The discussion revolves around the spectral emissions of a candle flame, specifically focusing on the blue part of the spectrum and the violet emission at 432 nm attributed to excited CH* molecules. Participants are questioning the reasons behind specific wavelengths of emissions and their relation to photon energies.
The discussion is active, with participants exploring the relationship between photon energies and spectral emissions. Some guidance has been offered regarding the use of the energy equation, but multiple interpretations and questions about the nature of spectral lines remain open.
Participants are considering the discrete nature of spectral lines and questioning the underlying reasons for the specific wavelengths observed in both CH* emissions and hydrogen's Balmer lines. There is an emphasis on understanding the differences in energy levels of atoms and molecules compared to everyday objects.
Kamakiri said:Homework Statement
In the spectrum of the blue part in a candle flame, there’s a violet emission at 432 nm due to excited CH* molecules (chemiluminescence). Why 432? Why not 400 or 500? There are emissions at 436, 475 and 520 nm too. Why these numbers?
2. The attempt at a solution
Is it because the energies of the photons emitted correspond to these wavelengths, as E = hc/λ?
I read about Balmer lines. The H-alpha spectral line of hydrogen gas is red, since the energy of the photons emitted correspond to 656.3 nm, as E = hc/λ. Is that right?Quantum Defect said:Why do the Balmer lines in the Hydrogen spectrum have the wavelengths that they have? If you know the answer to that question, you know why CH* emits at the wavelengths that it does.
Kamakiri said:I read about Balmer lines. The H-alpha spectral line of hydrogen gas is red, since the energy of the photons emitted correspond to 656.3 nm, as E = hc/λ. Is that right?