Ionisation energy

What is the definition of first ionisation energy (IE₁)?

The first ionisation energy is the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous ions.
Example:
Na(g) → Na⁺(g) + e⁻, IE₁ = +496 kJ mol⁻¹

What are the three main factors affecting ionisation energy?

• Atomic radius → Larger distance between nucleus & outer electron = weaker attraction = lower IE.

• Nuclear charge → More protons = stronger nuclear attraction between nucleus & electrons = higher IE.

• Electron shielding → Inner shells repel outer electrons, reducing effective nuclear charge = lower IE.

Why does first ionisation energy decrease down a group?

• Atomic radius increases (outer electrons further away).

• Shielding increases (more inner shells).

• Increased nuclear charge is outweighed by distance + shielding.
  ➡ Overall: Attraction decreases, so ionisation energy decreases.

Why does first ionisation energy increase across a period?

• Nuclear charge increases (more protons).

• Atomic radius decreases as the number of protons in the nucleus increases but shell number remains the same (no extra shells added).

• Shielding remains constant (same number of inner shells).
  ➡ Overall: Attraction increases, so ionisation energy increases.

Why is there a sharp drop in ionisation energy between the end of one period and the start of the next?

• New period = new shell added.

• Increased distance from nucleus.

• Increased shielding from inner shells.
  ➡ Attraction weaker, despite nuclear charge increasing.

why is a decrease in ionisation energy at aluminium in period 3 evidence of subshells

• the outermost electron in aluminium sits at a higher energy level subshell further from the nucleus 3p¹ compared to magensium which is at 3s²

• the atomic model niels bohr came up with didn’t explain subshells

how is a decrease in IE at sulfur in period 3 evidence for electron repulsion in an orbital

• phosphorus and sulfur both have outer electrons in the 3p orbital so shielding is the same

• however in phosphorus case the electron is being removed from a singly occupied shell whereas in sulfur it’s being removed from an orbital containing two electrons

• sulfur has a lower IE showing that the repulsion between two electrons in an orbital means that the electrons are easier to remove from shared orbitals

Why is the IE of boron lower than beryllium?

• Be configuration: 1s² 2s² → outer electron in 2s.

• B configuration: 1s² 2s² 2p¹ → outer electron in 2p.

• 2p orbital is higher in energy & further from nucleus.
  ➡ Easier to remove → lower IE (801 kJ mol⁻¹ vs 900 kJ mol⁻¹).

Why is the IE of oxygen lower than nitrogen?

• N configuration: 1s² 2s² 2p³ (half-filled stable subshell).

• O configuration: 1s² 2s² 2p⁴ → one 2p orbital contains a pair of electrons, causing repulsion.
  ➡ Easier to remove → lower IE (1314 kJ mol⁻¹ vs 1402 kJ mol⁻¹).

where are there minor dips in the increasing IE across periods?

•  Group 2 to 3 —>Be to B

• Group 5 to 6 —> P to S

• Group 5 to 6 —> N to O

What happens to successive ionisation energies of an element?

• Each successive electron is removed from an increasingly positive ion → requires more energy.

• As electrons are removed:

  • Less shielding

  • Stronger nuclear attraction

  • Proton-to-electron ratio increases so remaining electrons are held onto tightly
    ➡ IE values increase overall.

What do “big jumps” in successive ionisation energy graphs show?

• Big jumps = electron removed from a new, inner shell (much closer to nucleus).

• Small jumps = electron removed from the same shell (or subshell).

Example:

• Easy removal: outer 4s electron (paired, repulsion).

• Harder: second electron (from 4s⁺ ion).

• Large jump: third electron (closer 3p shell).

what are successive ionisation energies

• you can remove all the electrons from an atom leaving only the nucleus

• each time you remove an electron there’s a successive ionisation energy

what is second ionisation energy

• the energy needed to remove 1 electron from each ion in 1 mole of gaseous 1+ ions to form 1 mole of gaseous 2+ ions

how can you tell what group an element is in from a successive ionisation graph?

• count how many electrons are removed before the first big jump to find the big number