Metallic bonding
IB Chemistry SLΒ· 20 min read
1. Nature of Metallic Bondingβ β ββββ± 10 min
Unlike ionic bonding (where electrons are transferred) or covalent bonding (where electrons are shared between two atoms), metallic bonding involves delocalization of valence electrons across the entire solid lattice. Metals have low ionization energy, so they readily lose valence electrons to form positive cations.
Metallic bonding
The non-directional electrostatic attraction between positively charged metal ions and the delocalized sea of valence electrons that surrounds the ions
Example:
In solid sodium, each cation sits in a lattice surrounded by a sea of delocalized 3sΒΉ electrons.
Explain why valence electrons in magnesium are delocalized rather than bound to individual atoms
- 1
Magnesium is a group 2 metal with electron configuration . Metals have low ionization energy, so valence electrons are weakly attracted to the nucleus.
- 2
When magnesium atoms pack into a solid lattice, the 3s orbitals of adjacent atoms overlap across the entire lattice.
- 3
Overlapping orbitals create a continuous band of low-energy molecular orbitals, so electrons are free to move through the lattice instead of being bound to a single atom.
2. Physical Properties of Metals (Explained by Bonding)β β ββββ± 15 min
All characteristic physical properties of metals can be directly linked to their delocalized bonding and non-directional lattice structure:
- Electrical conductivity: Delocalized electrons move freely through the lattice when a potential difference is applied, so metals conduct electricity as solids and liquids.
- Thermal conductivity: Free delocalized electrons transfer kinetic energy quickly through the lattice, giving metals high thermal conductivity.
- Malleability and ductility: Non-directional bonding means layers of cations can slide past each other without breaking the electrostatic attraction to delocalized electrons.
- Luster: Delocalized electrons absorb and re-emit most visible light, giving metals a shiny appearance.
Explain why metals are malleable but ionic solids are brittle
- 1
In metals, bonding is non-directional electrostatic attraction between cations and delocalized electrons.
- 2
When a force is applied, layers of cations slide past one another. Delocalized electrons can move to re-establish electrostatic attraction between the new cation positions, so the structure does not break.
- 3
In ionic solids, bonding is attraction between fixed oppositely charged ions. When a force shifts layers, like charges align, creating repulsion that breaks the lattice apart.
Exam tip:
Always connect the property directly to delocalized electrons or non-directional bonding in exam answers, don't just state the property.
3. Trends in Metallic Bond Strengthβ β β βββ± 15 min
The strength of metallic bonding depends on two key factors: the charge of the metal cation, and the ionic radius of the cation. Higher cation charge increases electrostatic attraction to delocalized electrons, leading to stronger bonding. Smaller ionic radius also increases attraction, because cations are closer to the delocalized electrons. Stronger bonding requires more energy to break the lattice, so it correlates to higher melting points.
Explain why magnesium has a higher melting point than sodium
- 1
Magnesium forms cations, while sodium forms cations. Magnesium has a higher cation charge than sodium.
- 2
has a smaller ionic radius (72 pm) than (102 pm), because more protons pull on the same number of electron shells.
- 3
Higher charge and smaller radius create much stronger electrostatic attraction between ions and delocalized electrons, compared to sodium.
- 4
Stronger metallic bonding requires more thermal energy to break the lattice, so magnesium has a higher melting point than sodium.
4. Common Pitfalls
Wrong move:
Saying metallic bonding is attraction between neutral metal atoms
Why:
The metallic lattice is made of positive cations, not neutral atoms, and attraction is between cations and delocalized electrons
Correct move:
State that metallic bonding is electrostatic attraction between positively charged metal cations and a sea of delocalized valence electrons
Wrong move:
Claiming metals only conduct electricity when molten
Why:
This rule applies to ionic compounds, not metals
Correct move:
Metals conduct electricity in both solid and molten states, because delocalized electrons are mobile even in solid lattices
Wrong move:
Explaining higher melting point only by higher atomic mass
Why:
Melting point depends on bond strength, not mass, and many high mass metals have low melting points (e.g. mercury)
Correct move:
Explain melting point trends using cation charge and ionic radius, which determine the strength of electrostatic attraction
Wrong move:
Saying metallic bonds break when a metal is hammered into shape
Why:
Metallic bonds do not break when layers slide, they re-form in new positions
Correct move:
When force is applied, layers of cations slide past each other, and delocalized electrons re-adjust to maintain attraction so the structure stays intact
5. Quick Reference Cheatsheet
Factor | Effect on Bond Strength | Effect on Melting Point |
|---|---|---|
Higher cation charge | Stronger attraction, stronger bond | Higher |
Smaller cation radius | Stronger attraction, stronger bond | Higher |
Lower cation charge | Weaker attraction, weaker bond | Lower |
Larger cation radius | Weaker attraction, weaker bond | Lower |
Electrical conductivity | Delocalized mobile electrons | Conducts solid + liquid |
Malleability | Non-directional bonding allows sliding | Shape can be changed without breaking |
When this came up on past exams
AI-estimated based on syllabus patterns β cross-check with official past papers for accuracy. Use only as revision-focus signals.
- 2023 Β· 1
Identify property of metallic bonding
- 2022 Β· 2
Explain melting point trend in metals
What's Next
Understanding metallic bonding completes the three core types of chemical bonding covered in IB Chemistry SL. This foundation is essential to explain the properties of alloys (mixtures of metals with other elements), and to compare how different bonding types give different physical properties to solid materials. You will also build on this knowledge when studying periodic trends in metallic character across periods and groups, and when analyzing structure-property relationships in materials chemistry.