Metallic bonding
Metals are made of positive ions packed in a regular lattice. Each atom lets go of its outer (valence) electrons, and those electrons are no longer tied to one atom, they drift freely through the whole structure. We picture this as positive ions sitting in a sea of delocalised electrons.
- Metallic bond
- The strong electrostatic attraction between a lattice of positive metal ions and the sea of delocalised electrons that surrounds them.
"Delocalised" just means the electrons are not fixed in one place. Because the attraction pulls in every direction and holds the whole lattice together, metallic bonds are strong, which is why most metals have high melting points.
Why metals conduct and bend
The sea of electrons explains the two properties you are most often asked about.
Electrical and thermal conductivity. An electric current is a flow of charge. Because the delocalised electrons are free to move through the lattice, they can carry charge from one end of the metal to the other, so metals conduct electricity. Those same moving electrons also transfer energy quickly, so metals conduct heat well.
Malleability and ductility. Malleable means it can be hammered into shape; ductile means it can be drawn into wires. When a force is applied, the layers of positive ions slide over one another. The electron sea simply flows with them and keeps holding the lattice together, so the metal changes shape without shattering.
Explain, in terms of structure, why copper is used for electrical wiring but a block of solid ionic salt is not.
Giant covalent structures
In a giant covalent (macromolecular) structure, huge numbers of atoms are joined by a continuous network of strong covalent bonds. There are no separate molecules, the whole crystal is effectively one giant molecule. Breaking it apart means breaking many strong bonds, so these substances have very high melting and boiling points.
The three you must know are all built from just a few elements:
| Structure | Bonding detail | Key properties |
|---|---|---|
| Diamond (carbon) | Each carbon bonds to 4 others in a rigid 3D lattice | Extremely hard, very high melting point, does not conduct electricity |
| Graphite (carbon) | Each carbon bonds to 3 others in flat layers; one delocalised electron per atom | Soft and slippery, conducts electricity, high melting point |
| Silicon dioxide, SiO2 | Each silicon bonds to 4 oxygens, each oxygen to 2 silicons | Hard, very high melting point, does not conduct |
Diamond is hard because every atom is locked by four strong bonds with no weak points. Graphite is the interesting one: its atoms form layers, and the fourth outer electron of each carbon becomes delocalised. Those free electrons let graphite conduct electricity (unusual for a non-metal), while weak forces between the layers let them slide, which is why graphite is soft enough to use in pencils and as a lubricant. Silicon dioxide (the main compound in sand and quartz) behaves much like diamond because it shares the same kind of rigid 3D network.
Diamond and graphite are both just carbon
Same element, different arrangement. When you meet two forms of one element with different properties, the answer is always in the structure, not the atoms themselves.
Intermolecular forces
Simple molecular substances, such as water (H2O), carbon dioxide (CO2) and iodine (I2), are made of small molecules. Inside each molecule the atoms are joined by strong covalent bonds. But between the molecules there are only weak forces of attraction, called intermolecular forces.
- Intermolecular forces
- The weak forces of attraction acting between separate molecules, much weaker than the covalent bonds within a molecule.
When you melt or boil a simple molecular substance, you only need to overcome those weak intermolecular forces, not the strong covalent bonds. That takes little energy, so these substances have low melting and boiling points and many are liquids or gases at room temperature. A common mistake is to say melting "breaks the covalent bonds", it does not; the molecules stay whole and simply move apart.
They also do not conduct electricity, because the molecules are neutral overall: there are no free electrons and no ions to carry charge.
Common slip
Low melting point does not mean weak covalent bonds. The covalent bonds are strong; it is the forces between molecules that are weak. Always name which force is being overcome.
Comparing the structures
Almost every "explain the property" question comes down to identifying the structure first, then reasoning from it.
| Structure | Melting point | Conducts? | Why |
|---|---|---|---|
| Metallic | High | Yes (solid & liquid) | Delocalised electrons carry charge |
| Giant covalent | Very high | Only graphite | Continuous network of strong bonds |
| Simple molecular | Low | No | Weak intermolecular forces; no charged particles free to move |
Where this is assessed
Linking structure to properties is classic Criterion A (Knowing and understanding). When you interpret melting-point or conductivity data to deduce a structure, you are also showing Criterion B (Inquiring and designing) thinking.
Check yourself
1. Why can graphite conduct electricity but diamond cannot? +
In graphite each carbon bonds to only three others, so one outer electron per atom becomes delocalised and free to move, carrying charge. In diamond every carbon uses all four outer electrons in bonds, so there are no free electrons to carry a current.
2. Carbon dioxide is a gas at room temperature but silicon dioxide melts above 1600 °C. Explain the difference. +
CO2 is simple molecular: only weak intermolecular forces need to be overcome to separate the molecules, so it boils very easily. SiO2 is giant covalent: melting means breaking a continuous network of strong covalent bonds, which needs a huge amount of energy, so its melting point is much higher.
3. Why can a metal be bent into shape without breaking? +
The layers of positive ions can slide over each other when a force is applied, and the sea of delocalised electrons moves with them, keeping the lattice bonded together throughout. So the metal changes shape rather than shattering.
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