Atomic Trends

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Atomic radii: Size of the entire atom

1. The number of shells; within a group, the number of shells increases and therefore atomic radius generally increases in a group from top to bottom.

 

2. The increased nuclear attraction as you go through the period causes the atomic radius to decrease as electrons are pulled in closer to the nucleus

Effective nuclear charge - Zeff

: net charge of the nucleus after the attraction and repulsion of all the electrons are taken into account.

• as you work across a row of the periodic table, the number of protons increase but the number of shells remain the same and shielding does not occur frequently. Thus Zeff increases from LHS to RHS across the periodic table. 

 

• as you work down a column of the periodic table, both the number of protons increase and the number of shells too.  The result is a decrease in Zeff from top to bottom (due to increasing shielding by inner electrons. This shielding reduces the effective nuclear charge (Zeff), so outer electrons feel less attraction.

 

Ionisation energy - the amount of energy to remove 1 mole of electron from 1 mol of gaseous atoms

Ionisation energy generally decreases down a group as the outer electrons are on average situated further from the nucleus making it easier to remove 1 mole of electron from 1 mole of gaseous atem.

 

Ionisation energy generally increases across a period as Zeff increase therefore electron are drawn to and feel more  positive charge towards the nucleus therefore it is harder to remove electrons.

Electron affinity (ea) - the amount of energy associated with the addition of 1 mole of electrons into 1 mole of gaseous atoms.

 

1. Across a Period (↑ Electron Affinity) (becomes more negative)
   More protons increase nuclear attraction, pulling electrons closer. Smaller atomic radius strengthens this pull, making it easier to gain electrons.

2. Down a Group (↓ Electron Affinity)
   Larger atomic radius means weaker nuclear attraction. More inner electrons shield the nucleus, reducing its pull on new electrons. This makes elements lower in a group less likely to gain electrons.