Measurement Errors Deviation Causes in Absorbance Quiz

At high concentrations typically above 0.01 M the average distance between molecules decreases. This leads to molecular interactions that alter the electronic environment and charge distribution of the analyte which changes its ability to absorb light at the selected wavelength.

The Beer Lambert Law assumes a constant refractive index. However in concentrated solutions the refractive index changes significantly compared to the pure solvent. This physical change alters the molar absorptivity and leads to a nonlinear relationship between absorbance and concentration.

If an analyte exists in equilibrium between two forms with different colors a change in concentration may shift the equilibrium. This results in a non proportional change in the concentration of the light absorbing species leading to a curved calibration plot.

Stray light is any radiation reaching the detector that has not passed through the sample at the selected wavelength. It adds to the transmitted light intensity which the instrument interprets as a lower absorbance. This effect is most severe at high absorbance levels.

Instrumental deviations arise from the limitations of the spectrophotometer. The Law assumes monochromatic light. Using a wide slit or a source with multiple wavelengths causes the molar absorptivity to vary resulting in negative deviations from linearity.

Suspended particles or bubbles scatter light in all directions. Since less light reaches the detector the instrument assumes this light was absorbed by the analyte. This creates a falsely high absorbance reading that does not represent the actual chemical concentration.

The Law is strictly valid only for monochromatic light. Because the molar absorptivity is different for every wavelength in a broad beam the total absorbance is no longer directly proportional to the concentration of the analyte.

If a substance absorbs light and then re emits it as fluorescence the detector may pick up this emitted light as if it were transmitted light. This makes the sample appear less absorbent than it actually is leading to an underestimate of the concentration.

At the peak of a broad absorption band the molar absorptivity is relatively constant over a small range of wavelengths. This minimizes the errors caused by the finite slit width of the instrument and ensures the best adherence to the linear law.

Any process that changes the chemical identity of the analyte as its concentration changes will cause a deviation. Solvent evaporation changes the concentration itself but polymerization or ionization actually changes the nature of the absorbing species.

At high absorbance very little light reaches the detector. In this state even a tiny amount of stray light becomes a large percentage of the total detected signal causing the absorbance reading to flatten out or drop off significantly on the calibration curve.

As a rule of thumb the Law is most reliable when concentrations are below 0.01 M and absorbance values are between 0.1 and 1.0. Outside of these ranges physical and chemical interactions make the linear relationship less dependable.

If the sample cuvette and the blank cuvette have different path lengths or different optical purities the zeroing process will be imperfect. This adds a constant offset to every reading which can be seen as a non zero intercept on the calibration curve.

Diluting samples brings them into the linear range. Narrower slits provide light that is closer to being monochromatic. Filtering removes light scattering particles. Using a broader light source would actually increase deviations.

Negative deviations where the slope decreases as concentration increases are typically caused by instrumental factors like stray light or the use of light that is not perfectly monochromatic.