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Difference in energy between visible line spectrum and ultraviolet?

  • Thread starter cheekygeek
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Homework Statement


The emission spectrum of an unknown element contains two lines - one in the visible portion of the spectrum, and the other, ultraviolet. Based on the following figure and what you know about Niels Bohr's model of the atom, account for the difference in energy between these lines.

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Homework Equations


n/a


The Attempt at a Solution


Bohr suggests that colours in the line spectra are emitted by electrons. Electrons orbit the nucleus in very specific orbits, and each orbit has a distinct energy. When the electron absorbs energy from some external source, it jumps (transitions) to a higher orbit or energy level. The attraction of the nucleus eventually pulls the electron back to a lower energy level and when it does the energy that the electron absorbed is emitted. In this case the line which is in the visible portion of the spectrum is produced when the electron is transitioning from a high level of orbit to a lower level of orbit and in turn releases energy that features longer wavelengths, which means that the line is lower in energy compared to the line which is emitting on the ultraviolet portion of the spectrum. The line, which is on the ultraviolet region of the spectrum, indicates that the electron is jumping from the highest (higher than the previous one) energy level of orbit to the lowest energy level of orbit and thus releasing more energy and shorter wavelengths.


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Thank you in advance!
 

Answers and Replies

  • #2
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In this case the line which is in the visible portion of the spectrum is produced when the electron is transitioning from a high level of orbit to a lower level of orbit and in turn releases energy that features longer wavelengths, which means that the line is lower in energy compared to the line which is emitting on the ultraviolet portion of the spectrum.
That part looks a bit strange. The first comparison ("longer") has nothing to compare with.
The main idea is right - the energy difference is larger for the transition which emits UV radiation.

The line, which is on the ultraviolet region of the spectrum, indicates that the electron is jumping from the highest(higher than the previous one)
Usually, that's not the dominant effect. The energy of the lower state (the "target" of the transition) is more important.
 

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