How Bonding Type Predicts Melting Point

Here is a list of melting points that spans more than 3,000 degrees: methane −182 °C, water 0 °C, iodine 114 °C, sodium chloride 801 °C, silicon 1,414 °C, silicon dioxide about 1,710 °C, magnesium oxide about 2,852 °C, tungsten 3,422 °C. Could you sort those without the numbers, just from what the substances are made of? With one idea, you can come close.

The idea is this: melting point measures what has to be overcome to let particles slide past each other. A solid melts when its particles gain enough energy to escape their fixed positions. If only weak forces hold them there, a little heat is enough. If strong bonds hold them there, you need a lot. And what holds the particles in place depends on the type of bonding.

An analogy: getting a crowd moving

Imagine a hall full of people you want to get moving about. If they are strangers who just happen to be standing near each other, a gentle nudge sets them wandering. If they are holding hands in pairs, the pairs still wander easily, because nothing links one pair to the next. But if everybody is holding hands with every neighbour in a single web across the whole hall, nobody can move until you break many of those grips.

That picture maps onto the main structure types:

• Simple molecules are the pairs. Strong bonds inside each molecule, weak links between molecules. Easy to set moving.
• Giant structures (ionic, covalent network, metallic) are the single web. Strong bonds reach through the entire solid. Hard to set moving.

The crucial detail is that melting a simple molecular solid does not break its covalent bonds. It only overcomes the weak forces between molecules.

Simple molecular substances: low melting points

In methane, iodine, water and sulfur, each molecule is held together by strong covalent bonds, but molecules are attracted to one another only by intermolecular forces. These are much weaker, so melting points are low:

Substance	Main forces between molecules	Melting point
Neon (Ne atoms)	Weak dispersion	−249 °C
Argon (Ar atoms)	Dispersion	−189 °C
Methane, CH₄	Dispersion	−182 °C
Water, H₂O	Hydrogen bonding	0 °C
Iodine, I₂	Strong dispersion (many electrons)	114 °C
Sulfur, S₈	Dispersion	115 °C

Within this group, two things raise the melting point:

1. More electrons, bigger molecules. Dispersion forces grow with the number of electrons, so iodine (106 electrons per molecule) and S₈ melt far higher than methane (10 electrons).
2. Hydrogen bonding. Water molecules are small and light, yet water melts 182 degrees above methane, which has almost the same molar mass. That gap is the effect of hydrogen bonding.

Still, even the “high” values here are modest. Nothing in the simple molecular class of everyday substances comes close to ionic or network solids.

Ionic compounds: high melting points

In sodium chloride there are no molecules at all. Every Na⁺ is surrounded by six Cl⁻ and every Cl⁻ by six Na⁺, throughout the crystal. Melting means letting ions move while strong electrostatic attractions pull on them in every direction. Hence sodium chloride melts at 801 °C.

The strength of an ionic lattice depends mainly on two things, both captured by lattice energy:

• Ionic charge. Doubling the charge on both ions roughly quadruples the attraction.
• Ionic size. Smaller ions sit closer together and attract more strongly.

Magnesium oxide shows both effects at once. Mg²⁺ and O²⁻ carry double charges and are small. The result is a melting point of about 2,852 °C, more than three times that of NaCl. That is why magnesium oxide is used to line furnaces.

Giant covalent structures: very high melting points

In diamond, silicon and silicon dioxide, atoms are joined by covalent bonds into one continuous network. To melt the solid, covalent bonds themselves must break, and there are a vast number of them.

Substance	Melting point
Silicon, Si	1,414 °C
Silicon dioxide (quartz), SiO₂	about 1,710 °C
Diamond and graphite, C	do not melt at normal pressure; they turn straight to vapour at several thousand degrees

Carbon is the extreme. At atmospheric pressure it does not form a liquid at all; it sublimes. Liquid carbon exists only under high pressure. Compare silicon with carbon: both form the same diamond-type network, but Si–Si bonds are longer and weaker than C–C bonds, so silicon gives way at a much lower temperature. More on these networks in giant covalent structures.

Graphite is a useful warning. Between its layers there are only weak forces, which is why it is soft and slippery. But melting requires breaking the strong covalent bonds within the layers, so graphite’s resistance to heat is as extreme as diamond’s. Softness and melting point measure different things.

Metals: from liquid at room temperature to over 3,400 °C

Metals are giant structures of positive ions held together by a sea of delocalised electrons, as described in metallic bonding. Their melting points spread across a huge range, which makes them the least predictable group:

Metal	Melting point
Mercury, Hg	−39 °C
Caesium, Cs	28 °C
Gallium, Ga	30 °C
Sodium, Na	98 °C
Magnesium, Mg	650 °C
Aluminium, Al	660 °C
Iron, Fe	1,538 °C
Tungsten, W	3,422 °C

Some trends do hold. Across period 3, sodium (one delocalised electron per atom) melts at 98 °C, magnesium (two) at 650 °C and aluminium (three) at 660 °C. More delocalised electrons and a smaller, more highly charged ion usually mean stronger bonding, though the rise from magnesium to aluminium is small, a reminder that crystal structure matters too. Down group 1, larger atoms give weaker bonding: caesium melts in a warm hand.

The transition metals in the middle of the d-block, such as tungsten, reach the highest values because their d electrons join in the bonding. Tungsten has the highest melting point of any metal; see the element with the highest melting point for the full ranking.

Putting the pattern together

Structure type	What must be overcome	Typical melting points
Simple molecular	Weak intermolecular forces	Below about 300 °C, often below room temperature
Ionic	Electrostatic attraction between ions	Several hundred to nearly 3,000 °C
Giant covalent	Strong covalent bonds throughout	Well above 1,000 °C
Metallic	Metallic bonding	Very wide range, most above 500 °C

As a rule of thumb, a substance that is a gas or liquid at room temperature, or melts on a hot plate, is almost certainly simple molecular (mercury being the famous metallic exception). A substance that survives a Bunsen flame without melting is almost certainly a giant structure.

Common mistakes

“When a molecular substance melts, its covalent bonds break.” This is the most common error in exam answers. When ice melts, every H₂O molecule survives intact. Only the hydrogen bonds between molecules are disrupted. Write “overcome intermolecular forces”, never “break covalent bonds”, for simple molecules.

“Covalent means low melting point.” Only for simple molecules. Diamond and quartz are covalent and among the most heat-resistant materials known. Always distinguish simple molecular from giant covalent; see simple molecular vs giant covalent.

“Ionic bonds are stronger than covalent bonds, so ionic compounds melt higher.” Silicon dioxide melts above sodium chloride, and diamond will not melt at all at normal pressure. The comparison that matters is the structure, not a ranking of bond types.

“Metals always have high melting points.” Mercury, caesium and gallium are all liquid on a warm day or in the hand.

Mixing up melting and boiling. Both depend on the same forces, but boiling requires particles to escape completely. Quote the right one.

Key takeaways

• Melting point reflects the strength of whatever holds particles in fixed positions.
• Simple molecular solids melt low because only weak intermolecular forces are overcome; size and hydrogen bonding raise the value.
• Ionic solids melt high; higher charges and smaller ions raise it further (NaCl 801 °C, MgO about 2,852 °C).
• Giant covalent solids melt very high because covalent bonds must break (Si 1,414 °C, SiO₂ about 1,710 °C; carbon sublimes).
• Metals span the widest range, from mercury at −39 °C to tungsten at 3,422 °C.
• The same structure types explain solubility too: see how bonding type predicts solubility.