Factors Affecting Bond Strength, Trends in Melting Points, and Formation of Alloys
Definition
Alloy
An alloy is a mixture of a metal with one or more other elements, which can be either metals or non-metals
Factors Affecting Bond Strength in Metals
- The strength of metallic bonds arises from the electrostatic attraction between positively charged metal cations and the "sea" of delocalized electrons surrounding them.
- Two key factors influence this attraction: the charge of the cations and the radius of the cations.
Charge of Cations
- The charge of a metal cation directly affects the strength of the metallic bond.
- Cations with higher charges exert a stronger pull on the delocalized electrons, increasing the bond strength.
Example
- Sodium (Na) forms cations with a charge of $1+$.
- Magnesium (Mg) forms cations with a charge of $2+$.
- Aluminium (Al) forms cations with a charge of $3+$.
- As the charge increases from $1+$ to $3+$, the metallic bond becomes stronger.
- This explains why aluminium has a higher melting point than magnesium or sodium.
Radius of Cations
- The size of the cation plays an equally important role.
- Smaller cations allow the delocalized electrons to get closer to the nucleus, increasing the strength of the electrostatic attraction.
Example
- Sodium cations ($Na^+$) have a larger ionic radius than magnesium cations ($Mg^{2+}$).
- Magnesium cations are larger than aluminium cations ($Al^{3+}$).
- This trend helps explain why aluminium, with its small cation radius and high charge, forms the strongest metallic bonds among these three metals.
Common Mistake
- Don’t confuse the size of the neutral atom with the size of the cation.
- When a metal atom loses electrons to form a cation, its radius decreases significantly.
Trends in Melting Points of Metals
- The melting point of a metal reflects the strength of its metallic bonds.
- Stronger metallic bonds require more energy to break, leading to higher melting points.
- Let’s examine two key trends:
s-Block Metals: Lower Melting Points
- The s-block metals, including Group 1 (alkali metals) and Group 2 (alkaline earth metals), generally have lower melting points compared to other metals.
- This is due to:
- Larger cation radii: As you move down the group, the cations become larger, weakening the attraction between the cations and the delocalized electrons.
- Lower charges: Group 1 metals have a charge of $1+$, leading to weaker metallic bonds.
Example
Lithium ($Li$) has a melting point of 181°C, while potassium ($K$) melts at just 63°C.
Example
- Consider the melting points of Group 1 metals.
- As the ionic radius increases from lithium to potassium, the metallic bonds weaken, resulting in lower melting points.
p-Block Metals: Higher Melting Points
- In contrast, p-block metals (such as aluminium) tend to have higher melting points.
- Their smaller cation radii and higher charges result in stronger metallic bonds.
- Additionally, p-block metals contribute more delocalized electrons per atom to the "sea," further strengthening the bonds.
Example
Aluminium ($Al$) has a melting point of 660°C, significantly higher than sodium ($98°C$) or magnesium ($650°C$).
Self review
Can you explain why aluminium has a higher melting point than sodium, based on the factors affecting bond strength?
