Simple Molecular vs Giant Covalent Substances

Water, carbon dioxide and iodine are covalent substances. So are diamond and sand. Yet water boils at 100 °C and carbon dioxide turns to gas at −78 °C, while diamond survives temperatures above 3,500 °C. All five are held together by covalent bonds, so why the enormous difference? The answer is structure: whether the covalent bonds make small separate molecules or a giant network. This comparison sets out the two types side by side.

The key difference in one sentence

• In a simple molecular substance, atoms are joined by covalent bonds into small, separate molecules, which are attracted to each other only by weak intermolecular forces.
• In a giant covalent substance, covalent bonds link all the atoms into one continuous network, with no separate molecules.

Side by side

Feature	Simple molecular	Giant covalent
Particles	Separate molecules	One continuous network of atoms
Bonds within	Strong covalent bonds (inside each molecule)	Strong covalent bonds (throughout)
Forces between molecules	Weak intermolecular forces	None — there are no separate molecules
What must be overcome to melt/boil	Weak intermolecular forces	Many strong covalent bonds
Melting and boiling points	Low	Very high
State at room temperature	Often gas or liquid; some low-melting solids	Solid
Electrical conductivity	Doesn’t conduct (no charged particles free to move)	Usually doesn’t conduct — graphite and graphene are exceptions
Hardness (as solids)	Soft, often brittle	Usually very hard (graphite is soft)
Solubility in water	Varies — polar molecules may dissolve; non-polar don’t	Insoluble
Formula represents	A molecule (e.g. H₂O)	A ratio (e.g. SiO₂) or an element (C)
Examples	H₂, O₂, N₂, H₂O, CO₂, CH₄, NH₃, I₂, S₈, sugar, wax	Diamond, graphite, graphene, silicon, silicon dioxide, silicon carbide

What actually happens when they melt

This is the heart of the comparison, and the most common exam question.

Simple molecular: molecules separate

When ice melts or water boils, the O–H covalent bonds inside each water molecule don’t break. The molecules simply move apart, overcoming the intermolecular forces (here, hydrogen bonds) between them (see intermolecular forces).

Compare the energies for water:

• Boiling (overcoming intermolecular forces): about 41 kJ mol⁻¹.
• Breaking both O–H bonds in a mole of water molecules: about 2 × 463 = 926 kJ mol⁻¹.

The molecules survive boiling intact; steam is still H₂O molecules.

Giant covalent: bonds break

In diamond, there are no molecules to separate. To melt it, you must break the strong C–C covalent bonds themselves — about 347 kJ mol⁻¹ each, and each carbon atom has four of them (shared with neighbours, so two per atom). That’s why the melting point is so high.

Why simple molecular substances differ from each other

Not all simple molecular substances have the same melting point. The strength of the intermolecular forces depends on:

1. Size of the molecule: bigger molecules, with more electrons, have stronger London (dispersion) forces. Among the halogens: F₂ (gas, boils at −188 °C), Cl₂ (gas, −34 °C), Br₂ (liquid, 59 °C), I₂ (solid, sublimes readily and melts at 114 °C).
2. Polarity: polar molecules add permanent dipole–dipole forces (see how to tell if a molecule is polar).
3. Hydrogen bonding: molecules with O–H, N–H or F–H bonds attract each other strongly (see hydrogen bonding).

So water (small but hydrogen-bonded) is a liquid, while methane (similar size, only London forces) is a gas. Even so, all simple molecular substances melt far below giant covalent ones.

Polymers: in between

Polymers such as poly(ethene) are made of very long molecules — thousands of atoms each — but they’re still separate molecules held together by intermolecular forces. They’re simple molecular in type, but because the molecules are so large, the intermolecular forces add up, and polymers are solids that soften or melt at moderate temperatures (often 100–300 °C).

Exceptions and special cases

• Graphite and graphene are giant covalent but conduct electricity, because each carbon has one delocalised electron (see giant covalent structures).
• Graphite is soft, because its layers slide over each other.
• Iodine sublimes (turns straight from solid to gas) on gentle heating — a sign of weak intermolecular forces.
• Fullerenes (C₆₀) are forms of carbon made of separate molecules, so they’re simple molecular, unlike diamond and graphite (see allotropes of carbon).

How to identify the type from data

A typical exam question gives data and asks you to identify the structure:

Substance	Melting point	Conducts as solid?	Conducts as liquid?	Soluble in water?	Structure
A	−95 °C	No	No	No	Simple molecular
B	1,710 °C	No	No	No	Giant covalent
C	801 °C	No	Yes	Yes	Ionic
D	1,085 °C	Yes	Yes	No	Metallic

Key clues:

• Low melting point + no conduction → simple molecular.
• Very high melting point + no conduction in any state → giant covalent (unless it’s graphite).
• High melting point + conducts only when molten or dissolved → ionic (see properties of ionic compounds).
• Conducts as a solid and a liquid → metallic (or graphite).

Worked exam answer

Question: Silicon dioxide melts at about 1,700 °C. Carbon dioxide turns to gas at −78 °C. Both contain covalent bonds between a group 14 element and oxygen. Explain the difference.

Model answer: Carbon dioxide is a simple molecular substance made of separate O=C=O molecules. Within each molecule the atoms are held by strong covalent double bonds, but between molecules there are only weak intermolecular forces. Very little energy is needed to overcome these forces, so CO₂ changes state at a low temperature, and its covalent bonds are not broken.

Silicon dioxide is a giant covalent structure. Each silicon atom is covalently bonded to four oxygen atoms and each oxygen to two silicon atoms, in a continuous 3D network with no separate molecules. To melt it, many strong Si–O covalent bonds must be broken, which needs a very large amount of energy, so its melting point is very high.

Notice the structure of the answer: name each structure type, say which forces or bonds must be overcome, compare their strength, and link that to the energy needed and the temperature. The same pattern works for comparing iodine with diamond, or methane with graphite.

Common misconceptions

• “Covalent bonds are weak, so covalent substances melt easily.” Covalent bonds are strong; simple molecular substances melt easily because the forces between molecules are weak.
• “Boiling water breaks it into hydrogen and oxygen.” It separates whole H₂O molecules.
• “All giant covalent substances are hard non-conductors.” Graphite is soft and conducts.
• “Big molecules are giant covalent.” Polymers are very large molecules, but still separate molecules.

Key takeaways

• Simple molecular: separate molecules, weak forces between them → low melting points, non-conductors.
• Giant covalent: continuous network of strong covalent bonds → very high melting points, hard, insoluble, usually non-conductors.
• Melting a simple molecular substance overcomes intermolecular forces; melting a giant covalent one breaks covalent bonds.
• Graphite is the key exception: giant covalent but soft and conducting.

For the full picture, see properties of covalent compounds and covalent bonding explained.