What is UV-Visible Spectroscopy ?

Spectroscopy refers to the absorption or emission of light or other radiations by matter. When a substance or chemical absorbs light or radiation, they reach an excited state followed by de-excitation or emission of the light or radiation forming distinct spectra. UV-visible spectroscopy or  ultraviolet-visible refers to the absorption and emission in the ultraviolet and visible regions of the electromagnetic spectrum.

The interaction of electromagnetic radiation/ light and matter can be studied using spectroscopy or spectrophotometry. Spectroscopic techniques can be used to identify substances by analyzing the spectrum of radiation absorbed or emitted by them. Spectrophotometers are routinely used to estimate the concentration of coloured compounds, analysis of the structure of chemicals, molecular analysis of intermediates formed in an enzyme-catalyzed reaction and intermolecular bonding.

Spectrophotometer

A spectrophotometer is a device used to measure the amount of light passing through the substance of interest. ‘Spectro’ refers to the range of wavelengths that the light source produces. The UV-visible spectrophotometer is a variation of a colorimeter. A colorimeter is used to measure the transmittance and the absorbance of light passing through a coloured solution of the test sample. The colour is produced by the addition of specific reagents. The intensity of the colour is directly proportional to the concentration of the substance. The colorimeter uses a broad range of wavelengths produced using filters.

A UV-visible spectrophotometer uses a monochromatic light produced by the grating of a prism. The components of a UV-visible spectrophotometer include a light source with UV or visible light, a wavelength selector or a monochromator, a cuvette or sample container, and a digital display.

The Beer-Lambert law

Beer-Lambert law defines the relationship between the absorbance or attenuation of a light that is passed through a substance and the properties of the substance.  This was proposed in 1852 by August Beer, a German mathematician and chemist.  Beer’s law states that the absorbance of a substance dissolved in a solution is directly proportional to its concentration and path length. Hence, it can be applied only when there is a linear relationship. It is mathematically represented as follows.

$$A=\epsilon lc$$

where, A = absorbance, ε = molar extinction coefficient or molar absorptivity or absorption coefficient, l = path length and c = concentration.

Transmittance and absorption

Transmittance is the intensity of light that reaches the detector after passing through the sample (I) divided by the intensity of the incident light (I₀).

$$T=\frac{I}{I_0}$$

Here, T = transmittance, I = intensity of the light coming out of the sample and I₀ = intensity of the incident light.

Absorbance is equal to the negative log of transmittance and can be represented as follows: transmittance (T) and absorbance (A) can be expressed as:

$$A=-log\left(T\right)$$

Therefore,

$$A=\epsilon lc=log\left(\frac{I}{I_0}\right)$$

The concentration of a substance in a solution can be calculated if the absorption coefficient for a given wavelength and the thickness of the path length for light transmitted through the solution is known.

UV-Visible spectra

Sunlight or white light is perceived as a uniform homogenous colour. However, it is composed of a broad range of radiations with wavelengths in the ultraviolet (UV), visible and infrared (IR) portions of the spectrum. The colours of the visible spectrum can be separated by passing sunlight or white light through a prism which disperses it into its component colours.

Wavelength

Electromagnetic radiations like visible light is considered as a wave characterized by a wavelength or frequency. The distance between adjacent peaks or troughs is the wavelength. Visible wavelength ranges from 400 to 800 nm. Red and violet are the longest and shortest visible wavelengths respectively. Other colors of the spectrum, in the order of increasing wavelength, may be remembered using the mnemonic: VIBGYOR

VIBGYOR

Frequency

Frequency is the number of wave cycles that travel past a fixed point per unit of time and is designated in cycles per second, or hertz (Hz).

UV-Visible Excitation and Absorption Spectra

Molecular orbitals

In atomic theory, the atomic orbital (AO) is a mathematical function that describes the probability of finding an electron of an atom in a specific region around the nucleus. Each orbital of an atom is defined by three unique parameters namely the energy of the electron, angular momentum and the magnetic quantum number or the vector component of the angular momentum. The orbitals are sharp, principal, diffuse and fundamental and are simply referred to as s, p, d, and f orbitals.

Similar to the atomic orbitals, the shared electron pairs of covalently bonded atoms may be considered to occupy molecular orbitals (MO). In general, two atomic orbitals can be combined to form a molecular orbital. For example, in a hydrogen molecule, the two 1s atomic orbitals combine to form a sigma (s) bond molecular orbital with low energy and a second higher energy orbital known as an anti-bonding orbital. The bonding orbital contains two electrons of opposite spin, which forms the covalent bond. Similar sigma bonds may be formed by the overlapping between p-orbitals.

Another type of MO is known as the π-orbital, which is formed by the lateral overlap of two p-orbitals. These bonds are weaker than sigma bonds and are known as pi-bonds.

Many organic molecules undergo electronic transitions when they interact with high-energy radiations in the UV and visible range of the electromagnetic spectrum. When energy from the UV or visible light is absorbed by a molecule, one of its electrons jumps from the lower to the higher energy molecular orbital.For example, a hydrogen molecule at ground state would have both the electrons in the lower energy bonding orbital or the sigma MO (σ MO). This is known as the highest occupied molecular orbital (HOMO). When excited, the electrons transition to the anti-bonding orbital (σ* MO) or the lowest unoccupied molecular orbital (LUMO). This is known as a σ - σ* transition. Similarly, when molecules containing double bonds absorb light, they undergo a π - π* transition. In the non-bonding orbitals or n-orbitals, the n - n* transition energy gaps are lower than the σ - σ* transition. Hence the wavelength of light absorbed is longer. This allows an important application of UV-visible spectroscopy to study organic compounds with multiple double bonds or conjugated pi-systems. In these molecules, the energy gap in the n - n* transitions is much narrower than individual double bonds. Therefore, these molecules absorb light at higher wavelengths strongly. These molecules that absorb strongly at UV-visible region are known as chromophores. The n- π*  and the π - π* transitions have the lowest energy transitions and are achieved within the UV-visible spectrum.  In general, the promotion of an electron from the HUMO to the LUMO is an energetically favoured transition and the resulting species is said to be in an excited state.

Contexts and Applications

This topic is useful for various undergraduate and postgraduate courses, especially for, Bachelors and Masters in Chemistry.