what is the difference between excitation and emission spectroscopy?
Absorption:


Light with a specific frequency is selected
This frequency is shined (radiated) onto a molecule (sample)
The amount of light that shines through the molecule is measured, ie. the detector is behind the sample (180 degrees).
Step 1 is repeated with a different frequency.


When the sample absorbs a specific frequency, less light of that frequency will shine through to the detector. The machine records this and a peak appears in the spectrum for that frequency.

Frequency is usually measured in wavelength, and a wide range of wavelengths are applicable. In UV-VIS, this can be eg. 300-800 nm, so 800 nm is measured, then 799 nm, then 798 nm...down to 300 nm. In IR, you will go into micrometers. In X-ray, you will be around 10 nm I think.

Emission:


A light with a specific frequency is selected
The sample is irradiated with this frequency.
The amount of light that shines from the molecule is measured, ie. the detector is to the side of the sample (90 degrees).
The emitted light is measured for a range of frequencies
Step 1 is repeated with a different frequency.


In both processes, the sample absorbs light. The major difference here is the detector being located to the side of the sample. This allows it to ignore the light being shined onto the molecule, and instead measure light that is emitted from the sample. Think of the incident light as a column. In absorption spectroscopy, that column is absorbed by the sample and continues to the other side, where the detector can measure it. In emission spectroscopy, you want to avoid the column of light. Think of the sample like a light bulb that shines light in all directions. The detector in fluorescence is placed to the side, avoiding the column of light and picking up the light from the sample instead. The advantage to fluorescence is you can often selectively measure a moiety (part of the molecule) with superior sensitivity, whereas in absorption impurities or stray light are more prone to interfere with the signal and also leads to higher concentrations usually being necessary.

The reason the emission wavelength detected is different than the frequency used to irradiate it is due to Stoke's Shifts.

Stoke's Shift

Stoke's Shift requires basic knowledge of the difference between electronic and vibrational states. A molecule has electronic states, S=0, 1, 2, 3...etc. Within each electronic state is a set of vibrational states, v=1,2,3.... For example, two electrons can be in the 0 electronic state (termed ground state), but one electron can be in the v=2 vibrational state, while the other electron in v=4.

This is easier to picture if you look up Stoke's Shift, or specifically the diagram on Wikipedia's page for Fluorescence.

For example, in emission spectroscopy, an electron absorbs light, being excited electronically from S=0, v=1, to S=1, v=2. Vibrational relaxations will occur before electronic relaxations. That means the electron will first relax from S=1, v=2 to S=1, v=0. After this, it relaxes to S=0, emitting a photon in the process for detection. Because it relaxed its vibrational state first, energy was lost not in the form of light (thermodynamically). As a result, the photon finally emitted is at a lower energy than the photon that had been absorbed. That's the rudimentary explanation.