Properties of covalent compounds

Every physical property difference between covalent and ionic compounds traces back to one structural fact: what is doing the holding, and how strongly.

Inside the particle versus between particles. A covalent compound is built from a small, fixed number of atoms per molecule (CCl₄ has one carbon and four chlorines, no more, no less), joined by strong shared-electron-pair bonds. But separate molecules only brush against each other, held by comparatively feeble intermolecular forces. An ionic compound has no separate "molecules" at all — it is one continuous three-dimensional lattice of alternating positive and negative ions, each one gripped from every direction by its neighbours.

Why the ionic grip is so much stronger. The force between two charges follows Coulomb's law:

where: $F$ is the electrostatic force of attraction, $q_1$ and $q_2$ are the magnitudes of the charges on the two ions, $r$ is the distance between their centres, and $k$ is a constant. In an ionic crystal, every ion sits at distance $r$ from several oppositely-charged neighbours simultaneously, so the total attraction pulling the lattice together is enormous and acts equally in all directions. In a covalent compound, this strong electrostatic attraction is confined inside each molecule (holding the shared electron pairs in place); between molecules there is no such charge-charge attraction at all, only the weak forces left over.

What that difference produces. Melting or boiling a covalent compound only needs to separate molecules from each other — the strong bonds inside each molecule never break, so little energy is needed, hence low melting and boiling points. Melting an ionic compound means ripping apart the entire charged lattice, ion by ion, which needs far more energy — hence the high melting points of salts like sodium chloride. Solubility follows the same logic: polar water molecules can prise apart and surround individual ions, dissolving most ionic solids, but water usually cannot compete with the strong internal bonds of a non-polar covalent molecule, so most covalent compounds resist dissolving in water and dissolve instead in non-polar solvents where similar weak forces are at play.

There is no formula that predicts a melting point from first principles at this level, but the physical chain of reasoning is exact and always runs the same way once you know which type of bonding is present.

Start with what must be broken. Ask what has to be pulled apart to melt, boil, or dissolve the substance. For an ionic solid, the answer is always "the whole lattice" — every ion is locked to several oppositely-charged neighbours, so disrupting the solid means overcoming that Coulombic grip everywhere at once. For a covalent solid or liquid, the answer is always "only the links between molecules" — the far stronger bonds holding each molecule's own atoms together are never touched during melting or boiling.

Match the energy needed to the temperature required. Overcoming a lattice of strong, all-directional ionic attractions demands a great deal of energy, which shows up experimentally as a high melting or boiling point. Overcoming only weak intermolecular attractions demands comparatively little energy, so covalent compounds usually melt and boil at low temperatures — some, like CCl₄, are liquid or even gaseous at room temperature.

Apply the same logic to conductivity. Electric current needs charge carriers that are free to move. A solid ionic lattice has ions, but they are locked in fixed positions, so no current flows. Melt the lattice or dissolve it in water and the ions become mobile, so conduction switches on. A covalent compound never has ions at all — its particles are neutral molecules from the start — so, with rare exceptions such as acids that ionise in water, it fails to conduct electricity in any state: solid, liquid, or dissolved.

Apply it to solubility. Water is a polar solvent, good at pulling apart and surrounding charged ions, so it dissolves most ionic compounds readily. It is generally a poor solvent for non-polar covalent molecules, which dissolve instead in non-polar or organic solvents where similarly weak intermolecular forces make mixing easy. Exceptions like sugar dissolving in water happen only when the covalent molecule itself is strongly polar.

Sodium chloride melts at about 801 degrees Celsius, while carbon tetrachloride, CCl₄, melts at about minus 23 degrees Celsius. Explain the huge gap, and predict which of the two conducts electricity when melted.

Step 1. Identify the bonding in each. Sodium chloride forms by complete electron transfer between a metal and a non-metal, giving Na⁺ and Cl⁻ ions locked into a crystal lattice — it is electrovalent. Carbon tetrachloride forms by carbon sharing electron pairs with four chlorine atoms, giving discrete, neutral CCl₄ molecules — it is covalent.

Step 2. Compare what melting requires. Melting NaCl means overcoming the strong electrostatic attraction, given by $F = k\,q_1q_2/r^2$, acting between every Na⁺ and its several Cl⁻ neighbours throughout the entire lattice — a high-energy process. Melting CCl₄ means overcoming only the weak forces of attraction between separate CCl₄ molecules; the strong C–Cl covalent bonds inside each molecule stay fully intact.

Step 3. Read off the melting points. Because breaking an all-directional ionic lattice needs far more energy than separating loosely-held molecules, NaCl requires a temperature roughly 824 degrees Celsius higher than CCl₄ to melt — the 801 degree Celsius versus minus 23 degree Celsius gap is direct experimental evidence of ionic versus covalent bonding.

Step 4. Predict the conductivity. Once molten, NaCl's Na⁺ and Cl⁻ ions are free to move, so molten NaCl conducts electricity well. Molten CCl₄ still contains only neutral molecules, with no ions produced by melting, so it does not conduct electricity in the molten state either. Answer: the roughly 824 degree Celsius gap is due to breaking a strong, all-directional ionic lattice versus separating weakly-held covalent molecules; molten NaCl conducts, molten CCl₄ does not.

• Assuming every covalent compound is a gas or has an extremely low melting point — sugar and many polymers are solids at room temperature; the deciding factor is the strength of the forces present, not the physical state alone.
• Thinking that dissolving in water automatically means a substance conducts electricity — plain dissolving does not create ions; only substances that actually ionise, such as acids, release charge carriers in water.
• Believing solid ionic compounds conduct because "they contain ions" — those ions are fixed in the lattice and cannot move, so solid ionic compounds do not conduct.
• Forgetting that the covalent bonds inside a molecule are themselves strong — a covalent compound's low melting point is due to weak forces between molecules, never a weak bond within the molecule.

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