A level Chemistry AQA - Periodicity

Trends across a period - atomic radius

Decreases

Why does atomic radius decrease across a period?

Nuclear charge increases so there is a stronger electrostatic attraction between the nucleus of elements across a period and their outer electrons, so the outer electrons are drawn in closer to the nucleus

Trends across Period 3 - melting points

1. melting points of P3 metals increase

2. melting point of silicon is very high

3. melting points of P, S, Cl, Ar are low

4. The melting point of S is greater that P

Why do the melting points of metals increase across P3?

• The number of outer electrons increases

• More delocalised electrons.

• Greater electrostatic forces of attraction between the positive metal ions and the delocalised electrons.

• Ionic radius from Na→Al decreases so charge density increases.

Why is the melting point of silicon very high?

• Has a giant covalent lattice structure

• A lot of energy is required to break the many strong covalent bonds

Why are the melting point of P, S, Cl, Ar low?

• Are simple covalent molecules

• Little energy is required to overcome the weak van der Waals forces between molecules

Why is the melting point of S greater than P?

• S has more electrons per molecule than P or Cl2

• Therefore S has stronger van der Waals forces of attraction between molecules

• More energy is required to overcome those forces

Trends across a period - first ionisation energy

Increases

Why does first ionisation energy increase across a period?

• Nuclear charge increases

• Atomic radius decreases

• Outer electron is closer to the nucleus

• Shielding stays the same

• Stronger electrostatic forces of attraction between nucleus and outer electron

• More energy is required to lose outer electron

First IE in Group 3

Al (G3) has lower first ionisation energy than expected

Why does aluminium have a lower first IE than expected

• Electron being added into 3p subshell

• This subshell is further from the nucleus

First IE in Group 6

S has lower first IE than expected

Why does sulfur have a lower first IE than expected?

• Pairing of electrons in the p-subshell

• The electrons in the pair repel

• Less energy required to remove one of the electrons

First IE in group 1

Have the lowest first IE in each period

Why do elements in G1 have the lowest first IE in each period?

• Greatest atomic radius

• Lowest nuclear charge

First IE in Group 0

Highest first IE in each period

Why do elements in G0 have highest first IE in each period?

• smallest atomic radius

• Highest nuclear charge