Covalent Bonding and Molecular Substances
In 1926, Gilbert Lewis at UC Berkeley mapped the first covalent bond, showing two atoms sharing electrons rather than transferring them, explaining why water boils at 100 °C not −80 °C.
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Q1 · Water, oxygen, and glucose are all essential for life, what do you think these molecules have in common at the atomic level, and how might their atoms be held together?
Q2 · If ionic compounds form when electrons are transferred between atoms, why do you think some atoms prefer to share electrons instead?
● Know
- How covalent bonds form by sharing electron pairs
- The difference between single, double, and triple bonds
- The properties of molecular substances (low melting points, poor conductors)
● Understand
- Why molecular substances have LOW melting points (intermolecular forces break, not covalent bonds)
- Why covalent compounds generally do not conduct electricity
- The difference between intra- and inter-molecular forces
● Can do
- Draw dot-and-cross diagrams for simple molecules (H₂, O₂, H₂O, CH₄)
- Predict properties of molecular substances
- Explain the melting misconception for covalent compounds
Bend a copper wire back and forth in your hands: it flexes smoothly and never shatters, because the atoms inside can slide past each other while still held by the same sea of electrons that surrounds them. In a metallic solid, metal atoms release their valence electrons, becoming positive metal cations. These cations arrange themselves in a regular, repeating pattern, the metal lattice. The released valence electrons are not attached to any individual atom; instead they form a delocalised 'sea of electrons' that moves freely throughout the entire lattice. The metallic bond is the electrostatic attraction between the positive metal cations and the surrounding sea of negative electrons. Because the electrons are mobile and not fixed in directional bonds, this model explains most characteristic metal properties.
The sea of electrons model also explains metallic lustre. When light hits the metal surface, the free electrons can absorb the photon energy and re-emit it as light across the whole visible spectrum, this is why polished metals appear shiny and reflective. Silver reflects about 97% of visible light, which is why mirrors are made by vacuum-depositing a silver film onto glass. The electron sea is what gives metals their characteristic appearance as well as their electrical and thermal behaviour.
Copper (Cu) has 1 valence electron per atom. In a copper wire, roughly 8.5 × 10²⁸ free electrons per cubic metre form the 'sea'. When a voltage is applied, this sea drifts slowly (about 0.1 mm/s average drift velocity) but the electric signal travels at close to the speed of light, like pushing a hose already full of water.
Australia produces about 900,000 tonnes of copper per year, mostly from Olympic Dam in South Australia (BHP) and the Cadia Valley mine in NSW (Newcrest). The delocalised electron sea in copper is the reason this metal is chosen over cheaper alternatives for almost all electrical wiring in Australian homes and industry.
Both electrical and thermal conductivity in metals are explained by the delocalised electrons. For electrical conductivity: when a voltage is applied across a metal, the free electrons in the sea drift collectively from the negative terminal to the positive terminal, carrying electric charge. No ions need to move, only the electron sea. This is why metals conduct electricity as solids, unlike ionic compounds which need dissolved or molten ions.
For thermal conductivity: when one end of a metal is heated, the electrons in that region gain kinetic energy and move faster. They rapidly collide with electrons throughout the lattice, transferring kinetic energy, heat, through the metal far faster than atomic vibration alone could achieve. This is why a metal spoon left in a hot pot becomes hot all the way to the handle quickly, while a wooden spoon does not. Silver conducts better than iron because silver has more free electrons per atom available to carry both charge and heat, the electron sea is denser.
Silver (thermal conductivity 429 W/m·K) conducts heat roughly 8 times better than iron (50 W/m·K). In a silver ring worn on a finger, heat equalises to body temperature in about 2 seconds; an iron ring would take about 16 seconds. Jewellers prefer silver also because it can be cast, polished, and hallmarked, all properties related to its metallic bonding.
Aluminium (electrical conductivity 3.8 × 10⁷ S/m) is used for high-voltage transmission lines in Australia rather than copper, it is lighter for the same conductance, allowing longer spans between towers. The AEMO (Australian Energy Market Operator) network uses thousands of kilometres of aluminium transmission cable to carry power from Queensland coal plants and South Australian wind farms across the grid.
Malleability is the ability to be hammered into sheets; ductility is the ability to be drawn into wires. Both are explained by the sea of electrons model. When a force is applied to a metal, layers of cations slide past each other. In an ionic lattice, when layers shift, like charges align and repel, shattering the crystal. In a metal, the electron sea simply flows around the shifting cation layers, maintaining the attractive force throughout the deformation. The metal can change shape permanently without breaking because the bond is non-directional: it doesn't matter which cations are next to each other, the sea holds them all.
This is why metals can be shaped by rolling, pressing, drawing, and forging, all industrial processes that exploit metallic malleability. Copper can be drawn into wires with diameters less than 0.1 mm. Gold can be beaten into sheets 0.1 µm thick (thinner than a cell membrane). Iron can be rolled into steel beams kilometres long. In every case, the electron sea is the silent enabler, it is what makes metals the most versatile structural materials available to engineers.
BlueScope's hot rolling mill at Port Kembla presses steel slabs at 1200 °C into steel sheets 0.5 mm thick at 80 km/h. The metallic bonding allows the steel to be deformed by a factor of 200:1 in thickness without fracturing, directly because the electron sea flows around displaced cation layers.
BlueScope Steel's Port Kembla steelworks in NSW is Australia's largest steel producer, rolling around 2.6 million tonnes of hot-rolled steel per year. The malleability of steel (a metallic alloy of iron and carbon) is fundamental to every product BlueScope makes, from Colorbond roofing to heavy structural sections used in Australian construction.
Malleability is the ability to be hammered into , while ductility is the ability to be drawn into wires. Both are explained by the model. When a force is applied, layers of slide past each other. In a metal the electron sea flows around the shifting layers, maintaining the between particles. The metallic bond is , so the metal changes shape without breaking.
At the start of this lesson, you heard about copper wire, which bends, stretches, and hammers into shape without breaking, versus a ceramic cup that shatters when you drop it. Both are solid and strong, but their bonding is completely different. Metallic bonding explains why metals are ductile and conductive in a single elegant model.
Now that you've worked through the lesson, can you explain what "sea of delocalised electrons" means and how it accounts for both the conductivity and the bendability of metals? How does this compare to the ionic and covalent bonding you studied in earlier lessons?
Q1. Explain how a covalent bond forms between two hydrogen atoms. Draw a dot-and-cross diagram to support your explanation.
Q2. A student claims that when ice melts, covalent bonds in water molecules are broken. Evaluate this claim and provide a correct explanation of what actually happens.
Q3. Compare the structure and bonding of sodium chloride (ionic) and water (covalent). Use this comparison to explain why these substances have such different melting points.