harder to break.

Dative covalent (coordinate) bonds

In a dative covalent bond both electrons of the shared pair come from the same atom. It needs one atom with a lone pair and another with an empty space in its outer shell. In a displayed formula it is drawn as an arrow pointing from the atom that donates the lone pair to the atom that accepts it. Once formed, a dative bond is identical to any other covalent bond.

• Ammonium ion, NH₄⁺: NH₃ + H⁺ → NH₄⁺. The nitrogen's lone pair bonds to an H⁺ ion, which has no electrons. All four N–H bonds are then the same, and the ion is tetrahedral.
• Oxonium ion, H₃O⁺: H₂O + H⁺ → H₃O⁺. One of oxygen's two lone pairs forms the bond, so one lone pair is left.
• NH₃BF₃: the lone pair on the nitrogen of NH₃ fills the gap in the outer shell of boron in BF₃, written H₃N→BF₃.
• Aluminium chloride, Al₂Cl₆: at lower temperatures, two AlCl₃ molecules join. A lone pair on a chlorine of each molecule forms a dative bond to the aluminium of the other, giving each aluminium four bonding pairs.

3. Shapes of molecules and ions

The shape of a molecule is set by the electron pairs in the outer shell of its central atom. Pairs of electrons are regions of negative charge, so they repel one another and move as far apart as possible. This is electron-pair repulsion theory. Count the pairs and you know their arrangement; then the positions of the atoms give the shape.

• Lone pairs repel more strongly than bonding pairs. A lone pair is held by only one nucleus, so it sits closer to the central atom and spreads out more. The order of repulsion is lone pair–lone pair > lone pair–bonding pair > bonding pair–bonding pair.
• Each lone pair squeezes the bond angles by roughly 2.5°: 109.5° in CH₄, 107° in NH₃ and 104.5° in H₂O.
• A double or triple bond counts as one bonding region, because its electrons all point in the same direction. CO₂ has two bonding regions around carbon, so it is linear; each carbon in ethene, C₂H₄, has three, so the H–C–H angles are 120°.
• The shape is named from the atoms only, even though the lone pairs help to decide it. Ammonia has four pairs arranged tetrahedrally, but its shape is pyramidal.

With five or six pairs, lone pairs take the positions that keep them furthest from other pairs. That gives a few more shapes you can work out without needing to learn their names: ClF₃ (3 bonding, 2 lone) is T-shaped, SF₄ (4 bonding, 1 lone) is a distorted tetrahedron, XeF₂ and I₃⁻ (2 bonding, 3 lone) are linear, and BrF₅ (5 bonding, 1 lone) is square pyramidal.

4. Electronegativity and polarity

Electronegativity is measured on the Pauling scale. Fluorine is the most electronegative element, with a value of 4.0, followed by oxygen, nitrogen and chlorine.

• It increases across a period: the nuclear charge increases while the bonding electrons stay in the same shell with similar shielding, and the atoms get smaller, so the bonding pair is pulled in more strongly.
• It decreases down a group: the bonding pair is further from the nucleus and more shielded by inner shells.

Polar bonds and the bonding continuum

When two identical atoms are bonded, as in H₂ or Cl₂, the electrons are shared equally and the bond is non-polar. When the atoms have different electronegativities, the bonding pair is pulled towards the more electronegative atom. That atom gains a small negative charge, δ−, and the other is left with a small positive charge, δ+. The bond is a polar covalent bond and has a permanent dipole: in hydrogen chloride, H^δ+–Cl^δ−.

Pure covalent and pure ionic bonding are the two ends of a continuous scale. The bigger the difference in electronegativity, the more uneven the sharing. A small difference gives a slightly polar bond, and a very large difference (more than about 1.7) means the electron is effectively transferred and the bonding is mainly ionic. Most real compounds lie somewhere in between. Ionic compounds can also have some covalent character: a small, highly charged positive ion can distort, or polarise, the electron cloud of a large negative ion, pulling electron density into the space between them.

Polar and non-polar molecules

A molecule with polar bonds is not necessarily polar overall. Each bond dipole has a direction, and if the molecule is symmetrical the dipoles cancel out. CO₂ is linear, so its two C=O dipoles point in opposite directions and cancel; BF₃ (trigonal planar) and CCl₄ (tetrahedral) cancel in the same way, so these molecules are non-polar. Water is bent, so its two O–H dipoles add up to give a polar molecule with δ− on the oxygen side. Replacing one Cl in CCl₄ by H gives CHCl₃, which is no longer symmetrical and is polar.

A simple test: run a thin jet of a liquid from a burette past an electrically charged rod. A jet of a polar liquid such as water bends towards the rod, because the molecules turn so that their oppositely charged ends face it. This happens whether the rod is positive or negative. A jet of a non-polar liquid such as cyclohexane or tetrachloromethane is not deflected. The more polar the molecules, the greater the deflection.

5. Metallic bonding

In a solid metal, each atom releases its outer electrons into a shared "sea" that is spread through the whole structure. These electrons are delocalised: they do not belong to any one atom. What is left is a regular, closely packed lattice of positive ions held together by their attraction to the delocalised electrons. This is a giant metallic lattice.

Metallic bonding is stronger when there are more delocalised electrons per atom, when the ions have a greater charge, and when the ions are smaller. Magnesium has a higher melting temperature (650 °C) than sodium (98 °C): each Mg atom gives two electrons to the sea and forms a smaller 2+ ion, while Na gives one electron and forms a larger 1+ ion.

• High melting temperatures: lots of energy is needed to overcome the strong attraction between the ions and the delocalised electrons.
• Good electrical conductors, as solids and liquids: the delocalised electrons can move through the structure when a potential difference is applied.
• Good thermal conductors: the mobile electrons carry kinetic energy quickly through the metal.
• Malleable and ductile: all the ions are the same, so when a force makes layers of ions slide over each other, the surroundings of every ion stay the same and the bonding is not broken. The metal changes shape instead of shattering.

6. Giant covalent structures

In a giant covalent (macromolecular) lattice, every atom is joined to its neighbours by strong covalent bonds in a network that runs through the whole crystal. To melt it, many of these bonds must be broken, so the melting temperatures are very high. These substances are insoluble in water, because water molecules cannot provide the energy needed to break covalent bonds.

7. Simple molecular structures

Most covalent substances are made of separate small molecules. The bonds inside each molecule are strong, but the forces between molecules (London forces, permanent dipole forces and hydrogen bonds, covered in the topic on intermolecular forces) are weak. Melting or boiling only overcomes these intermolecular forces, not the covalent bonds, so simple molecular substances have low melting and boiling temperatures and many are gases or liquids at room temperature. They do not conduct electricity, because the molecules are neutral and no electrons are free to move.

• Iodine, I₂: a solid in which I₂ molecules are packed in a regular lattice, held only by London forces. It melts at 114 °C and gives a purple vapour when warmed.
• Ice, H₂O: each water molecule forms two covalent O–H bonds and takes part in hydrogen bonds to its neighbours, with each oxygen surrounded tetrahedrally by four hydrogens. This open arrangement holds the molecules further apart than in liquid water, which is why ice is less dense than water and floats.

8. Structure, bonding and properties

Quick checks

Interactive diagrams