Below is the energy level diagram representing the transitions made by an electr

Question

Below is the energy level diagram representing the transitions made by an electron in a hydrogen atom which result in the observed lines of both the absorption and emission spectra.

4 different energy photons are represented (approximate wavelengths are given in parentheses):

infrared (~ 10-4 m)
red (~ 10-6 m)
blue (~ 10-7 m)
ultraviolet (~ 10-8 m)

_________ Red photon absorption (~ 10-6 m)
_________ Blue photon emission (~ 10-7 m)
_________Highest energy emission
_________ Lowest energy absorption

Explanation / Answer

En = -13.6 /n^2 ----> electron present in in the n'th state

Electron in 1st state has highest energy

As the n increase ... the energy decreases

Highest energy emission = [E3 to E1 ] -----> h

The difference between ground state and its upper states will be higher than other ones

Lowest energy absorption = [E4 - E5] ----> a

The lowest energy line (which is the longest wavelength line) in the Balmer series appears in the red portion of the visible spectrum.

Blamer series -----> upper states to 2nd state transition

Visible spectrum is observed in Balmer series

The four lines in the question stands for four kinds of electron transition which are all in the range of visible light :
1) n=3 to n=2;(red)
2) n=4 to n=2;(green)
3) n=5 to n=2;(blue)
4) n=6 to n=2.(purple)

Red photon absorption - b

Blue photon emission - g

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