### Hydrogen spectrum

The emission spectrum of atomic hydrogen is divided into a number of spectral series, with wavelengths given by the Rydberg formula. These observed spectral lines are due to electrons moving between energy levels in the atom. The spectral series are important in astronomy for detecting the presence of hydrogen and calculating red shifts. Further series were discovered as spectroscopy techniques developed.

## Physics

In physics, the spectral lines of hydrogen correspond to particular jumps of the electron between energy levels. The simplest model of the hydrogen atom is given by the Bohr model. When an electron jumps from a higher energy to a lower, a photon of a specific wavelength is emitted.

The spectral lines are grouped into series according to Template:Mvar. Lines are named sequentially starting from the longest wavelength/lowest frequency of the series, using Greek letters within each series. For example, the 2 → 1 line is called "Lyman-alpha" (Ly-α), while the 7 → 3 line is called "Paschen-delta" (Pa-δ). Some hydrogen spectral lines fall outside these series, such as the 21 cm line; these correspond to much rarer atomic events such as hyperfine transitions.[1] The fine structure also results in single spectral lines appearing as two or more closely grouped thinner lines, due to relativistic corrections.[2] Typically one can only observe these series from pure hydrogen samples in a lab. Many of the lines are very faint and additional lines can be caused by other elements (such as helium if using sunlight, or nitrogen in the air). Lines outside of the visible spectrum typically cannot be seen in observations of sunlight, as the atmosphere absorbs most infra-red and ultraviolet wavelengths.The longest wavelength in the Lyman series of hydrogen 122nm.

## Rydberg formula

The energy differences between levels in the Bohr model, and hence the wavelengths of emitted/absorbed photons, is given by the Rydberg formula:[3]

$\left\{1 \over \lambda\right\} = R \left\left( \left\{1 \over \left(n^\prime\right)^2\right\} - \left\{1 \over n^2\right\} \right\right) \qquad \left\left( R = 1.097373 \times 10^7 \ \mathrm\left\{m\right\}^\left\{-1\right\} \right\right)$

where Template:Mvar is the upper energy level, Template:Mvar is the lower energy level, and Template:Mvar is the Rydberg constant.[4] Meaningful values are returned only when Template:Mvar is greater than Template:Mvar and the limit of one over infinity is taken to be zero.

## Series

All wavelengths are given to 3 significant figures.

### Lyman series (n′ = 1)

Main article: Lyman series
Template:Mvar λ (nm)
2 122
3 103
4 97.3
5 95.0
6 93.8
$\infty$ 91.2

The series is named after its discoverer, Theodore Lyman, who discovered the spectral lines from 1906–1914. All the wavelengths in the Lyman series are in the ultraviolet band.[5][6]

### Balmer series (n′ = 2)

Main article: Balmer series
Template:Mvar λ (nm)
3 656
4 486
5 434
6 410
7 397
$\infty$ 365

Named after Johann Balmer, who discovered the Balmer formula, an empirical equation to predict the Balmer series, in 1885. Balmer lines are historically referred to as "H-alpha", "H-beta", "H-gamma" and so on, where H is the element hydrogen.[7] Four of the Balmer lines are in the technically "visible" part of the spectrum, with wavelengths longer than 400 nm and shorter than 700 nm. Parts of the Balmer series can be seen in the solar spectrum. H-alpha is an important line used in astronomy to detect the presence of hydrogen.

### Paschen series (Bohr series) (n′ = 3)

Template:Mvar λ (nm)
4 1870
5 1280
6 1090
7 1005
8 954
9 923
10 902
11 887
$\infty$ 820

Named after the Austro-German physicist Friedrich Paschen who first observed them in 1908. The Paschen lines all lie in the infrared band.[8]

### Brackett series (n′ = 4)

Template:Mvar λ (nm)
5 4050
6 2620
7 2160
8 1940
9 1820
$\infty$ 1460

Named after the American physicist Frederick Sumner Brackett who first observed the spectral lines in 1922.[9]

### Pfund series (n′ = 5)

Template:Mvar λ (nm)
6 7460
7 4650
8 3740
9 3300
10 3040
$\infty$ 2280

Experimentally discovered in 1924 by August Herman Pfund.[10]

### Humphreys series (n′ = 6)

Template:Mvar λ (nm)
7 12400
8 7500
9 5910
10 5130
11 4670
$\infty$ 3280

Discovered by American physicist Curtis J. Humphreys.[11]

### Further (n′ > 6)

Further series are unnamed, but follow exactly the same pattern as dictated by the Rydberg equation. Series are increasingly spread out and occur in increasing wavelengths. The lines are also increasingly faint, corresponding to increasingly rare atomic events.

## Extension to other systems

The concepts of the Rydberg formula can be applied to any system with a single particle orbiting a nucleus, for example a He+ ion or a muonium exotic atom. The equation must be modified based on the system's Bohr radius; emissions will be of a similar character but at a different range of energies.

All other atoms possess at least two electrons in their neutral form and the interactions between these electrons makes analysis of the spectrum by such simple methods as described here impractical. The deduction of the Rydberg formula was a major step in physics, but it was long before an extension to the spectra of other elements could be accomplished.