Simple Molecular Structures

Simple molecular substances consist of small, discrete molecules held together by weak intermolecular forces. The properties of these substances are determined primarily by the strength of the intermolecular forces between molecules — not by the strength of the covalent bonds within them. Understanding this distinction is central to explaining the physical properties of simple molecular compounds.

What Are Simple Molecular Substances?

A simple molecular substance is composed of individual molecules, each containing a fixed, small number of atoms covalently bonded together. Examples include water (H₂O), carbon dioxide (CO₂), methane (CH₄), oxygen (O₂), iodine (I₂), glucose (C₆H₁₂O₆), and sulfur (S₈).

The key structural feature is that strong covalent bonds exist within each molecule, but only weak intermolecular forces (London, dipole–dipole, or hydrogen bonding) act between molecules. It is the weak intermolecular forces that must be overcome during melting and boiling — not the covalent bonds.

Low Melting and Boiling Points

Simple molecular substances generally have low melting and boiling points because the intermolecular forces between molecules are weak. It takes relatively little energy to separate the molecules from each other (to melt or boil the substance), even though the covalent bonds within each molecule are very strong.

Important distinction: When a simple molecular substance boils, the molecules themselves remain intact. The covalent bonds are not broken. Only the intermolecular forces are overcome. For example, when water boils at 100°C, the O–H covalent bonds within each H₂O molecule remain intact — it is the hydrogen bonds between molecules that are broken.

Many simple molecular substances are gases or liquids at room temperature. Those that are solids (like iodine, naphthalene, or sugar) have relatively low melting points compared to giant structures.

Boiling Points of Common Simple Molecular Substances

Substance	Formula	Mr	Intermolecular forces	bp / °C
Helium	He	4	London only	−269
Methane	CH₄	16	London only	−162
Propane	C₃H₈	44	London only	−42
Hydrogen chloride	HCl	36.5	London + dipole-dipole	−85
Ammonia	NH₃	17	London + dipole + H-bonding	−33
Water	H₂O	18	London + dipole + H-bonding	100
Ethanol	C₂H₅OH	46	London + dipole + H-bonding	78
Iodine	I₂	254	London only	184

Notice that iodine, with only London forces, has a higher boiling point than ammonia (which has hydrogen bonding). This is because I₂ has far more electrons (106 vs 10), making its London forces cumulatively very strong. Molecular size and electron count matter alongside the type of force.

Electrical Non-Conductivity

Simple molecular substances do not conduct electricity in any state — solid, liquid, or gas. This is because they have no free charged particles — no mobile ions and no delocalised electrons.

Even polar molecules, which have partial charges (δ+ and δ−), do not conduct because these partial charges are fixed within the molecule and cannot move independently. True electrical conduction requires either free-moving ions or delocalised electrons.

This is an important contrast with:

Ionic compounds — which conduct when molten or dissolved (mobile ions)
Metals — which conduct in the solid state (delocalised electrons)
Graphite — which conducts along layers (delocalised electrons)

Solubility Patterns

The solubility of simple molecular substances follows the general rule: "like dissolves like."

Polar molecules tend to dissolve in polar solvents (e.g., water). For example, ethanol and glucose dissolve in water because they can form hydrogen bonds or dipole–dipole interactions with water molecules, which compensate for the breaking of water–water hydrogen bonds.

Non-polar molecules tend to dissolve in non-polar solvents (e.g., hexane, cyclohexane). For example, iodine dissolves well in hexane because both are non-polar and interact through London forces.

Non-polar molecules do not dissolve well in water because dissolving them would break water's hydrogen bonds without forming sufficiently strong new interactions to compensate.

Iodine: A Typical Example

Iodine (I₂) is a classic example of a simple molecular solid and is frequently used in exam questions.

Structure: Discrete I₂ molecules arranged in a regular crystalline lattice, held together by London dispersion forces.

Properties:

Melting point: 114°C — relatively low, because only weak London forces between molecules need to be overcome
Sublimes easily — iodine converts directly from solid to purple vapour even below its melting point, showing how weak the intermolecular forces are
Does not conduct electricity — no free ions or delocalised electrons
Insoluble in water — non-polar molecules cannot form hydrogen bonds with water
Soluble in hexane — non-polar solute dissolves in non-polar solvent (like dissolves like)
Soluble in aqueous KI — forms the triiodide ion I₃⁻ through a chemical reaction, which is a different mechanism

Despite having only London forces, iodine is a solid at room temperature because it has a very large electron cloud (many electrons), making its London forces relatively strong for a simple molecular substance.

Comparing Structure Types from Properties

Being able to identify a structure type from physical properties is a key exam skill. Here is a systematic approach:

High Melting Point + Conducts When Solid

→ Metal (delocalised electrons)

High Melting Point + Conducts When Molten/Dissolved but NOT Solid

→ Giant ionic (mobile ions when melted)

Very High Melting Point + Never Conducts

→ Giant covalent (e.g., diamond, SiO₂)

Very High Melting Point + Conducts as Solid (layered)

→ Giant covalent — graphite (delocalised electrons in layers)

Low Melting Point + Never Conducts

→ Simple molecular (weak intermolecular forces, no free charges)

Worked Example: Identifying Unknown Substances

Four substances W, X, Y, Z have the following properties:

Substance	mp / °C	Solid conductivity	Molten conductivity	Soluble in water	Soluble in hexane
W	1710	No	No	No	No
X	114	No	No	No	Yes
Y	801	No	Yes	Yes	No
Z	1085	Yes	Yes	No	No

W: Very high mp, never conducts, insoluble in everything → giant covalent (SiO₂). All electrons localised in covalent bonds.

X: Low mp, never conducts, insoluble in water, soluble in hexane → simple molecular, non-polar (I₂). Only London forces; dissolves in non-polar hexane.

Y: High mp, conducts when molten, water-soluble → giant ionic (NaCl). Ions are mobile when molten; ion-dipole forces enable dissolution in water.

Z: High mp, conducts as solid, insoluble in water → metallic (Cu). Delocalised electrons provide conductivity in the solid state.

Predicting Structure from Properties

When given an unknown substance and asked to predict its structure, consider the following data points:

Melting/boiling point: Low → simple molecular. High → giant structure.
Electrical conductivity: Conducts as solid → metal or graphite. Conducts when molten → ionic. Never conducts → simple molecular or giant covalent.
Solubility in water: Soluble → ionic or polar molecular. Insoluble → non-polar molecular, giant covalent, or metal.
Solubility in non-polar solvents: Soluble → simple molecular (non-polar). Insoluble → ionic or giant structure.
Hardness: Hard → giant structure. Soft → simple molecular.

Common Misconceptions

"Simple molecular substances have weak bonds." This is wrong. The covalent bonds within the molecules are strong. The intermolecular forces between molecules are weak. Always be precise about which forces you are referring to.

"The boiling point of water is high because the O–H bond is strong." This is wrong. Boiling does not break covalent bonds. Water has a high boiling point because of strong hydrogen bonds between molecules. The O–H bond strength affects the stability of the water molecule itself, not the boiling point.

"Molecular substances cannot dissolve in water." This is wrong. Many polar molecular substances dissolve well in water (ethanol, glucose, HCl). It is non-polar molecular substances that are generally insoluble.

Why Structure Determines Properties

The central theme of this entire course is that structure determines properties. For simple molecular substances:

The covalent bonds within molecules determine chemical reactivity (which bonds break and form in reactions)
The intermolecular forces between molecules determine physical properties (melting point, boiling point, solubility)
The absence of free charged particles determines electrical behaviour (no conduction)

Being able to explain physical properties in terms of structure and bonding — using the correct terminology and making clear which forces are being discussed — is essential for full marks in A-Level Chemistry exams.

Common exam mistake: When asked "why does iodine have a low melting point?", students sometimes write "because it has weak covalent bonds." The I–I covalent bond is actually reasonably strong (151 kJ mol⁻¹). The low melting point is because the intermolecular London forces between I₂ molecules are weak. When iodine melts, the I–I bonds within each molecule remain intact — only the forces between molecules are overcome.

A-Level Deep Dive: Simple Molecular Structures

Spec mapping

Edexcel 9CH0 specification, Topic 2 — Bonding and Structure, sub-topic 2.7 sets out the simple molecular aspect: discrete covalent molecules held together by intermolecular forces, with low melting and boiling points and an inability to conduct electricity in any phase. This thread is examined far beyond Topic 2.

Cross-paper relevance:

Paper 1, Topic 4 (Inorganic chemistry — period 3): the chlorides Cl₂, P₄O₁₀, SO₃, SiCl₄ and (for contrast) MgCl₂ form a classic "structure determines property" sequence. Cl₂, P₄O₁₀, SO₃ and SiCl₄ are simple molecular and have low boiling points; MgCl₂ is ionic and has a very high melting point. Period 3 oxides and chlorides are the inorganic vehicle for testing simple molecular reasoning at A2.
Paper 2, Topics 6 onwards (Organic chemistry): essentially every organic compound — alkanes, alkenes, haloalkanes, alcohols, aldehydes, ketones, carboxylic acids, esters, amines, amides — is simple molecular. Boiling points, solubilities and physical states are all explained through chain length, branching, polarity and hydrogen bonding, all of which sit on the simple molecular framework laid down in 2.7.
Paper 3 (Synoptic / general principles): identification questions routinely give a table of physical data — melting point, electrical conductivity in solid and liquid states, solubility in water and in non-polar solvents — and ask candidates to assign each substance to a structural class. Low melting point, poor conductor in all phases and (often) solubility in non-polar solvents are the diagnostic signature of simple molecular.

(refer to the official Pearson Edexcel specification document for exact wording)

Worked example with full mark scheme

Question (7 marks):

(a) Iodine (I₂) sublimes at 184 °C. Sulfur (S₈) melts at 115 °C. Both are simple molecular. Explain, in terms of structure and bonding, why their melting/sublimation points differ in this way. (3)

(b) Solid iodine does not conduct electricity. Explain why. (2)

(c) Predict, with reasoning, whether tetrachloromethane (CCl₄) would dissolve in water and whether it would dissolve in hexane. (2)

Solution with full mark scheme:

(a) B1 — both substances consist of discrete molecules (I₂ and S₈) held together by London (dispersion) forces only, since both are non-polar.

M1 — comparing the strengths of London forces. The S₈ ring contains 8 sulfur atoms and a total of 128 electrons; an I₂ molecule contains 2 iodine atoms and 106 electrons. London force strength depends on the total number of electrons and on molecular polarisability.

A1 — concluding that I₂ has more polarisable electrons per molecule (large, diffuse outer shells) and that the contact area between S₈ rings is also large, so both have substantial London forces. The reason I₂ goes to the gas phase at a higher temperature than S₈ melts is that sublimation requires complete separation of molecules whereas melting only loosens them; the comparison "melting point of S₈ vs sublimation point of I₂" is not a like-for-like ranking of intermolecular force strength. A clean A1 answer states that the polarisability of I₂'s electron cloud, combined with the phase change being considered (sublimation rather than melting), accounts for the higher temperature.

Common error: writing "iodine has stronger covalent bonds than sulfur, so it sublimes higher". This loses the M and A marks because boiling/sublimation does not break covalent bonds — only intermolecular forces are overcome.

(b) B1 — solid iodine consists of discrete I₂ molecules. Within each molecule the bonding electrons are localised in the I–I covalent bond; there are no delocalised electrons free to move through the lattice.

B1 — there are also no ions present (iodine is a non-metal element existing as neutral molecules), so there are no mobile charge carriers of any kind. Hence solid I₂ is an electrical insulator.

(c) B1 — CCl₄ is non-polar overall: each C–Cl bond is polar, but the tetrahedral arrangement of four identical C–Cl bonds means the bond dipoles cancel by symmetry. Consequently CCl₄ cannot form hydrogen bonds with water, nor can it interact strongly with water's dipole. To dissolve in water, CCl₄ would have to disrupt water's hydrogen-bonded network without compensating energetic gain. So CCl₄ is essentially insoluble in water.

B1 — hexane (C₆H₁₄) is also non-polar; the only intermolecular force in pure hexane is London dispersion. CCl₄ and hexane molecules can interact via London forces of comparable strength, and there is no hydrogen-bonded network to disrupt. So CCl₄ dissolves readily in hexane ("like dissolves like").

Total: 7 marks (B5 M1 A1).

Specimen question modelled on the Edexcel 9CH0 Paper 1 format

Question (6 marks): The boiling points of the Group 7 (halogen) elements at standard atmospheric pressure are:

Halogen	F₂	Cl₂	Br₂	I₂
Boiling point / °C	−188	−34	59	184

Explain the trend in boiling points down the group. Your answer should refer to structure, type of intermolecular force, and the factors that determine its strength. (6)

Mark scheme decomposition by AO:

B1 (AO1) — identifying that all four halogens are simple molecular substances consisting of diatomic non-polar molecules (X₂).
B1 (AO1) — identifying that the only intermolecular force present is the London (dispersion) force, since the molecules are non-polar.
M1 (AO2) — explaining what determines London force strength: the number of electrons in the molecule (and equivalently the molecular mass), combined with the polarisability of the electron cloud — larger, more diffuse electron clouds polarise more easily and produce larger instantaneous and induced dipoles.
M1 (AO2) — applying this to the data: down the group, atomic number increases, so the number of electrons per X₂ molecule increases (F₂ has 18 electrons, Cl₂ has 34, Br₂ has 70, I₂ has 106), and the outermost electrons occupy progressively larger and more polarisable shells.
A1 (AO2) — concluding that London forces strengthen down the group, so more energy is required to separate the molecules, hence boiling point increases.
A1 (AO3) — synoptic point: the change in physical state at room temperature (F₂ and Cl₂ gases, Br₂ liquid, I₂ solid) is a direct consequence of London force strength scaling with electron count, not a change in bonding type — all four are still simple molecular with the same kind of intermolecular interaction.

Total: 6 marks split AO1 = 2, AO2 = 3, AO3 = 1. Edexcel routinely uses the halogen boiling-point trend to test whether candidates can articulate the one variable (London force strength) that changes monotonically while structure type stays the same.

Synoptic links

This sub-topic threads through the entire A-Level Chemistry course:

Topic 2.6 — Intermolecular forces. The melting points, boiling points, viscosities and surface tensions of simple molecular substances are determined entirely by intermolecular forces. Mastery of London, permanent dipole–dipole and hydrogen bonding (Topic 2.6) is the prerequisite for explaining any simple molecular property quantitatively. Without 2.6 in place, 2.7 reduces to memorisation.
Topic 4 — Inorganic chemistry, period 3. The chlorides and oxides of the period 3 elements span structural classes: NaCl and MgCl₂ are ionic giants; AlCl₃ shows borderline behaviour (dimerising to Al₂Cl₆ in the gas phase, simple molecular); SiCl₄, PCl₃, PCl₅, SCl₂ and Cl₂ are all simple molecular. Their boiling points, conductivities and reactions with water are direct exam targets that sit on a 2.7 foundation.
Topic 6 onwards — Organic chemistry. Every alkane, alkene, alkyne, alcohol, aldehyde, ketone, carboxylic acid, ester, amine and amide encountered at A-Level is simple molecular. The boiling-point ladder of straight-chain alkanes (methane gas → ethane gas → … → pentane liquid → hexadecane solid wax), the elevated boiling points of alcohols relative to alkanes of similar Mr (hydrogen bonding) and the solubility of small alcohols in water are all simple molecular reasoning in organic clothing.