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9.2.1: Electronic Spectra - Ultraviolet and Visible Spectroscopy - Organics

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    222266
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    Objectives

    After completing this section, you should be able to

    1. identify the ultraviolet region of the electromagnetic spectrum which is of most use to organic chemists.
    2. interpret the ultraviolet spectrum of 1,3-butadiene in terms of the molecular orbitals involved.
    3. describe in general terms how the ultraviolet spectrum of a compound differs from its infrared and NMR spectra.
    Key Terms

    Make certain that you can define, and use in context, the key term below.

    • ultraviolet (UV) spectroscopy
    • Molar absorptivity
    Study Notes

    Ultraviolet spectroscopy provides much less information about the structure of molecules than do the spectroscopic techniques studied earlier (infrared spectroscopy, mass spectroscopy, and NMR spectroscopy). Thus, your study of this technique will be restricted to a brief overview. You should, however, note that for an organic chemist, the most useful ultraviolet region of the electromagnetic spectrum is that in which the radiation has a wavelength of between 200 and 400 nm.

    UV-Visible Absorption Spectra

    To understand why some compounds are colored and others are not, and to determine the relationship of conjugation to color, we must make accurate measurements of light absorption at different wavelengths in and near the visible part of the spectrum. Commercial optical spectrometers enable such experiments to be conducted with ease, and usually survey both the near ultraviolet and visible portions of the spectrum. Ultraviolet-visible absorption spectroscopy provides much less information about the structure of molecules than do the spectroscopic techniques studied earlier (infrared spectroscopy, mass spectroscopy, and NMR spectroscopy) and mainly provides information about conjugated pi systems. For an organic chemist the most useful ultraviolet region of the electromagnetic spectrum involves radiation with a wavelength between 200 and 400 nm. UV/Vis absorption spectra also involve radiation from the visible region of the electromagnetic spectrum with wavelengths between 400 and 800 nm.

    A diagram highlighting the various kinds of electronic excitation that may occur in organic molecules is shown below. Of the six transitions outlined, only the two lowest energy ones, n to pi* and pi to pi* (colored blue) are achieved by the energies available in the 200 to 800 nm range of a UV/VIs spectrum. These energies are sufficient to promote or excite a molecular electron to a higher energy orbital in many conjugated compounds.

    electrns.gif

    When sample molecules are exposed to light having an energy that matches a possible electronic transition within the molecule, some of the light energy will be absorbed as the electron is promoted to a higher energy orbital. A UV/Vis spectrometer records the wavelengths at which absorption occurs, together with the degree of absorption at each wavelength. Absorbance, abbreviated 'A', is a unitless number which contains the same information as the 'percent transmittance' number used in IR spectroscopy. 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:

    \[A = \log \left (\dfrac{I_0}{I} \right) \]

    The resulting spectrum is presented as a graph of absorbance (A) versus wavelength, as in the isoprene spectrum shown below. Since isoprene is colorless, it does not absorb in the visible part of the spectrum and this region is not displayed on the graph. Notice 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.

    Typically, there are two things that are noted and recorded from a UV-Vis spectrum. The first is λmax, which is the wavelength at maximal light absorbance. As you can see, isoprene has λmax, = 222 nm. The second valuable piece of data is the absorbance at the λmax. In the isoprene spectrum the absorbance at the value λmax of 222 nm is about 0.8.

    isopren1.gif

    Molar absorptivity

    Molar absorptivity (\(epsilon\)) is a physical constant, characteristic of the particular substance being observed and thus characteristic of the particular electron system in the molecule. Molar absorptivities may be very large for strongly absorbing chromophores (>100,000) and very small if absorption is weak (10 to 100). The magnitude of \(epsilon\) reflects both the size of the chromophore and the probability that light of a given wavelength will be absorbed when it strikes the chromophore. Molar absorptivity (\(\epsilon\)) is defined via Beer's law as:

    \[\epsilon = \dfrac{A}{c\; l}\]

    where

    • \(A\) is the sample absorbance
    • \(c\) is the sample concentration in moles/liter
    • \(l\) is the length of light path through the sample in cm

    If the isoprene spectrum show above was obtained from a dilute hexane solution (c = 4 * 10-5 moles per liter) in a 1 cm sample cuvette, a simple calculation using the above formula indicates a molar absorptivity of 20,000 at the maximum absorption wavelength, symbolized as \(λ_{max}\).

    The only molecular moieties likely to absorb light in the 200 to 800 nm region are functional groups that contain pi-electrons and hetero atoms having non-bonding valence-shell electron pairs. Such light absorbing groups are referred to as chromophores. A list of some simple chromophores and their light absorption characteristics are provided below. The oxygen non-bonding electrons in alcohols and ethers do not give rise to absorption above 160 nm. Consequently, pure alcohol and ether solvents may be used for spectroscopic studies.

    Chromophore Example Excitation λmax, nm ε @ λmax Solvent
    C=C Ethene π → π* 171 15,000 hexane
    C≡C 1-Hexyne π → π* 180 10,000 hexane
    C=O Ethanal n → π*
    π → π*
    290
    180
    15
    10,000
    hexane
    hexane
    N=O Nitromethane

    n → π*
    π → π*

    275
    200
    17
    5,000
    ethanol
    ethanol
    C-X X=Br
    X=I
    Methyl bromide
    Methyl Iodide

    n → σ*
    n → σ*

    205
    255
    200
    360
    hexane
    hexane

    Electronic Transitions (cause of UV-Visible absorption)

    As previously noted, electronic transitions in organic molecules lead to UV and visible absorption. As a rule, energetically favored electron promotion will be from the highest occupied molecular orbital (HOMO) to the lowest unoccupied molecular orbital (LUMO), and the resulting species is called an excited state. The molecular orbital picture for the hydrogen molecule (H2) consists of one bonding σ MO, and a higher energy antibonding σ* 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).

    Hydroden has one lone pair and one unpaired electron in the sigma orbital and one unpaired electron in the sigma anti-bonding orbital.

    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.

    Ethene has one lone pair and one unpaired electron in the pi orbital and one unpaired electron in the pi anti-bonding orbital.

    The electronic transitions of both molecular hydrogen and ethene are too energetic to be accurately recorded. Where electronic transition 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. The MO diagram for 1,3-butadiene, the simplest conjugated system. 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 pi bonding, while the upper two are pi antibonding.

    1,3-butadiene had three lone pairs and one unpaired electron in the bonding orbital and one unpaired electron in the anti-bonding orbital.

    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.

    1,3,5-hexatriene has five lone pairs in the bonding orbital and one unpaired electron in the antibonding orbital.

    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.

    Bond line drawing of beta-carotene.

    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:

    4-methyl-3-penten-2-one has four lone pairs in the bonding orbital, three lone pairs and one unpaired electron in the n orbital, and one unpaired electron in the anti-bonding orbital.

    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.

    Use of UV/Vis Spectroscopy in Biological Systems

    The bases of DNA and RNA are good chromophores:

    Four RNA/DNA base structures. From left to right: adenine, guanine, cytosine, thymine.

    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).

    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’).

    Graph of DNA melting. Starts with strand of double stranded DNA (one strand red and one strand blue) at low absorbance and low temperature. Dotted line shows absorbance as temperature increases until the DNA becomes single stranded. Dashed line at 'melting temp' with melting DNA structure. Blue strand starts bending away from red strand.

    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.

    Later 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:

    Reaction: Reduced molecule plus NAD + goes to oxidized molecule plus 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 L*mol-1*cm-1. The figure below shows the spectra of both compounds superimposed, with the NADH spectrum offset slightly on the y-axis:

    Graph of absorbance vs wavelength comparing NADH (blue, top) and NAD+ (red, bottom). NADH starts at an absorbance of 0.6 and peaks to 1.1 around 260 nanometers. Second peak around 340 nanometers to an absorbance of 0.7. NAD+ starts at an absorbance of 0.4 and peaks to 0.8 around 260 nanometers. No second peak.

    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.

    Three aromatic amino acids. From left to right: phenylalanine, tyrosine, tryptophan.

    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.

    Exercise \(\PageIndex{1}\)
    1. 50 microliters of an aqueous sample of double stranded DNA is dissolved in 950 microliters of water. This diluted solution has a maximal absorbance of 0.326 at 260 nm. What is the concentration of the original (more concentrated) DNA sample, expressed in micrograms per microliter?
    2. What is the energy range for 300 nm to 500 nm in the ultraviolet spectrum? How does this compare to energy values from NMR and IR spectroscopy?
    3. Identify all isolated and conjugated pi bonds in lycopene, the red-colored compound in tomatoes. How many pi electrons are contained in the conjugated pi system?

    figE2-2-2.png

    Answer

    1) Using ε = A/c, we plug in our values for ε and A and find that c = 3.27 x 10-5M, or 32.7 mM.

    2)

    E = hc/λ

    E = (6.62 × 10−34 Js)(3.00 × 108 m/s)/(3.00 × 10−7 m)

    E = 6.62 × 10−19 J

    The range of 3.972 × 10-19 to 6.62 × 10-19 joules. This energy range is greater in energy than the in NMR and IR.

    3)

    lycopene_example.png

    Objective

    After completing this section, you should be able to use data from ultraviolet spectra to assist in the elucidation of unknown molecular structures.

    Study Notes

    It is important that you recognize that the ultraviolet absorption maximum of a conjugated molecule is dependent upon the extent of conjugation in the molecule.

    The Importance of Conjugation

    A comparison of the UV/Vis absorption spectrum of 1-butene, λmax = 176 nm, with that of 1,3-butadiene, λmax = 292 nm, clearly demonstrates that the effect of increasing conjugation is to shift toward longer wavelength (lower frequency, lower energy) absorptions. Further evidence of this effect is shown below. The spectrum on the left illustrates that conjugation of double and triple bonds also shifts the absorption maximum to longer wavelengths. From the polyene spectra displayed in the right it is clear that each additional double bond in the conjugated pi-electron system increases the absorption maximum about 30 nm. Also, the molar absorptivity (ε) roughly doubles with each new conjugated double bond. Spectroscopists use the terms defined in the table below when describing shifts in absorption. Thus, extending conjugation generally results in bathochromic and hyperchromic shifts in absorption.

    conjenyn.gifpolyene.gif

    Terminology for Absorption Shifts
    Nature of Shift Descriptive Term
    To Longer Wavelength Bathochromic
    To Shorter Wavelength Hypsochromic
    To Greater Absorbance Hyperchromic
    To Lower Absorbance Hypochromic

    Many other kinds of conjugated pi-electron systems act as chromophores and absorb light in the 200 to 800 nm region. These include unsaturated aldehydes and ketones and aromatic ring compounds. A few examples are displayed below. The spectrum of the unsaturated ketone (on the left) illustrates the advantage of a logarithmic display of molar absorptivity. The \(\pi \rightarrow \pi^*\) absorption located at 242 nm is very strong, with an ε = 18,000. The weak \(n \rightarrow \pi^*\) absorption near 300 nm has an ε = 100.

    enone2a.gif thioestr0.gif

    Benzene exhibits very strong light absorption near 180 nm (ε > 65,000) , weaker absorption at 200 nm (ε = 8,000) and a group of much weaker bands at 254 nm (ε = 240). Only the last group of absorptions are completely displayed because of the 200 nm cut-off characteristic of most spectrophotometers. The added conjugation in naphthalene, anthracene and tetracene causes bathochromic shifts of these absorption bands, as displayed in the chart below. All the absorptions do not shift by the same amount, so for anthracene (green shaded box) and tetracene (blue shaded box) the weak absorption is obscured by stronger bands that have experienced a greater red shift. As might be expected from their spectra, naphthalene and anthracene are colorless (with their absorptions in the UV range), but tetracene is orange (since its absorptions move into the visible range).

    polyarom.gif

    Looking at UV-Vis Spectra

    Below is the absorbance spectrum of an important biological molecule called nicotinamide adenine dinucleotide, abbreviated NAD+. This compound absorbs light in the UV range due to the presence of conjugated pi-bonding systems.

    UV-vis spectra of NAD plus where the max wavelength is at 260 nanometers.

    Below is the absorbance spectrum of the common food coloring Red #3. The extended system of conjugated pi bonds causes the molecule to absorb light in the visible range. Because the λmax of 524 nm falls within the green region of the spectrum, the compound appears red to our eyes (recalling the color wheel from Section 14.7).

    The max wavelength is at 524 nanometers

    Example 14.8.2

    How large is the π - π* transition in 4-methyl-3-penten-2-one?

    Solution

    Example 14.8.3

    Which of the following molecules would you expect absorb at a longer wavelength in the UV region of the electromagnetic spectrum? Explain your answer.

    image035.png

    Solution

    Exercise \(\PageIndex{1}\)

    Which of the following would show UV absorptions in the 200-300 nm range?

    14.8.png

    Answer

    B and D would be in that range.


    9.2.1: Electronic Spectra - Ultraviolet and Visible Spectroscopy - Organics is shared under a not declared license and was authored, remixed, and/or curated by LibreTexts.