Introduction

Spectroscopy is the study of the interactions between radiant energy and matter. We see colours as a consequence of absorption by organic and inorganic compounds. For organic chemists, a primary fact is that the structure of certain organic compounds determine the wavelengths at which the compounds will absorb energy, therefore spectroscopic techniques can be used to determine chemical structure of unknown compounds.


Electromagnetic Radiation

Electromagnetic Radiation is energy transmitted through space in the form of waves. Each type of radiation is characterised by its wavelength (), the distance from the crest of one wave to the next.



The entire spectrum of electromagnetic radiation is shown below:


The visible region of the spectrum is between 400 and 750 nm (1 nm = 10-9m), but this is only a small part of the spectrum. Wavelengths slightly shorter than those if the visible region are in the ultraviolet region, while slightly longer wavelengths are in the infrared region of the spectrum.
Radiation can also be characterised by its frequency, (), which is defined as the number of complete cycles per second (cps), also called Hertz (Hz), and can be depicted as follows:


Radiation of a higher frequency has more waves per second, so the wavelength is shorter. By their definintions, wavelength and frequency are inversely proportional, and can be expressed mathematically:


= frequency in Hz
c = 3 x 1010 cm/sec (speed of light)
= wavelength in cm

In Infrared spectroscopy, the frequency is expressed as a wave number (number of cycles per second or cm-1). Wavenumbers and wavelengths can be interconverted using the following equation:


Some common spectroscopy units and symbols are listed in the table below:

Symbol or Unit Definition
Frequency:
Definition in Hz (Cycles per second)
cm-1 Wavenumber: frequency reciprocal
Wavelength
Wavelength (usually expressed in units of meters, nm, cm etc.)
m micrometer, 10-6 m
nm nm, 10-9 m
Å angstrom, 10-10 m


Radiation is transmitted in energy packets called photons or quanta. The energy of a photon is inversely proportional to the wavelength. Shorter wavelength radiation has a higher energy, so a photon of ultraviolet light has more energy than a photon of visible light.
Molecules only absorb specific wavelength of electromagnetic radiation. Absorption of an ultraviolet photon (higher energy) can promote an electron to a higher energy orbital. Infrared energy does not contain suficient energy to promote an electron, but instead results in greater vibration of bonded atoms. The intensity of radiation is proportional to the number of photons, regardless of the energy. When a sample absorbs photons, whether infrared or ultraviolet, the number of transmitted photons must decrease and it is this change in intensity that is measured in absorption spectroscopy.


Spectral Features

Any spectrum of a compound is a graph of either wavelength or frequency, continuously changing over a small part of the electromagnetic spectrum, versus either percent transmission (%T) or absorbance (A).


Most infrared spectra record wavelength or frequency versus %T. If a sample does not absorb at a particular frequency, then it is displayed on the spectrum as 100 %T. When a compound absorbs radiation, the intensity of transmitted radiation decreases, and this results in a decrease of %T and appears as a dip in the spectrum, called an absorption peak or band. The portion of the spectrum at 100 %T is called the base line, recorded at the top of the spectrum.

Most visible and ultraviolet spectra are usually presented as graphs of A versus wavelength. In these cases the baseline is at the bottom of the spectrum, and absorption is recorded as an increase in signal. The general appearance of spectra using %T and A is shown below.




Absorption of Infrared Radiation

Atoms bonded by covalent bonds can oscillate or vibrate, in a manner similar to two balls attached by a spring. When molecules absorb Infrared (IR) radiation, the absorbed energy causes an increase in the amplitude of the vibrations of the bonded atoms. The molecule is then in an excited vibrational state. The absored energy is then dissipated as heat when the molecule returns to the ground state.

Different types of bonds (C-H, C-C, O-H etc...) absorb infrared energy at different characteristic wavelengths.

A bond within a molecule will undergo different types of oscillations, and so a molecule may absorb energy at more than one wavelength. For example, the O-H bond absorbs energy at approx. 3300 cm-1. Energy at this wavlength causes increased stretching vibrations of the O-H bond. There is also an absorption at approx. 1250 cm-1, witch leads to increased bending vibrations. These different types of vibration are called fundamental modes of vibration.




The amount of energy that a bond absorbs depends on the change in the bond moment (or dipole moment) as the bonded atoms vibrate, where a greater change in the bond moment results in the absorption of more energy. Nonpolar bonds do not absorb IR radiation, because there is no change in the bond moment when the atoms oscillate. Relativey nonpolar bonds such as C-C or C-H have weak absorption bands, whereas strogly polar bonds such as C=O have very strong absorption patterns.


The Infrared spectrum

An Infrared Spectrophotometer is the instrument used to measure the absorption of infrared radiation of a compound. A schematic diagram is shown below:




There is the light source, emmiting infrared light, which is split by mirrors into two beams, a reference and a sample beam. After passing through the reference cell containg solvent (or usually nothing), and the sample cell, the two beams are combined in a chopper (another mirror system). The alternating beam is diffracted by a grating that seperates the beam into its different wavelengths. The detector measures the difference in intensities of the two segments of the beam at each wavelength, and passes this information to the recorder to produce a spectrum.
A typical spectrum is displayed below:




The scales at the bottom are in wavnumbers - ranging from 4000 to approx. 400 cm-1. The frequency or wavelength of the minimum point (maximum absorption) of an absorption band is used to identify each band. The bands can be classified by intensity: strong, (s), medium,(m) and weak, (w). A weaker band overlapping another is called shoulder.


Interpretation of Infrared Spectra

Chemists have determined ranges of wavelengths of absorption for each functional group. Correlation Tables provide summaries of this information. A typical correlation table is shown below:




We can see that O-H and N-H stretching bands are found between 3700 and 3000 cm-1. If the infrared spectrum of an unknown compound shows a band in this region, then it may be reasonable to assume the the structure may contain one of these groups.
The region between 4000 and 1400 cm-1 is useful for the identification of the various structural groups, and shows mainly the stretching mode absorptions. At wavnumbers lower than 1400 cm-1, spectra can be very complicated due to the combination of bands arising from both stretching and bending modes. Determination of structure from this region alone is uncertain, but most organic compounds has its own unique absorption pattern in this region of the spectrum, and is called the fingerprint region.


In the following sections we will discuss several of the characteristic infrared absorptions.



Group assignments in Infrared Spectra


Carbon-Carbon and Carbon-Hydrogen

C-C single bonds, (bonds between sp3-hybridised carbon atoms) produce weak absorption bands in the Ir spectrum, and therefore are not that useful for structure determination. Bonds between sp2-hybridised carbon atoms (C=C) produce a variable strength absorption at approx. 1700-1600 cm-1, whereas aryl carbon-carbon bonds have bands at slightly lower frequencies. Carbon-carbon triple bonds between sp-hybridised carbon atoms show a weak but characteristic absorbance band at approx. 2250-2100 cm-1, in a region of the spectrum with little absorbance by other groups takes place.

Bond Region of Spectrum (cm-1)
sp2-hybridised - double bond (alkyl) 1700-1600
sp2-hybridised - double bond (aryl) 1600-1450
sp-hybridised - triple bond 2250-2100


Almost all organic compounds contain C-H bonds. The C-H stretch is seen at 3300-2800 cm-1. The stretching peaks are usefull in determining the hybridisation of the carbon atom.

Bond Region of Spectrum (cm-1)
sp3 -C-H (alkanes, alkyl) 3000-2800
sp2 -C-H (alkenes) 3300-3000
sp -C-H (alkynes) ~3300


The geminal coupling (two methyl groups on the same carbon), often exhibit a double bending peak at approx. 1385-1360 cm-1, but are not visible on all spectra and a single peak is sometimes observed.


Carbon-Halogen

The stretching absorption of the C-X haloalkane bond lies in the fingerprint region of the infrared spectrum and can range from 1430-500 cm-1. Without any further information it is not possible to identify a band in this region as the definite presence of a halogen in an organic compound.


Alcohols and Amines

Alcohols and amines exhibit O-H and N-H stretching absorptions between 3700-3000 cm-1, to the left of the C-H absorption. If two peaks are observed in the spectrum of a compound, the absorption is due to two hydrogens on the amine nitrogen (-NH2). One observed peak means that there is only one hydrogen on the nitrogen, corresponding to a secondary amine (>N-H). If no peak is observed it may be possible that a tertiary amine is present (R3N).

Alcohols and amines also exhibit C-O and C-N absorptions in the fingerprint region of the spectrum.

Bond Region of Spectrum (cm-1)
O-H or N-H 3700-3000
C-O or C-N 1300-900


Hydrogen bonding changes the position and appearance of of an absorption band. Hydrogen bonding in pure liquids can be extensive, giving a wide absorption band at approx. 3300 cm-1. When hydogen bonding is less extensive a weaker, but sharper peak, is observed. Absorption by N-H bonds is less intense than by O-H bonds, due to weaker and less polar hydrogen bonds in amines.


Ethers

Ethers have a characteristic C-O stretching band within the fingerprint region at approx. 1260-1050 cm-1. As the oxygen atom is electronegative, the stretching causes a large change in the bond dipole moment. Therefore the C-O absorption is usually very strong.


Carbonyl Compounds

Probably the most distinctive band in the infrared spectrum is the one due to the carbonyl stretching mode, a strong peak observed between 1820-1640 cm-1. The carbonyl group can be part of several functional groups, so the exact position in the spectrum, together with data from other absorptions is neccessary to deduce a possible structure. Different functional groups, along with their relevant carbonyl frequencies are discussed below:

Ketones produce the simplest spectra of the carbonyl compounds. If the ketone is aliphatic, then the stretching bands for C=O or C-H are strong. Other funtional groups may complicate the spectrum.

Aldehydes produce similar spectra to those of ketones. The important difference is that the aldehyde has a hydrogen atom bonded to the carbonyl carbon. This C-H bond has two characteristic stretching bands (to the right of the aliphatic C-H band) at approx. 2900-2820 cm-1 and 2780-2700 cm-1. Both are weak but sharp bands, though the band at ~2900 cm-1 may be swamped by the C-H bands.

Carboxylic acids show the typical C=O absorption as well as having a characteristic O-H band at approx. 3300 cm-1 and slopes inro the aliphatic C-H absorption band. The main reason for the difference in the bands for the carboxyl O-H rather than the alcohol O-H is due to carboxylic acids forming dimers.



Esters show both the typical carbonyl band and a C-O band. The C-O band, like in ethers, is observed at approx. 1300-1100 cm-1, but in esters the band is very strong and can be used to distinguish between esters and ketones.