Say that an atom has the following energy levels:
E1 = -9.20 eV
E2 = -6.20 eV
E3 = -3.70 eV
E4 = -2.40 eV
E5 = -1.60 eV
Suppose atoms in the E2 level are exposed to incident radiation of wavelength 413.3 nm. Would this radiation be absorbed by atoms in the E2 level?
I am looking for some clarification on this problem if possible...
The Attempt at a Solution
E2->3 = -2.5 eV
E2->4 = -3.8 eV
E = hf = 1240/413.3 = 3 eV (radiation energy)
The atom starts out in the n = 2 energy level. According to my professor, the atom would NOT absorb the radiation because it's not exactly 2.5 eV (to go up to n = 3).
1) Why is it that the photon has enough energy to go up to n = 3 but the atom won't absorb the photon? Is it just because of quantization of energy? It seems strange to me that it wouldn't absorb.
My second question is if the wavelength is 496 nm then the photon will be absorbed by the n = 2 atom. Then the atom could emit wavelengths from 3->1, 2->1, and 3->2.
3->1 = 5.5 eV
2->1 = 3 eV
3->2 = 2.5 eV
2) Why could the atom emit 3->1 since it takes 5.5 eV and the atom only absorbed 2.5 eV? As well as 2->1 takes 3 eV?