Inorganic chemistry

1. How does atomic radius change across Period 3?

1. Increases
2. Decreases
3. Stays the same
4. Increases then decreases

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B: Decreases. Increasing nuclear charge pulls the outer electrons in closer.

2. Why does melting point rise sharply from Na to Si across Period 3?

1. Structures change from simple molecular to giant metallic/covalent with stronger bonding
2. Atoms get smaller only
3. Elements become non-metals
4. Electronegativity decreases

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A: Structures change from simple molecular to giant metallic/covalent with stronger bonding. Stronger bonding requires more energy to overcome.

3. Why does melting point drop suddenly after silicon in Period 3?

1. These elements form giant structures
2. These elements are simple molecular or monatomic, with only weak forces between particles
3. They become metals
4. They gain extra protons

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B: These elements are simple molecular or monatomic, with only weak forces between particles. Only weak intermolecular or interatomic forces need to be overcome.

4. How does electronegativity change down a group?

1. Increases
2. Decreases
3. No change
4. Becomes negative

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B: Decreases. Greater atomic radius reduces the pull on a bonding pair from further atoms.

5. Why does first ionisation energy generally decrease down a group?

1. Nuclear charge decreases
2. Atomic radius and shielding increase, outweighing increased nuclear charge
3. Atoms become smaller
4. Electrons are lost from full shells

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B: Atomic radius and shielding increase, outweighing increased nuclear charge. The outer electron is further away and more shielded, so easier to remove.

6. What is the general trend in electronegativity across Period 3?

1. It decreases
2. It increases
3. It stays constant
4. It becomes negative

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B: It increases. Increasing nuclear charge with similar shielding increases the pull on bonding electrons.

7. Which Period 3 element has the highest first ionisation energy?

1. Chlorine
2. Argon
3. Sodium
4. Silicon

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B: Argon. Argon has the largest nuclear charge of the Period 3 elements, with a full, stable outer shell. Its outer electrons are held most strongly, so it has the highest first ionisation energy.

8. First ionisation energies (kJ mol⁻¹): Na 496, Mg 738, Al 578, Si 786. Why is Al lower than Mg?

1. Al has a larger atomic radius than Mg
2. Al's outer electrons are paired in one orbital
3. Al's outer electron is in a 3p subshell, higher in energy and shielded by 3s
4. Al has fewer protons than Mg

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C: Al's outer electron is in a 3p subshell, higher in energy and shielded by 3s. Al's outermost electron is in a 3p orbital. It is higher in energy and slightly shielded by the 3s electrons, so it is easier to remove despite the greater nuclear charge.

9. Why is the first ionisation energy of sulfur lower than that of phosphorus?

1. Sulfur has a larger nuclear charge
2. Sulfur has more shielding
3. Repulsion between paired electrons in a 3p orbital makes one easier to remove
4. Sulfur's outer electron is in the 3s subshell

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C: Repulsion between paired electrons in a 3p orbital makes one easier to remove. In S (3p⁴) two electrons share one 3p orbital. Their mutual repulsion makes one electron easier to remove than the unpaired 3p electron in P (3p³).

10. Why does silicon have a very high melting point?

1. Strong London forces between Si atoms
2. Strong covalent bonds throughout a giant lattice must be broken
3. Delocalised electrons in a metallic lattice
4. Strong intermolecular hydrogen bonds

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B: Strong covalent bonds throughout a giant lattice must be broken. Silicon has a giant covalent lattice. Melting it needs a large amount of energy to break many strong covalent bonds.