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Put four white or colourless solids on a bench: a spoonful of table salt, a lump of sugar, a pinch of sand and a strip of aluminium foil. At a glance they are all just “stuff that holds its shape”. Heat them, hit them with a hammer, try to pass a current through them or drop them in water, though, and they behave in four completely different ways. Salt survives the Bunsen flame, sugar melts and caramelises, sand barely notices, and the foil bends instead of shattering.
Those differences are not random. Almost every crystalline solid you will meet in a chemistry course falls into one of four families, and the family is decided by two questions: what particles sit at the points of the structure, and what holds those particles together. Answer those two questions and you can predict a surprising amount about a substance before you ever touch it.
The big comparison table
| Ionic | Molecular (simple covalent) | Covalent network (giant covalent) | Metallic | |
|---|---|---|---|---|
| Particles in the structure | Positive and negative ions | Separate small molecules | Atoms | Positive metal ions in a “sea” of delocalised electrons |
| What holds it together | Electrostatic attraction between oppositely charged ions, in every direction | Weak intermolecular forces between molecules (the bonds inside each molecule are strong) | Strong covalent bonds running through the whole structure | Attraction between the metal ions and the delocalised electrons |
| Typical elements involved | A metal plus a non-metal | Non-metals only | Mainly carbon, silicon, and some compounds such as SiO₂ | Metals only |
| Melting point | High (NaCl 801 °C) | Low (ice 0 °C, iodine about 114 °C) | Very high (silicon about 1,414 °C, quartz around 1,700 °C) | Wide range, usually high (mercury −39 °C, sodium 98 °C, tungsten about 3,422 °C) |
| Conducts as a solid? | No | No | No (graphite is the famous exception) | Yes |
| Conducts when molten or dissolved? | Yes | No (a few react with water to make ions) | Does not melt or dissolve easily | Yes, as a liquid |
| Solubility in water | Often soluble | Varies: polar ones may dissolve, non-polar ones usually do not | Insoluble | Insoluble (some react) |
| Mechanical behaviour | Hard but brittle | Soft, often crumbly | Very hard (diamond) or slippery layers (graphite) | Malleable and ductile |
| Everyday examples | Table salt, magnesium oxide, calcium carbonate | Ice, sugar, dry ice, sulfur, wax | Diamond, graphite, silicon, sand (SiO₂) | Copper, iron, aluminium, gold |
The melting points for elements above come from this site’s element data. The table is the whole story in miniature; the rest of this article explains why each column looks the way it does.
Ionic solids: a charged grid
An ionic solid is built from ions, not molecules. In sodium chloride, every Na⁺ is surrounded by six Cl⁻ ions and every Cl⁻ by six Na⁺ ions, and the pattern repeats in all three directions. There is no such thing as “one molecule of NaCl”; the formula just tells you the ions are present in a 1 : 1 ratio. You can explore how the lattice is arranged in ionic lattice structure.
Why the high melting point? To melt the solid you must let the ions slide past each other, which means partly overcoming the attraction of each ion to all its oppositely charged neighbours. That takes a lot of energy. Ions with bigger charges attract more strongly, which is why magnesium oxide (Mg²⁺ and O²⁻) melts at well over 2,000 °C while sodium chloride melts at 801 °C.
Why no conduction as a solid? Conduction needs charged particles that can move. The ions are charged but locked in place. Melt the solid or dissolve it in water and the ions are free to drift, so the liquid or solution conducts.
Why brittle? Tap an ionic crystal hard enough and one layer shifts by half an ion. Suddenly positive ions sit next to positive ions and negative next to negative. The layers repel and the crystal cleaves along a flat face.
Molecular solids: strong inside, weak between
Molecular solids are made of discrete molecules such as H₂O, CO₂, I₂, S₈ or sucrose. Each molecule is held together by strong covalent bonds, but the molecules are attracted to each other only by intermolecular forces: London dispersion forces, permanent dipole attractions and, where possible, hydrogen bonds (see intermolecular forces).
Why the low melting point? Melting a molecular solid does not break any covalent bonds. When ice melts, the water molecules are still H₂O; only the hydrogen bonds between them loosen. Because intermolecular forces are much weaker than ionic or covalent bonds, a small amount of heat is enough. Dry ice (solid CO₂) is so weakly held that at normal pressure it turns straight into gas at about −78 °C.
Why no conduction? Molecules are neutral and there are no free electrons. Even when melted, there is nothing charged to carry a current.
Trend worth knowing: bigger molecules with more electrons have stronger dispersion forces, so melting points rise with size. That is why fluorine and chlorine are gases at room temperature, bromine is a liquid and iodine is a solid.
This is the family where students most often go wrong, by saying that “covalent bonds are weak, so the melting point is low”. The covalent bonds are not the issue at all. The weak link is between molecules.
Covalent network solids: one giant molecule
In a covalent network (also called giant covalent or macromolecular) solid, every atom is joined to its neighbours by covalent bonds, and those bonds continue throughout the whole crystal. A diamond is, in effect, one enormous molecule of carbon atoms, each bonded to four others in a tetrahedral arrangement. Silicon adopts the same structure, and silicon dioxide builds a network in which every silicon is bonded to four oxygens and every oxygen bridges two silicons.
Why the very high melting point? There are no weak links. To melt the solid you must break strong covalent bonds, lots of them. Silicon melts at about 1,414 °C, and carbon in the form of diamond or graphite survives far higher temperatures still.
Why so hard (usually)? Pushing atoms out of place means bending or breaking covalent bonds in every direction, so diamond is the hardest natural material.
The graphite exception. Graphite is also a covalent network, but each carbon bonds to only three others, forming flat hexagonal sheets. The fourth electron from each atom is delocalised across the sheet, so graphite conducts electricity along the layers. The sheets are held to each other only by weak forces, so they slide, which is why graphite is soft and used in pencils and lubricants. Its melting point stays very high because melting would still require breaking the strong bonds within each sheet. For a closer look at the difference, see simple molecular vs giant covalent.
Metallic solids: ions in an electron sea
In a metal, each atom gives up its outer electrons to a shared pool. What remains is a regular array of positive ions surrounded by delocalised electrons that belong to the whole lump rather than to any single atom. The attraction between the ions and this electron sea is the metallic bond (metallic bonding has the full picture).
Why do metals conduct? The delocalised electrons can move through the structure. Apply a voltage and they drift towards the positive terminal. The same mobile electrons also carry heat efficiently, which is why a metal spoon in hot tea warms up quickly.
Why malleable? When a force pushes one layer of ions past another, the electron sea simply flows with them. There is no sudden repulsion as in an ionic crystal, so the metal bends or stretches rather than shattering.
Why such a wide range of melting points? The strength of metallic bonding depends on how many electrons each atom contributes and how small and highly charged the ions are. Sodium gives one electron from a fairly large atom and melts at about 98 °C. Transition metals such as tungsten contribute more electrons, and tungsten does not melt until about 3,422 °C. Mercury is liquid at room temperature, a reminder that “metallic solid” describes the bonding type, not a guaranteed solid state.
How to identify the type of solid
When an exam question hands you an unknown substance and a list of properties, work through these questions in order:
- Does it conduct as a solid? If yes, it is almost certainly metallic. (Check the special case of graphite.)
- Does it conduct when molten or dissolved, but not as a solid? That points to ionic.
- Is the melting point low (below a few hundred °C) and does it not conduct in any state? It is molecular.
- Is the melting point very high, with no conduction in any state and no solubility? It is a covalent network.
You can also reason from the formula. A metal combined with a non-metal usually forms an ionic solid. Non-metals alone usually form molecular solids, unless the element or compound is one of the few network formers (C, Si, SiO₂, SiC, boron). Pure metals and alloys are metallic.
When the neat boxes break down
Real substances do not always sit in one box. Aluminium chloride contains a metal and a non-metal, yet it sublimes at a fairly low temperature and has a lot of covalent character. Graphite conducts like a metal while being a covalent network. Some ionic compounds, such as calcium carbonate, are barely soluble in water. And many substances are not crystalline at all: glass and most plastics are amorphous solids, with no long-range order, so they soften gradually instead of melting at one sharp temperature.
The four types are best seen as reference points on a continuum of bonding, not rigid categories. The article on the bonding continuum develops that idea further.
When to use each model
- Use the ionic model when explaining high melting points, brittleness and conduction only in the molten or dissolved state for metal–non-metal compounds.
- Use the molecular model whenever a question mentions boiling points of small covalent substances, and always point to intermolecular forces.
- Use the covalent network model for diamond, graphite, silicon and silica, and stress that melting means breaking covalent bonds.
- Use the metallic model for conductivity, malleability and ductility of metals and alloys.
Key takeaways
- The four types are defined by their particles (ions, molecules, atoms, or metal ions with delocalised electrons) and the forces holding those particles together.
- Ionic solids have high melting points, are brittle, and conduct only when molten or dissolved.
- Molecular solids have low melting points because only weak intermolecular forces are overcome on melting; the covalent bonds inside the molecules stay intact.
- Covalent network solids have very high melting points because melting breaks strong covalent bonds; graphite is the conducting exception.
- Metallic solids conduct in the solid state and are malleable because their delocalised electrons can move and flow around shifting layers of ions.
- Treat the categories as guides: aluminium chloride, graphite and amorphous solids all show that real materials can blur the boundaries.
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