When electrons jump between energy levels they absorb or release light of very specific frequencies. Those frequencies show up as sharp lines — and the pattern of lines is direct evidence that electrons live in fixed energy levels.
📘 What you need to know
Electrons occupy fixed energy levels around the nucleus.
Absorbing energy makes an electron jump up a level; falling back down releases energy as light.
The frequency of light emitted matches the energy gap between the two levels.
A line emission spectrum shows only certain frequencies — evidence that electron energies are quantised.
Lines get closer together (converge) towards the high-energy end; the convergence limit relates to ionisation.
Absorption and emission
Electrons sit in energy levels around the nucleus, and they can move between them — but only by taking in or giving out a matching amount of energy.
Absorbing energy lifts an electron from a lower level to a higher one — the excited state (higher energy, further from the nucleus).
The electron then falls back down, releasing energy as light. The frequency emitted equals the frequency absorbed for that jump.
Absorbing energy pushes an electron out to a higher level; emission is the electron falling back and releasing light.
The line emission spectrum
If the emitted light is in the visible range, passing it through a diffraction grating splits it into a line emission spectrum — a set of sharp coloured lines on a dark background. Each line is one specific energy value, produced by one specific electron jump.
Every element has its own unique fingerprint of lines, because every element has its own unique set of energy levels. That’s how astronomers can tell which elements are in a distant star just from its light.
Convergence and ionisation
Look closely at a hydrogen spectrum and you’ll notice the lines get closer together towards the high-energy (blue/UV) end. This is called convergence.
The lines converge because the energy levels themselves get closer together as they go up.
At the convergence limit, the electron has enough energy to leave the atom completely — this corresponds to the ionisation energy.
Electron jumps down to n = 1 give UV (Lyman), to n = 2 give visible (Balmer), to n = 3 give infrared (Paschen). Levels converge at the top.
Which jump gives which light
The size of the drop decides the type of light released — bigger drops mean higher energy.
Falling to n = 1 → ultraviolet (largest gap, highest energy) — the Lyman series.
Falling to n = 2 → visible light — the Balmer series.
Falling to n = 3 → infrared (smallest gap, lowest energy) — the Paschen series.
WORKED EXAMPLE
Which electron transition in the hydrogen atom emits visible light? A: n = 1 → n = 2, B: n = 2 → n = 3, C: n = 2 → n = 1, D: n = 3 → n = 2.
Emission = electron falling to a lower levelso rule out A and B (those are jumps up, i.e. absorption).Visible light = falling to n = 2 (Balmer series)C falls to n = 1 → ultraviolet. D falls to n = 2 → visible.Answer: D (n = 3 → n = 2)
⚠️ Common mix-up
Absorption vs emission: absorption is the electron jumping up (energy in); emission is falling down (light out).
The series is named by the level the electron falls to, not the one it starts from.
Up next: Energy Levels, Sublevels & Orbitals — a closer look at how those energy levels are built from subshells and orbitals.
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