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Emission & Absorption Spectra

Emission & Absorption Spectra

11 min read Updated Mar 26, 2026

Every element has a unique set of energy levels. When electrons jump between these levels, they absorb or emit photons with very specific energies - and therefore very specific wavelengths. This creates a spectral “fingerprint” that is unique to each element.

Emission Spectra

When atoms are heated or electrically excited, electrons jump to higher energy levels. These excited states are unstable, so electrons quickly fall back to lower levels, releasing photons in the process. Each photon has a wavelength determined by the energy difference between the two levels:

The result is a line spectrum - a set of discrete, bright lines on a dark background. Each line corresponds to a specific electron transition. Because each element has a unique set of energy levels, each element produces a unique line spectrum.

The Hydrogen Emission Series

Hydrogen’s emission spectrum is the simplest and most important for the MCAT. The transitions are grouped into named series based on the final energy level:

Hydrogen energy levels and the lines they emit

Bohr model
energy (eV) electron free · ionised n = ∞ 0 eV n = 6 -0.38 eV n = 5 -0.54 eV n = 4 -0.85 eV n = 3 -1.51 eV n = 2 -3.40 eV n = 1 -13.60 eV 10.2 eV: the n = 1 → 2 gap alone That single step is three quarters of the whole 13.6 eV it takes to strip the electron off. Every rung above it shares the last quarter, which is why they pile up. Lyman lands on n = 1 · ultraviolet Balmer lands on n = 2 · visible Paschen lands on n = 3 · infrared What the eye actually sees A hot hydrogen tube emits only these four visible lines, on an otherwise black field. Hα 656 nm Hβ 486 nm Hγ 434 nm Hδ 410 nm 400 nm wavelength 700 nm line drop ΔE λ Hα 3 → 2 1.89 eV 656 nm Hβ 4 → 2 2.55 eV 486 nm Hγ 5 → 2 2.86 eV 434 nm Hδ 6 → 2 3.02 eV 410 nm Absorption is this in reverse Shine white light through cold hydrogen and the same four wavelengths go missing: dark lines on a continuous background, at exactly the positions of the bright emission lines above.
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Scroll sideways to see the whole map.

Why the rungs crowdEnergy goes as 1/n², so the n = 1 to n = 2 gap is 10.2 eV while n = 5 to n = 6 is under 0.2 eV. Every question that asks which transition emits the shortest wavelength is really asking which gap is biggest, and the biggest gaps are always at the bottom.

Reading a series off the ladderA series is named for where the electron lands, not where it starts. Landing on n = 1 means falling the whole way, so Lyman lines are ultraviolet. Landing on n = 2 gives the four visible Balmer lines. Landing on n = 3 gives infrared.

The sign conventionEnergies are negative because the electron is bound: zero is the free electron at n = ∞. Absorbing a photon moves the electron up toward zero, and it takes the full 13.6 eV to ionise hydrogen from the ground state.

The ladder and the spectrum are the same fact. Each coloured line on the right is drawn at the wavelength its arrow on the left produces, through λ = hc/ΔE. A bigger drop makes a bluer photon, which is why H-α from n = 3 is red and H-δ from n = 6 is violet.

| Series | Final level (nfn_{f}) | Transitions from | Region of spectrum |
|--------|-------------------|------------------|--------------------|
| Lyman | 1 | n = 2, 3, 4, … to 1 | Ultraviolet (UV) |
| Balmer | 2 | n = 3, 4, 5, … to 2 | Visible light |
| Paschen | 3 | n = 4, 5, 6, … to 3 | Infrared (IR) |

Key Energy Relationships

The largest energy transition in any series is from n = infinity to nfn_{f} (the series limit). The smallest is from nfn_{f} + 1 to nfn_{f}.

Within the Balmer series (visible light):

  • n = 3 to n = 2: red light (656 nm) - lowest energy visible transition
  • n = 4 to n = 2: blue-green (486 nm)
  • n = 5 to n = 2: blue-violet (434 nm)
  • n = 6 to n = 2: violet (410 nm) - highest energy visible transition

Remember: shorter wavelength = higher frequency = higher energy.

Absorption Spectra

An absorption spectrum is the reverse of an emission spectrum. When white light (containing all wavelengths) passes through a gas of atoms, the atoms absorb photons that match specific energy transitions. The result is a continuous spectrum with dark lines where specific wavelengths have been absorbed.

The wavelengths of the dark lines in an absorption spectrum match exactly the wavelengths of the bright lines in the emission spectrum for the same element. This makes sense - the same energy gaps are involved, whether the electron is jumping up (absorption) or falling down (emission).

| Feature | Emission Spectrum | Absorption Spectrum |
|---------|------------------|-------------------|
| Appearance | Bright lines on dark background | Dark lines on continuous background |
| Process | Electrons fall to lower levels | Electrons jump to higher levels |
| Energy | Released as photons | Absorbed from photons |
| Same wavelengths? | Yes - same element, same lines |

Continuous vs. Line Spectra

A continuous spectrum contains all wavelengths of light (like a rainbow) and is produced by hot, dense objects like the sun or an incandescent light bulb.

A line spectrum contains only specific wavelengths and is produced by excited atoms in the gas phase. The discrete lines reflect the quantized energy levels of the atoms.

Which hydrogen emission series produces visible light? What is the final energy level for transitions in this series?
Click to reveal answer

The Balmer series produces visible light. All transitions in the Balmer series end at n = 2 (nfn_{f} = 2). Transitions from n = 3, 4, 5, and 6 down to n = 2 produce red, blue-green, blue-violet, and violet light respectively.

How do emission and absorption spectra of the same element compare?
Click to reveal answer

They are complementary. The bright lines in the emission spectrum appear at exactly the same wavelengths as the dark lines in the absorption spectrum. Both correspond to the same electron transitions between the same energy levels - emission lines result from electrons falling down, absorption lines result from electrons jumping up.

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