Properties of Covalent Compounds

Covalent compounds are everywhere: the water you drink, the oxygen you breathe, the sugar in your tea, the petrol in a car, the plastic in a phone case. Most of them share a set of typical properties — low melting and boiling points, no electrical conductivity, and variable solubility — that are very different from ionic compounds. But a few covalent substances, like diamond and sand, break all these patterns. This article explains the typical properties, why they arise, and where the exceptions come from.

Two kinds of covalent substance

Covalent bonds form when non-metal atoms share pairs of electrons (see covalent bonding explained). Covalent substances come in two very different types:

• Simple molecular: small, separate molecules — H₂O, CO₂, CH₄, I₂, glucose.
• Giant covalent: a continuous network of covalent bonds — diamond, graphite, silicon dioxide.

Most of this article concerns simple molecular substances, which make up the vast majority of covalent compounds. The giant covalent exceptions are covered at the end (see also simple molecular vs giant covalent substances).

Strong bonds, weak forces

The key to understanding simple molecular substances is to distinguish two kinds of attraction:

1. Covalent bonds within each molecule — strong (typically 150–1,000 kJ mol⁻¹).
2. Intermolecular forces between molecules — weak (typically a few to about 40 kJ mol⁻¹) (see intermolecular forces).

Physical changes such as melting, boiling and dissolving only involve the weak forces between molecules. The covalent bonds stay intact.

Property 1: low melting and boiling points

Because only weak intermolecular forces need to be overcome, simple molecular substances melt and boil at low temperatures. Many are gases or liquids at room temperature.

Substance	Formula	Boiling point	State at 20 °C
Hydrogen	H₂	−253 °C	Gas
Methane	CH₄	−162 °C	Gas
Carbon dioxide	CO₂	sublimes at −78 °C	Gas
Ammonia	NH₃	−33 °C	Gas
Water	H₂O	100 °C	Liquid
Ethanol	C₂H₅OH	78 °C	Liquid
Iodine	I₂	184 °C (melts at 114 °C)	Solid
Sucrose	C₁₂H₂₂O₁₁	decomposes above about 186 °C	Solid

What makes some higher than others?

The strength of intermolecular forces depends on:

• Molecule size: bigger molecules have more electrons, giving stronger London (dispersion) forces. In the alkanes, boiling points rise steadily with chain length — methane is a gas, octane a liquid, and long-chain waxes are solids.
• Polarity: polar molecules add permanent dipole–dipole attractions (see how to tell if a molecule is polar).
• Hydrogen bonding: molecules with H attached to N, O or F form especially strong intermolecular attractions (see hydrogen bonding). Water boils at 100 °C, while methane — a molecule of similar size — boils at −162 °C.

Many simple molecular substances are also volatile: they evaporate easily, which is why you can smell perfume, petrol and ethanol.

Property 2: poor electrical conductivity

Simple molecular substances generally don’t conduct electricity in any state — solid, liquid or gas. Molecules have no overall charge, and there are no free electrons or ions to carry a current.

Pure water, for example, is a very poor conductor. Tap water conducts because it contains dissolved ions from minerals.

There’s an important exception: some covalent molecules react with water to form ions. Hydrogen chloride gas doesn’t conduct, but in water it ionises completely to form hydrochloric acid:

HCl(g) + H₂O(l) → H₃O⁺(aq) + Cl⁻(aq)

The resulting solution conducts well. Ammonia does something similar, forming NH₄⁺ and OH⁻, though only partly.

Property 3: solubility — “like dissolves like”

Covalent substances vary widely in solubility, following the rule “like dissolves like”:

• Polar molecules, especially those that can hydrogen-bond, tend to dissolve in polar solvents such as water. Examples: ethanol, sugar, ammonia, ethanoic acid.
• Non-polar molecules dissolve in non-polar solvents such as hexane, but not in water. Examples: iodine (in hexane it gives a purple solution), oils, fats, wax, petrol.

Why doesn’t oil dissolve in water? Water molecules are strongly attracted to each other by hydrogen bonds. A non-polar molecule can’t form hydrogen bonds, so pushing it between water molecules would break water–water attractions without replacing them. It’s more favourable for the oil and water to stay separate.

This explains everyday observations: salad dressing separates into layers, grease won’t wash off with water alone (you need soap or detergent, whose molecules have both a polar and a non-polar end), and sugar dissolves easily in tea.

Property 4: soft solids

Solid simple molecular substances — iodine, sulfur, wax, sugar — are usually soft or crumbly, because the molecules are held in the crystal only by weak forces. You can easily scratch a candle with a fingernail, for example.

Property 5: often low density and flammable

Many covalent compounds containing carbon and hydrogen are flammable: their molecules react with oxygen, forming stronger bonds in CO₂ and H₂O and releasing energy. This is why fuels such as methane, propane and petrol are covalent substances. (Water and CO₂, of course, don’t burn — they’re already fully oxidised.)

Giant covalent substances: the exceptions

A few covalent substances form continuous networks of covalent bonds with no separate molecules (see giant covalent structures). Their properties are completely different:

Property	Simple molecular	Giant covalent (e.g. diamond, SiO₂)
Melting point	Low	Very high (diamond above 3,500 °C; SiO₂ about 1,700 °C)
Hardness	Soft	Very hard (diamond, quartz)
Solubility	Varies	Insoluble in all solvents
Conductivity	No	No — except graphite and graphene, which conduct

The reason: melting a giant covalent substance means breaking strong covalent bonds, not weak intermolecular forces.

Covalent vs ionic compounds

Property	Simple covalent	Ionic
Particles	Molecules	Ions
Melting point	Low	High
Conducts when molten	No	Yes
Conducts when dissolved	Usually no (unless it forms ions)	Yes
Solubility in water	Only if polar	Often (not always)
Solubility in non-polar solvents	If non-polar	No

See properties of ionic compounds and ionic vs covalent bonds.

Worked example: ranking boiling points

Question: Put these in order of increasing boiling point and explain: methane (CH₄), ethanol (C₂H₅OH) and octane (C₈H₁₈).

Answer: methane (−162 °C) < ethanol (78 °C) < octane (about 126 °C).

Methane is a small, non-polar molecule with only weak London forces between molecules. Ethanol is small too, but its O–H group allows hydrogen bonding between molecules, so more energy is needed to separate them. Octane has no hydrogen bonding, but it’s a much larger molecule with many more electrons, so its London forces are strong enough to outweigh ethanol’s hydrogen bonds.

Size and hydrogen bonding both matter, and a large enough molecule can boil higher than a smaller hydrogen-bonded one.

Common misconceptions

• “Covalent bonds are weak because covalent compounds melt easily.” The bonds are strong; the forces between molecules are weak.
• “Boiling breaks covalent bonds.” It separates whole molecules.
• “All covalent substances are non-conductors.” Graphite conducts, and substances like HCl form conducting solutions.
• “All covalent substances have low melting points.” Giant covalent structures have some of the highest known.

Key takeaways

• Simple molecular substances have strong bonds within molecules but weak forces between them.
• This gives low melting and boiling points, volatility and softness.
• They don’t conduct because molecules are neutral and there are no free charges.
• “Like dissolves like”: polar molecules dissolve in water; non-polar ones don’t.
• Giant covalent substances are the exception: very high melting points, hard and insoluble.