The big idea: A neon sign glows its own particular red and a sodium street lamp its own orange — each element gives out only its own set of colours, because an atom can only hold a few allowed energies, its energy levels (we say the energy is quantised: it comes in fixed steps).
When an electron drops from a higher level to a lower one, the energy it loses leaves as a single packet of light — a photon.
Because only certain drops are allowed, an atom gives out only certain colours — a pattern of separate lines called a line spectrum.
An atom can only hold a few allowed energies — the discrete energy levels. When an electron DROPS from a higher level to a lower one, the energy it loses leaves as a single photon of light (a downward arrow = emission).
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New words, plainly: Quantised = only fixed values are allowed (like stairs, not a ramp).
Photon = one tiny packet of light energy.
Transition = an electron jumping between two levels.
Line spectrum = the set of separate lines (colours) an atom gives out or takes in.
Emission spectrum
- Electron drops to a lower level and gives out a photon
- You see bright coloured lines on a dark background
- Each line = one allowed energy gap in the atom
Absorption spectrum
- Electron absorbs a photon and jumps up to a higher level
- You see dark lines missing from a continuous rainbow
- The same gaps are missing as the emission lines — same atom, same energies
Why a line spectrum is a fingerprint: Every element has its own set of energy levels, so it has its own pattern of lines.
That is how we match a spectrum to an element — and how we know which energy-level diagram a spectrum came from.
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The energy a photon carries equals the gap between the two levels. The data booklet gives two ways to link that energy to the light:
- energy of the photon — equals the energy lost in the jump (J)
- Planck constant, 6.63 × 10⁻³⁴ J s (given)
- frequency of the emitted (or absorbed) light (Hz)
- energy of the photon — equals the energy of the jump (J)
- Planck constant, 6.63 × 10⁻³⁴ J s (given)
- speed of light, 3.00 × 10⁸ m s⁻¹ (given)
- wavelength of the light (m)
Bigger drop → shorter wavelength: From , a bigger energy gap means a bigger E, which means a shorter wavelength λ (and higher frequency).
So the biggest drop makes the shortest-wavelength line; the smallest drop makes the longest-wavelength line.
A BIG drop loses a lot of energy → a high-frequency, SHORT-wavelength photon. A SMALL drop loses little energy → a low-frequency, LONG-wavelength photon. Photon energy = the gap between the two levels.
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An electron in an atom drops between two levels and loses 3.0 × 10⁻¹⁹ J of energy. Find the wavelength of the photon it emits. (h = 6.63 × 10⁻³⁴ J s, c = 3.00 × 10⁸ m s⁻¹.)
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How this is tested — energy levels and spectra are almost always a Paper 1A multiple-choice question:
Paper 1A
- Match an emission-line pattern to the right set of energy levels (more gaps → more lines).
- Pick the longest-wavelength transition — the smallest energy drop.
- Count how many different wavelengths appear from level n to the ground state.
Paper 2
- Use E = hf or E = hc/λ to turn a level gap into a frequency or wavelength.
The classic trap: Thinking the biggest drop gives the longest wavelength — it's the opposite (biggest drop = shortest λ).
Counting distinct wavelengths: Each different gap between two levels makes one line. From level n down to the ground state, the number of different downward jumps is n(n − 1) ÷ 2.
For n = 3 that is 3 lines (3→1, 3→2, 2→1); for n = 4 it is 6 lines.
From n = 3 there are THREE different downward jumps (3→1, 3→2, 2→1). Each jump has its own energy gap, so each makes a photon of a different wavelength — three emission lines in the spectrum.
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In a particular atom an electron is excited to the third level (n = 3), and the electrons fall back to the ground state by every possible route. (a) How many different wavelengths can appear in the emission spectrum? (b) State which transition produces the longest-wavelength photon, and explain why.
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