Atomic Structure: Ionisation Energy

Atomic Structure: Ionisation Energy

1. Definition of ionisation energy

   • The minimum amount of energy needed to remove 1g of electron from 1 mol of gaseous atom to form 1 mol of gaseous ion

2. Trend of ionisation energy

   • down the group
     • I.E. decreases
       • Shielding effect
       • nuclear attraction decreases
       • size of atom increases
   • across the group
     • I.E. increases
       • # of protons increases
       • nuclear attraction increases
       • nuclear charge increases
       • electron energy level will be held tightly together

3. I.E. between shells

   • as the electrons get closer to the nucleus, I.E. increases
     • log of ionisaion energy:

4. Factors affecting I.E.

   • atomic radius
     • Smaller the atomic radius, the higher is the atomic energy
   • Nuclear charge
     • the higher the positive charge of the nuceus, stronger is it's attraction for the electrons
   • Orbital penetration
     • It's easier to remove electrons from p orbitals than from s orbitals, because the s orbitals penetrate towards the nucleus more closely than the p orbitals
   • Electron pairing
     • paired electrons are easier to remove than unpaired ones because repulsion between electrons in the same orbital is higher than repulsion between electrons in different orbitals
   • Shielding or screening effect of the inner orbitals
     • inner electron orbitals effectively shield the outermost electrons from the nucleus

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