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Metallic bonding is the third major type of bonding you need to understand. It is found in metals and alloys and explains their distinctive properties: high melting points, electrical and thermal conductivity, malleability, and ductility. The model of metallic bonding is fundamentally different from both ionic and covalent bonding.
In a metal, the atoms are packed closely together in a regular lattice arrangement. Each metal atom loses its outer shell electron(s), becoming a positive ion (cation). These electrons do not belong to any particular atom — they become delocalised, forming a "sea" of electrons that spreads throughout the entire structure.
The metallic bond is the strong electrostatic attraction between the positive metal ions and the delocalised electrons. This is the definition you need to learn precisely.
Notice the key features of this model:
The strength of metallic bonding depends on three interrelated factors:
1. Number of delocalised electrons per atom: The more electrons each atom contributes to the sea, the stronger the bonding. Sodium (Group 1) contributes 1 electron, magnesium (Group 2) contributes 2, and aluminium (Group 3) contributes 3. More delocalised electrons means greater charge density in the electron sea and stronger attraction.
2. Charge on the metal ion: Directly related to the number of electrons lost. Na⁺ has a 1+ charge, Mg²⁺ has 2+, and Al³⁺ has 3+. Higher charge means stronger electrostatic attraction to the electron sea.
3. Ionic radius: Smaller ions mean the positive charge is closer to the delocalised electrons, increasing the strength of attraction. As you go down a group, the ions get larger and the metallic bonding generally gets weaker.
Comparing the Period 3 metals illustrates how bonding strength changes:
| Metal | Ion | Electrons donated | Ionic radius / pm | Melting point / °C |
|---|---|---|---|---|
| Na | Na⁺ | 1 | 102 | 98 |
| Mg | Mg²⁺ | 2 | 72 | 649 |
| Al | Al³⁺ | 3 | 53 | 660 |
Going from sodium to aluminium:
All three factors work together to strengthen the metallic bonding from left to right across the period, which is reflected in the increasing melting points.
Note that aluminium's melting point (660°C) is only slightly higher than magnesium's (649°C) despite having a higher charge. This is because aluminium has a different crystal structure (face-centred cubic vs hexagonal close-packed for magnesium), which affects the efficiency of packing.
A large amount of energy is required to overcome the strong electrostatic attractions between the positive ions and the delocalised electrons. The stronger the metallic bonding (more delocalised electrons, higher charge, smaller ions), the higher the melting point.
Transition metals generally have very high melting points because they can contribute d electrons to the sea of delocalised electrons, in addition to their s electrons. Tungsten, for example, melts at 3422°C.
When a voltage is applied across a metal, the delocalised electrons can flow through the structure towards the positive terminal. This movement of charged particles is an electric current. The key point is that the electrons are already free to move — they do not need to be released from bonds first.
Metals conduct in the solid state (unlike ionic compounds, which must be melted or dissolved). Electrical conductivity decreases with increasing temperature because the positive ions vibrate more vigorously, scattering the delocalised electrons and impeding their flow.
The delocalised electrons also explain why metals are good thermal conductors. When one end of a metal is heated, the delocalised electrons in that region gain kinetic energy and move rapidly through the structure, transferring energy to other parts of the metal. The lattice vibrations of the positive ions also contribute, but electron movement is the dominant mechanism.
Malleability means the metal can be hammered into sheets. Ductility means it can be drawn into wires. These properties arise because the layers of metal ions can slide over each other without breaking the metallic bond. As layers shift, the delocalised electron sea simply adjusts to the new positions of the ions, maintaining the bonding throughout.
This is fundamentally different from ionic compounds, where displacing layers brings like charges together and causes shattering. In metals, there are no alternating charges — the positive ions can be in any arrangement relative to each other, and the electron sea holds them together regardless.
An alloy is a mixture of two or more metals (or a metal with a non-metal like carbon). Alloys are generally harder and have higher melting points than pure metals.
The reason is that the atoms of the different elements have different sizes. When a differently-sized atom is introduced into the lattice, it disrupts the regular arrangement of layers. This makes it harder for the layers to slide over each other, increasing hardness and reducing malleability.
Examples:
Going down Group 1 (Li → Na → K → Rb → Cs):
The result is that the metallic bond strength decreases because the larger ionic radius means the positive charge is further from the delocalised electrons. This is clearly reflected in the melting points: Li (181°C), Na (98°C), K (64°C), Rb (39°C), Cs (28°C).
| Feature | Ionic | Covalent (simple) | Metallic |
|---|---|---|---|
| Particles | Positive and negative ions | Atoms sharing electrons | Positive ions + delocalised electrons |
| Force | Electrostatic (ion–ion) | Electrostatic (nuclei–shared pair) | Electrostatic (ion–electron sea) |
| Directional? | No | Yes | No |
| Structure | Giant lattice | Molecules or giant network | Giant lattice |
| Melting point | High | Low (molecules) / very high (giant) | Generally high |
| Conductivity | Molten/dissolved only | None (usually) | Solid and molten |
| Malleability | Brittle | N/A (molecules) / brittle (giant) | Malleable and ductile |
A substance has a melting point of 649°C, conducts electricity as a solid and as a liquid, is malleable, and is insoluble in water. Identify the bonding type.
This is a metal. The melting point of 649°C matches magnesium.
The sea-of-electrons model is useful but has limitations. It does not explain why some metals are better conductors than others, or why transition metals have particularly high melting points. A more sophisticated treatment uses band theory, where the atomic orbitals of all the atoms in the metal merge to form continuous energy bands. However, the simple model is sufficient for A-Level and explains the key trends effectively.
This lesson sits inside Edexcel 9CH0 Topic 2 — Bonding and Structure, with the metallic-bonding content concentrated in sub-topic 2.3 (the metallic bonding model and the physical properties it explains: electrical conductivity, thermal conductivity, malleability, ductility, and high melting point). Although Topic 2 introduces the model, metallic bonding resurfaces repeatedly across the linear specification. Paper 1 (Advanced Inorganic and Physical Chemistry) Topic 4 uses the same model to explain the rise in melting points across Period 3 from Na to Al, and to contrast that rise with the abrupt switch to giant covalent (Si) and simple molecular (P, S, Cl) structures. Paper 1 Topic 15 revisits metallic bonding for the d-block: transition-metal melting points and densities are interpreted in terms of additional d-electron contributions to the delocalised sea. Paper 3 (General and Practical Principles) examines this content synoptically, typically by asking candidates to compare ionic, giant covalent and metallic structures from a data table, then justify physical-property differences from the bonding model. The connections to Topic 1 (ionisation energies — why metals lose electrons easily in the first place) and Topic 3 (redox reactivity — how readily those outer electrons are surrendered) are also routinely tested (refer to the official Pearson Edexcel specification document for exact wording).
Question (7 marks): The melting points of selected Period 3 elements are: Na 98 °C, Mg 650 °C, Al 660 °C, Si 1414 °C, P 44 °C. Explain the pattern in melting points across these five elements, with reference to bonding and structure throughout. (7)
Solution with mark scheme:
Step 1 — Na to Mg to Al: a metallic-bonding argument.
Sodium, magnesium and aluminium are all metals with giant metallic lattices consisting of positive ions in a delocalised sea of electrons. Across this sub-section of the period:
All three factors strengthen the electrostatic attraction between cations and delocalised electrons, so progressively more energy is required to overcome the metallic bond.
M1 — explicitly identifies metallic bonding in Na, Mg, Al. M1 — references at least two of {cation charge, number of delocalised electrons, ionic radius}. A1 — links these to a stronger metallic bond and hence a higher melting point Na → Al.
Step 2 — the jump to silicon.
Silicon is not a metal: it is a giant covalent (macromolecular) solid with each Si atom covalently bonded to four others in a tetrahedral lattice. To melt silicon, strong covalent bonds throughout the lattice must be broken, requiring far more energy than overcoming metallic attractions, hence the very high melting point of 1414 °C.
B1 — identifies Si as giant covalent and refers to breaking covalent bonds.
A common slip here is to say "Si melts higher than Al because metallic bonding in Si is stronger". This is wrong: silicon does not exhibit metallic bonding at all. Examiners explicitly penalise this misclassification.
Step 3 — the collapse at phosphorus.
Phosphorus exists as discrete P₄ molecules held together in the solid only by weak London (dispersion) forces between molecules. The covalent bonds inside each P₄ tetrahedron are not broken on melting — only the intermolecular forces are. These forces are far weaker than either metallic or giant covalent bonding, so the melting point collapses to 44 °C.
M1 — identifies P as simple molecular (P₄). A1 — states that only weak London forces are overcome on melting (the intramolecular covalent bonds remain intact).
Total: 7 marks (M3 A2 B1 + 1 implicit linking mark).
A perfect script would also note that Mg → Al is a much smaller rise than Na → Mg (only ~10 °C), because aluminium's face-centred-cubic packing differs from magnesium's hexagonal close packing, partly offsetting the bonding-strength gain.
Question (6 marks): Aluminium is used for overhead high-voltage power lines, while copper is used for household wiring. Both metals are excellent electrical conductors. Using your knowledge of metallic bonding, structure, and the physical properties of metals, explain why each metal is preferred for its specific application. (6)
Mark scheme decomposition by AO:
| AO | Marks | Earned by |
|---|---|---|
| AO1 | 2 | Stating that both metals conduct electricity via delocalised electrons free to move through the giant metallic lattice; recognising both are malleable and ductile because cation layers can slide past one another while still held by the delocalised electron sea. |
| AO2 | 3 | Applying the bonding model to context: aluminium has lower density (lighter cations, less massed lattice) so overhead cables sag less under their own weight over long spans; copper has higher conductivity per unit cross-section (more delocalised electrons per unit volume in its denser lattice) so thinner wires suffice indoors where weight is not the constraint. |
| AO3 | 1 | Evaluating a trade-off: aluminium is cheaper and lighter but a poorer conductor than copper, so overhead cables use a thicker aluminium core (often reinforced with steel) — a choice that makes engineering sense only because pylons can support the bulk. Indoors, copper's higher conductivity-per-volume wins because runs are short, weight is irrelevant, and cable bulk inside walls matters. |
Total: 6 marks (AO1 = 2, AO2 = 3, AO3 = 1). Note the AO3 mark depends on an explicit trade-off — a candidate who simply lists properties of each metal without comparing the engineering constraints will score AO1 + AO2 only.
Metallic bonding connects to several other parts of the 9CH0 specification:
Topic 4 — Inorganic Chemistry (Group 1 and 2 trends): the decrease in melting point down Group 1 (Li 181 °C → Cs 28 °C) and the more complex trend down Group 2 follow directly from the metallic-bonding model. Each atom still donates the same number of electrons (1 in Group 1, 2 in Group 2), the cation charge is unchanged, but the ionic radius increases down the group, so the delocalised electrons are further from the nuclear charge and the metallic bond weakens. This is the same argument used in this lesson, transplanted into a vertical group context.
Topic 4 — Period 3 structure-type sweep: the synoptic question "explain Period 3 melting points from Na to Ar" requires distinguishing metallic (Na, Mg, Al), giant covalent (Si), and simple molecular (P₄, S₈, Cl₂, Ar atoms) structures. The metallic-bonding model accounts for only the first three; getting the structure-type assignment right is the critical first step.
Topic 15 — d-block transition metals: transition metals exhibit enhanced metallic bonding because both 4s and 3d electrons can become delocalised, contributing more electrons per atom to the sea. This is why W melts at 3422 °C and why d-block densities far exceed those of s-block neighbours. The simple model from this lesson is the foundation; Topic 15 just extends the electron count.
Topic 12 — Acid–base equilibria of hydrated metal cations: ions such as [Fe(H₂O)₆]³⁺ are weakly acidic. The acidity depends on the charge density of the bare cation (charge ÷ ionic radius) — exactly the same property that controls metallic-bond strength. A small, highly charged cation polarises coordinated water strongly and bonds delocalised electrons strongly: one underlying property, two surface manifestations.
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