Metallic bonding and giant metallic lattices

A metal is not made from separate metal molecules. In a solid metal, many positive metal ions are arranged in a giant metallic lattice. The outer-shell electrons from the metal atoms are delocalised, so they are not attached to one particular atom and can move through the lattice.

The word strong matters. A lot of energy is needed to overcome metallic bonding, so many metals have relatively high melting and boiling points compared with simple molecular substances.

The word mobile also matters. The positive metal ions vibrate about fixed positions in the solid lattice, but the delocalised electrons can move. These mobile electrons carry charge, so metals conduct electricity in the solid and liquid states.

At this point, keep the model simple. You do not need details of cubic packing, hexagonal packing or unit cells for this lesson.

Worked reasoning: why aluminium has a high melting point for a metal

Aluminium has a giant metallic lattice. Metallic bonding is the strong electrostatic attraction between positive metal ions and delocalised electrons. A large amount of energy is needed to overcome these attractions, so aluminium has a high melting point compared with simple molecular elements such as phosphorus or chlorine.

Giant covalent lattices

A giant covalent lattice is a network of atoms joined by strong covalent bonds. The structure extends throughout the solid, so there are no small separate molecules to pull apart when it melts.

Diamond, graphite, graphene and silicon are the OCR examples for this lesson.

Diamond

In diamond, each carbon atom forms covalent bonds to four other carbon atoms. This creates a rigid three-dimensional giant covalent lattice. There are no delocalised electrons that can move through the structure, so diamond does not conduct electricity.

Silicon

Silicon forms a giant covalent lattice similar in idea to diamond: silicon atoms are bonded to other silicon atoms by strong covalent bonds throughout the solid. Silicon has a very high melting point because many strong covalent bonds must be broken before the lattice can melt. You do not need a semiconductor explanation here.

Graphite

In graphite, each carbon atom forms covalent bonds to three other carbon atoms in layers. One outer-shell electron from each carbon atom becomes delocalised. These delocalised electrons can move through the layers, so graphite conducts electricity. Graphite is still a giant covalent structure because there are strong covalent bonds throughout each layer.

Graphene

Graphene is a single layer of carbon atoms in a giant covalent network. Like graphite, each carbon atom is bonded to three other carbon atoms and has delocalised electrons. This bonding model explains why graphene is strong and conducts electricity.

Explaining properties from structure

For this topic, a good explanation uses three questions in order:

1. What particles are present?
2. What forces or bonds hold the particles together?
3. Which particles or electrons can move?

That sequence works for melting point, boiling point, solubility and electrical conductivity.

Melting and boiling points

Giant metallic lattices usually have high melting and boiling points because metallic bonding is strong. Giant covalent lattices usually have very high melting and boiling points because strong covalent bonds must be broken throughout the network.

Solubility

Giant metallic and giant covalent lattices are generally insoluble in common solvents such as water. The solvent would have to overcome many strong attractions or bonds in the giant lattice. Some metals react with water or acids, but that is not the same as simply dissolving.

Electrical conductivity

Electrical conductivity needs mobile charged particles. In metals, the mobile charged particles are delocalised electrons. In graphite and graphene, delocalised electrons also carry charge. In diamond, there are no mobile delocalised electrons, and the atoms are not charged ions, so it does not conduct electricity.

Melting point trends across Periods 2 and 3

Across Periods 2 and 3, the structure changes from giant metallic to giant covalent to simple molecular. The melting point trend follows that structural change.

For the metals at the start of a period, melting points are explained by metallic bonding. Across the Period 3 metals, sodium, magnesium and aluminium, metallic bonding gets stronger because there are more delocalised electrons and stronger attraction between the positive ions and the delocalised electrons. Aluminium has a higher melting point than sodium for this reason.

The giant covalent elements have very high melting points. Carbon in Period 2 and silicon in Period 3 form giant covalent lattices, so strong covalent bonds must be broken. Silicon has the highest melting point in Period 3 for this reason.

The later non-metals have simple molecular lattices. Nitrogen, oxygen, fluorine and chlorine exist as small molecules such as N2, O2, F2 and Cl2. Phosphorus and sulfur are often treated as P4 and S8 in this melting-point context. Argon and neon are monatomic. These substances have low melting points because only weak induced dipole-dipole interactions between molecules or atoms are overcome when they melt.

Important contrast

When P4, S8 or Cl2 melts, the covalent bonds inside the molecules are not broken. The weak intermolecular forces between molecules are overcome. This is a common place to lose marks.

Worked example: explaining Al, Si, P4 and S8 data

Suppose a question gives these melting points:

A strong answer would reason like this:

Aluminium has a giant metallic lattice. Strong metallic bonding is the electrostatic attraction between positive metal ions and delocalised electrons, so a lot of energy is needed to overcome it. Silicon has a giant covalent lattice, with each silicon atom bonded into a network by strong covalent bonds, so even more energy is needed to melt it. P4 and S8 are simple molecular substances. They have much lower melting points because only weak induced dipole-dipole interactions between molecules are overcome. S8 has a higher melting point than P4 because S8 has more electrons and stronger induced dipole-dipole interactions.

Notice the order of the explanation: structure type, bonding or force, then energy needed. That is the pattern to repeat.

Exam precision and common traps

The safest explanation is a chain:

structure -> particles -> bonding or force -> mobility or energy needed -> property

Use exact terms. Say "positive ions and delocalised electrons" for metallic bonding. Say "strong covalent bonds between atoms" for giant covalent lattices. Say "weak induced dipole-dipole interactions between molecules" for simple molecular elements.

Avoid these traps:

• Do not say that a simple molecular substance has a low melting point because it has weak covalent bonds. Its covalent bonds are strong; its intermolecular forces are weak.
• Do not say that graphite conducts because it has free ions. It conducts because it has delocalised electrons.
• Do not say that diamond is one molecule in the ordinary small-molecule sense. It is a giant covalent lattice of many carbon atoms.
• Do not explain Period 2 or Period 3 melting-point trends using ionisation energy, atomic radius alone or reactivity. The required explanation is structure and bonding.
• Do not add metal packing details. They are outside the boundary for this lesson.

Mini worked answer: fixing a weak response

Weak response: "Silicon has a high melting point because it has strong intermolecular forces."

Improved response: "Silicon has a giant covalent lattice. Strong covalent bonds between silicon atoms must be broken to melt the lattice, so a large amount of energy is needed."

The improved version names the structure, names the bond, and links it to energy needed.