Metallic Bonding and Alloys
Quick Notes
- Metallic bonding: lattice of positive metal ions in a “sea of delocalised electrons.”
- Delocalised electrons explain electrical/thermal conductivity, malleability, ductility, and luster of metals.
- Melting/boiling points: high due to strong electrostatic attraction between ions and electrons in the lattice.
- Conductivity: delocalised electrons carry charge in solid and molten states.
- Malleable/ductile: layers of ions can slide without breaking metallic bonding.
- Alloys: retain metallic bonding but modify properties.
- Interstitial alloy: small atoms fit in gaps between metal atoms, making the solid harder/stronger (e.g., steel = Fe + C).
- Substitutional alloy: similar-sized atoms replace each other in lattice (e.g., brass = Cu + Zn).
Full Notes
What is Metallic Bonding?
Metallic bonding is a strong electrostatic attraction between positive metal ions (cations) and a “sea” of delocalised electrons.
Metal atoms lose their outer electrons easily; these electrons become delocalised and form a shared pool across the entire structure. The positive ions are arranged in a fixed lattice, held together by attraction to the delocalised electrons.
Example:Structure of Sodium
Each sodium atom loses one outer electron, forming Na⁺ ions. These electrons form a mobile electron cloud around the lattice.
Properties Explained by Metallic Bonding
- High melting/boiling points: strong attractions between cations and delocalised electrons require large energy to break.
- Electrical and thermal conductivity: free electrons carry charge and transfer energy in both solid and molten states. Example: copper wiring.
- Malleability and ductility: layers of ions can slide over each other without disrupting bonding. Example: gold jewellery.
- Luster: mobile electrons reflect light, giving metals a shiny surface.
Models of Metallic Solids
In a model, metal cations are arranged in a regular lattice with delocalised electrons shown as a surrounding electron cloud.
Example:Model of bonding in magnesium
Alloys
Alloys are mixtures containing at least one metal. They retain metallic bonding but often have altered properties.
Interstitial Alloys
Form when small atoms (often nonmetals like carbon) fit into the spaces (interstices) between larger metal atoms in the lattice.
The atoms have significantly different radii and this means layers can no longer easily slide over each other, giving a harder, stronger material
Example: Steel (Fe + C) – carbon atoms block movement of iron layers.
Substitutional Alloys
Atoms of similar size replace each other in the lattice. The structure remains uniform, and properties are an average of the components.
Example: Brass (Cu + Zn) – zinc atoms replace copper atoms in the lattice.
Worked Example
Explain why steel is stronger than pure iron.
- Steel is an interstitial alloy – small carbon atoms fit between iron atoms.
- This blocks the movement of iron layers, making it harder for them to slide.
- Conclusion: Increased strength and hardness compared to pure iron.
If asked to model or explain the structure of a metallic solid or alloy, be sure to refer to delocalized electrons, positive metal ions, and the type of alloy (interstitial or substitutional). Be specific about the relative size of the atoms involved.
Summary
- Metallic bonding = lattice of metal cations + delocalised electrons.
- Explains high conductivity, malleability, ductility, and luster.
- Alloys modify properties but retain metallic bonding.
- Interstitial = small atoms fill gaps (harder/stronger).
- Substitutional = similar-sized atoms replace each other (mixed properties).