The Beer–Lambert Law Explained
The Beer–Lambert law is the principle that turns a spectrophotometer from a light meter into a quantitative analytical instrument. It explains why measuring absorbance lets you calculate exactly how much of a substance is in a solution.
The equation
The law states that absorbance (A) is directly proportional to concentration:
A = ε l c
- A — absorbance (unitless)
- ε — molar absorptivity, a constant unique to each substance at a given wavelength
- l — path length, the distance the light travels through the sample (typically 1 cm)
- c — concentration of the absorbing substance
Because ε and l are constant for a given measurement, absorbance rises in a straight line with concentration. Measure the absorbance and you can read off the concentration.
Building a calibration curve
In practice, you measure a series of standards of known concentration and plot absorbance against concentration. The resulting straight line — the calibration curve — is then used to determine unknown samples. A high correlation coefficient (R² approaching 0.9999) confirms an excellent linear fit and reliable quantification.
Where the law breaks down
The linear relationship holds well at low to moderate absorbance, but deviations appear when:
- Concentration is too high — molecules interact and the response is no longer linear.
- Stray light is significant — unwanted light reaching the detector flattens the curve at high absorbance.
- The sample scatters or fluoresces — adding signal that is not true absorbance.
For the most accurate results, work within the instrument’s validated linear range, keep stray light low, and dilute highly concentrated samples into the reliable region of the curve.