The Lambert-Beer Law introduces the concept of absorbance (A) of a sample as $A=log\frac{I}{I_0}$. Where $I_0$ represents the intensity of the incident light and I the intensity of the light passing through the cell. We can also express absorbance as a function of cuvette length and solute concentration. \begin{equation} A=log\frac{I_0}{I}=\epsilon\cdot c\cdot l \end{equation} Where $l$ is the length of the cuvette in cm, $c$ represents the concentration of solute in mol/l and $\epsilon$ is the molar absorptivity (molar extinction coefficient) measured in l/mol.cm.

For a given concentration and cuvette length, the molar absorptivity determines whether the intensity of the band (absorbance) is high or low. It is very common to represent $log\epsilon$ in ordinates instead of the absorbance, in abscissas the wavelength is represented. To see the importance of the molar absorptivity coefficient, we will compare its value in the $\pi \rightarrow\pi^{\ast}$ transition of 1,3-butadiene ($\lambda =217\;nm$), which presents a $\epsilon =21000\;l/mol.cm$ ($log\epsilon=4.32$), with the transition $n\rightarrow \pi^{\ast}$ of acetone ($\lambda =280\ ;nm$) which presents $\epsilon =12\;l/mol.cm$ ($log\epsilon =1.08$). In the case of 1,3-butadiene an intense band is observed while in acetone it corresponds to a very low intensity band (forbidden transition). In general, those with a molar absorptivity of less than 100 l/mol.cm are considered prohibited transitions.