IB DP Chemistry - S2.3.2 Strength of metallic bonds- Study Notes - New Syllabus - 2026, 2027 & 2028
IB DP Chemistry – S2.3.2 Strength of metallic bonds – Study Notes – New Syllabus
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- IB DP Chemistry 2025 SL- IB Style Practice Questions with Answer-Topic Wise-Paper 1
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Structure 2.3.2 — Strength of Metallic Bonds and Trends in Melting Points (s- and p-block metals)
Structure 2.3.2 — Strength of Metallic Bonds and Trends in Melting Points (s- and p-block metals)
Metallic Bond Strength Depends On:
- Charge on the metal ion: Higher charge → stronger electrostatic attraction with delocalized electrons.
- Size (radius) of the metal ion: Smaller ions → stronger attraction between nucleus and sea of electrons due to higher charge density.
- Number of delocalized electrons: More electrons → stronger bond.
Explanation of Melting Point Trends:
s-block metals (Groups 1 and 2):
Group 1 (e.g. Na, K, Rb):
- Each atom loses 1 electron → 1 delocalized electron per ion.
- Metallic bonds are relatively weak due to low charge and large ionic radius.
- Melting points decrease down the group as ion size increases → weaker attraction.
Group 2 (e.g. Mg, Ca, Sr):
- Each atom loses 2 electrons → 2 delocalized electrons per ion.
- Higher charge and more electrons → stronger bonding than Group 1 metals.
- Melting points are higher than Group 1 and generally decrease down the group.
p-block metals (e.g. Al, Ga, In, Tl):
- More variation in structure and bonding.
- Aluminium has very strong metallic bonding due to:
- High charge \( \text{Al}^{3+} \)
- Small radius
- 3 delocalized electrons per atom
- Melting point trend is less regular due to changes in atomic structure and presence of partially filled p-orbitals.
Summary Table: Melting Point Trends
| Group | Example Metals | Charge | Ion Size Trend | Melting Point Trend |
|---|---|---|---|---|
| Group 1 | Li, Na, K | \( +1 \) | Increases down group | Decreases |
| Group 2 | Mg, Ca, Sr | \( +2 \) | Increases down group | Decreases |
| p-block | Al, Ga, In, Tl | Mostly \( +3 \) | Increases irregularly | Irregular trend |
Key Concept:
Stronger metallic bonds → higher melting points. Strength depends on charge, radius, and number of delocalized electrons.
Example
Explain why aluminium has a much higher melting point than sodium.
▶️Answer/Explanation
Aluminium forms \( \text{Al}^{3+} \) ions and contributes 3 delocalized electrons per atom.
Sodium forms \( \text{Na}^{+} \) ions and contributes only 1 delocalized electron.
Aluminium has a smaller ion with higher charge → stronger metallic bonds → much higher melting point.
Example
Explain why aluminium has a higher melting point than sodium and magnesium.
▶️Answer/Explanation
Aluminium forms \( \text{Al}^{3+} \) ions and releases 3 delocalized electrons per atom.
Sodium forms \( \text{Na}^+ \) with 1 electron; magnesium forms \( \text{Mg}^{2+} \) with 2 electrons.
Higher charge and more electrons → stronger attraction in aluminium → higher melting point.
Example
Why does calcium have a lower melting point than magnesium?
▶️Answer/Explanation
Both Ca and Mg are Group 2 metals and form \( \text{M}^{2+} \) ions with 2 delocalized electrons each.
However, calcium has a larger atomic radius → lower charge density → weaker attraction between cations and delocalized electrons.
Therefore, calcium has a lower melting point than magnesium.
Example
Predict which metal has stronger metallic bonding: potassium (K) or beryllium (Be)?
▶️Answer/Explanation
Potassium is Group 1: forms \( \text{K}^+ \) with 1 delocalized electron.
Beryllium is Group 2: forms \( \text{Be}^{2+} \) with 2 delocalized electrons and is much smaller in size.
Be has higher cation charge and greater electron density → stronger metallic bonding than K.
Example
Why does lithium have a higher melting point than cesium?
▶️Answer/Explanation
Both are Group 1 metals and form \( \text{M}^+ \) ions with 1 delocalized electron.
Lithium has a much smaller radius → higher charge density → stronger attraction to electrons.
Thus, lithium has stronger metallic bonds → higher melting point than cesium.