Atomic Spectra: Emission and Absorption Spectra Explained

An atomic spectrum is the pattern of colours or wavelengths of light that an atom emits or absorbs. When atoms of an element are heated or passed through an electric discharge, they give out light only at certain specific wavelengths. Passing that light through a prism splits it into bright lines on a dark background. This is an emission spectrum.

This matters because every element has its own unique set of lines, like a fingerprint. Astronomers use these lines to find what stars are made of, and scientists use them to identify unknown substances in laboratories without touching them.

Key Concepts

In Bohr's model, electrons move only in fixed energy levels around the nucleus. The ground state is the lowest level. When an atom absorbs energy, an electron jumps to a higher level called an excited state. When it falls back, it emits a photon whose energy equals the difference between the two levels.

• Emission spectrum: bright coloured lines on a dark background, produced when electrons fall from higher to lower levels.
• Absorption spectrum: dark lines on a continuous rainbow background, produced when white light passes through a cool gas and electrons jump up, absorbing specific wavelengths.
• Spectral series: groups of lines like the Lyman series in ultraviolet, Balmer series in visible light, and Paschen series in infrared for hydrogen.

Formulas

The photon energy is E = hν, where h is Planck's constant (6.626 × 10⁻³⁴ J s) and ν is frequency in hertz (s⁻¹). Since c = νλ, we also write E = hc/λ, where c = 3 × 10⁸ m/s and λ is wavelength in metres.

For hydrogen, the Rydberg formula gives 1/λ = R (1/n₁² − 1/n₂²), where R = 1.097 × 10⁷ m⁻¹, n₁ is the lower level and n₂ is the higher level. This assumes a single-electron atom like hydrogen.

Worked Example

Find the wavelength of light emitted when an electron in hydrogen falls from level 3 to level 2 (the red Balmer line).

1. Use 1/λ = R (1/2² − 1/3²) = 1.097 × 10⁷ × (1/4 − 1/9).
2. 1/4 − 1/9 = 5/36 = 0.1389.
3. 1/λ = 1.097 × 10⁷ × 0.1389 = 1.524 × 10⁶ m⁻¹.
4. λ = 1 / (1.524 × 10⁶) = 6.56 × 10⁻⁷ m = 656 nm, which is red light.

Common Misconceptions

• Students think atoms emit all colours like a bulb. In reality, each element emits only specific lines.
• Many confuse emission and absorption spectra. Absorption lines occur at exactly the same wavelengths as emission lines for the same element.
• Some think electrons orbit like planets on smooth paths. In Bohr's model they jump between fixed energy levels without passing through in-between energies.

Real-World Applications

Atomic spectra identify elements in stars and nebulae, detect sodium and mercury in street lamps, analyse metal composition in factories, and test for lead or other impurities in water and food samples.

In short, emission spectra show bright lines when atoms release light, absorption spectra show dark lines when atoms take in specific wavelengths, and both act as unique fingerprints of elements.