Atomic spectra arise from the interaction of atoms with electromagnetic radiation. When atoms absorb or emit energy, electrons transition between different energy levels, resulting in characteristic spectral lines.
The sodium D-line (589.0 and 589.6 nm) originates from ( 3p \rightarrow 3s ) transition. With a low-pressure sodium lamp, the line is Doppler-broadened (( \sim 3 ) GHz at 500 K). With a tunable diode laser locked to the D-line, one can:
This exemplifies how a simple atomic transition becomes a rich testing ground with laser technology.
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Atoms emit or absorb light at discrete wavelengths because their electrons occupy quantized energy levels. When an electron transitions between levels ( E_i ) (initial) and ( E_f ) (final), the photon energy is:
[ \Delta E = |E_i - E_f| = h\nu = \frachc\lambda ] Atomic spectra arise from the interaction of atoms
where ( h ) is Planck’s constant, ( \nu ) frequency, ( c ) speed of light, and ( \lambda ) wavelength. The result is a line spectrum — unique to each element.
Key concepts in atomic spectra include:
For hydrogen-like atoms, the Rydberg formula gives wavelengths: This exemplifies how a simple atomic transition becomes
[ \frac1\lambda = RZ^2 \left( \frac1n_1^2 - \frac1n_2^2 \right) ]
where ( R ) is the Rydberg constant and ( Z ) the atomic number.