Why can metals conduct and bend?
Lesson 1.8.3 listed what metals have in common: they are shiny, conduct electricity and heat, and can be shaped. Gold can be hammered into leaf hundreds of times thinner than a hair, and copper can be drawn into thin wire. Yet lesson 2.4.1 showed that ionic crystals shatter when struck. Both are built from positively charged particles, so why don’t metals shatter? The answer lies in the special kind of bonding inside a metal.
Metallic bonding: the sea of electrons
Metal atoms have few valence electrons (usually 1 to 3) and low ionization energies, so they let go of them easily. In a metal, those released valence electrons no longer belong to any one atom; they move freely through the whole piece of metal, shared by all the metal cations. This is called the sea of electrons电子气. The strong attraction between the metal cations and the sea of electrons is the metallic bond金属键.
Compared with the last two lessons:
| Ionic bond | Covalent bond | Metallic bond | |
|---|---|---|---|
| Particles | Cations and anions | Atoms | Metal cations and free electrons |
| Where the electrons go | Transferred from metal to nonmetal | Shared by two atoms | Shared by the whole metal |
| Usually found | Between reactive metals and nonmetals | Between nonmetal atoms | In metals and alloys |
Explaining metallic properties
- Electrical conductivity: with a voltage applied, the free electrons drift in one direction, carrying a current.
- Thermal conductivity: the moving electrons keep colliding with the cations, carrying energy quickly from hot regions to cold ones.
- Malleability and ductility: under force, layers of metal atoms can slide past each other, but the sea of electrons holds every layer together, so the metallic bonding isn’t broken. Ionic crystals are different: once a layer of ions shifts, like charges meet, repel, and the crystal shatters.
- Luster: the free electrons absorb visible light of every wavelength and re-emit it, so most metals look silvery and shiny.
Start with the metal: connect the battery, then hit it with the hammer. Then switch to the ionic crystal NaCl, do the same two things, and compare:
The strength of metallic bonding
The larger the charge on the metal cations (the more valence electrons) and the smaller the cations, the stronger their attraction to the sea of electrons. Stronger metallic bonding means higher melting and boiling points and a harder metal. Boiling separates the metal atoms completely, so the boiling point is the better guide to bond strength.
| Metal | Valence electrons | Melting point | Boiling point |
|---|---|---|---|
| Na | 1 | 97.8 °C | 883 °C |
| Mg | 2 | 650 °C | 1090 °C |
| Al | 3 | 660.3 °C | 2519 °C |
Down a group, the metal atoms get larger and the metallic bonding weakens. In lesson 2.3.6, the alkali-metal melting points fell from 180.5 °C for lithium to 28.4 °C for cesium.
Alloys
An alloy合金 is a metallic material made by melting two or more metals (or a metal and a nonmetal) together. AP sorts alloys into two types by structure:
- In a substitutional alloy替代型合金, atoms of similar size replace some of the host metal’s atoms in the lattice. In brass, zinc atoms replace some copper atoms.
- In an interstitial alloy间隙型合金, much smaller atoms fit into the gaps between the host atoms. In steel, carbon atoms sit between the iron atoms.
Either way, the extra atoms break up the neat layers and make them harder to slide, so alloys are usually harder than the pure metal. The effect is especially strong in interstitial alloys: steel is much harder than pure iron.
Example 1: Copper versus salt
Why can a lump of copper be hammered into a sheet, while a salt crystal shatters?
- Bonding in copper. Copper is a metal: its cations sit in a sea of electrons, and metallic bonds have no fixed direction.
- What hammering does to copper. After layers slide, they are still surrounded by the sea of electrons, and the cations still attract it, so the metallic bonding is intact; the copper just changes shape.
- Bonding in salt. In salt, Na⁺ and Cl⁻ alternate, held by ionic bonds, with every ion surrounded by opposite charges.
- What hammering does to salt. Once a layer slides by one position, Na⁺ faces Na⁺ and Cl⁻ faces Cl⁻; the like charges repel, and the crystal splits along that layer.
- Conclusion. The difference is direction: metallic bonding has none, so shifted layers still hold together; ionic bonding needs opposite charges side by side, so a shift turns attraction into repulsion.
Example 2: Comparing Na, Mg, and Al
Sodium, magnesium, and aluminum are all in period 3. Compare the strength of their metallic bonding, and check with data.
- What sets the strength. The more charge on the metal cations (the more electrons they add to the sea) and the smaller they are, the more strongly cations and electrons attract.
- Compare charges. Na, Mg, and Al have 1, 2, and 3 valence electrons, forming Na⁺, Mg²⁺, and Al³⁺.
- Compare sizes. Across a period, the ions get smaller: Na⁺ 102 pm, Mg²⁺ 72 pm, Al³⁺ 54 pm.
- Conclusion. Both factors point the same way: Na < Mg < Al.
- Check with data. The boiling points are 883 °C, 1090 °C, and 2519 °C. The melting points follow the same order, though Mg and Al differ by only 10 °C; melting also depends on the crystal structure.
- Copper (Cu)
A metal: copper cations in a sea of electrons.
- Sodium chloride (NaCl)
A reactive metal and a reactive nonmetal: Na⁺ and Cl⁻ ions.
- Carbon dioxide (CO₂)
Two nonmetals forming molecules with covalent C=O bonds.
- Iron (Fe)
A metal, held together by metallic bonding.
- Magnesium oxide (MgO)
Mg²⁺ and O²⁻ ions held by ionic bonds.
- Water (H₂O)
The O–H bonds are covalent.
- Brass (a copper–zinc alloy)
An alloy keeps a metallic structure with metallic bonding.
- Potassium bromide (KBr)
K⁺ and Br⁻ ions held by ionic bonds.
- Chlorine (Cl₂)
Two chlorine atoms share a pair of electrons: a covalent bond.
Common mistakes
Mistake 1: “A metallic bond is a pair of electrons shared by two neighboring metal atoms.” A covalent shared pair belongs to just two atoms; the free electrons in a metal belong to the whole piece and move among all the cations.
Mistake 2: “The electrons in a metal move only when a current flows.” Free electrons are always moving randomly. A voltage adds a slow overall drift in one direction on top of that random motion, and that drift is the current.
Mistake 3: “An alloy is a compound of several metals.” Alloys are generally mixtures: their components don’t come in a fixed ratio, and they can include nonmetals, such as the carbon in steel.
Quick check
Which fact can the sea-of-electrons model explain?
The free electrons in a metal drift under a voltage, carrying a current. Dissolving salt involves ions, nitrogen's stability comes from its strong N≡N bond, and water's high boiling point comes from hydrogen bonding (lesson 2.4.8).
Which metal has the highest boiling point?
Aluminum has 3 valence electrons, and Al³⁺ has the largest charge and smallest radius, so its metallic bonding is strongest and its boiling point highest (2519 °C).
Summary
- The valence electrons released by metal atoms form a sea of electrons shared by all the cations; the strong attraction between the cations and that sea is the metallic bond.
- The sea of electrons explains why metals conduct electricity and heat, are malleable and ductile, and are shiny. Under force, ionic crystals shatter as like charges meet, while metal layers slide without breaking.
- The higher the cation charge and the smaller the cation, the stronger the metallic bonding. Alloys are substitutional or interstitial, and are usually harder than the pure metal.
On to AP (for later)
AP U2.4 uses particle diagrams of metals and alloys: a pure metal is a neat array of identical spheres; a substitutional alloy replaces some spheres with others of similar size; an interstitial alloy has small spheres tucked between the large ones. A common question asks why an interstitial alloy such as steel is harder and less malleable than the pure metal: the small atoms block the layers from sliding.