Infrared Spectroscopy

- mirrored from UCLA
http://www.chem.ucla.edu/~pang/chem11cl_net/11cl_home.html -

Introduction

Adapted from : R. L. Pecsok L. D. Shields, Modern Methods of Chemical Analysis (Wiley, New York, 1968); and A.T. Schwartz et al., Chemistry in Context (American Chemical Society, Washington, DC 1994).

Spectroscopy is the study of the interaction of electromagnetic radiation with matter. There are many forms of spectroscopy, each contributing useful information to identify substances and to determine various characteristics of their structure.

A portion of the electromagnetic spectrum is shown in Figure 1, along with the names associated with various regions of the electromagnetic spectrum. Our eyes can detect only a very limited range of wavelengths, the visible spectrum between about 300 and 800 nm.

Figure 1 The Electromagnetic Spectrum

Atoms and molecules can absorb electromagnetic radiation, but only at certain energies (wavelengths). The diagram in Figure 2 illustrates the relationships between different energy levels within a molecule. The three groups of lines correspond to different electronic configurations. The lowest energy, most stable electron configuration is the ground state electron configuration. Certain energies in the visible and uv regions of the spectrum can cause electrons to be excited into higher energy orbitals; some of the possible absorption transitions are indicated by the vertical arrows. Very energetic photons (uv to x-ray region of the spectrum) may cause an electron to be ejected from the molecule (ionization). Photons in the infrared region of the spectrum have much less energy than photons in the visible or uv regions of the electromagnetic spectrum. They can excite vibrations in molecules. There are many possible vibrational levels within each electronic state. Transitions between the vibrational levels are indicated by the vertical arrows on the left side of the diagram. Microwave radiation is even less energetic than infrared radiation. It cannot excite electrons in molecules, nor can it excite vibrations; it can only cause molecules to rotate. Microwave ovens are tuned to the frequency that causes molecules of water to rotate, and the ensuing friction causes heating of water-containing substances. Figure 3 illustrates these three types of molecular responses to radiation.

Figure 2 Energy Levels in Molecules

Figure 3 Molecular responses to radiation

What do we mean by molecular vibrations? Picture a diatomic molecule as two spheres connected by a spring. When the molecule vibrates, the atoms move towards and away from each other at a certain frequency. The energy of the system is related to how much the spring is stretched or compressed. The vibrational frequency is proportional to the square root of the ratio of the spring force constant to the masses on the spring. The lighter the masses on the spring, or the tighter (stronger) the spring, the higher the vibrational frequency will be. Similarly, vibrational frequencies for stretching bonds in molecules are related to the strength of the chemical bonds and the masses of the atoms. Molecules differ from sets of spheres-and-springs in that the vibrational frequencies are quantized. That is, only certain energies for the system are allowed, and only photons with certain energies will excite molecular vibrations. The symmetry of the molecule will also determine whether a photon can be absorbed.

The number of vibrational modes (different types of vibrations) in a molecule is 3N-5 for linear molecules and 3N-6 for nonlinear molecules, where N is the number of atoms. So the diatomic molecule we just discussed has 3 x 2 - 5 = 1 vibration: the stretching of the bond between the atoms. Carbon dioxide, a linear molecule, has 3 x 3 - 5 = 4 vibrations. These vibrational modes, shown in Figure 4, are responsible for the "greenhouse" effect in which heat radiated from the earth is absorbed (trapped) by CO2 molecules in the atmosphere. The arrows indicate the directions of motion. Vibrations labeled A and B represent the stretching of the chemical bonds, one in a symmetric (A) fashion, in which both C=O bonds lengthen and contract together (in-phase), and the other in an asymmetric (B) fashion, in which one bond shortens while the other lengthens. The asymmetric stretch (B) is infrared active because there is a change in the molecular dipole moment during this vibration. To be "active" means that absorption of a photon to excite the vibration is allowed by the rules of quantum mechanics. [Aside: the infrared "selection rule" states that for a particular vibrational mode to be observed (active) in the infrared spectrum, the mode must involve a change in the dipole moment of the molecule.] Infrared radiation at 2349 (4.26 um) excites this particular vibration. The symmetric stretch is not infrared active, and so this vibration is not observed in the infrared spectrum of CO2. The two equal-energy bending vibrations in CO2 (C and D in Figure 4) are identical except that one bending mode is in the plane of the paper, and one is out of the plane. Infrared radiation at 667 (15.00 um) excites these vibrations. Aside: another way of illustrating the out-of-plane mode is to place circled + or - signs on the atoms, signifying motion above of below the plane of the paper, respectively. Thought question: Why do you think it takes more energy (shorter wavelengths, higher frequencies) to excite the stretching vibration than the bending vibration?

Figure 4 Vibrations of CO2.

In addition to bond stretching and bond bending, more complicated molecules vibrate in rocking and twisting modes, which arise from combinations of bond bending in adjacent portions of a molecule. (These are sketched in the handout you received in lecture.) Torsions involve changes in dihedral angles. This type of mode is analogous to twisting the lid off the top of a jar. No bonds are stretched, and no bond angles change, but the spatial relationship between the atoms attached to each of two adjacent atoms will change. The torsional mode for ethane is illustrated below.

Without going into details at this point, we can note some general trends. The stronger the bond, the more energy will be required to excite the stretching vibration. This is seen in organic compounds where stretches for triple bonds such as C[[equivalence]]C and C[[equivalence]]N occur at higher frequencies than stretches for double bonds (C=C, C=N, C=O), which are in turn at higher frequencies than single bonds (C-C, C-N, C-H, O-H, or N-H). The heavier an atom, the lower the frequencies for vibrations that involve that atom. The characteristic regions for common infrared stretching and bending vibrations are given in Figure 5. Further details are given in the tables at the end of this handout.

How are infrared spectra obtained, and what do they look like? An infrared spectrometer consists of a glowing filament that generates infrared radiation (heat), which is passed through the sample to be studied. A detector measures the amount of radiation at various wavelengths that is transmitted by the sample. This information is recorded on a chart, where the percent of the incident light that is transmitted through the sample (% transmission) is plotted against wavelength in microns (um) or the frequency (). Remember that energy is inversely proportional to wavelength. If we define wavenumber (a.k.a. "reciprocal centimeters") = 1/ (), we have a parameter that is directly proportional to energy. Figure 6 shows the infrared spectrum of a gaseous sample of carbon dioxide. Note that the intensity of the transmitted light is close to 100% everywhere except where the sample absorbs: at 2349 (4.26 um) and at 667 (15.00 um).

Figure 5 Chart of Characteristic Vibrations

Figure 6 Infrared spectrum of Carbon Dioxide

Because the IR spectrum of each molecule is unique, it can serve as a signature or fingerprint to identify the molecule. This feature, along with the fact that it is a non-destructive technique, have made infrared spectroscopy a valuable method in chemical analysis. Areas in which it is used extensively include pharmaceutical analysis, quality control in industrial processes, environmental chemistry, geology and astronomy. One difficulty, however, is that the infrared (IR) spectra of molecules with more than a few atoms can be very complicated. How do we know what vibration each absorption band in the IR spectrum corresponds to? There are really three possible answers.

 1. It is possible to perform elaborate chemical calculations that allow us to develop pictures of each vibrational mode. How accurate these calculations are depends on the method used. As part of the laboratory assignment, you will examine the results of two different methods, as described further below. As with all modeling, these calculations are subject to a variety of errors, and so the quality of the results varies widely with the method and the complexity of the molecule being studied.

2. In many cases, it is not important to know the exact nature of each vibration. Rather, we might just want to whether certain functional groups (e.g. -COOH, -NH2, etc.) are present in the molecule. It turns out that the some molecular vibrations can be approximately described just in terms of the motions of a few of the atoms, while the other atoms move only slightly or not at all. This approximation is called "functional group analysis". It is particularly useful as a tool for qualitative analysis of organic molecules, and for monitoring the progress of organic reactions.

3) In other cases, we may not even care what the modes are! We may just want to obtain a spectrum of our sample, and compare it to a library of spectra of known compounds, in order to identify our sample. This procedure is common in environmental and forensic analyses.

As part of your laboratory and take-home assignment, you will participate in the first two types of analyses.

Laboratory Assignment

1. Computer exercises:

A. Using the Spartan program.

Introduction: The Spartan[TM] molecular modeling program uses quantum mechanics to solve the Schrodinger wave equation for the molecule of interest. Once the positions of the atoms and the molecular wavefunction are known, the various molecular vibrations and their frequencies are calculated. The Spartan program can use several different quantum mechanical methods to determine the molecular geometry (the lowest energy arrangement of the atoms in molecule) and the vibrational frequencies. All of the methods involve solving the Schrödinger equation for a system with many electrons; it is a formidable task, because it requires calculating many very complicated integrals. Ab initio ("from first principles") calculations do just that, without any approximations. If we provide a sufficiently exact description of the wavefunction on each atom and accurately account for their interactions in the molecule, the results (wavefunction, geometry, vibrational frequencies) will be very accurate. Poorer or simpler descriptions of the atomic wavefunctions and/or the interactions between electrons in the molecule will cause the results to be less accurate. Bond lengths may differ somewhat from actual (experimental) values, and vibrational frequencies will also be in so-so agreement with experiment. As part of the lab, you will examine the results of an "STO-3G" level ab initio calculation of the geometry and vibrations of CO2.

As you might imagine, ab initio calculations can be very time consuming. For molecules having more than a few second-row atoms, accurate calculations can take hours or days, even on the most powerful computers! Simplifying the calculations by setting some of the integrals equal to zero or to constants results in great time savings, often with little loss of accuracy. The parameters are set to give good agreement with experiment for a test set of molecules. This type of calculation is called "semi-empirical". Note that the calculated frequencies may deviate from experimental values by as much as 200 .

Activities on the Silicon Graphics, or Hewlett-Packardworkstation: The four molecules you will study are carbon dioxide, acetic acid, propionamide, and allyl benzoate. The calculations have already been done, and have been stored in the computer. You will examine the results. The files you will look at are:

* carbon_dioxide_STO3G

* carbon dioxide_PM3

* acetic_acid_IR

* propionamide_IR

* allyl benzoate_IR

* First, write the structure of each molecule, and calculate the number of normal modes. (You can check organic chemistry textbooks, or the Merck Index, for the structures of compounds with which you are not familiar.)

* After opening the Spartan file, select "VIBRATIONS" from the output menu, and make a list of the vibrational frequencies (round to the nearest integer!). Clicking on a frequency in the table will cause the corresponding vibrational mode to be displayed. You can use the mouse to rotate the molecule or to move it around on the screen to make it easier to see the displacements of the atoms. For each mode, identify the atoms involved and the types of vibration (stretch, bend, rock, torsion). (You can do this by sketching the molecule, and using arrows to show the direction of motion.) Do this for all of the modes of carbon dioxide and acetic acid, and for all the modes at wavenembers higher than 2100 and for every fifth one below 2100 for propionamide. For allyl benzoate, identify the range of frequencies that arise from the different types of vibration, (i.e. C-H stretch).

* Compare the vibrational frequencies and modes of CO2 calculated by the two methods, and with the experimental data given earlier in this handout.

Questions: What types of modes occur at the highest frequencies? At the lowest? Are there similar modes in the different molecules? If so, describe several of them. If not, suggest a reason.

2. Post-lab exercise

You will be given an infrared spectrum of an "unknown" molecule, along with its molecular formula.

(a) Using the method described below, calculate the number of sites of unsaturation (see explanation below).

(b) Draw several different Lewis structures consistent with the molecular formula. (For each arrangement of atoms, be sure that you draw the best Lewis structure!)

(c) Using the attached data tables, complete the IR diagnostic worksheet for the spectrum of your unknown. Propose a structure for the unknown that is consistent with your analysis.

Asides:

1) The data tables give characteristic experimental frequencies that are typical for the vibrations of different functional groups. Because they are based on experimental data, just like the spectrum of

your unknown, you can assume that the frequency ranges are accurate.

2) The tables make the approximation that the observed bands in the IR spectra arise from vibrational modes involving the motion of only a few atoms, e.g. a -NH2 group. But you know from the calculations you did in lab that real vibrational modes are much more complicated, and may involve many atoms on a molecule. Thus, there are assumptions and approximations involved in using the functional group tables!

3) The Spartan program provides convenient graphics that allow us to visualize vibrational modes. In reality, the amplitude of the motions are much smaller than were shown on the screen! Actual bond stretching vibrations involve a change in bond lengths of only, at most, 1-2% (i.e. about 0.01 Å!), while actual angle bending distortions are <= 5deg..

3) For more information on computational chemistry methods, such as the ab initio and semi-empirical methods used here, you will find a tutorial on the World Wide Web. Using Netscape or MOSAIC, go to http://www.nih.gov/molecular_modeling/gateway.html.

Determining Sites of Unsaturation

In organic compounds, the term saturated means that all the atoms have single bonds and that if the compound is a hydrocarbon, the general formula is . If there are any double or triple bonds, or if there is a ring in the molecular structure, then the compound is said to have sites of unsaturation. For example, hexane, (I) is saturated. Hexene (II) and cyclohexane(III), both isomers with the formula , have two fewer hydrogen atoms, or one site of unsaturation each.

When other atoms besides carbon and hydrogen are present, a formula can be used to determine the number of sites of unsaturation.

where the summation is over all atoms in the molecular formula. The valence is the number of bonds that the atom forms. Thus, the valence for C = 4; N = 3; O = 2; For H, Cl and the other halogens, the valence = 1.

Examples:

(1) Calculate the number of sites of unsaturation in the molecule .

Answer # = {[(2(4-2)+4(1-2)+2(2-2)]+2}/2 = 1

(2) Calculate the number of sites of unsaturation in .

Answer: # = {[(6(4-2)+6(1-2)]}/2 = 4

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