[horseb0x]

I noticed that the hydrogen spectral lines are grouped into 6 series and given a value for n. I also noticed that each series was named after its discoverer but "coincidentally?" falls into a specific region of the EM spectrum so the Lyman series (n=1) of lines are all in the UV region, the Balmer series (n=2) in the visible region, the Paschen series (n=3) the IR region etc. Firstly is this "n" the principle quantum number? If so what have these series' got to do with the different energy shells of the Bohr model? For example what has the balmer series got to do with the 2nd energy shell? Finally what is it about this correlation that causes the lines of each series to appear where they do. For example why do the lines all appear in the UV region when n=1 but lie in the visible region when n=2?

[opti384]

When an electron in an excited state moves to the 1st energy shell (Lyman series), it emits UV rays.

[rabolisk]

First, the 'n' refers to the principal quantum number, n. When a hydrogen's electron is excited from n=1 to some other n ≥ 1, it absorbs some specific frequency of light. When it relaxes back down, it emits a light (visible or not) of some wavelength. Lyman series is when an electron moves from n ≥ 1 to n = 1. Balmer series is when an electron moves from n ≥ 2 to n = 2, and so on. The wavelength of the emitted EM wave can be calculated by the Rydberg equation, the derivation of which is beyond the scope of your chemistry. One thing to notice is that a UV ray, with a higher frequency and shorter wavelength, corresponds to a higher energy than a visible light. This means that the transition associated with the Lyman series (to n=1) corresponds to greater energy difference than the transition associated with the Balmer series (to n=2). Consult your textbook and the following link for more.

http://en.wikipedia.org/wiki/Hydrogen_spectral_series

[horseb0x]

Thanks a lot. That pretty much answers all my questions but a few more came to mind now. The number of spectral lines should significantly decrease as you move up n because there are less higher energy levels for electrons to fall from. For example shouldn't there be more lines in the Lyman series than the Balmer series because in the Balmer series the transition from 2 to 1 is an absorption whereas in the Lyman series its an emission. Is that how it is? Also would I be right in assuming that the larger the atom, the greater the amount of spectral lines? For example I assume Uranium with its 92 electrons and 7 energy shells should emit way more spectral lines than hydrogen because it has way more ground energy levels for electrons to fall back down to?

[jeffrey.struss]

You are forgetting there are theoretically an infinite number of levels.

I know, citing wikipedia and all that, but here we go. It lists the first 9 or so lines.

http://en.wikipedia.org/wiki/Lyman_series
http://en.wikipedia.org/wiki/Balmer_series

As you can see the Balmer series goes on in to UV.


Also, in the Balmer series the transition is going from 1 to 2 (not as you said from 2 to 1). Going UP in number is an absorption, going down is emission. Since the balmer series starts at 2 it can't go down to 1. n is always >2 for the balmer series.