1The idea in one paragraph

Real bonds are rarely purely ionic, purely covalent or purely metallic. A bond between two atoms sits somewhere on a continuum, and two numbers tell you where: how different the two atoms' electronegativities are, and how high they are on average. Plot those two numbers and you land on a bonding triangle with ionic, covalent and metallic at its corners; your position predicts how the substance behaves. Materials scientists use the same thinking in reverse. They mix metals to make alloys harder than either metal alone, link small molecules into long chains to make polymers with the properties they want, and combine different materials into composites that do what no single one can.

2Bonding is a continuum

The three models are useful because they are simple, and simple models are ideals. Consider the bond in hydrogen chloride. The pair is shared, so it is covalent; but chlorine pulls it strongly, so the bond has a large δ+ and δ−, some of the character of an ionic bond. Or consider aluminium chloride: from a metal and a non-metal, so "ionic" by the S2.1 rule of thumb, yet it turns to vapour at about 180 °C, which no ionic lattice does. Most real bonds are a blend.

Bonding is a continuum: most bonds have partly ionic, partly covalent and partly metallic character.

The bonding triangle (a version is in the data booklet) turns that blend into a position. The bottom axis is the average electronegativity of the two elements; the vertical axis is the difference in electronegativity, Δχ. Figure 1 draws it with examples.

The three corners are the three pure models, and the reasoning behind each is short.

• Top: ionic. A large Δχ means one atom takes the electrons almost completely. Caesium fluoride, from the least and the most electronegative elements, is at the apex.
• Bottom left: metallic. Both atoms have low electronegativity, so neither holds its electrons tightly and the electrons are delocalized. Pure metals sit on the base, where Δχ = 0.
• Bottom right: covalent. Both atoms have high electronegativity, so both hold on to electrons, and they share them. Cl₂ and F₂ sit on the base at the right.

In between, the character changes gradually. Moving up from the bottom-right corner, bonds become more polar, then polar covalent, then predominantly ionic. The dashed lines in Figure 1 mark rough boundaries only; above about Δχ = 1.8 bonding is mostly ionic, but nothing special happens at that value.

3Placing a compound on the triangle

The guide asks you to determine the position of a binary compound (two elements) from electronegativity data. It is two sums. Take magnesium oxide, with χ(Mg) = 1.3 and χ(O) = 3.4.

Do the same for aluminium chloride, χ(Al) = 1.6 and χ(Cl) = 3.2: average 2.4, Δχ 1.6. It sits below the ionic zone, in the polar covalent region, which is exactly why it behaves as a covalent substance. For a pure element, Δχ is zero and the average is just its own electronegativity: silicon sits on the base at 1.9, between the metallic and covalent corners, as a metalloid should.

You do not calculate a percentage of ionic character; the guide says so. You read the region off the triangle and say what kind of bonding dominates.

4From position to properties

Once you know which corner a substance is nearest, the properties follow from the model for that corner.

The chlorides of period 3 show the continuum at work. Following them from sodium to phosphorus, the average electronegativity rises and Δχ falls, so the bonding moves steadily from ionic towards covalent, and the melting points collapse.

The triangle has limits, and the guide wants you to know them. It uses only two numbers per compound, so it cannot tell a covalent network from covalent molecules: SiO₂ (average 2.65, Δχ 1.5) sits near AlCl₃ and H₂O, yet melts above 1600 °C because it is a network. It says nothing about structure, only about the bond. Discrete labels such as "ionic" and "covalent" are useful precisely because they are simplifications, and the continuum is the more honest description.

Composites show the same design thinking at a larger scale. A composite is made of two or more materials with different properties, combined so that each makes up for the other's weakness. Concrete is mostly ionic and covalent-network solids: very strong when squeezed, but it cracks when pulled or bent. Steel is metallic: strong when pulled. Set steel bars inside concrete and you have reinforced concrete, which resists squeezing and bending and holds up bridges and buildings that neither material could hold alone.

5Alloys

An alloy is a mixture of a metal with other metals or with non-metals. Some common examples:

You do not have to learn specific alloys; you do have to explain their properties.

Why alloys form at all. Metallic bonding is non-directional: the cations are held by the electron sea around them, not by bonds to particular neighbours. So a cation of a different element, or a small non-metal atom, can take a place in the lattice or sit in the gaps between cations without breaking the bonding. The alloy still has delocalized electrons, so it still conducts and is still metallic.

Why alloys are mixtures, not compounds. The composition can vary continuously: brass can be 60% copper or 70% copper. There is no fixed ratio and no new chemical formula, so an alloy is a mixture, even though it is held together by metallic bonding.

Why alloys are harder and stronger. In a pure metal all the cations are the same size, so the layers are regular and slide over each other easily, which is what makes a pure metal soft and malleable. Atoms of a different size distort the layers, as Figure 2 shows. The distorted layers can no longer slide past each other easily, so a larger force is needed to change the metal's shape: the alloy is harder and stronger, and less malleable.

Other properties can change too. Chromium in stainless steel forms a thin oxide layer on the surface that stops the iron beneath from rusting. The bonding is the same idea; the design is choosing what to add.

6Polymers and plastics

A polymer is a very large molecule, a macromolecule, made of many small repeating subunits called monomers, joined by covalent bonds. A chain may contain thousands of monomer units.

Polymers are everywhere in living things. Natural polymers include proteins (from amino acids), starch and cellulose (from glucose), DNA (from nucleotides) and natural rubber. Synthetic polymers include poly(ethene), PVC, PTFE, polystyrene, nylon and PET. Plastics are synthetic polymers that can be moulded into shape.

Figure 3 shows the structure behind the properties of plastics: long chains, with strong covalent bonds along each one and weaker intermolecular forces between chains.

Properties can be tuned by changing the structure. Longer chains mean more intermolecular forces per chain, so a higher melting point and a stronger plastic. Chains with few branches pack closely, so their intermolecular forces are stronger and the plastic is denser and more rigid; branched chains cannot pack closely, giving a softer, lower-density plastic. Polar side groups add dipole–dipole forces: the C–Cl bonds in PVC make it stiffer than poly(ethene). Plastics whose chains contain ester or amide links can be broken down by hydrolysis, which is one route to biodegradable plastics (Structure 3.2).

7Addition polymers

An addition polymer forms when monomers containing a C=C double bond join together. In each monomer the double bond breaks (strictly, its π bond does), and the two carbon atoms each use the freed electron to bond to a neighbouring monomer. Nothing else is lost, so the polymer is the only product and the atom economy is 100%. Figure 4 shows ethene becoming poly(ethene).

The polymer is too long to draw, so you draw its repeating unit: the smallest section that repeats along the chain, in square brackets with a subscript n and a bond sticking out through each bracket. For an addition polymer made from one monomer, the repeating unit is always the two carbons of the old C=C with everything attached to them, now joined by a single bond.

The same rule works for any alkene. Figure 5 shows three more.

The monomer structures are given in the exam, or you deduce them, so the skill is the drawing. Three checks catch almost every mistake.

1. The main chain contains only the two carbon atoms of the C=C. Side groups such as CH₃ or Cl hang off the chain; they are never in it.
2. There is no double bond in the repeating unit.
3. There are bonds through both brackets, showing the chain continues.

To go backwards, from a section of polymer to its monomer, find the two-carbon unit that repeats, draw it once, and put the C=C back. A chain reading –CH₂–CH(CH₃)–CH₂–CH(CH₃)– has the repeating unit –CH₂–CH(CH₃)–, so the monomer is CH₂=CH–CH₃, propene.