There is a significant difference between the physical properties of NaCl and Cl₂, as shown in the table below, which results from the difference between the ionic bonds in NaCl and the covalent bonds in Cl₂.

Some Physical Properties of NaCl and Cl₂

Each Na⁺ ion in NaCl is surrounded by six Cl^- ions, and vice versa, as shown in the figure below. Removing an ion from this compound therefore involves breaking at least six bonds. Some of these bonds would have to be broken to melt NaCl, and they would all have to be broken to boil this compound. As a result, ionic compounds such as NaCl tend to have high melting points and boiling points. Ionic compounds are therefore solids at room temperature.

Cl₂ consists of molecules in which one atom is tightly bound to another, as shown in the figure above. The covalent bonds within these molecules are at least as strong as an ionic bond, but we don't have to break these covalent bonds to separate one Cl₂ molecule from another. As a result, it is much easier to melt Cl₂ to form a liquid or boil it to form a gas, and Cl₂ is a gas at room temperature.

The difference between ionic and covalent bonds also explains why aqueous solutions of ionic compounds conduct electricity, while aqueous solutions of covalent compounds do not. When a salt dissolves in water, the ions are released into solution.

	H2O
NaCl(s)		Na+(aq) + Cl-(aq)

These ions can flow through the solution, producing an electric current that completes the circuit. When a covalent compound dissolves in water, neutral molecules are released into the solution, which cannot carry an electric current.

	H2O
C12H22O11(s)		C12H22O11(aq)

When two chlorine atoms come together to form a covalent bond, each atom contributes one electron to form a pair of electrons shared equally by the two atoms, as shown in the figure below. When a sodium atom combines with a chlorine atom to form an ionic bond, each atom still contributes one electron to form a pair of electrons, but this pair of electrons is not shared by the two atoms. The electrons spend most of their time on the chlorine atom.

Ionic and covalent bonds differ in the extent to which a pair of electrons is shared by the atoms that form the bond. When one of the atoms is much better at drawing electrons toward itself than the other, the bond is ionic. When the atoms are approximately equal in their ability to draw electrons toward themselves, the atoms share the pair of electrons more or less equally, and the bond is covalent. As a rule of thumb, metals often react with nonmetals to form ionic compounds or salts, and nonmetals combine with other nonmetals to form covalent compounds. This rule of thumb is useful, but it is also naive, for two reasons.

• The only way to tell whether a compound is ionic or covalent is to measure the relative ability of the atoms to draw electrons in a bond toward themselves.
• Any attempt to divide compounds into just two classes (ionic and covalent) is doomed to failure because the bonding in many compounds falls between these two extremes.

The first limitation is the basis of the concept of electronegativity. The second serves as the basis for the concept of polarity.

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