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Introduction

Infrared spectroscopy is a powerful tool utilized across many different chemistry disciplines. As the name suggests, this type of spectroscopy utilizes light from the infrared region of the electromagnetic spectrum (10¹² – 10¹⁴ Hz). IR spectroscopy induces vibrations within a molecule that subsequently change the dipole moment. When the dipole moment is changed, IR light is absorbed, causing a band to show up on the spectrum.

Metal carbonyls are metal complexes with ligands in which carbons are bound to oxygen. In the case of this article, the discussion will revolve around terminal metal carbonyls, where the carbonyl ligand is simply carbon monoxide.

Generally, IR spectroscopy doesn't give great details about molecular structure or connectivity. This is especially true in organic chemistry, where IR spectroscopy is mainly used to detect functional groups. However, in inorganic and organometallic chemistry it can actually be quite useful in determining molecular geometry and symmetry. Depending on the number of ligands and their orientation, one can expect to see varying patterns in the spectra.

Molecular symmetry and geometry is important in inorganic chemistry. Different geometries in molecules can lead to different chemical and physical properties. Understanding the difference in geometry, symmetry, and point group can then be applied to explain why these varying properties arise. Once the differences are understood, chemists can use the information to their advantage when it comes to synthesis, analysis, application, and so on.

IR Activity

In order for IR bands to appear in the spectra, there needs to be a change in the dipole moment. The dipole moment is essentially a measure of polarity within a molecules (think of differences in electronegativity of elements that make up any given molecule). When IR radiation hits a molecule, it will cause the bonds to vibrate, rotate, stretch, and bend. If these movements result in a change of dipole moment, IR bands will appear in the spectra because the light is being absorbed, so it won't be detected by the spectrometer. These changes in dipole are directly related to the position of molecules in a given sample. These chemical bonds are modeled using Hooke's Law (think of two balls connected by a spring). Hooke's Law is used to understand the energy between to bodies connected through some force. In the molecular case, it is the bond strength that acts as the "spring." When IR light hits the molecule, the force constant (bond strength) won't change, but the position of the molecules will. This will then cause the dipole moment to change, as it is dependent on the position of the molecules. Once the dipole moment changes, a peak will show. Depending on the position of the peak on the spectrum, scientists can then learn about the strength and nature of that particular bond.

As a quick way to see if a molecule will be IR active, one can look at the symmetry within a molecule. As a qualitative way of doing this, use VSEPR or the point group in order to visualize the molecule. Then draw different bending and stretching modes of the atoms (include dipole for the overall molecule). If the dipole moment changes, that bending/stretching mode will be IR active. This is a good, quick way to predict IR peaks, but usually scientists have to work in the reverse order. They first look at a spectra, then they fit bending and stretching modes to what is seen in spectroscopic data.

Below, in the figure, there are examples of the types of stretching and bending modes that one can expect to see occurring with a molecule. Each stretch and bend has its own unique name so it can be easily identified.

Reference

The above method for finding IR modes works well for simple molecules, but becomes difficult as the complexity of the molecule increases. For these systems, more quantitative methods need to be used. These qualitative methods are rooted in group theory and linear algebra. Knowing the point groups and stretching mode operators is imperative to fully understanding IR spectroscopy on a theoretical level, but isn't entirely necessary for the application of it.

Using CO

Scientists often use carbon monoxide (CO) when doing IR studies on metal complexes. Carbon monoxide gives large peaks that are very pronounced and easy to pick out from other stretches in spectra.

Using CO also lets scientists understand the electronic nature of metal centers in coordination chemistry. CO is a special ligand because it is a σ donor and π acceptor. The carbon donates its lone pair of electrons to the metal in a sigma fashion. The pi system in the CO allows it to accept electrons into the π* orbital from the metal, which strengthens the M-C bond but weakens the C-O bond. This phenomenon is known as backbonding.

Scientists have also been able to study metal carbonyl compounds and their reactions using IR spectroscopy coupled with isotopic labeling. Earlier in the article, it was established that scientists model chemical bonds using Hooke's Law. Using this model, mass will effect the stretching of bonds, so by substituting isotopes for either/both carbon and oxygen in the complex, scientists can learn more about the nature of complexes, their properties, and reactions. In general, substituting heavier isotopes at either carbon (¹²C for ¹³C) or oxygen (¹⁶O for ¹⁸O) will lower the stretching frequencies.

Fluctuations in the IR Peaks

The metal can affect where stretches show up in the IR spectra. Referring to the previous section, it is know that metals can put electrons into the π* orbital of the CO ligand. If a metal is more electron rich, it will be more willing to donate electrons back into CO. If the metal is more electron poor, the there will be little to no backbonding. The table below describes this relationship.

Reference

As shown in the table above, the nature of the metal center plays an important role in where CO stretches show. Anionic, or negatively charged, metal centers will be more electron donating, which will push electron density into the π* orbitals of the CO, which will reduce the C-O bond strength, which reduces the stretching frequency. As the metal center becomes more cationic, or electropositive in nature, there are less electrons to donate into the π* orbital, so the C-O bond is stronger and has a higher stretching frequency.

The metal can affect where IR bands of carbonyls show up, but this also happens when changing ligands on a metal carbonyl complex. Ligands that are electron donating will have a similar affect on IR peaks as an anionic metal. Electron withdrawing ligands with act like cationic metal centers. The table below shows how electron donating versus electron withdrawing ligands affects CO stretching.

Reference

Symmetry from IR Peaks

It has been established earlier that the symmetry of a given molecule gives a specific number of peaks in an IR spectrum. Since it can be rather difficult to figure out what are IR active versus IR inactive modes using simple methods, it is easier to look at a table and fit the number of peaks seen to the spectrum in question. Below is a table that gives the number of peaks of carbonyl ligands given, give the number of CO ligands as well as the coordination number. Keep in mind, these peaks will all show up in relatively the same area as described above. It should also be noted, this technique will work for any ligands where there are IR active modes. Carbonyls have just been a useful ligand because of the prominent peaks.

Reference



The table gives a lot of information about the geometry of carbonyl molecules given the number of stretches. Coupling IR experiments with different ones to find out how many carbonyls a compound actually has gives a wealth of information about any given metal carbonyl.

It should also be noted from above that one can tell which isomer is present in a given sample. Isomers have different and unique properties, while having the same make up. It is always important to understand why these differences arise, and spectroscopy is an important tool scientists use.