Brian C. Smith

As part of our ongoing review of infrared spectroscopic theory, we will discuss why different functional groups have different peak heights and widths. We find that Beer’s law determines peak height, and that the absorptivity in Beer’s law depends upon a mysterious quantity called the dipole moment, and is in large part responsible for why infrared peaks vary in size, from large to non-existent. Infrared peak widths can vary from a few wavenumbers to thousands of wavenumbers in width. We explore the role of intermolecular interactions in this phenomenon. We will wrap up by discussing the importance of integrating peak position, height, and width information to properly interpret spectra.

When I first introduced infrared spectral interpretation theory many columns ago (1), it was interspersed with discussions of the interpretation of specific functional groups. This approach made sense at the time, but, by not presenting all the theory of a piece, some important concepts that you need to know to interpret spectra successfully might have gotten short shrift. Hence the need for this Grand Review.

There are three important pieces of information in an infrared spectrum; peak positions, peak heights, and peak widths. All these pieces of information can be used to distinguish different functional groups from each other, and I have emphasized repeatedly in my columns the importance of integrating all these pieces of information to interpret spectra. It then makes sense for you to learn the reasons why different functional groups have different peak position, heights, and widths.

In my last column, I discussed why different functional groups have different peak positions (2). Here, we will discuss why different functional groups have different peak heights and widths. We will find that Beer’s law determines peak height, and that the absorptivity in Beer’s law depends upon a mysterious quantity called the 

, and is in large part responsible for why infrared peaks vary in size from large to nonexistent. Infrared peak widths can vary from a few wavenumbers to thousands of wavenumbers in width. We will explore the role of intermolecular interactions in this phenomenon.

When plotted in absorbance, the peak heights in an infrared spectrum are proportional to concentration, allowing infrared spectra to be used to determine the concentration of chemical species in samples (3). The relationship between absorbance and concentration is summarized in Beer’s law (3):

 equals absorbance, ɛ equals absorptivity, 

The pathlength is simply the thickness of sample seen by the infrared beam. The absorptivity is the proportionality constant between absorbance and concentration. The absorptivity also has physical meaning; it is a constant for a given pure molecule at a given wavenumber. For example, the absorptivity of pure acetone at 1716 cm

 (if not stated, assume all peak positions are in cm

 going forward) under standard conditions of temperature and pressure is a constant. This means that, under these conditions and for a fixed pathlength, a sample of pure acetone will always absorb the same amount of light at 1716.

Since different molecules have different absorptivities, this quantity must somehow be related to chemical structure. Consider the electronic structure of the CO

, the electronegativity difference between the carbon and the oxygen atoms causes the electrons to be shared unevenly (4), leading to a partial positive charge on the carbon, and partial negative charges on the oxygens. In Figure 1, a partial positive charge is represented by a δ+, and a partial negative charge by a δ-.

A dipole moment can be defined as two charges separated by a distance. Since chemical bonds, such as the ones in CO

, can consist of charges separated by a distance, chemical bonds have what are called 

. The bond dipole is a vector quantity; it has a magnitude and direction. The arrows in Figure 1 show the dipole moment vectors for CO

. Each arrow’s length represents the magnitude of the dipole moment for that bond, while the arrow head denotes its direction.

The magnitude of a bond dipole is simply the size of the charges multiplied by their distance apart, and can be calculated as such:

The bond dipole is a measure of charge asymmetry, and bonds with large charges held far apart having bigger dipole moments than bonds with small charges held close together. Thus, polar bonds, such as O-H and C=O, have large charge asymmetry, and therefore large dipole moments. Covalent bonds such as C-C bonds have small change asymmetry, and therefore small dipole moments.

The net dipole moment for a molecule is the sum of the bond dipoles for that molecule. Note in Figure 1 that, for CO

 at equilibrium, the two bond dipoles are equal in magnitude and opposite in direction. When added together, these two dipole moments cancel each other, giving a net dipole of zero.

So what does all this dipole moment stuff have to do with molecular vibrations and peak heights? Consider the symmetric and asymmetric stretching peaks of CO

During the symmetric stretch, the length of the bonds changes, but, since they stretch in phase with each other, the two bonds are always the same length, and always point in opposite directions. We say the symmetry of the molecule is maintained during the symmetric stretch (hence its name), and further meaning that the net dipole of the molecule is always zero at all points during this stretch. We can then say that, for the symmetric stretch of CO

, with respect to the bond length, dx, is zero. In other words:

An inspection of the infrared spectrum of CO

 shows the intensity of its symmetric stretching peak is zero, which means there is no peak corresponding to the symmetric stretch in CO

’s spectrum. Notice that there is a connection between dμ/dx being zero and the peak intensity being zero.

 undergoes an asymmetric stretch, as seen in Figure 2, the bond’s lengths are different at the same point in time. This means at points during this vibration, the two bond dipoles have different values, and that the net dipole is non-zero during this vibration. In this case, we say the symmetry of the molecule is broken, hence the name 

. Since the two bond dipoles values vary during the asymmetric stretch, this means the net dipole during this vibration has non-zero values. In other words:

. In fact, dμ/dx is quite large for this vibration, since we are stretching large dipole moments asymmetrically. The asymmetric stretching peak in the spectrum of CO

, and is quite intense, as seen in Figure 3. In fact, it is the intensity of this absorbance that makes CO

You may have seen 2350 peak before in sample spectra measured with a Fourier-transform infrared (FT-IR) spectrometer. This peak occurs when the CO

 concentration inside the spectrometer increases between the measurement of the background and sample spectra (please consult my book on FT-IR to learn how to avoid this problem [5]).

Note that dμ/dx for the asymmetric stretch of CO

 is large, and the vibration’s corresponding peak is large. What, then, is the connection between dμ/dx and intensity, and how does this new discovery of ours fit in with Beer’s law?

I will skip the derivation here, but quantum mechanics tells us that (6):

where ɛ equals absorptivity, dμ equals change in dipole moment during a vibration, and dx equals change in bond length during a vibration.

Therefore, the absorptivity depends upon the change in dipole moment with respect to bond length during a vibration squared. In other words, the absorptivity depends upon how the electronic structure of a molecule or functional group changes during a vibration. It is this equation that links peak height to molecular structure, and is why different functional groups have different peak heights.

A vibration such as the symmetric stretch of CO

 for which dμ/dx = 0 and the resultant peak intensity is zero is called an 

. On the other hand, vibrations such as the asymmetric stretch of CO

 for which dμ/dx ≠ 0 and the resultant peak’s intensity is non-zero are called infrared active peaks. One of the reasons that the number of normal modes a molecule has is greater than its number of infrared peaks is that not all vibrations are infrared active.

Polar functional groups such as C-O or C=O have large values of 

, hence have vibrations with large values of dμ/dx, and generally have intense peaks. On the other hand, for non-polar functional groups, such as C-C bonds that have small values of μ and their vibrations have small values of dμ/dx, their peaks are generally small. This means that infrared spectroscopy is better at detecting some functional groups than others, and that polar functional groups are generally the easiest to see. In other words, infrared spectroscopy is not a democracy, and, from the point of view of an infrared spectrometer, not all functional groups are created equal. This actually limits us as to what functional groups infrared spectroscopy can see. Fortunately, other molecular spectroscopy techniques, such as Raman scattering, are excellent at seeing the functional groups infrared (IR) cannot see, and thus can take up the slack.

In addition to peak positions and heights, peak widths can also be used to distinguish functional groups. Infrared peak widths vary greatly across the spectrum of known molecules. Figure 4 shows the infrared spectrum of liquid water.

Infrared peaks are generally narrow at their tops, and wider at their bottoms. How, then, do we measure peak widths in a standard fashion? One way spectroscopists do this is to measure what is called the full width half maximum (FWHM) of a peak (6). To measure a FWHM, take the height of a peak, go to where it is half this height, and then measure its width there. For example, if a peak is 1.0 absorbance unit (AU) high, we would measure this peak’s FWHM at 0.5 AU. The FWHM os the O-H stretching peak of water at 3300 cm

, as illustrated in Figure 4. This is an unusually broad peak for an infrared spectrum.

Figure 5 shows an example of a spectrum with narrow peaks; it is of the molecule benzonitrile.

In Figure 5, the FWHM of the C≡N stretching peak at 2228 is less than 10 cm

. The barely discernible red dot shows the FWHM. This peak is 50 times narrower than the FWHM of the O-H stretching peak of water. Why?

Consider Figure 6, which shows the interactions between neighboring water molecules, known as intermolecular interactions. The O-H bond in water is polar, because of the electronegativity differences between oxygen and hydrogen. This means, in an O-H bond, the hydrogen has a partial positive charge on it, and that the oxygen has a partial negative charge on it. Because of this, individual water molecules coordinate with each other with the partially positive hydrogen of one molecule interacting with the partially negative oxygen on a neighboring molecule, as seen in Figure 6. These two molecules form a weak, but very real, chemical bond called a hydrogen bond. These are represented by the dashed lines in Figure 6. Hydrogen bonds are an example of a strong intermolecular interaction, and note that water has broad peaks. In fact, any functional group that engages in hydrogen bonding, including alcohols, carboxylic acids, and amines, will have broadened peaks as a result of hydrogen bonding.

As noted, benzonitrile has narrow peaks. Why? The bonds in this molecule are non-polar, so there is no hydrogen bonding, and the interactions between neighboring molecules is weak, consisting of what are called van der Waals bonds (6). This shows us that molecules with weak intermolecular interactions have narrow peaks. Ultimately, then, the peak widths in infrared spectra are determined by the strength of intermolecular interactions; this is why peak width differences can be used to distinguish functional groups from each other. An example of this is seen in Figure 7.

The spectrum on the left in Figure 7 shows the N-H stretching peaks of a primary amine and on the right is the O-H stretching peak of an alcohol. Note these peaks all appear around 3300. The FWHM of the N-H stretching peaks is about 70 cm

, whereas the FWHM of the O-H stretch is about 500 cm

. Even though these two functional groups have peaks in similar places, the difference in their peak widths can be used to distinguish them from the other.

The difference in peak width seen here exists because the polarity of an N-H bond is lower than that of an O-H bond; as a result, N-H hydrogen bonds are weaker than those of O-H bonds, meaning that amines have weaker intermolecular interactions than alcohols.

The pieces of information available to us in an infrared spectrum include peak positions, heights, and widths. All these pieces of information can be used to distinguish functional groups. We covered peak positions in the last column. Here, we saw that Beer’s law determines peak size, and that the absorptivity depends upon the change in dipole moment with respect to bond length during a vibration. We also saw that peak widths are determined by the strengths of intermolecular interactions. We now know why integrating peak position, height, and width information to interpret spectra is important.

Quantitative Spectroscopy: Theory and Practice

 (Academic Press, Boston, Massachusetts, 2002).

 (Cornell University Press, Ithaca, New York, 1938).

Fundamentals of Fourier Transform Infrared Spectroscopy

 (CRC Press, Boca Raton, Florida, 3rd Ed., 2011).

 (John Wiley and Sons, New York, New York, 3rd Ed., 1996).

, PhD, is the founder and CEO of Big Sur Scientific, a maker of portable mid-infrared cannabis analyzers. He has over 30 years experience as an industrial infrared spectroscopist, has published numerous peer-reviewed papers, and has written three books on spectroscopy. As a trainer, he has helped thousands of people around the world improve their infrared analyses. In addition to writing for 

, Dr. Smith writes a regular column for its sister publication 

 and sits on its editorial board. He earned his PhD in physical chemistry from Dartmouth College. He can be reached at: 

Articles

Integrating peak position, height, and width information is important to properly interpret infrared spectra. But do you understand why peak heights and widths vary? We explain.

The Grand Review II: Why Do Different Functional Groups Have Different Peak Heights and Widths?

Over the last five years and 30+ columns, we have covered a lot of territory. It is now appropriate to step back and take a look at where we have been. “IR Spectral Interpretation Workshop” columns can be divided into five broad categories: theory, functional groups containing the C-H bond, those containing the C-O bond, those with the C=O bond, and those with organic nitrogen compounds. In the first installment of “The Grand Review,” we discuss why different functional groups have peaks with different positions.

What a long, strange, and hopefully interesting and useful trip it has been for you, my column readers. Since January 2015, 

 has published over 30 installments of my “Infrared Spectral Interpretation Workshop”columns. Over these years, we have discussed the theory behind IR spectral interpretation, a process to follow which maximizes the probability of obtaining the right answer, interpretation aids such as spectral subtraction and library searching, and how to interpret the spectra of functional groups containing C-H, C-O, C=O, and N-H bonds.

There is still plenty of new material to be covered, but given the time that has elapsed and the amount of ground we have covered, it is time to step back and take a look at where we’ve been. The next several columns will contain a “Grand Review” of the major concepts you need to understand to interpret infrared spectra properly.

When I first started writing this column, one of the challenges was trying to get the theory discussion out of the way as quickly as possible so we could get to interpreting spectra. This meant that interpretation theory and practice were mixed in the same article. While this made sense at the time, it meant discussion of crucial theoretical concepts was spread over multiple installments, which, from a pedagogic point of view, is not optimal. This “Grand Review” gives me the chance to rectify this situation by introducing infrared spectral interpretation theory as a whole. In this first installment, we will review the theory behind peak positions.

Why Different Functional Groups Have Different Peak Positions

There is a lot of information to be gleaned from infrared spectra, but perhaps the most important is the molecular structural information given to us by peak positions (1). A large part of what I have taught you are the diagnostic group wavenumber peak positions for common functional groups, where we assign peaks in a spectrum to a specific vibration of a given functional group. How do we know what functional groups absorb where? Part of this is from measuring infrared spectra of molecules with known chemical structures and looking for correlations (1–5); the other is from theory (6–9).

Because the absorption of infrared light excites molecular vibrations, we need to understand these vibrations to make sense of infrared spectra. The constituent vibrations of a molecule (or any physical object, for that matter) are referred to as its 

 (1). An illustration of the symmetric and asymmetric stretching normal modes of carbon dioxide is shown in Figure 1.

 molecule is symmetric because both C=O bonds are the same length. During the symmetric stretching vibration, the two oxygen atoms move in opposite directions, as illustrated in Figure 1, and the two C=O bonds have the same length at all points during the vibration, preserving the symmetry of the molecule. On the other hand, during the asymmetric stretch, the two oxygen atoms move in different directions; at many points during the vibration the bond lengths differ from each other, and the symmetry of the molecule is broken.

To understand the relationship between molecular vibrations and peak positions, consider the generic diatomic molecule shown in Figure 2, which is a ball and spring model of a molecule, where individual atoms are represented by balls and chemical bonds by springs. The atoms in Figure 2 have masses 

, respectively. A real-world example of such a diatomic molecule would be hydrogen chloride (HCl), where atom 1 would be chlorine, and atom 2 hydrogen.

Diatomic molecules have only one vibration (1), the stretching of their single chemical bond, and therefore have only one infrared peak. Our goal is to derive an equation that will tell us the wavenumber of the vibrational peak in the spectrum of a diatomic molecule, like the one shown in Figure 2. To do this, we will assume when this molecule vibrates, it does so in what is called a harmonic fashion. This means that when the chemical bond stretches, the two atoms move in phase with each other, with the same amplitude, and pass through their equilibrium positions at the same time (1). Real-world harmonic systems include a swinging pendulum and a rotating wheel; many systems whose motion repeats can be characterized mathematically as a harmonic oscillator. Most molecular vibrations are much more complicated than the harmonic motion assumed here, but assuming harmonic motion makes the math easier.

To begin, we apply physics to the situation. We define the reduced mass of our diatomic molecule as such:


By using a single quantity to represent the mass of a two-mass system, the mathematics of the derivation is made easier.

The second piece of physics we have to consider is Hooke’s Law (1), which describes the behavior of springs, and vibrating chemical bonds, as such (equation 2):

, is a measure of a spring’s stiffness, and when applied to chemical bonds, it is a measure of bond strength.

Anyone who has ever stretched a rubber band has an intuitive grasp of Hooke’s Law. 

 is the amount of force needed to stretch and maintain the rubber band at a given length, and is equal to the amount of force the rubber band is putting on the person trying to stretch it (the restoring force). Experience shows it takes more force to stretch a rubber band a long distance than a short distance, which is why the variable 

 appears in Hooke’s Law. Experience also shows it takes more force to stretch a stiff rubber band than a weak one, which is why the force constant 

If one takes the reduced mass, Hooke’s Law, applies one of Newton’s Laws, and solves the resulting differential equation, this result is obtained (1):

Equation 3 is the equation we have been looking for—it predicts the position in wavenumbers of the single vibrational peak for a generic diatomic molecule. Note there are several constants in the equation, including the value of π and the speed of light, and that the two variables are the force constant, MR, and reduced mass, 

Equation 3 is perhaps the most important equation in infrared spectroscopy. It relates components of a molecule’s structure, bond strength, and atomic mass, to what is called a 

, the peak position in wavenumber. If two molecules have different molecular structures, they will have different spectra (1–5), and this is a direct result of equation 3; molecules with different structures have different bond strengths or atomic masses, and therefore different peak positions. As we have seen throughout the last five years, peak positions in an infrared spectrum can be used to determine the structures of unknown molecules. Equation 3 gives us the correlation between chemical structure and peak position; equation 3 is why infrared spectra are useful.

 is in the numerator in equation 3, as the force constant goes up the peak position increases. This means strong chemical bonds have higher peak positions than weak chemical bonds. The force constant is determined by a chemical bond’s electronic structure, which is how the electrons and nuclei are arranged with respect to each other in a bond (1). When comparing peak position shifts between molecules caused by changes in 

, the phenomenon is called an electronic effect; a change in electronic structure caused the bond strength to change, causing 

 to change, and resulting in a change in peak position.

An interesting example of an electronic effect is the position for carbon-carbon single bond, (C-C), double bond (C=C), and triple bond (C

C) stretching peaks. C-C stretching peaks typically fall from 1400 to 1000 cm

 (going forward, assume all peak positions are in cm

units even if not denoted), as shown in the spectrum of hexane in Figure 3.

In Figure 3 the C-C stretching peaks are a series of small peaks just to the right of the CH

Carbon-carbon double bond, or C=C stretching peaks, fall at a higher wavenumber than C-C stretches, from 1680 to 1630, because a double bond is stronger, has a higher force constant, and therefore a higher peak position than a C-C bond. An illustration of a C=C stretching peak is shown in the spectrum of 1-Octene in Figure 4.

The peak labeled B at 1641 is an example of a C=C stretching peak.

C bonds are even stronger than C=C bonds, and therefore have a stretching peak position typically near 2100. An example of a C

C stretching peak is shown in the spectrum of 3,3-dimethyl-1-butyne in Figure 5.

In summary then, we have seen a classic example of an electronic effect, where a change in bond strength causes C-C, C=C, and C

C stretching peaks to appear from 1400 to 1000, 1680 to 1630, and ~2100, respectively.

The denominator in equation 3 contains the reduced mass, therefore molecules with heavy atoms in them have lower wavenumber peaks than molecules with light atoms in them. When peak positions change because of a change in atomic mass, it is called a 

. A classic example of a mass effect is the wavenumber region where peaks of organic and inorganic molecules fall. As we have seen, organic functional groups contain light elements such as C, H, N, and O. In general, most of the peaks from these functional groups fall above 400. Figures 3, 4, and 5 illustrate this well.

Inorganic molecules are composed of heavier metal atoms such as Na, K, Ca, and Pb. Many of the peaks from inorganic functional groups fall below 400. This is because metal atoms tend to weigh more than non-metal atoms. In fact, the reason that most FT-IRs can’t scan below 400 is that is where the inorganic material KBr, often used in beamsplitters, begins to absorb (1). An example of an inorganic with strong far infrared peaks is seen in Figure 6, the far infrared spectrum of calcium carbonate (CaCO

Note that the spectrum in Figure 6 is plotted from 600

, well into the far infrared that is typically defined as being from 400

In summary, we have seen how changes in electronic structure and atomic mass impact peak positions. Much of what we know about why different functional groups have different peak positions can be understood utilizing these concepts (1).

Equation 3 is straightforward to understand, and lays a good foundation for understanding the connection between infrared peak positions and chemical structure. However, it is based on assumptions that are not always correct. In reality, most molecular vibrations are not harmonic, but anharmonic, which means their vibrational motion is much more complicated that what we have assumed here. In this case, we use the anharmonic oscillator model to describe these vibrations. The net effect is that equation 3 is modified with extra terms to take anharmonic motion, or anharmonicity, into account (7–9). This being said, equation 3 is still approximately correct, and serves to explain much of what we observe in terms of why different functional groups absorb at different wavenumbers.

The correlation between molecular structure and peak positions is a large part of what makes infrared spectroscopy useful. Using the Harmonic Oscillator model, Equation 3 was derived , relating peak positions to force constant and reduced mass. When changes in force constant cause a change in peak position, it is called an electronic effect. When changes in atomic mass cause a change in peak position, it is called a mass effect. Much of what we know about why different functional groups have different peak positions can be understood using the concepts of electronic effect and mass effect. The limitations of the harmonic oscillator model were also discussed.

Infrared Spectral Interpretation: A Systematic Approach

The Infrared Spectra of Complex Molecules 

(Chapman and Hall, London, United Kingdom, 1975).

D. Lin-Vien, N. Colthup, W. Fately, and J. Graselli, 

The Handbook of Infrared and Raman Characteristic Frequencies of Organic Molecules

 (Academic Press, Boston, Massachusetts, 1991).

Introduction to Infrared and Raman Spectroscopy

 (Academic Press, Boston, Massachusetts,1990).

Infrared and Raman Characteristic Group Frequencies

 (Wiley and Sons, New York, New York, 3rd Ed., 2001).

E. B. Wilson, J.C. Decius, and P. C. Cross, 

Molecular Vibrations: The Theory of Infrared and Raman Vibration Spectra

 (Dover Publications, New York, New York, 1955).

(Wiley and Sons, New York, New York, 1996).

is the founder and CEO of Big Sur Scientific, a maker of portable mid-infrared cannabis analyzers. He has over 30 years experience as an industrial infrared spectroscopist, has published numerous peer-reviewed papers, and has written three books on spectroscopy. As a trainer, he has helped thousands of people around the world improve their infrared analyses. In addition to writing for 

Dr. Smith writes a regular column for its sister publication 

and sits on its editorial board. He earned his PhD in physical chemistry from Dartmouth College. He can be reached at: 

Articles in this column have addressed five main areas: theory, functional groups containing the C-H bond, those containing the C-O bond, those with the C=O bond, and those with organic nitrogen compounds. Here, we review the concepts.

The Grand Review I: Why Do Different Functional Groups Have Different Peak Positions?

Our final look at the infrared spectra of organic nitrogen functional groups will be of the nitro, or NO

, group. This group has two strong and characteristic infrared features, making it easy to identify. However, attaching a NO

 group to a molecule scrambles its electronic structure, which is why these compounds are explosive, and why nitro groups make some of the infrared interpretation rules we learned irrelevant.

Over the several years that I have been writing these articles, there have been large sections of material devoted to the spectroscopy of functional groups that contain specific chemical bonds. When I began discussing organics containing nitrogen, I stated that the stretching and bending of N-H bonds would be important in determining whether or not there was nitrogen present in a molecule (1). We subsequently used N-H stretching and bending peaks to analyze the spectra of amines (2,3), amides (4,5), imides (6), and urethanes (7).

The nitro group is going to be the odd man out in this discussion. This functional group consists of a nitrogen atom with no hydrogens, but with two oxygens and a carbon attached, as seen in Figure 1.

, and that the carbon atom singly bonded to the nitro nitrogen is called the 

. Depending on whether the alpha carbon is saturated, or part of an aromatic ring, nitro molecules can be divided into saturated and aromatic nitro compounds.

 group is unusual. Normally, oxygen atoms form two chemical bonds (8). However, there is also a C-N bond in the nitro group, as seen in Figure 1. Given that nitrogen normally forms three bonds (8), how do we apportion the electrons in a nitro group to keep these compounds from falling apart?

There are three bonding electrons shared between the two oxygens in the nitro group, giving what are essentially two “bonds and a half,” as seen in Figure 1. The dashed lines in Figure 1 represent the half bonds. These NO bonds are similar to the carboxylate C-O bonds and a half discussed in a previous column (9). Neither oxygen atom has its full complement of two full chemical bonds, making the nitro groups unstable. Saturated nitro compounds such as nitroalkanes, otherwise known as rocket fuel (8), are rarely analyzed by infrared spectroscopy because they are liable to detonate during analysis. We will not study them further.

However, nitro groups attached to benzene rings can be relatively stable, assuming not too many nitro groups are attached. The chemical structure of trinitrotoluene, a commonly used explosive known as 

The presence of the three nitro groups destabilizes the benzene ring, resulting in TNT’s explosive properties. Di-nitrotoluene and mono-nitrotoluene are stable. We will restrict our discussion to non-explosive aromatic nitro compounds.

The Infrared Spectroscopy of the Nitro Functional Group

The infrared spectrum of an aromatic nitro compound, meta-nitrotoluene, is seen in Figure 3. Recall that highly polar bonds have intense infrared features due to the large change in dipole moment with respect to bond length, 

, during a vibration (10). Oxygen is more electronegative than nitrogen; therefore, the N-O bonds in the nitro group are relatively polar, and as a result, their asymmetric and symmetric stretching peaks are unusually large. The details of these vibrations are shown in Figure 4.

 stretch typically falls from 1550 to 1500 cm

, and is seen in Figure 3 labeled A at 1527 cm

 units, even if not labeled as such). The symmetric stretch is seen in Figure 3 labeled B at 1350 cm

, and, in general, this peak appears from 1390 cm

. Note how peaks A and B in Figure 3 are the two most intense peaks in the spectrum, and they stick down like eye teeth in the middle of the spectrum. The combination of a pair of intense peaks in these wavenumber ranges is unique, making the presence of a nitro group in a sample easy to spot.

The nitro group also exhibits a scissors bending vibration, similar to that of the methylene group (11). This gives rise to a medium intensity peak from 890 cm

. It can be seen in Figure 3 labeled C at 881 cm

The good news about the nitro group is that it has two strong infrared bands that are easy to spot. The bad news is that whenever the nitro group is attached to a benzene ring, it makes it difficult to determine the substitution pattern on the benzene ring. Recall that the benzene ring out-of-plane C-H bend band, in combination with the presence or absence of the aromatic ring-bending band at 690 cm

, can be used to determine the substitution pattern on a benzene ring (12). The presence of a nitro group makes it difficult to apply these rules. This is caused by the unique electronic structure of the nitro group, and how it interacts electronically with the benzene ring. Suffice it to say, one may need to use an analytical technique other than infrared spectroscopy to determine the substitution pattern on nitro-substituted aromatic rings.

Note that the structure of meta-nitrotoluene does contain a methyl group. We have learned that the diagnostic pattern for a methyl group includes saturated carbon asymmetric and symmetric stretches near 2962 cm

 (13,14). Note in Figure 3 that the saturated C-H stretches fall at 2926 cm

. Under normal circumstances, we would interpret these two peaks as the asymmetric and symmetric stretches of methylene groups, because they have peaks at 2926 cm

 (14). In this case, that interpretation is wrong, because the nitro groups scramble the electronic structure of the molecule, throwing off these peak positions. If you notice the nitro peaks first, it will give you a heads up that the saturated C-H stretches may be troublesome. Normally we can depend on the methyl group umbrella mode to indicate that the saturated C-H stretches could be an issue. However, in this case, the intense nitro symmetric stretching peak at 1350 cm

 is sitting on top of it. Unfortunately, then there is nothing in this spectrum clearly showing the presence of the methyl group. As stated above, the nitro group throws off some of the interpretation rules we have learned.

The diagnostic pattern for the presence of a nitro group in a sample then is a pair of intense peaks at about 1550 cm

, along with the scissor peak around 850 cm

. Table I summarizes the group wavenumbers for the nitro group.

The nitro group consists of a nitrogen atom with two oxygens and one carbon attached. The two nitrogen-oxygen bonds are “bonds and half” that destabilize nitro compounds, making their analysis an explosive proposition. The large dipole moments for the NO bonds give two strong peaks around 1550 cm

 as a result of the asymmetric and symmetric stretching of the NO

functional group. This is an unusual pattern, and is easy to spot. There is also a scissoring peak around 850 cm

. Nitro groups tend to scramble the electronic structure of molecules, making the interpretation of their spectra problematic.

Your Next Infrared Spectral Interpretation Challenge

Using everything you have learned in this and previous columns, determine the functional groups present in the Figure i spectrum and try to determine the chemical structure of this compound. Remember that inclusion of a peak position in the table does not necessarily mean that it will be useful in the structure determination.

Because this is a particularly tricky problem, answer these questions about the spectrum in this order to guide you. In all cases, be sure to justify your answer.

2. If so, is there a benzene ring present?

3. Are there any other functional groups present?

 Dr. Smith writes a regular column for its sister publication

We review one final organic nitrogen functional group, representing explosive compounds. This is the nitro group, which requires a different set of interpretation rules.

Organic Nitrogen Compounds X: Nitro Groups, An Explosive Proposition

The urethane functional group is found almost exclusively in polymer form, known as 

. These ubiquitous polymers make urethanes an economically important functional group. Urethanes are made using diisocyanates. How to identify diisocyanates, and distinguishing among different urethane types using infrared spectroscopy, are explained in this installment of the column. 

		
		The urethane functional group is found almost exclusively in polymer form, known as 

. These ubiquitous polymers are used as wood finishes and make up polyurethane foam, also known as foam rubber. The structural framework of a urethane is shown in Figure 1.

	Note that this functional group contains a carbonyl bond with an oxygen atom and nitrogen atom attached to it. The latter is why urethanes are included in this series on organic nitrogen compounds. On the left-hand side of Figure 1, the urethane looks like an ester because of the O-C=O linkage (1), but on the right-hand side of Figure 1, the N-C=O linkage looks like an amide (2,3). As we will see, the infrared spectra of urethanes combine features of both esters and amides.

 if one carbon and two hydrogens are attached to the nitrogen atom, as 

 if two carbon atoms and one hydrogen atom are attached to the nitrogen, and as 

 if three carbon atoms are attached to the nitrogen. Figure 1 shows the structure of a secondary urethane. Note that it contains two C-N bonds and one N-H bond.

	Polyurethanes are made by reacting alcohols with a functional group called a 

. The structural framework of a diisocyanate is seen in Figure 2.

	Note that a diisocyanate contains two isocyanate -N=C=O linkages, hence the “di” in the name. A molecule with one -N=C=O unit would simply be called an 

. When urethanes and alcohols are polymerized to make polyurethanes, there can be isocyanate end groups that occasionally appear in the spectra
	of polyurethanes.

	The infrared spectrum of a diisocyanate is shown in Figure 3.

	Note that the spectrum is dominated by the peak labeled A at 2277 cm

 (assume all peak positions are in wavenumbers going forward, even if not labeled as such). This peak is due to the asymmetric stretch of the -N=C=O group, which normally falls from 2280 to 2240. This peak is the perfect example of a group wavenumber, because it is intense and shows up in a part of the spectrum where very few other functional groups absorb (5). Throughout the over 30 installments of this column, you may have noticed that most of the time we focus our attention from 3500 to 2800, and from 2000 to 400. That is because very few functional groups absorb between 2800 and 2000. The only other infrared peak we have seen that absorbs near diisocyanates is the C≡C stretch of disubstituted alkynes, which fall from 2260 to 2190 (6). These two functional groups are easily distinguished, because the diisocyanate -N=C=O asymmetric stretching peak is broader and stronger than a typical C≡C stretch. This is because of the high polarity of the -N=C=O group compared to that of an alkyne triple bond (7). The strength and unusual position of the -N=C=O stretching peak are why it sometimes appears in polyurethane spectra, even though it is in low concentration.

	Based on the structural framework of urethanes shown in Figure 1, we can predict to a point what urethane spectra will look like. Primary and secondary urethanes contain N-H bonds, and so have N-H stretches that are medium in width and intensity, and fall in the same wavenumber range as the N-H stretching peaks of amides and amines (2–4). The C=O bond in urethanes means there should be a strong C=O stretching peak (8) from around 1700, and the C-O bond means there should be a strong peak between 1300 and 1000 (9).

	The infrared spectra of primary urethanes behave as expected. Their asymmetric and symmetric NH

 stretches are found from 3450 to 3400 and 3240 to 3200, respectively. Primary urethanes have two NH stretching peaks, as do primary amines and amides (2–4), because, in all these cases, the primary nitrogen contains two N-H bonds, hence two N-H stretching peaks (3). The carbonyl stretch of primary urethanes falls at 1740 to 1730, and there is an NH

 scissors (in-plane bending) peak from 1630 to 1610.

	Unfortunately, the C=O stretching peak of primary urethanes falls in the same region as other carbonyl-containing functional groups, such as ketones, aldehydes, esters, and carboxylic acids (8,9). However, these functional groups do not contain nitrogen, hence there are no N-H stretching peaks, making distinguishing primary urethanes from these compounds easy. But primary amides have two NH

stretches, and a C=O stretch (3), as do primary urethanes. Fortunately, the primary urethane C=O stretch is at 1740 to 1730, whereas, for primary amides, it falls from 1680 to 1630 (10). This difference in C=O stretching peak position means primary amides and urethanes can be distinguished from each other. The diagnostic group wavenumber pattern, then, for primary urethanes is two NH stretching peaks, a C=O peak around 1730, and an NH

	The infrared spectrum of a secondary urethane, a “polyurethane/polyether foam,” better known as foam rubber, is seen in Figure 4.

	Because this polymer is a secondary urethane, as seen in the figure, it has a single N-H bond, and, hence, a secondary nitrogen. The spectral signature of a secondary nitrogen, as we saw for secondary amides and amines (2–4), is a single N-H stretching peak. In Figure 4, this peak is labeled A at 3326. In general, for secondary urethanes, this peak falls from 3340 to 3250. Note that the NH stretching peak in Figure 4 is of medium intensity and height, consistent with many of the other NH stretches we have seen (11).

	The carbonyl stretch in Figure 4 labeled B is split, with peaks at 1720 and 1705. In the past, a double carbonyl stretch like this indicated the presence of two separate carbonyl bonds in a functional group, such as for acid anhydrides (12) and imides (13). However, this is clearly not the case here; urethanes do not contain two carbonyl bonds. The cause of this split carbonyl peak is that some urethane C=O groups engage in hydrogen bonding, whereas others don’t. The 1720 peak is from non-hydrogen bonded, or “unassociated,” carbonyls, whereas the 1705 peak is from the hydrogen bonded, or “associated,” carbonyl groups (10). We have already studied the effects of hydrogen bonding on peaks, such as for alcohols (9). The tendency for peaks upon hydrogen bonding is to broaden and move to lower wavenumbers, as we have seen here.

Here we delve into the interpretation of another organic nitrogen compound, known as polyurethane, a ubiquitous polymer used for wood finishes and foam rubber.

Brian C. Smith

Articles

The Grand Review II: Why Do Different Functional Groups Have Different Peak Heights and Widths?

The Grand Review I: Why Do Different Functional Groups Have Different Peak Positions?

Organic Nitrogen Compounds X: Nitro Groups, An Explosive Proposition

Organic Nitrogen Compounds IX: Urethanes and Diisocyanates