3.1.3.4 - Bonding and Physical Properties

Structure, melting, and conductivity

Solids with a crystal structure have particles arranged in a regular pattern. Those particles may be ions, atoms, or molecules, and the attractions between them are what give the solid its physical properties.

When a solid melts, energy is absorbed because the particles need more freedom to move. What must be overcome depends on the structure:

• in giant ionic, metallic, and macromolecular crystals, strong attractions or bonds within the giant structure must be overcome
• in molecular crystals, only intermolecular forces between molecules are overcome

The energy absorbed on melting is the enthalpy change of melting, or fusion. When a liquid vaporises, even more energy is needed because the particles must become much farther apart. The energy absorbed on boiling is the enthalpy change of vaporisation. Stronger attractions between particles therefore usually mean both a higher melting temperature and larger energy changes for melting and vaporisation.

The reverse changes release energy. Freezing and condensing are exothermic because attractions form again as particles become more ordered. Freezing need not bring particles closer together: water expands when it freezes.

Two simple checks unlock most of the patterns in this topic:

• how much energy is needed to separate the particles?
• are there any mobile ions or delocalised electrons present?

Those two checks explain most melting-point and conductivity patterns in this topic.

At the melting or boiling point, energy can change the arrangement and separation of particles while temperature stays constant. The energy changes the attractions between particles rather than increasing their average kinetic energy during the phase change.

Giant ionic and metallic crystals

Start with sodium chloride and magnesium because both are giant structures with high melting temperatures, but they conduct differently.

Sodium chloride is a giant ionic lattice of $\ce{Na^+}$ and $\ce{Cl^-}$ ions. Each positive ion is surrounded by negative ions and vice versa, so there are strong electrostatic attractions throughout the crystal. A large amount of energy is needed to separate these ions, so sodium chloride has a high melting point.

In the solid, sodium chloride does not conduct electricity because the ions are fixed in place. Once molten, or dissolved in water, the ions can move and carry charge.

The diagram below shows a simple 2D slice of the sodium chloride lattice, with the alternating pattern continuing in all directions through the crystal.

[DIAGRAM: asset_name: 1.3.4 - Bonding and Physical Properties - Diagram 1; asset_slug: 1.3.4 - Bonding and Physical Properties - Diagram 1; recommended_method: deterministic_code; description: A lattice slice contains exactly eight Na+ and eight Cl− ions, alternating in rows and columns without covalent bond lines.]

Magnesium is a metallic crystal. Each magnesium atom contributes its outer electrons to a shared pool, leaving a lattice of positive ions.

The strong attraction between $\ce{Mg^{2+}}$ ions and delocalised electrons gives magnesium a high melting temperature. Magnesium conducts electricity in both the solid and liquid states because the delocalised electrons are mobile.

This is the key comparison:

• sodium chloride conducts only when ions can move
• magnesium conducts because electrons can move even while the metal is solid

The diagram below shows the metallic lattice model for magnesium, where mobile delocalised electrons move through the whole structure.

[DIAGRAM: asset_name: 1.3.4 - Bonding and Physical Properties - Diagram 2; asset_slug: 1.3.4 - Bonding and Physical Properties - Diagram 2; recommended_method: deterministic_code; description: A metallic magnesium fragment contains nine Mg2+ ions and eighteen delocalised electrons.]

For conductivity, always identify the particle that moves. Writing only "it has charges" is too vague.

Molecular crystals: iodine and ice

Iodine and ice are both molecular crystals. That means they contain discrete molecules arranged in a regular pattern. The covalent bonds inside each molecule are strong, but the forces between molecules are much weaker than the bonds in giant ionic, metallic, or macromolecular structures.

Solid iodine contains $\ce{I2}$ molecules. Each molecule has a strong covalent bond, but neighbouring molecules are held together only by London forces. The iodine molecules are large, so these forces are strong enough to make iodine a solid, but they are still much weaker than the forces in giant structures. That is why iodine melts relatively easily and can sublime: when iodine changes state, the intermolecular forces are overcome, not the $\ce{I-I}$ covalent bonds.

Ice is also a molecular crystal, but its molecules are $\ce{H2O}$. In ice, each water molecule is hydrogen bonded to four neighbours in an open arrangement.

Hydrogen bonds are stronger than ordinary London forces for molecules of similar size, which is why water has unusually high melting and boiling temperatures for a small molecule. Even so, ice still melts at a much lower temperature than sodium chloride or diamond because melting ice only requires intermolecular forces to be overcome, not a giant lattice of strong bonds.

Neither iodine nor pure ice conducts electricity. They have no mobile ions and no delocalised electrons. The fact that they are covalent substances does not make them conductors; conductivity depends on mobile charged particles.

The diagram below compares the two molecular crystals directly: iodine is held together by London forces between $\ce{I2}$ molecules, whereas ice has an open lattice held together by hydrogen bonds between $\ce{H2O}$ molecules.

[DIAGRAM: asset_name: 1.3.4 - Bonding and Physical Properties - Diagram 3; asset_slug: 1.3.4 - Bonding and Physical Properties - Diagram 3; recommended_method: deterministic_code; description: A two-dimensional ice fragment contains two hydrogens per oxygen and selected dashed O–H···O hydrogen bonds.]

For ice in particular, remember what does not happen on melting: the $\ce{O-H}$ covalent bonds inside water molecules remain. The structure changes because hydrogen bonds are disrupted and rearranged.

Macromolecular crystals: diamond and graphite

Diamond and graphite are both macromolecular forms of carbon. They are made only of carbon atoms, but the atoms are arranged differently, so their properties are very different.

In diamond, each carbon atom forms four covalent bonds in a tetrahedral arrangement. The result is a rigid three-dimensional network of strong covalent bonds. A very large amount of energy is needed to break enough of these bonds to separate the atoms. Carbon does not normally melt at atmospheric pressure: graphite sublimes, and diamond can transform before subliming. The exam comparison of very high melting/sublimation temperatures expresses the strong covalent network. Diamond does not conduct electricity because all four outer electrons of each carbon atom are used in covalent bonds, so there are no mobile charge carriers.

The diagram below shows the local bonding pattern in diamond: each carbon atom is bonded to four others in a tetrahedral arrangement.

[DIAGRAM: asset_name: 1.3.4 - Bonding and Physical Properties - Diagram 4; asset_slug: 1.3.4 - Bonding and Physical Properties - Diagram 4; recommended_method: deterministic_code; description: A diamond network fragment shows four covalent bonds per carbon and continuing bonds beyond the fragment.]

In graphite, each carbon atom forms three covalent bonds. That creates flat hexagonal layers. The fourth outer electron from each carbon is delocalised within the layer, so graphite can conduct electricity along the layers.

The covalent bonds within each layer are still very strong, so graphite also requires a very high temperature to change to separated gaseous particles. However, the layers are held together only by weak forces, so they can slide over each other. That is why graphite is soft even though it has a giant covalent structure.

The diagram below shows why graphite behaves differently from diamond: the layers contain delocalised electrons, but only weak forces act between neighbouring layers.

[DIAGRAM: asset_name: 1.3.4 - Bonding and Physical Properties - Diagram 5; asset_slug: 1.3.4 - Bonding and Physical Properties - Diagram 5; recommended_method: deterministic_code; description: Graphite has three bonds per carbon within hexagonal sheets, weak forces between layers and electrons delocalised within the sheets.]

The comparison only becomes convincing when you explain both sides: why graphite conducts and why diamond does not.

Drawing and using the patterns

This topic often becomes much easier once you match the observable property to the particle model, whether you are sketching a structure or interpreting melting point and conductivity data.

Small drawing rules that keep the chemistry accurate:

• $\ce{NaCl}$: draw alternating $\ce{Na^+}$ and $\ce{Cl^-}$ ions in a repeating lattice; do not draw separate NaCl molecules
• $\ce{Mg}$: draw a regular array of positive ions with a shared sea of electrons, not paired $\ce{Mg-Mg}$ bonds
• $\ce{I2}$: draw pairs of iodine atoms as molecules, with weak forces between molecules
• ice: draw $\ce{H2O}$ molecules with dashed hydrogen bonds between molecules in an open arrangement
• diamond: show each carbon bonded to four others in a 3D network
• graphite: show hexagonal layers, three bonds per carbon, and delocalised electrons within the layers

When asked for a specified number of particles, count them before adding bonds or charges. A 4-by-4 sodium chloride slice has eight $\ce{Na+}$ and eight $\ce{Cl-}$, with opposite charges horizontally and vertically adjacent. The slice represents a three-dimensional lattice; the border does not mean the real crystal stops there. For a local diamond fragment, draw four bonds around the central carbon and show that the outer carbons also bond onwards, rather than implying a five-carbon molecule.

An unknown solid can be investigated using small samples and safe heating, solubility and conductivity tests. Compare conductivity of the solid with that of its melt or aqueous solution where safe and feasible. Rinse electrodes between solutions to avoid contamination and test the solvent alone as a control. High melting temperature plus conductivity only when molten strongly supports ionic bonding; low melting temperature and no conductivity support molecular structure. Insolubility alone cannot distinguish a metal from a giant covalent solid, and graphite is the important conducting giant-covalent exception. Use several observations together.

A quick classification summary pulls the whole topic together:

The pattern behind the table matters more than memorising isolated facts. High melting point usually means strong attractions throughout a giant structure. Conductivity depends on whether ions or electrons can actually move.

If you keep linking visible properties back to the particle arrangement, this topic becomes much easier to reason through and much harder to memorise incorrectly.
For a low-melting unknown, a small dry sample in a capillary can be heated in a melting-point apparatus. Heat slowly near melting and record the onset and completion temperatures. A suitable water bath provides gentle heating only within its temperature range; failure to melt below 100 °C does not prove giant bonding. Avoid attempting to melt an unknown at high temperature without an approved method and hazard assessment.