Ultraviolet - Visible Spectroscopy (UV)

Ultraviolet - Visible

Spectroscopy (UV)

Introduction to Ultraviolet -

Visible Spectroscopy

(UV)

Background Theory

Absorption of ultraviolet and visible radiation

Absorption of visible and ultraviolet (UV) radiation is

associated with excitation of electrons, in both atoms

and molecules, from lower to higher energy levels.

Since the energy levels of matter are quantized, only light

with the precise amount of energy can cause transitions

from one level to another will be absorbed.

The possible electronic transitions that light might cause are

In each possible case, an electron is excited from a full

(low energy, ground state) orbital into an empty

(higher energy, excited state) anti-bonding orbital.

Each wavelength of light has a particular energy

associated with it. If that particular amount of energy is

just right for making one of these electronic transitions,

then that wavelength will be absorbed.

The larger the gap between the energy levels, the greater

the energy required to promote the electron to the higher

energy level; resulting in light of higher frequency,

and therefore shorter wavelength, being absorbed.

All molecules will undergo electronic excitation following

absorption of light, but for most molecules very high

energy radiation (in the vacuum ultraviolet, <200 nm) is

required. Consequently, absorption of light in the UV-visible

region will only result in the following transitions:

Therefore in order to absorb light in the region from

200 - 800 nm (where spectra are measured), the molecule

must contain either

V bonds or atoms with non-bonding

orbitals. A non-bonding orbital is a lone pair on, say,

oxygen, nitrogen or a halogen.

π bonds are formed by sideways overlap of the half-filled

p orbitals on the two carbon atoms of a double bond.

The two red shapes shown in the diagram below for

ethene are part of the same π bonding orbital.

Both of the electrons are found in the resulting π bonding

orbital in the ground state.

1.Also known as UV-visible spectrophotometry

2.For further reading about bonding refer to http://www.chemguide.co.uk/analysis/uvvisiblemenu.html#top

Molecules that contain conjugated systems,

i.e. alternating single and double bonds, will have their

electrons delocalised due to overlap of the p orbitals in the

double bonds. This is illustrated below for buta-1,3-diene.

Benzene is a well-known example of a conjugated

system. The Kekulé structure of benzene consists of

alternating single double bonds and these give rise to the

delocalised

V system above and below the plane of the

carbon – carbon single bonds.

As the amount of delocalisation in the molecule increases

the energy gap between the π

bonding orbitals and π

anti-bonding orbitals gets smaller and therefore light of

lower energy, and longer wavelength, is absorbed.

Although buta-1,3-diene absorbs light of a longer

wavelength than ethene it is still absorbing in the UV

region and hence both compounds are colourless.

However, if the delocalisation is extended further the

wavelength absorbed will eventually be long enough to

be in the visible region of the spectrum, resulting in a

highly coloured compound. A good example of this is the

orange plant pigment, beta-carotene – which has

11 carbon-carbon double bonds conjugated together.

Beta-carotene absorbs throughout the UV region but

particularly strongly in the visible region between 400

and 500 nm with a peak at 470 nm.

Groups in a molecule which consist of alternating single

and double bonds (conjugation) and absorb visible light

are known as chromophores.

Transition metal complexes are also highly coloured,

which is due to the splitting of the d orbitals when

the ligands approach and bond to the central metal ion.

Some of the d orbitals gain energy and some lose

energy. The amount of splitting depends on the central

metal ion and ligands.

The difference in energy between the new levels affects

how much energy will be absorbed when an electron is

promoted to a higher level. The amount of energy will

govern the colour of light which will be absorbed.

For example, in the octahedral copper

complex, [Cu(H

]

yellow light has sufficient energy to

promote the d electron in the lower

energy level to the higher one.

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INTRODUCTION 3ULTRAVIOLET  VISIBLE SPECTROSCOPY UV

It is possible to predict which wavelengths are likely to

be absorbed by a coloured substance. When white light

passes through or is reflected by a coloured substance,

a characteristic portion of the mixed wavelengths is

absorbed. The remaining light will then assume the

complementary colour to the wavelength(s) absorbed.

This relationship is demonstrated by the colour wheel

shown on the right. Complementary colours are

diametrically opposite each other.

UV-visible spectrometers can be used to measure the

absorbance of ultra violet or visible light by a sample,

either at a single wavelength or perform a scan over a

range in the spectrum. The UV region ranges from 190

to 400 nm and the visible region from 400 to 800 nm.

The technique can be used both quantitatively and

qualitatively. A schematic diagram of a UV-visible

spectrometer is shown above.

The light source (a combination of tungsten/halogen

and deuterium lamps) provides the visible and near

ultraviolet radiation covering the 200 – 800 nm.

The output from the light source is focused onto the

diffraction grating which splits the incoming light into

its component colours of different wavelengths,

like a prism (shown below) but more efficiently.

For liquids the sample is held in an optically flat, transparent

container called a cell or cuvette. The reference cell or

cuvette contains the solvent in which the sample is

dissolved and this is commonly referred to as the blank.

For each wavelength the intensity of light passing

through both a reference cell (I

) and the sample cell (I)

is measured. If I is less than I

, then the sample has

absorbed some of the light.

The absorbance (A) of the sample is related to I and I

according to the following equation:

The detector converts the incoming light into a current,

the higher the current the greater the intensity. The chart

recorder usually plots the absorbance against wavelength

(nm) in the UV and visible section of the electromagnetic

spectrum. (Note: absorbance does not have any units).

UV-Visible Spectrometer

MONOCHROMATOR

BEAM

SPLITTER

SAMPLE

CELL

REFERENCE

CELL

RATI O

DETECTORS

PRISM

SOURCE

Red

Orange

Yellow

Green

Blue

Violet

UV Region

Visible Region Colour Wheel

Blue

Violet

Red

Orange

Green

Yellow

800 nm - Visible Region

400 nm190 nm

620 nm

590 nm

570 nm

495 nm

450 nm

If a substance

absorbs here...

...it appears

as this colour

Red 620-750 nm

Orange 590-620 nm

Yellow 570-590 nm

Green 496-570 nm

Blue 450-495 nm

Violet 380-450 nm

400 nm

Red

Orange

Yellow

Green

Blue

Violet

UV Region

Visible Region Colour Wheel

Blue

Violet

Red

Orange

Green

Yellow

800 nm - Visible Region

400 nm190 nm

620 nm

590 nm

570 nm

495 nm

450 nm

If a substance

absorbs here...

...it appears

as this colour

Red 620-750 nm

Orange 590-620 nm

Yellow 570-590 nm

Green 496-570 nm

Blue 450-495 nm

Violet 380-450 nm

400 nm

UV-Visible Spectrum

The diagram below shows a simple UV-visible absorption

spectrum for buta-1,3-diene. Absorbance (on the vertical axis)

is just a measure of the amount of light absorbed.

One can readily see what wavelengths of light are

absorbed (peaks), and what wavelengths of light are

transmitted (troughs). The higher the value, the more

of a particular wavelength is being absorbed.

The absorption peak at a value of 217 nm, is in the ultra-violet

region, and so there would be no visible sign of any light being

absorbed making buta-1,3-diene colourless. The wavelength

that corresponds to the highest absorption is usually referred

to as “lambda-max” (

lmax).

The spectrum for the blue copper complex shows that the

complementary yellow light is absorbed.

The Beer-Lambert Law

According to the Beer-Lambert Law the absorbance is

proportional to the concentration of the substance in

solution and as a result UV-visible spectroscopy can also

be used to measure the concentration of a sample.

The Beer-Lambert Law can be expressed in the form of

the following equation:

A =

ecl

Where

A = absorbance

l = optical path length, i.e. dimension of the cell

or cuvette (cm)

c = concentration of solution (mol dm

-3

)

e = molar extinction, which is constant for a

particular substance at a particular

wavelength (dm

mol

-1

)

If the absorbance of a series of sample solutions of

known concentrations are measured and plotted

against their corresponding concentrations, the plot of

absorbance versus concentration should be linear

if the Beer-Lambert Law is obeyed. This graph is known

as a calibration graph.

A calibration graph can be used to determine the

concentration of unknown sample solution by measuring

its absorbance, as illustrated below.

Since the absorbance for dilute solutions is directly

proportional to concentration another very useful

application for UV-visible spectroscopy is studying

reaction kinetics. The rate of change in concentration of

reactants or products can be determined by measuring

the increase or decrease of absorbance of coloured

solutions with time. Plotting absorbance against time one

can determine the orders with respect to the reactants and

hence the rate equation from which a mechanism for the

reaction can be proposed.

INTRODUCTION 4ULTRAVIOLET  VISIBLE SPECTROSCOPY UV

Maximum absorption at this wavelength

absorbance

1.0

0.8

0.6

0.4

0.2

200 220 240 260 280 300

wavelength (nm)

max=217nm

absorbance

350 400 450 500 550 600 650 700

wavelength (nm)

Modern Applications of UV Spectroscopy

UV-visible spectroscopy is a technique that readily allows one

to determine the concentrations of substances and therefore

enables scientists to study the rates of reactions, and determine

rate equations for reactions, from which a mechanism can

be proposed. As such UV spectroscopy is used extensively

in teaching, research and analytical laboratories for the

quantitative analysis of all molecules that absorb ultraviolet

and visible electromagnetic radiation.

Other applications include the following:

• In clinical chemistry UV-visible spectroscopy is used

extensively in the study of enzyme kinetics.

Enzymes cannot be studied directly but their activity can

be studied by analysing the speed of the reactions which

they catalyse. Reagents or labels can also be attached to

molecules to permit indirect detection and measurement

of enzyme activity. The widest use in the field of clinical

diagnostics is as an indicator of tissue damage.

When cells are damaged by disease, enzymes leak into

the bloodstream and the amount present indicates the

severity of the tissue damage. The relative proportions of

different enzymes can be used to diagnose disease,

say of the liver, pancreas or other organs which

otherwise exhibit similar symptoms.

• UV-visible spectroscopy is used for dissolution testing

of tablets and products in the pharmaceutical industry.

Dissolution is a characterisation test commonly used by

the pharmaceutical industry to guide formulation design

and control product quality. It is also the only test that

measures the rate of in-vitro drug release as a function of

time, which can reflect either reproducibility of the product

manufacturing process or, in limited cases,

in-vivo drug release.

• In the biochemical and genetic fields UV-visible

spectroscopy is used in the quantification of DNA

and protein/enzyme activity as well as the thermal

denaturation of DNA.

• In the dye, ink and paint industries UV-visible spectroscopy

is used in the quality control in the development and

production of dyeing reagents, inks and paints and the

analysis of intermediate dyeing reagents.

• In environmental and agricultural fields the quantification

of organic materials and heavy metals in fresh water can be

carried out using UV-visible spectroscopy.

INTRODUCTION 5ULTRAVIOLET  VISIBLE SPECTROSCOPY UV