Ionisation Energy

2026 Syllabus Objectives

By the end of this topic, you should be able to:

  1. Define and use the term first ionisation energy (IE)
  2. Construct equations to represent first, second and subsequent ionisation energies
  3. Identify and explain the trends in ionisation energies across a period and down a group of the Periodic Table
  4. Identify and explain the variation in successive ionisation energies of an element
  5. Understand that ionisation energies are due to the attraction between the nucleus and the outer electron
  6. Explain the factors influencing the ionisation energies of elements in terms of nuclear charge, atomic/ionic radius, shielding by inner shells and sub-shells, and spin-pair repulsion
  7. Deduce the electronic configurations of elements using successive ionisation energy data
  8. Deduce the position of an element in the Periodic Table using successive ionisation energy data

1. What is Ionisation Energy?

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Ionisation is the process of removing an electron from an atom or ion.

Ionisation energy is a measure of how much energy is needed to remove an electron from a gaseous atom or ion. It tells us how strongly an atom holds onto its electrons.

Think of it like this: the nucleus (positive charge) pulls on the electrons (negative charge) like a magnet. The stronger the pull, the more energy you need to remove an electron.

Key points:

  • Ionisation energy is always positive (you must put energy in)
  • It is an endothermic process (energy is absorbed)
  • This is because you need to overcome the force of attraction between the negative electron and the positive nucleus
  • Units: kJ mol⁻¹ (kilojoules per mole)
  • Measured under standard conditions: 298 K and 101 kPa

2. First Ionisation Energy

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Definition

First ionisation energy (IE₁) is the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions.

General Equation Format

The general equation is:

X(g) → X⁺(g) + e⁻

Where:

  • X = any element
  • (g) = gaseous state (this is essential!)
  • e⁻ = electron

Examples

Sodium: Na(g) → Na⁺(g) + e⁻ ΔH = +496 kJ mol⁻¹

Calcium: Ca(g) → Ca⁺(g) + e⁻ ΔH = +590 kJ mol⁻¹

Fluorine: F(g) → F⁺(g) + e⁻ ΔH = +1680 kJ mol⁻¹

Notice that fluorine has a much higher first ionisation energy than sodium. This is because fluorine holds its electrons much more tightly.

Important Notes

  • Always include state symbols (g) in your equations
  • The electron being removed is the outermost electron (the one furthest from the nucleus)
  • Different elements have different first ionisation energies depending on their electronic structure

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