Contents: 7 sections
Syllabus points
- Describe ionic, covalent, dative covalent and metallic bonding.
- Use electronegativity to explain bond polarity.
- Predict the shapes of simple molecules and ions, and their bond angles, using electron pair repulsion.
- Describe hydrogen bonding, permanent dipole and van der Waals forces, and explain their effect on physical properties.
- Explain the anomalous properties of water in terms of hydrogen bonding.
The three strong bonds
| Bond | Formed between | Held together by |
|---|---|---|
| Ionic | Metal and non-metal | Electrostatic attraction between oppositely charged ions |
| Covalent | Non-metal and non-metal | Shared pair of electrons attracted to both nuclei |
| Metallic | Metal atoms | Attraction between positive ions and delocalised electrons |
A dative covalent (or co-ordinate) bond is a covalent bond in which both electrons come from the same atom. Once formed it is identical to any other covalent bond. The standard examples are NH₄⁺, where the lone pair on nitrogen bonds to H⁺, and Al₂Cl₆.
Electronegativity and polarity
Electronegativity is the ability of an atom to attract the bonding electrons in a covalent bond.
It increases across a period, because nuclear charge rises while the atomic radius falls, and decreases down a group, because the bonding pair is further from the nucleus and better shielded. The most electronegative elements are fluorine, oxygen, nitrogen and chlorine.
When two different atoms are bonded, the electrons sit closer to the more electronegative one. This produces a permanent dipole, written δ+ and δ-.
A molecule with polar bonds is not necessarily a polar molecule. CO₂ is non-polar even though each C=O bond is polar, because the molecule is linear and the two dipoles point in opposite directions and cancel. H₂O is polar because it is bent, so the dipoles do not cancel. Symmetry decides it.
Shapes of molecules
The rule is short: electron pairs repel and arrange themselves as far apart as possible, and lone pairs repel more strongly than bonding pairs.
The order of repulsion is lone/lone > lone/bonding > bonding/bonding. Each lone pair squeezes the bond angle by roughly 2.5°.
| Electron pairs | Lone pairs | Shape | Bond angle |
|---|---|---|---|
| 2 | 0 | Linear | 180° |
| 3 | 0 | Trigonal planar | 120° |
| 4 | 0 | Tetrahedral | 109.5° |
| 4 | 1 | Pyramidal | 107° |
| 4 | 2 | Bent | 104.5° |
| 5 | 0 | Trigonal bipyramidal | 120° and 90° |
| 6 | 0 | Octahedral | 90° |
The ammonia and water sequence is worth learning as a set, because it shows the rule working. Methane has four bonding pairs and no lone pairs, so 109.5°. Ammonia has three bonding pairs and one lone pair, so 107°. Water has two bonding pairs and two lone pairs, so 104.5°.
To work out a shape: count the electron pairs around the central atom, decide how many are lone pairs, then read off the arrangement.
Intermolecular forces
These are far weaker than covalent bonds and act between molecules. They decide melting point, boiling point and solubility, not the strength of the substance's own bonds.
Van der Waals forces (induced dipole to induced dipole) exist between all molecules. Electrons move at random, so at any instant a molecule has a temporary dipole, which induces one in its neighbour. They get stronger as the number of electrons increases, which is why boiling points rise down Group 17 and along the alkanes.
Permanent dipole forces act between polar molecules and are stronger than van der Waals forces of comparable size.
Hydrogen bonding is the strongest of the three. It occurs when hydrogen is bonded to nitrogen, oxygen or fluorine and is attracted to a lone pair on the N, O or F of another molecule. Those three elements are highly electronegative and small, so the δ+ on hydrogen is large and the approach is close.
Why water is anomalous
Water boils at 100 °C, far above what its Mᵣ of 18 would suggest, because breaking apart the liquid means breaking many hydrogen bonds.
Ice is less dense than water, which is unusual and biologically important. In ice each water molecule forms four hydrogen bonds in a rigid tetrahedral lattice that holds the molecules further apart than in the liquid. Ice therefore floats, and ponds freeze from the top down.
Water also has a high surface tension and a high specific heat capacity for the same reason.
Structure and properties
| Structure | Example | Melting point | Conducts? |
|---|---|---|---|
| Giant ionic | NaCl | High | Only when molten or aqueous |
| Giant covalent | Diamond, SiO₂ | Very high | No (graphite is the exception) |
| Simple molecular | I₂, CO₂ | Low | No |
| Giant metallic | Mg | High | Yes, solid or liquid |
A simple molecular substance melts at a low temperature because only the intermolecular forces are broken, not the covalent bonds inside the molecules. Saying that the covalent bonds break when iodine melts is a serious error, and a common one.
Graphite conducts because each carbon uses only three of its four outer electrons in bonding, leaving one delocalised electron per atom free to move.
Common mistakes
- Saying covalent bonds break when a simple molecular solid melts. Only the intermolecular forces break.
- Calling CO₂ polar. The bonds are polar; the linear molecule is not.
- Saying hydrogen bonding occurs whenever hydrogen is present. It needs H bonded to N, O or F.
- Forgetting that lone pairs count when working out a shape, or forgetting they repel more.
- Writing the ammonia bond angle as 109.5°. One lone pair reduces it to 107°.
- Describing metallic bonding as attraction between atoms. It is between positive ions and delocalised electrons.