Calculating the Hydrogen Spectrum: Understanding Energy Levels and Transitions"

[Badrakhandama]

Hi,
I have a question concerning a hydrogen spectrum:

There are 4 energy levels drawn:
Starting from the top

-1.4 * 10^-19 Joules
-2.4 * 10^-19 Joules
-5.4 * 10^-19 Joules
-21.8 * 10^19 Joules

1st question:
Going from the second level to the third from the top (from -2.4 to -5.4) gives rise to a red line int he hydrgen emission spectrum. Calculate the wavelength. I have done this, and found it to be 6.63 * 10^-7 metres

2nd Question: This is where I am confounded!

Draw two arrows on the diagram labelled IR and the ther UK, to show the transitions giving rise to lines in the infrared and ultraviolet regions. Explain how you made your choice for both...

I know that Infrared usually has a wavelength of 10^-5 metres, and UV has 10^-8 metres.
From this, I tried to find what transition is required:

E2-E1 = Planck's COnstant * Frequency (Velocity/Wavelength).


But I do not think this is correct. HELP ANYBODY!

Thank You.

 

[Stonebridge]

You're doing fine.
The UV transition will involve a larger energy change, and the IR transition a smaller one than the calculation you did originally. The options are limited with those 4 levels you have so it shouldn't be hard to chose one each that fits.

 

[Badrakhandama]

But if UV transition requires a greater energy change, then why is it nt frm -1.4 to -5.4 * 10^-19?

Thanks

 

[Stonebridge]

All transitions to -5.4 (The Balmer Series) are in the visible.
The answer is that the transition you mention, although a larger one, is still not enough to give UV.
All transitions to -21.8 are UV (The Lyman Series), and all transitions to -2.4 are IR (Paschen Series).