Question

a) Is the energy absorption associated with bands in an infrared spectrum of higher or lower energy than the lines appearing in a visible line spectrum. Explain?
b) Identify the type of energy transition occuring in a molecule that causes a band to appear in an infrared spectrum.
c) Identify the type of energy transition occuring in an atom that causes a line to appear in a visible line spectrum.

Answer 1

In the electronmagnetic spectrum, the em radiations are arranged in order of increasing wavelengths or decreasing order of frequencies.

We know E = h$\nu$ and $\nu$=c/$\lambda$

Thus, E =hc/$\lambda$ where E is the energy associated with the electromagnetic radiation, $\lambda$ is the wavelength,$\nu$ is frequency of radiation and c is a constant (speed of light)

Please have a look at the attached image:

I have drawn an approximate range of frequencies of electromagnetic radiation (not to scale) to give a clear idea

a) Since the frequency of infrared radiations is lower than that of visible radiations and energy is directly proportional to frequency (as shown earlier ) thus energy associated with bands Of IR spectrum are lower than those in visible spectrum.

b) IR spectrum is based upon absorption of IR radiations by the mlecule in accordance with the quantisation of their Vibrational energy levels. (stretching and bending )

c)The electronic transition taking place in visible region are of following types:

$\sigma \rightarrow \sigma ^{*}$

$n \rightarrow \sigma ^{*}$

$n \rightarrow\pi ^{*}$

$\pi \rightarrow\pi ^{*}$

to get a general idea, please see the image:

ト* オケ 72 TL

(please note that the image is not drawn to scale, it is just to provide you with a rough idea)