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A Short Introduction to Ultraviolet-Visible Spectroscopy
This is an introduction to ultraviolet-visible spectroscopy.
Ultraviolet-visible or UV-vis spectroscopy has been used by chemists for a long time and was
really one of the first spectroscopic techniques developed. It turns out to be also the least
informative for structural analysis. So what? Why do we bother, you might ask? The answer
is that UV-vis spectroscopy tells us about the extent of conjugation in a compound, whether
we have a conjugated diene, an aromatic compound or something much more conjugated and
coloured.
In the electromagnetic spectrum you find that radiofrequencies, microwaves and infrared
radiation all have longer wavelengths than visible light. Ultraviolet light and X-rays are found
at the other end of the electromagnetic spectrum and have even shorter wavelengths. The
visible range spans from 800 nanometres (this is red) to 400 nanometres (which is blue or,
more precisely, violet). The UV region extends from 400 nanometres down to 200
nanometres.
If you wish to record a UV-vis spectrum, you will need a special UV cuvette. The more
expensive sort is made out of quartz. It is transparent not only to visible light but also to UV
light. A typical sample for a UV-vis measurement consists of a quite dilute solution
containing roughly about 1 mg of compound in 100 mL of solvent. The sample is filled into a
UV cuvette, and the cuvette is then placed inside a UV spectrophotometer. In addition, you
will need a reference which tends to be a second cuvette filled with only the solvent.
Whereas the reference does not absorb light, our sample may do so. The intensity I of the light
that passes through the sample is therefore smaller than the intensity I0 of the light going
through the reference. The UV spectrophotometer then calculates the absorbance A which is
defined as the logarithm of I0 over I.
The absorbance also equals the product of the molar absorptivity epsilon () times the
concentration c of the sample times the pathlength l of the UV cell. This is called the "BeerLambert Law". Both the molar absorptivity and the pathlength are constants. Incidentally, the
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pathlength of the most common UV cell is exactly 1 cm. Because the absorbance is, in effect,
proportional to the concentration of the sample, a UV-vis measurement therefore allows us to
determine the concentration of a sample provided that we know . Chemists and biochemists
often use UV-vis to measure concentrations, particularly when analyte concentrations become
very small.
While infrared light makes molecules vibrate, ultraviolet light has sufficient energy to cause
electronic transitions in conjugated compounds. When light is absorbed by a conjugated pisystem (-system), an electron is promoted from the a -orbital to the next higher energy
state, which happens to be an antibonding pi star (*) orbital. We call this a  to * transition.
With increasing conjugation less energy is required to see such a transition.
A UV-vis spectrophotometer scans the wavelengths in the visible and the UV region. It then
produces a plot of absorbance against wavelength, which we call a UV-vis spectrum. The
visible region is of interest only if our sample is coloured to the eye. If the sample is not and
appears "colourless", then we will be content with recording a UV spectrum between 400
nanometres and 200 nanometres.
You see here a UV spectrum with a single absorption peak. Other compounds will have more
than one absorption peak. Some may even show fine structure due to vibrational transitions,
but in most cases absorption peaks tend to be broad and any fine structure is usually blurred
out. The information you get from a UV spectrum is :
1) The wavelength at which the compound shows maximum absorbance, also called
lamda max (max). Note that some compounds may have several such maxima.
And
2) The molar absorptivity  which tells us whether an absorption peak is strong or
weak. No UV spectrophotometer can measure absorbances >2 accurately so that
strong and weak absorption peaks often have to be recorded using different sample
concentrations. Since  values can vary by several orders of magnitude, most
textbooks therefore provide you with UV-vis spectra where log  is plotted against
the wavelength.
A typical value for max for a conjugated diene is 220 nanometres, whereas for a benzene it is
about 260 nanometres. More conjugated -systems like this sunscreen or anthracene will have
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a max above 300 nanometres. We call the part of the molecule responsible for UV-vis
absorptions a chromophore.
Many substituents have an effect on max. An alkyl group, for example, causes only a small
shift by about 5 nm, whereas an additional conjugated C=C double bond moves the
absorption maximum 30 nm to higher wavelengths. We call this a "red shift". The opposite
would be a "blue shift", when the absorption peak moves towards smaller wavelengths.
Benzene, for example, has its highest absorption wavelength at 254 nm. A benzene with an
amino substituent absorbs only a little higher, at 280 nm. However, there are some
combinations of substituents which give rise to an unusually large "red shift". This is the case
when an electron-donating and an electron-withdrawing substituent are placed in para
position to each other on a benzene ring. So, while the amino group in 4-nitroaniline donates
electron density to the benzene ring, the nitro group at the other end withdraws it which gives
rise to an additional resonance structure. Such a push-pull arrangement between an electrondonating and an electron-withdrawing group results in a considerable red shift. "Push-pull"
systems are a common feature of many dyes.
You may have noticed that we did not dwell at all on the UV spectra of simple alkenes. The
reason is simple: the max of an alkene is 180 nanometres and cannot be observed in solution.
Even solvents without any double bonds, such as ethanol or water, start to absorb UV light at
around 200 nanometres.
Now, let's look at another special case: aldehydes and ketones. Like alkenes, carbonyl
compounds show a  to * star transition for the C=O double bond at around 180 nanometres,
which is well below the usual solvent cut-off. However, an electron from a lone pair at the
oxygen of an aldehyde or ketone can also get promoted into the * star orbital of the  bond.
Such an absorption is called an n to pi star (n to *) transition. The n to * star transition
requires less energy than a  to * transition and, consequently, causes a red shift in
absorption by about 80 – 100 nanometres. So, in the case of an ,-unsaturated carbonyl
compound you would expect to see two UV absorbances: one at 230 nanometres for the  to
* transition of the conjugated -system and one at around 310 nanometres for the n to *
transition involving the lone pair from the oxygen of the carbonyl group. Note that the  to *
transition is strong, whereas the n to * transition is very weak.
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More often than not, we tend to use UV spectroscopy without even thinking much about it.
Let's take the case of thin layer chromatography or TLC. After having developed a TLC,
chemists often use a UV lamp to visualise compounds on the TLC plate. Similarly, highperformance liquid chromatography or HPLC. The UV detector in an HPLC detects
compounds with a chromophore. It even quantifies how much compound we have, thus
becoming a useful tool not only in structural but also in quantitative analysis.
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