Now everything pays off. Melting point, volatility, solubility, conductivity — all of it comes from two questions: is this a molecule or a network? and what forces sit between the molecules?
📚 What you need to know
For molecular substances, melting and boiling break intermolecular forces, never covalent bonds.
Stronger intermolecular forces → higher melting and boiling point → less volatile.
Force strength depends on the number of electrons, the polarity, and the type of force present.
Like dissolves like: polar dissolves in polar, nonpolar dissolves in nonpolar.
Molecular covalent substances do not conduct — no ions, no free electrons. Some polar ones (HCl) conduct after ionising in water.
Giant covalent substances have very high melting points and are insoluble. Only graphite and graphene conduct.
Melting and boiling point
The single most important sentence in this topic: when a simple molecular substance melts or boils, the covalent bonds stay exactly where they are. Only the weak forces holding one molecule to the next are overcome.
Those forces are weak, which is why so many covalent substances are gases or liquids at room temperature. Compare that with a giant covalent structure, where melting really does mean breaking covalent bonds — hence the enormous difference in melting points.
Heating separates the molecules but leaves each one whole. The green dashes are what your energy is actually paying for.
What decides how strong the forces are
Three things, and you should check them in this order:
🧩 Comparing two molecular substances
Type of force first. Hydrogen bonding beats dipole–dipole, which beats London forces. If one substance can hydrogen bond and the other cannot, that usually settles it.
Then the number of electrons. If both have the same type of force, the one with more electrons has stronger London forces.
Then the shape. For isomers with identical electron counts, the longer, straighter molecule packs better and boils higher.
Here is all three ideas in one comparison. These molecules have almost the same number of electrons, so London forces are similar — the difference is entirely in what else they can do:
Almost identical electron counts, so the difference in boiling point comes entirely from the extra types of force available.
WORKED EXAMPLE
Place CH3CH2CH2OH, CH3COCH3 and CH3CH2CH2CH3 in order of increasing boiling point and explain.
Step 1 — check the electron counts32–34 electrons each, so London forces are about the same. Not the deciding factor.Step 2 — identify the strongest force in eachButane: nonpolar → London only.Propanone: polar C=O → London + dipole–dipole.Propan-1-ol: has an O–H → London + dipole–dipole + hydrogen bonding.butane < propanone < propan-1-ol
Volatility
Volatile just means “evaporates easily”. It is the flip side of boiling point — weak intermolecular forces mean a low boiling point and high volatility. That is why petrol and nail varnish remover evaporate in seconds while water takes hours.
Solubility
The rule is “like dissolves like”, and the reason is energy. Dissolving something means pulling the solvent particles apart to make room, then slotting the solute in. That only works if the new solute–solvent attractions are strong enough to pay for the ones you had to break.
Polar in polar. A polar solute can form dipole–dipole attractions or hydrogen bonds with a polar solvent. Ethanol and water mix in any proportion for exactly this reason.
Nonpolar in nonpolar. Only London forces are involved on both sides, and they match up fine. Iodine dissolves happily in hexane.
Polar in nonpolar (or the reverse) fails. A nonpolar solute cannot form strong enough attractions to compensate for breaking up the hydrogen bonding in water, so it stays separate.
Ethanol can join water’s hydrogen bonding network; hexane cannot, so it is pushed out into its own layer.
Size matters too. Small alcohols like ethanol (C2H5OH) dissolve completely in water because the –OH group dominates. As the carbon chain grows, the nonpolar hydrocarbon part takes over and solubility drops sharply — hexan-1-ol (C6H13OH) is only slightly soluble. Same functional group, very different behaviour.
Giant covalent substances are insoluble in everything. No solvent can supply the energy needed to break a continuous network of covalent bonds.
Electrical conductivity
To conduct, a substance needs charged particles that can move. Most covalent substances have neither free electrons nor ions, so the answer is usually no:
Simple molecular — do not conduct in any state. The molecules are neutral and all the electrons are tied up in bonds.
Exception: some polar molecules ionise in water. HCl is a gas that does not conduct, but dissolve it and it splits into H+ and Cl−, which do conduct.
Giant covalent — do not conduct, because every outer electron is used in bonding.
Exception:graphite and graphene, where each carbon has one delocalised electron free to move.
Summary table
Nonpolar molecular
Polar molecular
Giant covalent
Forces to overcome
London only
London + dipole–dipole (± H bonding)
Covalent bonds
Melting / boiling point
Low
Low to moderate
Very high
Volatility
High
Moderate
Very low
Soluble in water?
No
Often yes
No
Soluble in nonpolar solvents?
Yes
Sometimes
No
Conducts?
No
No, unless it ionises in water
No, except graphite & graphene
Example
Br2, CH4
H2O, HCl
Diamond, SiO2
WORKED EXAMPLE
A substance X melts at 1450 °C and conducts electricity poorly both as a solid and when molten. Deduce its structure.
Very high melting point → must be a giant structureThat leaves ionic, metallic or giant covalent.Does not conduct when molten → rules out ionicMolten ionic compounds have free-moving ions.Does not conduct as a solid → rules out metallicgiant covalent, like diamond or SiO₂
💡 Exam tip
Never write “the covalent bonds break” when explaining a boiling point. Write “the intermolecular forces are overcome”.
Name the specific force — London, dipole–dipole or hydrogen bonding — not just “intermolecular forces”.
For solubility answers, say what the solute forms with the solvent. “It is polar” alone is not an explanation.
Conductivity answers must mention mobile charged particles: free electrons or free ions.
⚠️ Common mix-up
Low melting point does not mean weak covalent bonds. The bonds in CH4 are strong; the forces between the molecules are not.
“Bigger molecule” is not an explanation. Say more electrons, which is what actually strengthens London forces.
HCl does not conduct as a gas. It only conducts once dissolved, because that is when ions appear.
Don’t assume every polar substance dissolves in water. It also has to be small enough for the polar part to win.
Up next: Chromatography — a practical technique built entirely on the ideas from these last few pages, since separation depends on which intermolecular forces each component can form.
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