Light Through Matter
When a beam of light passes through a colored solution, some photons are absorbed by dissolved molecules. The amount of absorption depends on three factors: how many molecules the light encounters (concentration), how far it travels through the solution (path length), and how strongly each molecule absorbs at that wavelength (molar absorptivity). The Beer-Lambert law combines these into a single elegant equation: A = εcl.
Absorbance and Transmittance
Transmittance is the fraction of light that makes it through the sample — a perfectly clear solution has T = 1 (100%), while an opaque one has T ≈ 0. Absorbance is the negative logarithm of transmittance, converting the exponential decay into a linear relationship with concentration. This linearity is what makes spectrophotometry so powerful for quantitative chemical analysis.
The Molar Absorptivity Spectrum
Every molecule has a unique absorption spectrum — a plot of ε versus wavelength that acts as a molecular fingerprint. Chromophores (light-absorbing groups) like conjugated double bonds, aromatic rings, and metal d-d transitions each contribute characteristic absorption bands. The wavelength of maximum absorption (λ_max) and the peak ε value together identify and quantify unknown compounds in solution.
Practical Spectrophotometry
In the laboratory, spectrophotometers measure absorbance across a range of wavelengths. By preparing standard solutions of known concentration and measuring their absorbance, you build a calibration curve. Unknown concentrations are then determined by interpolation. This principle underlies clinical blood tests, environmental water monitoring, forensic analysis, and countless other analytical applications. This simulator lets you explore how changing each variable affects the transmitted light intensity.