Bonding and physical properties

Diamond can cut glass, graphite rubs onto paper, sodium chloride shatters when struck, and copper can be drawn into wires. These substances behave very differently even though all of them are ultimately built from atoms and electrons. The reason lies in how their particles are arranged and held together.

Bonding describes the attractive forces acting between particles. Structure describes the larger arrangement produced by those particles and attractions. Neither idea should be considered alone: the physical properties of a substance emerge from the combination of both.

This topic therefore gives us a general method for turning microscopic models into macroscopic explanations. Instead of memorising isolated properties, we shall learn to predict them by asking what particles are present, what holds them together and whether any charged particles are free to move.

From bonding and structure to physical properties

A chemical bond is an electrostatic attraction that holds particles together. The particles involved depend on the type of bonding.

• Ionic bonding is the electrostatic attraction between oppositely charged ions.
• Metallic bonding is the electrostatic attraction between positive metal ions and delocalised electrons.
• Covalent bonding is the electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms.

A structure describes how large numbers of these particles are arranged. A substance may consist of a repeating giant structure or of separate, individual molecules.

What does the word lattice mean?

A lattice is a regular, repeating arrangement of particles in a crystalline solid. The particles may be ions, atoms or, in some solids, molecules.

The term is particularly useful when discussing giant ionic and giant metallic solids because the repeating arrangement is central to their structural models. We can also describe a molecular crystal as a molecular lattice when referring specifically to the ordered arrangement of molecules in the solid.

However, simple molecular is the broader structural description. A molecular substance consists of separate molecules whether it is solid, liquid or gaseous. The word lattice specifically refers to an ordered solid arrangement.

A four-step method for explaining a physical property

A reliable structure–property explanation can usually be built in four stages.

1. State the structure: identify whether the substance is ionic, metallic, giant covalent or simple molecular.
2. Identify the particles: state whether the structure contains ions, atoms, molecules or delocalised electrons.
3. Identify the relevant attraction or mobile particle: describe the bonds, intermolecular forces or charge carriers involved.
4. Connect the microscopic model to the property: explain how much energy is required or whether charged particles can move.

Giant ionic lattices

An ionic compound consists of positive ions and negative ions arranged in a regular, repeating three-dimensional lattice. Each ion is surrounded by ions of opposite charge, and strong electrostatic attractions act throughout the structure.

Ionic bonding is therefore not confined to one isolated pair of ions. The formula of an ionic compound, such as NaCl or MgO, gives the simplest ratio of ions within the lattice rather than representing a separate molecule.

Why do ionic compounds have high melting and boiling points?

In a solid ionic lattice, the ions are held in fixed positions by strong electrostatic attractions between oppositely charged ions.

To melt the substance, sufficient energy must be supplied to overcome enough of these attractions for the ions to move relative to one another. Because the attractions are strong and extend throughout the lattice, a large amount of energy is required.

Ionic compounds therefore generally have high melting points and very high boiling points. Some decompose before a normal boiling point can be measured.

Exam-style explanation: Sodium chloride has a giant ionic lattice containing oppositely charged Na⁺ and Cl⁻ ions. A large amount of energy is required to overcome the strong electrostatic attractions between the ions, so sodium chloride has a high melting point.

Why does MgO have a much higher melting point than NaCl?

Not all ionic lattices have equally strong attractions. A useful qualitative guide comes from Coulombic attraction:

The symbols q₁ and q₂ represent the ionic charges, while r represents the distance between the centres of the ions.

• NaCl contains Na⁺ and Cl⁻ ions.
• MgO contains Mg²⁺ and O²⁻ ions.
• The charge product is much greater in magnitude for Mg²⁺ and O²⁻.
• Mg²⁺ and O²⁻ also have relatively high charge densities.

The electrostatic attractions in MgO are therefore much stronger than those in NaCl. More energy is required to separate the ions, so MgO has a much higher melting point.

Why do ionic solids not conduct electricity?

Electrical conduction requires charged particles that can move through a substance. An ionic solid certainly contains charged particles, but its ions are held in fixed positions within the rigid lattice.

The ions can vibrate about their lattice positions, but they cannot move through the solid and transport charge from one place to another. Solid ionic compounds therefore do not conduct electricity.

Why do molten ionic compounds conduct?

When an ionic compound melts, the rigid lattice breaks down. The ions are no longer restricted to fixed lattice positions and can move through the liquid.

These mobile positive and negative ions carry electric charge, so the molten substance conducts electricity.

Why can aqueous ionic compounds conduct?

If an ionic compound dissolves in water, its ions become separated from the lattice and surrounded by water molecules. These hydrated ions can move through the solution and therefore carry charge.

Why are ionic crystals brittle?

In an undisturbed ionic lattice, positive ions are positioned next to negative ions. Suppose a sufficiently large force shifts one layer relative to another.

Ions of the same charge can then become aligned:

Strong electrostatic repulsion now acts between neighbouring positive ions and between neighbouring negative ions. The lattice splits rather than changing shape smoothly, so ionic crystals are brittle.

Are all ionic compounds soluble in water?

No. Water is polar and can form strong ion–dipole attractions with ions, so many ionic compounds dissolve in water. However, separating ions from an ionic lattice also requires energy.

Whether dissolution is favourable depends on the overall balance of interactions and energy changes. Some ionic substances therefore dissolve readily, whereas others, such as AgCl and BaSO₄, are only sparingly soluble.

Giant metallic lattices

A metal consists of a regular lattice of positive metal ions surrounded by delocalised electrons. These electrons originate from the outermost occupied shells of the metal atoms and are not associated with one particular ion.

Metallic bonding is the strong electrostatic attraction between the positive metal ions and the negatively charged delocalised electrons.

Why do metals conduct electricity?

Metals contain delocalised electrons that can move throughout the structure. When a potential difference is applied, the electrons undergo a net movement and carry electric charge.

Exam-style explanation: Copper contains mobile delocalised electrons. When a potential difference is applied, these electrons move through the giant metallic structure and carry electric charge.

Metals can therefore conduct while solid. Molten metals also conduct because the delocalised electrons remain mobile.

Why do metals conduct heat?

Mobile delocalised electrons can transfer kinetic energy rapidly through the structure. Vibrations of the closely packed metal ions also contribute to thermal energy transfer.

Metals are therefore generally good thermal conductors.

Why are metals malleable and ductile?

A malleable material can be hammered or rolled into sheets, whereas a ductile material can be drawn into wires.

When a force is applied to a metal, layers of positive metal ions can slide past one another. After the ions change position, they are still surrounded by and attracted to the sea of delocalised electrons.

Metallic bonding is therefore maintained as the structure changes shape. The metal bends or stretches instead of immediately fracturing.

Why do many metals have high melting points?

Melting requires sufficient energy to weaken the strong electrostatic attractions between positive metal ions and delocalised electrons. Many metals therefore have relatively high melting points.

However, metallic bond strength varies substantially from one metal to another. It is therefore inaccurate to say that every metal has a high melting point. Mercury, for example, is liquid at room temperature.

Metallic bonding tends to become stronger when:

• more electrons are delocalised per atom;
• the positive charge on the metal ion is greater; and
• the metal ion is smaller, giving a higher charge density.

These ideas help explain why magnesium generally has stronger metallic bonding than sodium.

Giant covalent structures

A giant covalent structure contains atoms joined by covalent bonds throughout an extended network. There are no separate, individual molecules within the structure.

Because strong covalent bonds extend throughout the structure, a large amount of energy is generally needed to separate the atoms. Giant covalent substances therefore tend to have very high melting points.

Diamond: a rigid three-dimensional network

Diamond is an allotrope of carbon. Each carbon atom forms four covalent bonds with four neighbouring carbon atoms in a tetrahedral arrangement.

The covalent bonds extend in three dimensions to produce a rigid giant network.

• Very high melting point: many strong C–C covalent bonds must be broken to separate the atoms.
• Extremely hard: strong directional covalent bonds extend throughout a rigid three-dimensional network.
• Does not conduct electricity: all four outer-shell electrons of each carbon atom are localised in covalent bonds, so there are no mobile charged particles.

Why is diamond hard? Each carbon atom is covalently bonded to four other carbon atoms in a rigid three-dimensional giant structure. The many strong covalent bonds resist deformation.

Graphite: strong layers with weak forces between them

Graphite is another allotrope of carbon. Each carbon atom forms three covalent bonds with three neighbouring carbon atoms, producing flat hexagonal layers.

Strong covalent bonds hold the carbon atoms together within each layer. Between neighbouring layers, however, there are much weaker instantaneous dipole–induced dipole (id-id) forces. These are one type of van der Waals force.

• Very high melting point: strong C–C covalent bonds within the layers must be broken to separate the carbon atoms.
• Soft and slippery: weak id-id forces between neighbouring layers allow the layers to slide over one another.
• Conducts electricity: each carbon atom contributes one electron to a delocalised electron system within the layer.

These delocalised electrons can move along the layers and carry electric charge. Electrical conduction is therefore much easier parallel to the layers than perpendicular to them.

What van der Waals force acts between graphite layers?

The weak attractions between neighbouring graphite layers are instantaneous dipole–induced dipole forces, usually shortened to id-id forces.

These arise because the electron distribution in one region can fluctuate momentarily, producing an instantaneous dipole. This dipole can distort the electron distribution in a neighbouring region and induce another dipole. The resulting opposite partial charges attract.

You may also encounter the name London dispersion forces in textbooks or other resources. This is not a different force. London dispersion forces are another name for instantaneous dipole–induced dipole forces.

Why does graphite conduct but diamond does not?

A carbon atom has four outer-shell electrons.

• In diamond, all four are used in four localised covalent bonds.
• In graphite, only three are used in localised covalent bonds.
• The remaining electron becomes delocalised within the layer.

Graphite therefore contains mobile delocalised electrons, whereas diamond does not. This difference in electronic structure explains their different electrical conductivities.

Silicon dioxide as a giant covalent structure

Solid silicon dioxide, SiO₂, also forms a giant covalent network. Each silicon atom is covalently bonded to four oxygen atoms, while each oxygen atom is covalently bonded to two silicon atoms.

The formula SiO₂ therefore gives the 1:2 ratio of silicon atoms to oxygen atoms within the extended structure. The solid does not consist of separate SiO₂ molecules.

Silicon dioxide has a high melting point because many strong Si–O covalent bonds must be broken to disrupt the network. It does not conduct electricity because it has no mobile ions or delocalised electrons.

Simple molecular substances

A simple molecular substance consists of separate molecules. Strong covalent bonds hold the atoms together within each molecule, while intermolecular forces act between neighbouring molecules.

At this point, it is useful to connect this chapter directly to the intermolecular forces studied earlier. In Reviewston, the main intermolecular attractions are organised as follows:

• Van der Waals forces include instantaneous dipole–induced dipole (id-id) forces and permanent dipole–permanent dipole (pd-pd) forces.
• Hydrogen bonding is treated separately as a particularly strong intermolecular attraction involving suitable O–H, N–H or F–H bonds.

How do the intermolecular forces fit together?

The type of intermolecular attraction present depends on the molecule.

In many comparison questions, we focus on the most important difference between the substances rather than listing every attraction present.

For example, when comparing CH₄ and H₂O, both substances experience id-id forces. However, H₂O also forms hydrogen bonds. It is this additional strong attraction that is especially important in explaining why water has a much higher boiling point than methane.

Why do simple molecular substances usually have lower melting and boiling points?

Intermolecular forces are generally much weaker than the ionic, metallic or covalent bonds that hold giant structures together.

Less energy is therefore required to separate the molecules. Simple molecular substances consequently tend to have lower melting and boiling points than giant structures.

When iodine melts, the I–I covalent bonds within I₂ molecules are not broken. Energy is used to overcome some of the instantaneous dipole–induced dipole forces between the separate I₂ molecules.

Why do different molecular substances have different boiling points?

A higher boiling point generally indicates that more energy is required to separate the molecules and therefore that the intermolecular attractions are stronger overall.

When comparing substances, first ask what intermolecular attractions are present.

• All molecules experience id-id forces.
• Polar molecules additionally experience pd-pd attractions.
• Molecules with suitable O–H, N–H or F–H bonds may additionally form hydrogen bonds.

The strength of id-id forces also depends on factors such as the size and polarisability of the electron cloud and the amount of contact between neighbouring molecules.

Larger, more polarisable molecules generally experience stronger id-id attractions because their electron clouds are more easily distorted. Molecular shape can also matter because greater surface contact allows stronger overall attractions between neighbouring molecules.

It is therefore too simplistic to say that all simple molecular substances have extremely low melting and boiling points. Large molecular substances can experience substantial intermolecular attractions and may be solids at room temperature.

Why do simple molecular substances usually not conduct electricity?

Most simple molecular substances consist of electrically neutral molecules. Their electrons are localised in covalent bonds or lone pairs, and they do not contain mobile ions or delocalised electrons.

They therefore do not normally conduct electricity in either the solid or liquid state.

A molecular substance may nevertheless produce ions when it reacts with water. For example, hydrogen chloride forms H₃O⁺ and Cl⁻ ions in aqueous solution:

The resulting solution conducts electricity because these ions are mobile. It is not the movement of neutral HCl molecules that carries charge.

Solubility of molecular substances

For a substance to dissolve, favourable attractions must form between the solute particles and solvent molecules. A useful first guideline is that substances tend to dissolve in solvents with which they can form similar or sufficiently strong intermolecular attractions.

• Polar molecules are generally more likely to dissolve in polar solvents.
• Non-polar molecules are generally more likely to dissolve in non-polar solvents.
• Molecules capable of hydrogen bonding with water may be relatively soluble in water, provided their non-polar regions are not too large.

Iodine, for example, is non-polar. Its molecules interact mainly through id-id forces. Iodine is therefore only slightly soluble in water but dissolves more readily in suitable non-polar solvents where favourable id-id attractions can form with the solvent molecules.

Why is ice less dense than liquid water?

Ice is a molecular solid. Each H₂O molecule contains strong O–H covalent bonds, while hydrogen bonds act between neighbouring water molecules.

Water molecules also experience van der Waals forces, as all molecules do. However, it is the extensive hydrogen-bonded arrangement that is especially important when explaining the unusual structure of ice.

In ordinary ice, the water molecules form an open, approximately tetrahedral hydrogen-bonded network. This arrangement contains considerable empty space.

When ice melts, some hydrogen bonds are disrupted and the rigid open structure partially collapses. The water molecules can move closer together on average.

For the same mass of H₂O, the more open structure of ice occupies a larger volume. Ice therefore has a lower density than liquid water and floats on it.

Comparing the four main structure types

Deducing structure from physical properties

Examination questions often provide physical-property data for an unfamiliar substance and ask you to deduce its likely structure. The goal is not to match one memorised property to one structure. Look for a pattern of evidence.

Evidence for a giant ionic structure

• High melting point.
• Does not conduct electricity as a solid.
• Conducts when molten.
• Conducts in aqueous solution if sufficiently soluble.
• Usually brittle.

Evidence for a giant metallic structure

• Conducts electricity as a solid.
• Conducts when molten.
• Usually conducts heat well.
• Malleable and ductile.
• Often has a moderate or high melting point.

Evidence for a giant covalent structure

• Very high melting point.
• Usually does not conduct electricity.
• Often very hard.
• Generally insoluble in common solvents.
• Graphite is an important conducting and soft exception.

Evidence for a simple molecular structure

• Usually relatively low melting and boiling points.
• Normally does not conduct as a solid or liquid.
• May be a gas, liquid or soft solid at room temperature.
• Solubility depends strongly on intermolecular interactions.

Worked example: identifying an ionic substance

Substance X is a brittle solid with a high melting point. It does not conduct electricity as a solid but conducts when molten.

The high melting point suggests strong attractions in a giant structure. The substance contains charged particles because the liquid conducts. However, those particles cannot move while X is solid.

Conclusion: X is likely to have a giant ionic lattice. Its ions are fixed in the solid but become mobile when the lattice melts.

Worked example: identifying a metal

Substance Y conducts electricity as a solid, can be hammered into sheets and can be drawn into wires.

Solid-state conduction suggests mobile charged particles, while malleability and ductility show that layers can change position without the structure fracturing.

Conclusion: Y is likely to have a giant metallic structure containing mobile delocalised electrons.

Worked example: why high melting point alone is insufficient

Substance Z has a very high melting point and does not conduct electricity as a solid.

These observations alone do not distinguish an ionic solid from a non-conducting giant covalent solid. Further information is required.

If Z conducts when molten, an ionic structure is likely because its ions have become mobile. If there are still no mobile charged particles, a giant covalent structure may be more likely.

Common mistakes and why they are wrong

Bonding and physical-properties checklist