Spectrometry Basics

Study Notes

Spectrometry analyzes molecules by observing their interaction with light, employing electromagnetic radiation.
Energy in electromagnetic (EM) spectrum is propagated in waveforms, utilizing various light forms like UV, visible, and infrared (IR) radiation.

Wavelength represents the distance between wave peaks, measured in micrometers (um) or nanometers (nm).
Wavelengths characteristic of specific regions:
- UV: 200-380 nm
- Visible: 380-780 nm
- Near IR: 780-3000 nm
- Medium IR: 3-15 um
- Far IR: 15-300 um

Energy in the EM spectrum is quantized, occurring in bundles called photons or quanta.
Planck’s constant describes the spectral density of electromagnetic radiation from a black body at thermal equilibrium at temperature T, where energy content determines chemical absorption or transmission.
Absorbance refers to light energy absorbed by photons, while transmission indicates the light exiting the sample.

Molecules exist in specific energy states and cannot have arbitrary energy levels.
Energy absorption moves molecules from an initial state to a higher energy state.
Internal energy (E) of a molecule at a given temperature is the sum of electronic (Ee), vibrational (Ev), and rotational (Er) energies, plus translational energy (Et), expressed as E = Ee + Ev + Er + Et.
Translational energy does not apply in spectrometry; Ee is relevant in UV-visible regions, while Ev and Et pertain to the IR region.

Molecules respond primarily to UV (185-380 nm) and visible (380-780 nm) light.
Electronic excitation occurs in these spectra, affecting non-bonding (n) electrons and sigma (σ) or pi (π) bonds.
Absorption of radiant energy elevates molecules to higher energy states in antibonding orbitals, termed σ* or π* for sigma and pi electrons, respectively.