Spectroscopic data undoubtedly provides a great deal of useful information about organic molecules. Competently deriving structural information from such data therefore, is a requisite skill for many undergraduates studying chemistry. Interpreting Organic Spectra covers the basic principles of spectroscopy in as non-mathematical a way as possible. It assumes no previous knowledge of spectroscopy and avoids excessive theory, approaching the topic as an exercise in pattern recognition. Hence the main focus of the book is in the provision of a variety of spectra for the student to interpret. Students are able to pace their progress by gaining confidence on the simpler spectra, and applying techniques learned to tackle more complex examples. As an introduction to the subject, it is ideal for A-level students as well as chemistry undergraduates and will prove to be a very useful reference tool for teachers and lecturers.

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Interpreting Organic Spectra

By David Whittaker

The Royal Society of Chemistry

Chapter 1 Infrared Spectroscopy, 1, 
Problems in Interpreting Infrared Spectra, 7, 
Chapter 2 Mass Spectrometry, 19, 
Problems in Interpreting Mass Spectra, 28, 
Chapter 3 Problems in Interpreting Infrared Spectra and Mass Spectra, 39, 
Chapter 4 Ultraviolet Spectroscopy, 51, 
Problems in Interpreting Ultraviolet Spectra, 55, 
Chapter 5 13C Nuclear Magnetic Resonance Spectroscopy, 59, 
Problems in Interpreting 13C NMR Spectra, 72, 
Chapter 6 Problems in Interpreting Infrared Spectra and 13C Nuclear Magnetic Resonance Spectra, 83, 
Chapter 7 Problems in Interpreting Mass Spectra and 13C Nuclear Magnetic Resonance Spectra, 95, 
Chapter 8 Problems in Interpreting Infrared Spectra, Mass Spectra, Ultraviolet Spectra and 13C Nuclear Magnetic Resonance Spectra, 107, 
Chapter 9 1H Nuclear Magnetic Resonance Spectroscopy, 149, 
Problems in Interpreting 1H NMR Spectra, 163, 
Chapter 10 Problems in Interpreting Infrared and 1H Nuclear Magnetic Resonance Spectra, 169, 
Chapter 11 Problems in Interpreting Mass Spectra and 1H Nuclear Magnetic Resonance Spectra, 181, 
Chapter 12 Problems in Interpreting Infrared Spectra, Mass Spectra, Ultraviolet Spectra, 13C Nuclear Magnetic Resonance Spectra and 1H Nuclear Magnetic Resonance Spectra, 193, 
Chapter 13 Difficult Problems in Interpreting Spectra, 235, 
Chapter 14 Answers to Problems, 257, 
Subject Index, 261,

CHAPTER 1

Infrared Spectroscopy

Infrared spectroscopy is the best means of identifying functional groups in a molecule. It involves measuring the absorption by a substance of radiation in the region from 4000 to 600 cm-1. The range covers roughly 2.5–16 �m (1 �m = 10-4 cm), but the frequency scale is now used universally.

Every bond in a molecule vibrates, resulting in a change in its dipole moment. This change in dipole moment provides a mechanism for the absorption of radiation. The energy of the vibration is such that the radiation absorbed is in the infrared region, which is of lower frequency, and hence lower energy, than visible light. Consequently, every bond in a molecule has an absorption peak in the infrared spectrum or the Raman spectrum of the molecule. Every substance therefore has its own unique infrared spectrum, so that we can identify any organic material by comparing its infrared spectrum with that of a known sample. In addition, each different functional group, such as O–H, C–H or C=C, absorbs within a narrow range of frequencies so that we can identify a functional group present in a molecule by the presence of an absorption band in a particular range of the infrared spectrum.

The frequency at which a bond absorbs radiation depends on the masses of the atoms forming the bond. The bonds which absorb radiation at the upper end of the frequency range are those which involve a light atom, hydrogen, with a heavier atom, such as nitrogen, oxygen or carbon. Hydrogen, being light, vibrates strongly and rapidly, so we see a strong, high energy absorption. As we move to lower frequencies, we come to the vibrations of two heavier atoms, such as C–N, C–O and C–C, and finally, at the lowest energies, we find the vibrations of the heavier atoms bonded to very heavy atoms, such as C–Cl, C–Br and C–I. Within this general trend, multiple bonds absorb higher energy radiation than single bonds, so the C [equivalent to] C bond absorbs at higher frequency than the C=C bond.

Before we study an infrared spectrum, we should consider how it is obtained. The sample to be studied is usually examined in one of four ways:

(1) As a pure liquid.

(2) As a mull. This method is used for solids. The solid is ground finely using a pestle and mortar, and mixed with a small amount of a liquid hydrocarbon (liquid paraffin) and run as a liquid. This has the disadvantage that the spectrum of the liquid hydrocarbon is super imposed on the spectrum of the sample.

(3) The compound is finely ground, mixed with K Br, compressed and run as a disc. This avoids adding extra absorption, but is time consuming.

(4) The compound is dissolved in a suitable solvent, and run as a solution. The spectrum cannot be observed in regions where the solvent absorbs.

Now let us consider an infrared spectrum. The spectrum of phenylamine, PhNH2 (also known as aniline), is shown in Figure 1.1.

The spectrum records the amount of radiation transmitted at each frequency, so the maximum absorption occurs when the least light is transmitted, and the recorder line is closest to the bottom of the spectrum. Positions of maximum absorption are difficult to measure accurately from the spectrum, so most spectrometers record them automatically, and print them alongside the spectrum, giving the % of radiation transmitted (%T) alongside. The machine can be set to record frequencies of all or only the stronger peaks.

Phenylamine has an -NH2 group, which characteristically absorbs in the region from 3500 to 3250 cm-1. The spectrum actually has two peaks in this range, at 3426 and 3352 cm-1. These do NOT represent a peak for each N-H bond (the bonds are indistinguishable so must have identical absorption frequencies): they result from the in-phase and out-of-phase vibrations of the N–H bonds:

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If the N–H bond is part of an amide, hydrogen bonding moves the absorption to lower frequency. If we have a secondary amine, with only one N–H bond, we obtain a single peak, which can easily be confused with the O–H bond.

A genuine O-H bond is shown in the next spectrum, that of butan-2-0l, CH3CH2CH(OH)CH3. This is shown in Figure 1.2. The O–H bond absorbs radiation in the range 3700–3200 cm-1, and we find that butan-2-01 has a peak at 3338 cm'. The O–H peak, like the N–H peak but unlike other peaks in the spectrum, is a broad rounded peak rather than the usual sharp spiked peaks. This is a result of intermolecular hydrogen bonding. The hydrogen atom on any particular oxygen atom is probably attached to another oxygen atom by hydrogen bonding, so bond vibrations vary over a frequency range, and the broad peak which we see is an envelope covering many absorptions at slightly different frequencies. If we dilute the sample of butan-2-01 with solvent which cannot form a hydrogen bond, such as dichloromethane, we get the spectrum shown in Figure 1.3, in which the broad O–H bond peak is reduced, and a small sharp absorption peak has appeared at higher frequency, resulting from the non-hydrogen bonded O–H bond. At greater dilution of the alcohol dissolved in dichloromethane, the original O–H peak has almost vanished, and the non-hydrogen bonded peak has further increased in size, as shown in Figure 1.4.

The O–H bond of a carboxylic acid shows even stronger hydrogen bonding than that of an alcohol, since the acid is so strongly hydrogen bonded as to exist in the dimeric form in the pure liquid:

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As a result, the carboxylic acid O–H group has an extremely broad absorption peak in the range 3200–2200 cm-1. Combined with a carbonyl peak in the range 1725–1680 cm-1, this makes identification of a carboxylic acid from its infrared spectrum an easy task. A typical carboxylic acid spectrum...

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