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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 1 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 2 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. 3 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. 4