Line Spectra of Hydrogen Atom

Hydrogen shows many spectral lines because of the following reasons. 

• Its electron can occupy many levels (n = 1, 2, 3.). 

• Each line in the hydrogen spectrum actually represents the transition from higher to lower energy levels. 

• These lines are grouped by the final energy level (Lyman n'=1, Balmer n'=2, etc.) 

• Higher energy levels are closer together. They tend to create more possible transitions.

Hydrogen produces a line spectrum because electrons exist only in discrete, quantised energy levels. When electrons jump between these fixed energy states, they emit photons with specific energies (E = h? ). This creates distinct spectral lines instead of continuous wavelengths. Bohr's model explains this through quantised orbits. Classical physics, on the other hand, would predict a continuous spectrum.

The hydrogen emission spectrum contains several spectral series, each named after its discoverer.

• Lyman series (n' = 1): To ground state, visible only in ultraviolet region
• Balmer series (n' = 2): Transitions to second level, appearing in visible region
• Paschen series (n' = 3): Moved to third level, visible in the infrared region
• Brackett series (n' = 4): Transitions to fourth level, appearing in the far infrared region
• Pfund series (n' = 5): Transitions to fifth level, showing in the infrared region
• Humphreys series (n' = 6): Transitions to sixth level, appearing in the infrared region