Ultraviolet and visible spectroscopy

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While interaction with infrared light causes molecules to undergo vibrational transitions, the shorter wavelength, higher energy radiation in the UV (200-400 nm) and visible (400-700 nm) range of the electromagnetic spectrum causes many organic molecules to undergo electronic transitions. What this means is that when the energy from UV or visible light is absorbed by a molecule, one of its electrons jumps from a lower energy to a higher energy molecular orbital.

Electronic transitions

Let’s take as our first example the simple case of molecular hydrogen, H2.  As you may recall from section 2.1A, the molecular orbital picture for the hydrogen molecule consists of one bonding σ MO, and a higher energy antibondingσ* MO.  When the molecule is in the ground state, both electrons are paired in the lower-energy bonding orbital – this is the Highest Occupied Molecular Orbital (HOMO).  The antibonding σ* orbital, in turn, is the Lowest Unoccupied Molecular Orbital (LUMO).

If the molecule is exposed to light of a wavelength with energy equal to ΔE, the HOMO-LUMO energy gap, this wavelength will be absorbed and the energy used to bump one of the electrons from the HOMO to the LUMO – in other words, from the σ to the σ* orbital. This is referred to as a σ - σ* transition. ΔE for this electronic transition is 258 kcal/mol, corresponding to light with a wavelength of 111 nm.

When a double-bonded molecule such as ethene (common name ethylene) absorbs light, it undergoes a π - π* transition.  Because π- π* energy gaps are narrower than σ - σ* gaps, ethene absorbs light at 165 nm - a longer wavelength than molecular hydrogen.

The electronic transitions of both molecular hydrogen and ethene are too energetic to be accurately recorded by standard UV spectrophotometers, which generally have a range of 220 – 700 nm.  Where UV-vis spectroscopy becomes useful to most organic and biological chemists is in the study of molecules with conjugated pi systems.  In these groups, the energy gap for π -π* transitions is smaller than for isolated double bonds, and thus the wavelength absorbed is longer.  Molecules or parts of molecules that absorb light strongly in the UV-vis region are called chromophores.

Let’s revisit the MO picture for 1,3-butadiene, the simplest conjugated system (see section 2.1B).  Recall that we can draw a diagram showing the four pi MO’s that result from combining the four 2pz atomic orbitals. The lower two orbitals are bonding, while the upper two are antibonding.

Comparing this MO picture to that of ethene, our isolated pi-bond example, we see that the HOMO-LUMO energy gap is indeed smaller for the conjugated system. 1,3-butadiene absorbs UV light with a wavelength of 217 nm.

As conjugated pi systems become larger, the  energy gap for a π - π* transition becomes increasingly narrow, and the wavelength of light absorbed correspondingly becomes longer.   The absorbance due to the π - π* transition in 1,3,5-hexatriene, for example, occurs at 258 nm, corresponding to a ΔE of 111 kcal/mol.

In molecules with extended pi systems, the HOMO-LUMO energy gap becomes so small that absorption occurs in the visible rather then the UV region of the electromagnetic spectrum.  Beta-carotene, with its system of 11 conjugated double bonds,  absorbs light with wavelengths in the blue region of the visible spectrum while allowing other visible wavelengths – mainly those in the red-yellow region - to be transmitted. This is why carrots are orange.

The conjugated pi system in 4-methyl-3-penten-2-one gives rise to a strong UV absorbance at 236 nm due to a π - π* transition.  However, this molecule also absorbs at 314 nm.  This second absorbance is due to the transition of a non-bonding (lone pair) electron on the oxygen up to a π* antibonding MO:

This is referred to as an n - π* transition.  The nonbonding (n) MO’s are higher in energy than the highest bonding p orbitals, so the energy gap for an n - π* transition is smaller that that of a π - π* transition – and thus the n - π* peak is at a longer wavelength.  In general, n - π* transitions are weaker (less light absorbed) than those due to π - π* transitions.

 

 

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Looking at UV-vis spectra

We have been talking in general terms about how molecules absorb UV and visible light – now let's look at some actual examples of data from a UV-vis absorbance spectrophotometer. The basic setup is the same as for IR spectroscopy: radiation with a range of wavelengths is directed through a sample of interest, and a detector records which wavelengths were absorbed and to what extent the absorption occurred.  Below is the absorbance spectrum of an important biological molecule called nicotinamide adenine dinucleotide, abbreviated NAD+ (we'll learn what it does in section 16.4)  This compound absorbs light in the UV range due to the presence of conjugated pi-bonding systems.

You’ll notice that this UV spectrum is much simpler than the IR spectra we saw earlier: this one has only one peak, although many molecules have more than one.  Notice also that the convention in  UV-vis spectroscopy is to show the  baseline at the bottom of the graph with the peaks pointing up.  Wavelength values on the x-axis are generally measured in nanometers (nm) rather than in cm-1 as is the convention in IR spectroscopy. 

Peaks in UV spectra tend to be quite broad, often spanning well over 20 nm at half-maximal height.  Typically, there are two things that we look for and record from a UV-Vis spectrum..  The first is λmax, which is the wavelength at maximal light absorbance.  As you can see, NAD+has λmax, = 260 nm.  We also want to record how much light is absorbed at λmax.  Here we use a unitless number called absorbance, abbreviated 'A'.  This contains the same information as the 'percent transmittance' number used in IR spectroscopy, just expressed in slightly different terms.  To calculate absorbance at a given wavelength, the computer in the spectrophotometer simply takes the intensity of light at that wavelength before it passes through the sample (I0), divides this value by the intensity of the same wavelength after it passes through the sample (I), then takes the log10 of that number:

                   A = log I0/I

You can see that the absorbance value at 260 nm (A260) is about 1.0 in this spectrum. 

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Here is the absorbance spectrum of the common food coloring Red #3:

Here, we see that the extended system of conjugated pi bonds causes the molecule to absorb light in the visible range.  Because the  λmaxof 524 nm falls within the green region of the spectrum, the compound appears red to our eyes.

Now, take a look at the spectrum of another food coloring, Blue #1:

Here, maximum absorbance is at 630 nm, in the orange range of the visible spectrum, and the compound appears blue. 

Applications of UV spectroscopy in organic and biological chemistry

UV-vis spectroscopy has many different applications in organic and biological chemistry.  One of the most basic of these applications is the use of the Beer - Lambert Law to determine the concentration of a chromophore.  You most likely have performed a Beer – Lambert experiment in a previous chemistry lab.  The law is simply an application of the observation that, within certain ranges, the absorbance of a chromophore at a given wavelength varies in a linear fashion with its concentration: the higher the concentration of the molecule, the greater its absorbance.   If we divide the observed value of A at λmax by the concentration of the sample (c, in mol/L), we obtain the molar absorptivity, or extinction coefficient (ε), which is a characteristic value for a given compound. 

                   ε = A/c

The absorbance will also depend, of course, on the path length - in other words, the distance that the beam of light travels though the sample.  In most cases, sample holders are designed so that the path length is equal to 1 cm, so the units for molar absorptivity are  mol * L-1cm-1.  If we look up the value of e for our compound at λmax, and we measure absorbance at this wavelength, we can easily calculate the concentration of our sample.   As an example, for NAD+ the literature value of ε at 260 nm is 18,000 mol * L-1cm-1.   In our NAD+spectrum we observed A260 = 1.0, so using equation 4.4 and solving for concentration we find that our sample is 5.6 x 10-5 M.  

 

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The bases of DNA and RNA are good chromophores:

Biochemists and molecular biologists often determine the concentration of a DNA sample by assuming an average value of  ε = 0.020 ng-1×mL for double-stranded DNA at its λmax of 260 nm (notice that concentration in this application is expressed in mass/volume rather than molarity:  ng/mL is often a convenient unit for DNA concentration when doing molecular biology).

 

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Because the extinction coefficient of double stranded DNA is slightly lower than that of single stranded DNA, we can use UV spectroscopy to monitor a process known as DNA melting.   If  a short stretch of double stranded DNA is gradually heated up, it will begin to ‘melt’, or break apart, as the temperature increases (recall that two strands of DNA are held together by a specific pattern of hydrogen bonds formed by ‘base-pairing’).

As melting proceeds, the absorbance value for the sample increases, eventually reaching a high plateau as all of the double-stranded DNA breaks apart, or ‘melts’.  The mid-point of this process, called the ‘melting temperature’, provides a good indication of how tightly the two strands of DNA are able to bind to each other. 

In section 16.8 we will see how the Beer - Lambert Law and UV spectroscopy provides us with a convenient way to follow the progress of many different enzymatic redox (oxidation-reduction) reactions.  In biochemistry, oxidation of an organic molecule often occurs concurrently with reduction of nicotinamide adenine dinucleotide (NAD+, the compound whose spectrum we saw earlier in this section) to NADH:

Both NAD+ and NADH absorb at 260 nm.  However NADH, unlike NAD+, has a second absorbance band with λmax = 340 nm and ε = 6290 mol L-1cm-1.  The figure below shows the spectra of both compounds superimposed, with the NADH spectrum offset slightly on the y-axis:

By monitoring the absorbance of a reaction mixture at 340 nm, we can 'watch' NADH being formed as the reaction proceeds, and calculate the rate of the reaction.

UV spectroscopy is also very useful in the study of proteins.  Proteins absorb light in the UV range due to the presence of the aromatic amino acids tryptophan, phenylalanine, and tyrosine, all of which are chromophores. 

Biochemists frequently use UV spectroscopy to study conformational changes in proteins - how they change shape in response to different conditions.  When a protein undergoes a conformational shift (partial unfolding, for example), the resulting change in the environment around an aromatic amino acid chromophore can cause its UV spectrum to be altered.

Infrared spectroscopy

Covalent bonds in organic molecules are not rigid sticks – rather, they behave more like springs.  At room temperature, organic molecules are always in motion, as their bonds stretch, bend, and twist.  These complex vibrations can be broken down mathematically into individualvibrational modes, a few of which are illustrated below.

The energy of molecular vibration is quantized rather than continuous, meaning that a molecule can only stretch and bend at certain 'allowed' frequencies.  If a molecule is exposed to electromagnetic radiation that matches the frequency of one of its vibrational modes,  it will in most cases absorb energy from the radiation and jump to a higher vibrational energy state - what this means is that the amplitude of the vibration will increase, but the vibrational frequency will remain the same.  The difference in energy between the two vibrational states is equal to the energy associated with the wavelength of radiation that was absorbed.  It turns out that it is the infrared region of the electromagnetic spectrum which contains frequencies corresponding to the vibrational frequencies of organic bonds.

Let's take 2-hexanone as an example.  Picture the carbonyl bond of the ketone group as a spring.  This spring is constantly bouncing back and forth, stretching and compressing, pushing the carbon and oxygen atoms further apart and then pulling them together.  This is thestretching mode of the carbonyl bond.  In the space of one second, the spring 'bounces' back and forth 5.15 x 1013 times - in other words,  the ground-state frequency of carbonyl  stretching for a the ketone group is about 5.15 x 1013 Hz. 

If our ketone sample is irradiated with infrared light, the carbonyl bond will specifically absorb light with this same frequency, which by equations 4.1 and 4.2 corresponds to a wavelength of 5.83 x 10-6 m and an energy of 4.91 kcal/mol.  When the carbonyl bond absorbs this energy, it jumps up to an excited vibrational state.

The value of ΔE - the energy difference between the low energy (ground)  and high energy (excited) vibrational states - is equal to 4.91 kcal/mol, the same as the energy associated with the absorbed light frequency.  The molecule does not remain in its excited vibrational state for very long, but quickly releases energy to the surrounding environment in form of heat, and returns to the ground state.

With an instrument called an infrared spectrophotometer, we can 'see' this vibrational transition.  In the spectrophotometer, infrared light with frequencies ranging from about 1013 to 1014 Hz  is passed though our sample of cyclohexane.  Most frequencies pass right through the sample and are recorded by a detector on the other side.

Our 5.15 x 1013 Hz carbonyl stretching frequency, however, is absorbed by the 2-hexanone sample, and so the detector records that the intensity of this frequency, after having passed through the sample, is something less than 100% of its initial intensity.

The vibrations of a 2-hexanone molecule are not, of course, limited to the simple stretching of the carbonyl bond.  The various carbon-carbon bonds also stretch and bend, as do the carbon-hydrogen bonds, and all of these vibrational modes also absorb different frequencies of infrared light.

The power of infrared spectroscopy arises from the observation that different functional groups have different characteristic absorption frequencies.   The carbonyl bond in a ketone, as we saw with our 2-hexanone example, typically absorbs in the range of  5.11 -  5.18 x 1013Hz, depending on the molecule.   The carbon-carbon triple bond of an alkyne, on the other hand, absorbs in the range  6.30 - 6.80 x 1013Hz.   The technique is therefore very useful as a means of identifying which functional groups are present in a molecule of interest.  If we pass infrared light through an unknown sample and find that it absorbs in the carbonyl frequency range but not in the alkyne range, we can infer that the molecule contains a carbonyl group but not an alkyne.

Some bonds absorb infrared light more strongly than others, and some bonds do not absorb at all. In order for a vibrational mode to absorb infrared light, it must result in a periodic change in the dipole moment of the molecule.  Such vibrations are said to be infrared active. In general, the greater the polarity of the bond, the stronger its IR absorption.  The carbonyl bond is very polar, and absorbs very strongly.  The carbon-carbon triple bond in most alkynes, in contrast, is much less polar, and thus a stretching vibration does not result in a large change in the overall dipole moment of the molecule. Alkyne groups absorb rather weakly compared to carbonyls.

Some kinds of vibrations are infrared inactive.  The stretching vibrations of completely symmetrical double and triple bonds, for example, do not result in a change in dipole moment, and therefore do not result in any absorption of light (but other bonds and vibrational modes in these molecules do absorb IR light).

Now, let's  look at some actual output from IR spectroscopy experiments.  Below is the IR spectrum for 2-hexanone.

There are a number of things that need to be explained in order for you to understand what it is that we are looking at.  On the horizontal axis we see IR wavelengths expressed in terms of a unit called wavenumber (cm-1), which tells us how many waves fit into one centimeter.   On the vertical axis we see ‘% transmittance’, which tells us how strongly light was absorbed at each frequency (100% transmittance means no absorption occurred at that frequency).  The solid line traces the values of % transmittance for every wavelength – the ‘peaks’ (which are actually pointing down) show regions of strong absorption.  For some reason, it is typical in IR spectroscopy to report wavenumber values rather than wavelength (in meters) or frequency (in Hz).  The ‘upside down’ vertical axis, with absorbance peaks pointing down rather than up, is also a curious convention in IR spectroscopy.  We wouldn’t  want to make things too easy for you!

 

 

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The key absorption peak in this spectrum is that from the carbonyl double bond, at 1716 cm-1 (corresponding to a wavelength of 5.86 mm, a frequency of 5.15 x 1013 Hz, and a ΔE value of 4.91 kcal/mol).  Notice how strong this peak is, relative to the others on the spectrum:  a strong peak in the 1650-1750 cm-1 region is a dead giveaway for the presence of a carbonyl group.  Within that range, carboxylic acids, esters, ketones, and aldehydes tend to absorb in the shorter wavelength end (1700-1750 cm-1), while conjugated unsaturated ketones and amides tend to absorb on the longer wavelength end (1650-1700 cm-1).

The jagged peak at approximately 2900-3000 cm-1 is characteristic of tetrahedral carbon-hydrogen bonds.  This peak is not terribly useful, as just about every organic molecule that you will have occasion to analyze has these bonds.  Nevertheless, it can serve as a familiar reference point to orient yourself in a spectrum. 

You will notice that there are many additional peaks in this spectrum in the longer-wavelength 400 -1400 cm-1 region.  This part of the spectrum is called the fingerprint region.  While it is usually very difficult to pick out any specific functional group identifications from this region, it does, nevertheless, contain valuable information.  The reason for this is suggested by the name: just like a human fingerprint, the pattern of absorbance peaks in the fingerprint region is unique to every molecule, meaning that the data from an unknown sample can be compared to the IR spectra of known standards in order to make a positive identification.  In the mid-1990's, for example, several paintings were identified as forgeries because scientists were able to identify the IR footprint region of red and yellow pigment compounds that would not have been available to the artist who supposedly created the painting (for more details see Chemical and Engineering News, Sept 10, 2007, p. 28).

Now, let’s take a look at the IR spectrum for 1-hexanol.

As you can see, the carbonyl peak is gone, and in its place is a very broad ‘mountain’ centered at about 3400 cm-1.  This signal is characteristic of the O-H stretching mode of alcohols, and is a dead giveaway for the presence of an alcohol group.  The breadth of this signal is a consequence of hydrogen bonding between molecules.

In the spectrum of octanoic acid we see, as expected, the characteristic carbonyl peak, this time at 1709 cm-1.

We also see a low, broad absorbance band that looks like an alcohol, except that it is displaced slightly to the right (long-wavelength) side of the spectrum, causing it to overlap to some degree with the C-H region.  This is the characteristic carboxylic acid O-H single bond stretching absorbance.

The spectrum for 1-octene shows two peaks that are characteristic of alkenes: the one at 1642 cm-1 is due to stretching of the carbon-carbon double bond, and the one at 3079 cm-1 is due to stretching of the s bond between the alkene carbons and their attached hydrogens.

Alkynes have characteristic IR absorbance peaks in the range of 2100-2250 cm-1 due to stretching of the carbon-carbon triple bond, and terminal alkenes can be identified by their absorbance at about 3300 cm-1, due to stretching of the bond between the sp-hybridized carbon and the terminal hydrogen.

It is possible to identify other functional groups such as amines and ethers, but the characteristic peaks for these groups are considerably more subtle and/or variable, and often are overlapped with peaks from the fingerprint region.  For this reason, we will limit our discussion here to the most easily recognized functional groups, which are summarized in table 1 in the tables section at the end of the text.

As you can imagine, obtaining an IR spectrum for a compound will not allow us to figure out the complete structure of even a simple molecule, unless we happen to have a reference spectrum for comparison.  In conjunction with other analytical methods, however, IR spectroscopy can prove to be a very valuable tool, given the information it provides about the presence or absence of key functional groups. IR can also be a quick and convenient way for a chemist to check to see if a reaction has proceeded as planned.  If we were to run a reaction in which we wished to convert cyclohexanone to cyclohexanol, for example, a quick comparison of the IR spectra of starting compound and product would tell us if we had successfully converted the ketone group to an alcohol 

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Chemical equivalence

If all protons in all organic molecules had the same resonance frequency in an external magnetic field of a given strength, the information in the previous paragraph would be interesting from a theoretical standpoint, but would not be terribly useful to organic chemists.  Fortunately for us, however, resonance frequencies are not uniform for all protons in a molecule.  In an external magnetic field of a given strength, protons in different locations in a molecule have different resonance frequencies, because they are in non-identical electronic environments.  In methyl acetate, for example, there are two ‘sets’ of protons.  The three protons labeled H_a have a different - and easily distinguishable – resonance frequency than the three H_b protons, because the two sets of protons are in non-identical environments: they are, in other words, chemically nonequivalent. 

On the other hand, the three H_a protons are all in the same electronic environment, and are chemically equivalent to one another.  They have identical resonance frequencies. The same can be said for the three H_b protons.

The ability to recognize chemical equivalancy and nonequivalency among atoms in a molecule will be central to understanding NMR.  In each of the molecules below, all protons are chemically equivalent, and therefore will have the same resonance frequency in an NMR experiment.　

You might expect that the equitorial and axial hydrogens in cyclohexane would be non-equivalent, and would have different resonance frequencies.  In fact, an axial hydrogen is in a different electronic environment than an equitorial hydrogen.  Remember, though, that the molecule rotates rapidly between its two chair conformations, meaning that any given hydrogen is rapidly moving back and forth between equitorial and axial positions.  It turns out that, except at extremely low temperatures, this rotational motion occurs on a time scale that is much faster than the time scale of an NMR experiment. 

In this sense, NMR is like a camera that takes photographs of a rapidly moving object with a slow shutter speed - the result is a blurred image.  In NMR terms, this means that all 12 protons in cyclohexane are equivalent.

Each the molecules in the next figure contains two sets of protons, just like our previous example of methyl acetate, and again in each case the resonance frequency of the H_a protons will be different from that of the H_b protons. 

Notice how the symmetry of para-xylene results in there being only two different sets of protons.

Most organic molecules have several sets of protons in different chemical environments, and each set, in theory, will have a different resonance frequency in ¹H-NMR spectroscopy.

When stereochemistry is taken into account, the issue of  equivalence vs nonequivalence in NMR starts to get a little more complicated.  It should be fairly intuitive that hydrogens on different sides of asymmetric ring structures and double bonds are in different electronic environments, and thus are non-equivalent and have different resonance frequencies. In the alkene and cyclohexene structures below, for example, H_a is trans to the chlorine substituent, while H_b is cis to chlorine.

What is not so intuitive is that diastereotopic hydrogens (section 3.10) on chiral molecules are also non-equivalent:

However, enantiotopic and homotopic hydrogens are chemically equivalent.

 

Example

How many different sets of protons do the following molecules contain? (count diastereotopic protons as non-equivalent).

The origin of the NMR signal

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NMR-active nuclei

The basis for nuclear magnetic resonance is the observation that many atomic nuclei spin about an axis and generate their own magnetic field, or magnetic moment.  For reasons that are outside the scope of this text, only those nuclei with an odd number of protons and/or neutrons have a magnetic moment.  Fortunately for chemists, several common nuclei, including hydrogen (¹H), the ¹³C isotope of carbon, the ¹⁹F isotope of fluorine, and the ³¹P isotope of phosphorus, all have magnetic moments and therefore can be observed by NMR – they are, in other words, NMR-active. Other nuclei - such as the common ¹²C and ¹⁶O isotopes of carbon and oxygen - do not have magnetic moments, and are essentially invisible in NMR. Other nuclei such as deuterium (²H) and nitrogen (¹⁴N) have magnetic moments and are NMR-active, but the nature of their magnetic moments is such that these nuclei are more difficult to analyze by NMR.   In practice it is ¹H, ¹³C, ¹⁹F, and ³¹P that are most often observed by NMR spectroscopy. In this chapter, we will develop our understanding of the principles behind NMR spectroscopy by focusing our attention first on the detection of protons in ¹H-NMR experiments (in discussion about NMR, the terms 'hydrogen' and 'proton' are used interchangeably). Much of what we learn, however, will also apply to the detection and analysis of other NMR-active nuclei, and later in the chapter we will shift our attention to NMR experiments involving ¹³C and ³¹P atoms.

Nuclear precession, spin states, and the resonance condition

When a sample of an organic compound is sitting in a flask on a laboratory benchtop, the magnetic moments of its hydrogen atoms are randomly oriented.   When the same sample is placed within the field of a very strong magnet in an NMR instrument (this field is referred to by NMR spectroscopists as the applied field, abbreviated B₀ ) each hydrogen will assume one of two possible spin states.  In what is referred to as the +½  spin state, the hydrogen's magnetic moment is aligned with the direction of B₀, while in the -½ spin state it is alignedopposed to the direction of B₀. 

Because the +½ spin state is slightly lower in energy, in a large population of organic molecules slightly more than half of the hydrogen atoms will occupy this state, while slightly less than half will occupy the –½ state.  The difference in energy between the two spin states increases with increasing strength of B₀.This last statement is in italics because it is one of the key ideas in NMR spectroscopy, as we shall soon see.

At this point, we need to look a little more closely at how a proton spins in an applied magnetic field.  You may recall playing with spinning tops as a child.  When a top slows down a little and the spin axis is no longer completely vertical, it begins to exhibit precessional motion, as the spin axis rotates slowly around the vertical. In the same way, hydrogen atoms spinning in an applied magnetic field also exhibit precessional motion about a vertical axis.  It is this axis (which is either parallel or antiparallel to B₀) that defines the proton’s magnetic moment.  In the figure below, the proton is in the +1/2 spin state.

The frequency of precession (also called the Larmour frequency, abbreviated ω_L) is simply the number of times per second that the proton precesses in a complete circle.  A　proton`s precessional frequency increases with the strength of B₀.

If a proton that is precessing in an applied magnetic field is exposed to electromagnetic radiation of a frequency ν that matches its precessional frequency ω_L, we have a condition called resonance.  In the resonance condition, a proton in the lower-energy +½ spin state (aligned with B₀) will transition (flip) to the higher energy –½ spin state (opposed to B₀).  In doing so, it will absorb radiation at this resonance frequency ν = ω_L.  This frequency, as you might have already guessed, corresponds to the energy difference between the proton’s two spin states. With the strong magnetic fields generated by the superconducting magnets used in modern NMR instruments, the resonance frequency for protons falls within the radio-wave range, anywhere from 100 MHz to 800 MHz depending on the strength of the magnet.