Hofstra Horizons Research

Magnetic Resonance Imaging and Nuclear Magnetic Resonance Spectroscopy: Studying Hydrogen in Different Environments

by Nanette M. Wachter

Magnetic Resonance Imaging and Nuclear Magnetic Resonance Spectroscopy: Studying Hydrogen in Different Environments

Most of us know someone who has had an MRI scan, or we have undergone the diagnostic procedure ourselves. MRI, magnetic resonance imaging, is a technique used by health professionals to acquire internal images of a patient without contrast agents or ionizing radiation. This noninvasive medical technique was developed in the early 1970s when scientists realized that a spectroscopic technique discovered in the mid-1940s, nuclear magnetic resonance (NMR) spectroscopy, could be applied to the analysis of tissues and tumors. Chemists and physicists employ NMR spectroscopy to characterize molecular structure and probe chemical dynamics.

How does magnetic resonance enable a chemist to visualize a molecule and a physician to diagnose a tumor? Both techniques use electromagnetic radiation in the radio frequency range to perturb nuclear spin. The subatomic particles of nuclei, protons and neutrons (collectively known as nucleons), produce a total angular momentum referred to as the nuclear spin. The most abundant element, hydrogen, has one proton in its nucleus and, except for its heavier isotopes (i.e., deuterium and tritium), no neutrons. Because it has an odd number of protons and no neutrons, a hydrogen nucleus has a spin that is easier to study than those of some other elements. A spinning particle is associated with a magnetic moment, that is, it behaves like a magnet and can be affected by a magnetic field. Nearly all organic and biological molecules contain hydrogen atoms. Since the hydrogen nucleus (which is a proton) possesses a magnetic moment, it can be observed by placing it in a magnetic field. When hydrogen is placed in a magnetic field, the field imposes a torque on the nuclear magnetic moment so that its spinning nucleus rotates, or precesses, about the applied magnetic field.1

Figure 1. Precession of nuclear magnetization, M. Ma represents the magnetization produced by population of nuclei with a-spin state, while Mß represents the magnetization produced by population of nuclei with ß-spin state.

Figure 1. Precession of nuclear magnetization, M. Ma represents the magnetization produced by population of nuclei with a-spin state, while Mß represents the magnetization produced by population of nuclei with ß-spin state.

The rate of precession depends on the type of nucleus (i.e., the number of nucleons) and on the magnetic field strength. It is often helpful to use vectors to demonstrate the direction of the magnetization with respect to the direction of the applied magnetic field (Figure 1). Hydrogen nuclei can precess in one of two directions, clockwise or counterclockwise, relative to an applied static magnetic field. Nuclei spinning clockwise (the alpha spin state) will act like little magnets and produce a bulk magnetization (M) aligned with the magnetic field. The direction of magnetization of those spinning counterclockwise (the beta spin state) will be opposed to the applied magnetic field and have higher energy than those aligned with the field. If energy is supplied that matches the difference between the energies of the two spin states, a quantum of energy will be absorbed and cause the nucleus to change its spin state (direction of precession). The absorption of energy is termed resonance, and so the technique is called nuclear magnetic resonance, or NMR.

How do chemists acquire structural information about molecules from NMR? NMR can be used to identify the connections of atoms in molecules because variations in the electronic environments of protons, such as the types of elements bonded to the hydrogen, affect the precessional frequencies of the nuclei. This phenomenon causes a shift in the frequency of the proton with respect to a reference and is referred to as the signal’s chemical shift. The resonant frequencies for hydrogen nuclei, however, are dependent on the strength of the applied magnetic field as well as the molecular environment of the proton. In order to compare NMR signals acquired on different NMR spectrometers, the frequencies are reported relative to the strength of the magnet used. Modern superconducting magnets typically employ field strengths ranging from 2,100 to 14,100 gauss (7-14.1 tesla), which corresponds to proton resonant frequencies of 300 MHz to 600 MHz. By convention, chemists report proton frequencies as a ratio of the signal’s chemical shift (in hertz) relative to the strength of the magnet (in megahertz) as parts per million. The NMR signals for hydrogen in organic molecules typically appear between 0 and 12 ppm, regardless of the strength of the magnet used. The region of the NMR spectrum closer to 0 ppm is referred to as “upfield,” while signals with higher chemical shift values are said to be “downfield.”

Figure 2. Ball-and-stick models of phenol, 2’-hydroxyacetophenone (2HA) and 4’-hydroxyacetophenone (4HA), respectively. Gray spheres represent carbon atoms, red spheres represent oxygen, and white spheres represent hydrogen.

Figure 2. Ball-and-stick models of phenol, 2’-hydroxyacetophenone (2HA) and 4’-hydroxyacetophenone (4HA), respectively. Gray spheres represent carbon atoms, red spheres represent oxygen, and white spheres represent hydrogen.

NMR as a tool for probing resonance-assisted hydrogen bonds. The three-dimensional structure of many biological molecules is dependent on hydrogen bonding interactions; for example, alpha helices and beta sheets in a protein’s tertiary structure or base-pairing in the DNA double helix. In much smaller molecules, the formation of inter- or intramolecular hydrogen bonds also has pronounced effects on their physical properties and reactivity. Regardless of the size of the molecule, intramolecular hydrogen bonds can form only if the molecular configuration can adopt conformation that allow the hydrogen bond donor group (typically an O-H or N-H substituent) to approach an acceptor atom on the molecule, such as another oxygen or nitrogen atom. Let’s consider three organic compounds capable of forming hydrogen bonds: phenol, 2’-hydroxyacetophenone (2HA) and 4’-hydroxyacetophenone (4HA), shown as ball-and-stick molecular models in Figure 2.

All three compounds have hydroxyl (O-H) groups and can therefore form intermolecular hydrogen bonds between the hydroxyl group on one molecule and an oxygen atom on another molecule. Only 2HA, however, is capable of forming an intramolecular hydrogen bond between the hydroxyl group and the oxygen on the carbonyl (C=O) group attached to the benzene ring (Figure 3).

NMR analysis of dilute chloroform solutions of each compound shows distinctly different signals for the hydroxyl protons. At room temperature, the chemical shift for the hydrogen attached to the oxygen of phenol appears at 4.8 ppm, while that of 4HA appears at 7.7 ppm, and the signal for the hydroxyl hydrogen in 2HA appears much further downfield at 12.2 ppm! What do these drastic differences in chemical shift imply? As noted earlier, hydrogen chemical shifts typically fall between 0 and 12 ppm, although values below 0 and above 12 ppm are not uncommon. The electrical environment of the nucleus is dependent on the molecular framework (i.e., the connectivity of the elements) and on the types of chemical bonds involved. Changes in the number of electrons near the nucleus alter the effective magnetic field felt by the proton. Signals closer to 0 ppm indicate that hydrogen experiences more electronic shielding from the magnetic field, while protons that give rise to NMR signals further downfield are said to be deshielded and encounter stronger effective magnetic fields. The large chemical shift observed for 2HA’s hydroxyl proton’s signal is partly attributed to weakening of the O-H bond by the proximal oxygen of the carbonyl (C=O) group.

Let’s back up a moment and examine an even simpler molecule, water. Small amounts of water in chloroform (a typical nonpolar organic solvent) produce a signal at 1.5 ppm for the hydrogen nuclei of water. The signal is therefore upfield of even the hydroxyl proton of phenol. In the water molecule, the oxygen atom is surrounded by two hydrogen atoms (H-O-H). Phenol, on the other hand, has a benzene ring attached to its oxygen atom. The benzene ring is made up of carbon atoms that have loosely held electrons hovering above and below the plane of the ring like two doughnuts (Figure 4). The electron network outside the plane of the molecule occupies pi bonds.

Figure 3. Representations of intermolecular hydrogen bonding in phenol (3a), and intramolecular hydrogen bond formation in 2HA (3b).

Figure 3. Representations of intermolecular hydrogen bonding in phenol (3a), and intramolecular hydrogen bond formation in 2HA (3b).

When an oxygen atom is attached to benzene, it can push its electrons into the pi bonds of the ring. Oxygen, therefore, becomes less basic, or, another way of looking at it, more acidic (i.e., better at donating a proton). As a result, the hydrogen of a hydroxyl group on a benzene ring is more easily dissociated from the oxygen. Interactions with oxygen atoms on other phenol molecules in solution, and even with the chlorinated solvent, weaken the O-H bond so that the hydrogen nucleus experiences less electronic shielding and appears further downfield. These weak intermolecular interactions are dependent on the chemical environment and temperature.2 If the solvent used to prepare the solution has a hydrogen bond acceptor atom (e.g., oxygen or nitrogen), then solute-solvent hydrogen bonds will compete with solute-solute hydrogen bonding. Variations in the number and types of intermolecular hydrogen bonds cause the signal to broaden. Since the solvent we are using in the current example, chloroform, does not contain oxygen or nitrogen atoms, it is not expected to compete with solute hydrogen bonding interactions. Temperature, however, can have a pronounced effect on molecular dynamics. The energy of a system is dependent on the temperature of its environment. Lowering the temperature of the system slows molecular motion and, consequently, enhances hydrogen bond interactions. As a result, the bond between the proton and hydrogen bond donor atom is weakened, and the chemical shift of the proton moves further downfield as the temperature is decreased. Chemical shifts that show little dependency on temperature variations imply that the hydrogen associated with that signal has less sensitivity to conformational or environmental changes.3 On the other hand, strong temperature dependence reflects a proton’s susceptibility to alterations in the molecule’s environment.

Figure 4. Delocalization of pi electrons above and below the plane of the benzene ring.

Figure 4. Delocalization of pi electrons above and below the plane of the benzene ring.

Imagine you are on a boat at sea. If you’re on a large cruise ship, you are less likely to be aware of the fluctuating conditions at sea. However, if you’re in a 20-foot sailboat, you’re much more susceptible to environmental factors. The cruise ship model applies to hydrogen in less polarized bonds, such as a carbon-hydrogen bond, while the sailor represents hydrogen in a more polarized O-H or N-H bond. By varying the temperature (i.e., weather conditions), we can explore the susceptibility of the hydrogen to external factors.

NMR investigations of dilute chloroform solutions of 2HA and 4HA at temperatures ranging from +50 C to -50 C reveal that the chemical shift for the hydroxyl proton in both compounds moves downfield (higher chemical shift values) as the temperature is decreased. This suggests that the hydrogen, or rather, the proton (the hydrogen without its electron), is moving away from the oxygen. Some other species with electrons must be attracting the proton; otherwise, why would the positive charge move away from the relatively electron-rich environment the oxygen provides? If we plot the chemical shift with respect to temperature, it is immediately apparent that the inverse relationship between chemical shift and temperature is linear. The graph shown in Figure 5 illustrates the relationship between the hydroxyl proton’s chemical shift and temperature for chloroform solutions of 2HA. Although a linear inverse relationship between chemical shift and temperature is observed for both 2HA and 4HA, the slopes of the lines for each graph are quite different. The magnitude of the slope for the dilute chloroform solution of 4HA is much steeper (-30 ppb/C) than the slope observed for the chloroform solution of 2HA (-1.9 ppb/C).4

Figure 5. Plots of hydroxyl proton chemical shift versus temperature (Kelvin) for chloroform solutions of 2HA.

Figure 5. Plots of hydroxyl proton chemical shift versus temperature (Kelvin) for chloroform solutions of 2HA.

It should be noted that the large temperature dependency for the chemical shift of the hydroxyl proton in chloroform solutions of 4HA is sensitive to the concentration of the solution. At high concentrations (five times greater), the slope’s magnitude is reduced to -20 ppb/C. When similar concentrations of 4HA are prepared in a more polar solvent (acetone) that is capable of hydrogen bonding with 4HA, the slope decreases to -8 ppb/C and there is little sensitivity to concentration changes.

The significantly small magnitude of the slope for the plot of the 2HA’s hydroxyl proton’s chemical shift versus temperature suggests that the hydrogen is much less sensitive to environmental changes. Moreover, the chemical shift is not concentration dependent and the slope does not vary with increasing concentration. These observations lead us to conclude that the hydrogen is involved in a more stable hydrogen bond than 4HA’s hydroxyl proton (recall the cruise ship).

As noted earlier, 2HA can form an intramolecular hydrogen bond with the carbonyl oxygen adjacent to the benzene ring. Earlier we saw that the oxygen in phenol can push its electron density into the pi bonds of the aromatic ring. The carbonyl group also contains pi electrons. However, in the carbonyl group the pi electrons delocalize from the carbonyl carbon to the oxygen, and since the carbonyl carbon is directly attached to the benzene ring, the pi electrons from the benzene ring can move out of the ring, or delocalize, toward the carbonyl group. Electron delocalization is referred to as “resonance.” This is unfortunate, as it is the same terminology used to describe the energy flow when nuclei transition between spin states. On the one hand, we’re observing energy frequencies necessary to excite nuclei. In the extended pi bonded systems, we are trying to describe the electronic structure and properties of the loosely held network of electrons outside the plane of the molecule. Organic chemists typically use multiple structures, referred to as resonance contributing structures, to represent electron delocalization. Figure 6 depicts resonance structures for 2HA.

The resonance structures help us conceptualize why the hydroxyl proton in 2HA is so far downfield (further removed from the hydroxyl oxygen and more exposed to the magnetic field), and why it shows much less sensitivity to temperature variations. 4HA has to seek out other molecules to form intermolecular hydrogen bonds. 2HA, conversely, has a hydrogen bond acceptor (the C=O oxygen) suitably located near the donor O-H group. Moreover, the basicity (anionic character) of the acceptor oxygen is enhanced by electron delocalization (resonance) from the pi system of the adjacent benzene ring. Resonance-assisted hydrogen bonding (RAHB) applies to systems in which electron delocalization is responsible for decreasing the electron density of the donor and increasing the anionic character of the acceptor.5

Figure 6. Resonance forms of 2HA depicting charge localization due to delocalization of pi and unshared electrons.

Figure 6. Resonance forms of 2HA depicting charge localization due to delocalization of pi and unshared electrons.

In my laboratory, we are exploring the relationship between hydrogen bonding and electron delocalization by developing more extensively conjugated compounds capable of forming intramolecular hydrogen bonds. One of the compounds we have synthesized, 4-N,N-dimethylamino-2’-hydroxychalcone (DMAHC, Figure 7), is a donor-acceptor system similar to 2HA but with a more extensive pi bond network. Additionally, there is a nitrogen atom attached to the second benzene ring on DMAHC. Nitrogen is better than oxygen at contributing electron density to an adjacent pi system (i.e., the benzene ring). As depicted in Figure 7, electron density from nitrogen donated into the ring is further “pushed” along the carbon chain to the carbonyl oxygen (the hydrogen acceptor atom). Empirically, electron delocalization can be observed by the human eye as well as NMR. DMAHC is much more colorful than most organic molecules. Colorful molecules have loosely held electrons. NMR analysis of DMAHC reveals that the chemical shift for the hydroxyl proton of DMAHC is even further downfield than that observed for 2HA, and in chloroform solutions appears at 13.3 ppm at room temperature.6

Figure 7. Resonance structures for DMAHC depicting variations in electronic structure and charge distribution.

Figure 7. Resonance structures for DMAHC depicting variations in electronic structure and charge distribution.

Perhaps of more significance, the signal for the hydroxyl proton in acetone solutions of DMAHC is found even further downfield, 13.5 ppm at ambient temperature. While acetone is capable of hydrogen bonding with the hydroxyl proton, there is no evidence that the solvent is interfering with intramolecular hydrogen bonding in DMAHC. In fact, it appears as if the polar solvent actually enhances intramolecular hydrogen bonding by stabilizing polarization in the molecule. Computational investigations of the electronic character performed by other researchers support the growing contribution of the zwitterionic form of the molecule.7

Why is it important to explore the contribution of hydrogen bonding to electron delocalization? Purine and pyrimidine rings found in DNA also contain conjugated pi electron networks with hydrogen bond donor and acceptor groups. The stability of the DNA double helix is attributed to hydrogen bond interactions between purine and pyrimidine bases, and stacking interactions between the adjacent base pairs.8

RAHB between the N-H donors and the oxygen or nitrogen acceptors of the purine and pyrimidine bases enhance polarization of the conjugated ring systems and may also contribute to base stacking electrostatic interactions important for the thermal stability of DNA. Since biological molecules exist in aqueous environments in living systems, it is important that we study hydrogen bond dynamics in solution.

Footnotes

  1. Claridge, Timothy D. W. (2009). High-Resolution NMR Techniques in Organic Chemistry, 2nd ed. Oxford: Elsevier, Ltd.
  2. Schneider, W. G., Bernstein, H. J. and Pople, J. A. (1958). Proton Magnetic Resonance Chemical Shift of Free and Associated Hydride Molecules. J. Chemical Physics, 28(4), 601-7.
  3. Gung, B. W., MacKay, J.A., and Zhu, Z. (1999). J. Organic Chemistry, 64, 700-6.
  4. Wachter, N. M. (2009). VT NMR Investigations of Hydrogen Bonding in Phenolic Systems. 237th National Meeting of the American Chemical Society, Salt Lake City, Utah.
  5. Jeffrey, G. A. (1997). An Introduction to Hydrogen Bonding. New York: Oxford University Press.
  6. Wachter-Jurcsak, N., and Detmer, C. A. (1999). 1H NMR Investigations of Resonance-Assisted Hydrogen Bonding in 4-(Dimethylamino)-2’-hydroxychalcone. Organic Letters, 1(5), 795-8.
  7. Garcia-Viloca, M., Gonzalez-Lafont, A., and Lluch, J. M. (2001). The 1H NMR chemical shift for the hydroxyl proton of 4-(dimethylamino)-2’-hydroxychalcone in chloroform: A theoretical approach to its inverse dependence on the temperature. Organic Letters, 3(4), 589-592.
  8. (a) Yakovchuk, P., Protozanova, E., and Frank-Kamenetskii, M.D. (2006). Base-stacking and base-pairing contributions into the thermal stability of the DNA double helix. Nucleic Acids Res., 34(2), 564-574. (b) Bloomfield, V. A., Crothers, D. M., Tinoco, I. (2000). Nucleic Acids: Structure, Properties and Function, Sausalito, CA: University Science Books. (c) Hobza, P., and Sandorfy, C. (1987). Non-empirical calculations on all 29 DNA base pairs. J. American Chemical Society, 109, 1302-7.

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