{"id":1798,"date":"2017-10-10T15:24:50","date_gmt":"2017-10-10T15:24:50","guid":{"rendered":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/?post_type=chapter&#038;p=1798"},"modified":"2018-10-05T19:51:26","modified_gmt":"2018-10-05T19:51:26","slug":"structure-determination-in-conjugated-systems-uv","status":"publish","type":"chapter","link":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/chapter\/structure-determination-in-conjugated-systems-uv\/","title":{"raw":"Structure Determination in Conjugated Systems UV","rendered":"Structure Determination in Conjugated Systems UV"},"content":{"raw":"<div class=\"elm-header\">\r\n<div class=\"elm-header-custom\">\r\n<div class=\"textbox learning-objectives\">\r\n<h3>Objectives<\/h3>\r\n<div id=\"elm-main-content\" class=\"elm-content-container\">\r\n<div>\r\n<div id=\"skills\">\r\n\r\nAfter completing this section, you should be able to\r\n<ol>\r\n \t<li>identify the ultraviolet region of the electromagnetic spectrum which is of most use to organic chemists.<\/li>\r\n \t<li>interpret the ultraviolet spectrum of 1,3-butadiene in terms of the molecular orbitals involved.<\/li>\r\n \t<li>describe in general terms how the ultraviolet spectrum of a compound differs from its infrared and NMR spectra.<\/li>\r\n<\/ol>\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<div id=\"elm-main-content\" class=\"elm-content-container\">\r\n<div>\r\n<div class=\"textbox key-takeaways\">\r\n<h3>Key Terms<\/h3>\r\nMake certain that you can define, and use in context, the key term below.\r\n<ul>\r\n \t<li>ultraviolet (UV) spectroscopy<\/li>\r\n<\/ul>\r\n<\/div>\r\n<\/div>\r\n<div id=\"note\">\r\n<div class=\"textbox\">\r\n<h3 class=\"boxtitle\">Study Notes<\/h3>\r\nUltraviolet 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.\r\n\r\n<\/div>\r\n<\/div>\r\n<img class=\"internal\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05155009\/wave.gif\" alt=\"\" width=\"269px\" height=\"88px\" \/>\u00a0\u00a0\u00a0\u00a0 <img class=\"internal\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05155011\/vispect2.gif\" alt=\"\" width=\"384px\" height=\"121px\" \/>\r\n<ul>\r\n \t<li><strong>Violet:<\/strong> \u00a0 400 - 420 nm<\/li>\r\n \t<li><strong>Indigo:<\/strong> \u00a0 420 - 440 nm<\/li>\r\n \t<li><strong>Blue:<\/strong> \u00a0 440 - 490 nm<\/li>\r\n \t<li><strong>Green:<\/strong> \u00a0 490 - 570 nm<\/li>\r\n \t<li><strong>Yellow:<\/strong> \u00a0 570 - 585 nm<\/li>\r\n \t<li><strong>Orange:<\/strong> \u00a0 585 - 620 nm<\/li>\r\n \t<li><strong>Red:<\/strong> \u00a0 620 - 780 nm<\/li>\r\n<\/ul>\r\nWhen white light passes through or is reflected by a colored substance, a characteristic portion of the mixed wavelengths is absorbed. The remaining light will then assume the complementary color to the wavelength(s) absorbed. This relationship is demonstrated by the color wheel shown below. Here, complementary colors are diametrically opposite each other. Thus, absorption of 420-430 nm light renders a substance yellow, and absorption of 500-520 nm light makes it red. Green is unique in that it can be created by absoption close to 400 nm as well as absorption near 800 nm.\r\n\r\n<img class=\"internal aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05155013\/colwheel.gif\" alt=\"\" width=\"248px\" height=\"184px\" \/>\r\n\r\nEarly humans valued colored pigments, and used them for decorative purposes. Many of these were inorganic minerals, but several important organic dyes were also known. These included the crimson pigment, kermesic acid, the blue dye, indigo, and the yellow saffron pigment, crocetin. A rare dibromo-indigo derivative, punicin, was used to color the robes of the royal and wealthy. The deep orange hydrocarbon carotene is widely distributed in plants, but is not sufficiently stable to be used as permanent pigment, other than for food coloring. A common feature of all these colored compounds, displayed below, is a system of <strong>extensively conjugated $$\\pi$$-electrons<\/strong>.\r\n<div id=\"section_1\">\r\n<h3 class=\"editable\">The Electromagnetic Spectrum<\/h3>\r\nThe visible spectrum constitutes but a small part of the total radiation spectrum. Most of the radiation that surrounds us cannot be seen, but can be detected by dedicated sensing instruments. This <strong>electromagnetic spectrum<\/strong> ranges from very short wavelengths (including gamma and x-rays) to very long wavelengths (including microwaves and broadcast radio waves). The following chart displays many of the important regions of this spectrum, and demonstrates the inverse relationship between wavelength and frequency (shown in the top equation below the chart).\r\n<p style=\"text-align: center\"><img class=\"internal aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05155015\/emspec.gif\" alt=\"\" width=\"629\" height=\"286\" \/><\/p>\r\nThe energy associated with a given segment of the spectrum is proportional to its frequency. The bottom equation describes this relationship, which provides the energy carried by a photon of a given wavelength of radiation.\r\n<p style=\"text-align: center\"><img class=\"internal aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05155017\/equation.gif\" alt=\"\" width=\"486px\" height=\"38px\" \/><\/p>\r\nTo obtain specific frequency, wavelength and energy values use this calculator.\r\n\r\n<\/div>\r\n<div id=\"section_2\">\r\n<h3 class=\"editable\">UV-Visible Absorption Spectra<\/h3>\r\nTo 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.\r\n\r\nThe visible region of the spectrum comprises photon energies of 36 to 72 kcal\/mole, and the near ultraviolet region, out to 200 nm, extends this energy range to 143 kcal\/mole. Ultraviolet radiation having wavelengths less than 200 nm is difficult to handle, and is seldom used as a routine tool for structural analysis.\r\n\r\n<img class=\"internal aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05155019\/electrns.gif\" alt=\"\" width=\"438px\" height=\"215px\" \/>\r\n\r\nThe energies noted above are sufficient to promote or excite a molecular electron to a higher energy orbital. Consequently, absorption spectroscopy carried out in this region is sometimes called \"electronic spectroscopy\". A diagram showing the various kinds of electronic excitation that may occur in organic molecules is shown on the left. Of the six transitions outlined, only the two lowest energy ones (left-most, colored blue) are achieved by the energies available in the 200 to 800 nm spectrum. 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 <strong>excited state<\/strong>.\r\n\r\nWhen 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. An optical spectrometer records the wavelengths at which absorption occurs, together with the degree of absorption at each wavelength. 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. <strong>Absorbance<\/strong> usually ranges from 0 (no absorption) to 2 (99% absorption), and is precisely defined in context with spectrometer operation.\r\n\r\n<\/div>\r\n<div id=\"section_3\">\r\n<h3 class=\"editable\">Electronic transitions<\/h3>\r\nLet\u2019s take as our first example the simple case of molecular hydrogen, H<sub>2<\/sub>.\u00a0 As you may recall from section 2.1A, the molecular orbital picture for the hydrogen molecule consists of one bonding \u03c3 MO, and a higher energy antibonding \u03c3* MO.\u00a0 When the molecule is in the ground state, both electrons are paired in the lower-energy bonding orbital \u2013 this is the Highest Occupied Molecular Orbital (HOMO).\u00a0 The antibonding \u03c3* orbital, in turn, is the Lowest Unoccupied Molecular Orbital (LUMO).\r\n\r\n<img class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05155022\/image023.png\" alt=\"image024.png\" width=\"535\" height=\"212\" \/>\r\n\r\nIf the molecule is exposed to light of a wavelength with energy equal to <strong>\u0394<\/strong>E, 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 \u2013 in other words, from the \u03c3 to the \u03c3* orbital. This is referred to as a <strong>\u03c3 <\/strong><strong>- <\/strong><strong>\u03c3<\/strong><strong>* transition<\/strong>. <strong>\u0394<\/strong>E for this electronic transition is 258 kcal\/mol, corresponding to light with a wavelength of 111 nm.\r\n\r\nWhen a double-bonded molecule such as ethene (common name ethylene) absorbs light, it undergoes a <strong>\u03c0<\/strong><strong> - <\/strong><strong>\u03c0<\/strong><strong>* transition.\u00a0 <\/strong>Because \u03c0- \u03c0* energy gaps are narrower than \u03c3 <strong>-<\/strong> \u03c3<strong>* <\/strong>gaps, ethene absorbs light at 165 nm - a longer wavelength than molecular hydrogen.\r\n\r\n<img class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05155024\/image025.png\" alt=\"image026.png\" width=\"597\" height=\"162\" \/>\r\n\r\nThe 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 \u2013 700 nm.\u00a0 Where UV-vis spectroscopy becomes useful to most organic and biological chemists is in the study of molecules with conjugated pi systems.\u00a0 In these groups, the energy gap for \u03c0 -\u03c0* transitions is smaller than for isolated double bonds, and thus the wavelength absorbed is longer.\u00a0 Molecules or parts of molecules that absorb light strongly in the UV-vis region are called <strong>chromophores<\/strong>.\r\n\r\nNext, we'll consider the 1,3-butadiene molecule. From valence orbital theory alone we might expect that the C<sub>2<\/sub>-C<sub>3<\/sub> bond in this molecule, because it is a sigma bond, would be able to rotate freely.\r\n\r\n<img class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05155027\/fig2-2-3.png\" alt=\"fig2-2-3.png\" width=\"158\" height=\"139\" \/>\r\n\r\nExperimentally, however, it is observed that there is a significant barrier to rotation about the C<sub>2<\/sub>-C<sub>3<\/sub> bond, and that the entire molecule is planar.\u00a0 In addition, the C<sub>2<\/sub>-C<sub>3<\/sub> bond is 148 pm long,\u00a0 shorter than a typical carbon-carbon single bond (about 154 pm), though\u00a0 longer than a typical double bond (about 134 pm).\r\n\r\nMolecular orbital theory accounts for these observations with the concept of <strong>delocalized \u03c0 bonds<\/strong>.\u00a0 In this picture, the four <em>p<\/em> atomic orbitals combine mathematically to form four pi molecular orbitals of increasing energy.\u00a0 Two of these - the bonding pi orbitals - are lower in energy than the <em>p<\/em> atomic orbitals from which they are formed, while two\u00a0 - the antibonding pi orbitals - are higher in energy.\r\n\r\n<img class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05155030\/fig2-2-4.png\" alt=\"fig2-2-4.png\" width=\"511\" height=\"440\" \/>\r\n\r\nThe lowest energy molecular orbital, pi<sub>1<\/sub>, has only constructive interaction and zero nodes. Higher in energy, but still lower than the isolated <em>p<\/em> orbitals, the pi<sub>2 <\/sub>orbital has one node but two constructive interactions - thus it is still a bonding orbital overall. Looking at the two antibonding orbitals,\u00a0 pi<sub>3<\/sub>* has two nodes and one constructive interaction, while pi<sub>4<\/sub>* has three nodes and zero constructive interactions.\r\n\r\nBy the <em>aufbau<\/em> principle, the four electrons from the isolated 2<em>p<\/em><sub>z<\/sub> atomic orbitals are placed in the bonding pi<sub>1<\/sub> and pi<sub>2<\/sub> MO\u2019s.\u00a0 Because pi<sub>1<\/sub> includes constructive interaction between C<sub>2<\/sub> and C<sub>3<\/sub>, there is a degree, in the 1,3-butadiene molecule, of pi-bonding interaction between these two carbons, which accounts for its shorter length and the barrier to rotation. The valence bond picture of 1,3-butadiene shows the two pi bonds as being isolated from one another, with each pair of pi electrons \u2018stuck\u2019 in its own pi bond.\u00a0 However, molecular orbital theory predicts (accurately) that the four pi electrons are to some extent delocalized, or \u2018spread out\u2019, over the whole pi system.\r\n\r\n<img class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05155033\/fig2-2-5.png\" alt=\"fig2-2-5.png\" width=\"408\" height=\"164\" \/>\r\n\r\n<a class=\"external\" title=\"http:\/\/wps.prenhall.com\/wps\/media\/objects\/340\/348272\/Instructor_Resources\/Chapter_15\/Text_Images\/FG15_02.JPG\" href=\"http:\/\/wps.prenhall.com\/wps\/media\/objects\/340\/348272\/Instructor_Resources\/Chapter_15\/Text_Images\/FG15_02.JPG\" target=\"_blank\" rel=\"external nofollow noopener\"><strong>space-filling view<\/strong><\/a>\r\n\r\n1,3-butadiene is the simplest example of a system of <strong>conjugated pi<\/strong><strong> bonds<\/strong>.\u00a0 To be considered conjugated, two or more pi bonds must be separated by only one single bond \u2013 in other words, there cannot be an intervening <em>sp<sup>3<\/sup><\/em>-hybridized carbon, because this would break up the overlapping system of parallel <em>p<\/em> orbitals.\u00a0 In the compound below, for example,\u00a0 the C<sub>1<\/sub>-C<sub>2<\/sub> and C<sub>3<\/sub>-C<sub>4<\/sub> double bonds are conjugated, while the C<sub>6<\/sub>-C<sub>7<\/sub> double bond is <strong>isolated<\/strong> from the other two pi bonds by <em>sp<sup>3<\/sup><\/em>-hybridized C<sub>5<\/sub>.\r\n\r\n<img class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05155035\/fig2-2-6.png\" alt=\"fig2-2-6.png\" width=\"256\" height=\"143\" \/>\r\n\r\nA very important concept to keep in mind is that <em>there is an inherent thermodynamic stability associated with conjugation.<\/em> This stability can be measured experimentally by comparing the\u00a0 <strong>heat of hydrogenation<\/strong> of two different dienes.\u00a0 (Hydrogenation is a reaction type that we will learn much more about in chapter 15: essentially, it is the process of adding a hydrogen molecule - two protons and two electrons - to a p bond). When the two <em>conjugated<\/em> double bonds of 1,3-pentadiene are 'hydrogenated' to produce pentane, about 225 kJ is released per mole of pentane formed.\u00a0\u00a0 Compare that to the approximately 250 kJ\/mol released when the two <em>isolated<\/em> double bonds in 1,4-pentadiene are hydrogenated, also forming pentane.\r\n\r\n<img class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05155038\/fig2-2-7.png\" alt=\"fig2-2-7.png\" width=\"361\" height=\"224\" \/>\r\n\r\nThe conjugated diene is lower in energy: in other words, it is more stable.\u00a0 In general, conjugated pi bonds are more stable than isolated pi bonds.\r\n\r\nConjugated pi systems can involve oxygen and nitrogen atoms as well as carbon.\u00a0 In the metabolism of fat molecules, some of the key reactions involve alkenes that are conjugated to carbonyl groups.\r\n\r\n<img class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05155040\/fig2-2-8.png\" alt=\"fig2-2-8.png\" width=\"113\" height=\"54\" \/>\r\n\r\nIn 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.\u00a0 Beta-carotene, with its system of 11 conjugated double bonds,\u00a0 absorbs light with wavelengths in the blue region of the visible spectrum while allowing other visible wavelengths \u2013 mainly those in the red-yellow region - to be transmitted. This is why carrots are orange.\r\n\r\n<img class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05155043\/image031.png\" alt=\"image032.png\" width=\"572\" height=\"136\" \/>\r\n<div>\r\n<div id=\"exercise\">\r\n<div class=\"textbox exercises\">\r\n<h3>Exercise 2.2.1<\/h3>\r\nIdentify all conjugated and isolated double bonds in the structures below. For each conjugated pi system, specify the number of overlapping <em>p<\/em> orbitals, and how many pi electrons are shared among them.\r\n\r\n<img class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05155045\/figE2-2-1.png\" alt=\"figE2-2-1.png\" width=\"346\" height=\"60\" \/>\r\n\r\n<\/div>\r\n<div class=\"textbox exercises\">\r\n<h3 class=\"boxtitle\">Exercise 2.2.2<\/h3>\r\nIdentify all isolated and conjugated pi bonds in lycopene, the red-colored compound in tomatoes.\u00a0 How many pi electrons are contained in the conjugated pi system?\r\n\r\n<img class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05155048\/figE2-2-2.png\" alt=\"figE2-2-2.png\" width=\"583\" height=\"87\" \/>\r\n\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<div id=\"section_4\">\r\n<div class=\"textbox exercises\">\r\n<div id=\"section_4\">\r\n<h3 class=\"editable\">Exercises<\/h3>\r\n<div id=\"s61719\">\r\n<div id=\"section_26\">\r\n\r\nWhat 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?\r\n<h3>Solution<\/h3>\r\n[reveal-answer q=\"18576\"]Show Answer[\/reveal-answer]\r\n[hidden-answer a=\"18576\"]\u00a0E = hc\/\u03bb E = (6.62 \u00d7 10<sup>-34<\/sup> Js)(3.00 \u00d7 10<sup>8<\/sup> m\/s)\/(3.00 \u00d7 10<sup>-7<\/sup>\u00a0m) E = 6.62 \u00d7 10<sup>-19<\/sup> J The range of 3.972 \u00d7 10<sup>-19<\/sup> to 6.62 \u00d7 10<sup>-19<\/sup>\u00a0joules. This energy range is greater in energy than the in NMR and IR.[\/hidden-answer]\r\n\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<div id=\"section_5\">\r\n<h3 class=\"editable\">Contributors<\/h3>\r\n<ul>\r\n \t<li><a class=\"external\" title=\"http:\/\/science.athabascau.ca\/staff-pages\/dietmark\" href=\"http:\/\/science.athabascau.ca\/staff-pages\/dietmark\" target=\"_blank\" rel=\"external nofollow noopener\">Dr. Dietmar Kennepohl<\/a> FCIC (Professor of Chemistry, <a class=\"external\" title=\"http:\/\/www.athabascau.ca\/\" href=\"http:\/\/www.athabascau.ca\/\" target=\"_blank\" rel=\"external nofollow noopener\">Athabasca University<\/a>)<\/li>\r\n \t<li>Prof. Steven Farmer (<a class=\"external\" title=\"http:\/\/www.sonoma.edu\" href=\"http:\/\/www.sonoma.edu\" target=\"_blank\" rel=\"external nofollow noopener\">Sonoma State University<\/a>)<\/li>\r\n \t<li><a title=\"Organic_Chemistry_With_a_Biological_Emphasis\" href=\"https:\/\/chem.libretexts.org\/Textbook_Maps\/Organic_Chemistry_Textbook_Maps\/Map%3A_Organic_Chemistry_with_a_Biological_Emphasis_(Soderberg)\" rel=\"internal\">Organic Chemistry With a Biological Emphasis <\/a>by\u00a0<a class=\"external\" title=\"http:\/\/facultypages.morris.umn.edu\/~soderbt\/\" href=\"http:\/\/facultypages.morris.umn.edu\/%7Esoderbt\/\" target=\"_blank\" rel=\"external nofollow noopener\">Tim Soderberg<\/a>\u00a0(University of Minnesota, Morris)<\/li>\r\n<\/ul>\r\n<\/div>\r\n<\/div>","rendered":"<div class=\"elm-header\">\n<div class=\"elm-header-custom\">\n<div class=\"textbox learning-objectives\">\n<h3>Objectives<\/h3>\n<div id=\"elm-main-content\" class=\"elm-content-container\">\n<div>\n<div id=\"skills\">\n<p>After completing this section, you should be able to<\/p>\n<ol>\n<li>identify the ultraviolet region of the electromagnetic spectrum which is of most use to organic chemists.<\/li>\n<li>interpret the ultraviolet spectrum of 1,3-butadiene in terms of the molecular orbitals involved.<\/li>\n<li>describe in general terms how the ultraviolet spectrum of a compound differs from its infrared and NMR spectra.<\/li>\n<\/ol>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<div id=\"elm-main-content\" class=\"elm-content-container\">\n<div>\n<div class=\"textbox key-takeaways\">\n<h3>Key Terms<\/h3>\n<p>Make certain that you can define, and use in context, the key term below.<\/p>\n<ul>\n<li>ultraviolet (UV) spectroscopy<\/li>\n<\/ul>\n<\/div>\n<\/div>\n<div id=\"note\">\n<div class=\"textbox\">\n<h3 class=\"boxtitle\">Study Notes<\/h3>\n<p>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.<\/p>\n<\/div>\n<\/div>\n<p><img decoding=\"async\" class=\"internal\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05155009\/wave.gif\" alt=\"\" width=\"269px\" height=\"88px\" \/>\u00a0\u00a0\u00a0\u00a0 <img decoding=\"async\" class=\"internal\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05155011\/vispect2.gif\" alt=\"\" width=\"384px\" height=\"121px\" \/><\/p>\n<ul>\n<li><strong>Violet:<\/strong> \u00a0 400 &#8211; 420 nm<\/li>\n<li><strong>Indigo:<\/strong> \u00a0 420 &#8211; 440 nm<\/li>\n<li><strong>Blue:<\/strong> \u00a0 440 &#8211; 490 nm<\/li>\n<li><strong>Green:<\/strong> \u00a0 490 &#8211; 570 nm<\/li>\n<li><strong>Yellow:<\/strong> \u00a0 570 &#8211; 585 nm<\/li>\n<li><strong>Orange:<\/strong> \u00a0 585 &#8211; 620 nm<\/li>\n<li><strong>Red:<\/strong> \u00a0 620 &#8211; 780 nm<\/li>\n<\/ul>\n<p>When white light passes through or is reflected by a colored substance, a characteristic portion of the mixed wavelengths is absorbed. The remaining light will then assume the complementary color to the wavelength(s) absorbed. This relationship is demonstrated by the color wheel shown below. Here, complementary colors are diametrically opposite each other. Thus, absorption of 420-430 nm light renders a substance yellow, and absorption of 500-520 nm light makes it red. Green is unique in that it can be created by absoption close to 400 nm as well as absorption near 800 nm.<\/p>\n<p><img decoding=\"async\" class=\"internal aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05155013\/colwheel.gif\" alt=\"\" width=\"248px\" height=\"184px\" \/><\/p>\n<p>Early humans valued colored pigments, and used them for decorative purposes. Many of these were inorganic minerals, but several important organic dyes were also known. These included the crimson pigment, kermesic acid, the blue dye, indigo, and the yellow saffron pigment, crocetin. A rare dibromo-indigo derivative, punicin, was used to color the robes of the royal and wealthy. The deep orange hydrocarbon carotene is widely distributed in plants, but is not sufficiently stable to be used as permanent pigment, other than for food coloring. A common feature of all these colored compounds, displayed below, is a system of <strong>extensively conjugated $$\\pi$$-electrons<\/strong>.<\/p>\n<div id=\"section_1\">\n<h3 class=\"editable\">The Electromagnetic Spectrum<\/h3>\n<p>The visible spectrum constitutes but a small part of the total radiation spectrum. Most of the radiation that surrounds us cannot be seen, but can be detected by dedicated sensing instruments. This <strong>electromagnetic spectrum<\/strong> ranges from very short wavelengths (including gamma and x-rays) to very long wavelengths (including microwaves and broadcast radio waves). The following chart displays many of the important regions of this spectrum, and demonstrates the inverse relationship between wavelength and frequency (shown in the top equation below the chart).<\/p>\n<p style=\"text-align: center\"><img loading=\"lazy\" decoding=\"async\" class=\"internal aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05155015\/emspec.gif\" alt=\"\" width=\"629\" height=\"286\" \/><\/p>\n<p>The energy associated with a given segment of the spectrum is proportional to its frequency. The bottom equation describes this relationship, which provides the energy carried by a photon of a given wavelength of radiation.<\/p>\n<p style=\"text-align: center\"><img decoding=\"async\" class=\"internal aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05155017\/equation.gif\" alt=\"\" width=\"486px\" height=\"38px\" \/><\/p>\n<p>To obtain specific frequency, wavelength and energy values use this calculator.<\/p>\n<\/div>\n<div id=\"section_2\">\n<h3 class=\"editable\">UV-Visible Absorption Spectra<\/h3>\n<p>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.<\/p>\n<p>The visible region of the spectrum comprises photon energies of 36 to 72 kcal\/mole, and the near ultraviolet region, out to 200 nm, extends this energy range to 143 kcal\/mole. Ultraviolet radiation having wavelengths less than 200 nm is difficult to handle, and is seldom used as a routine tool for structural analysis.<\/p>\n<p><img decoding=\"async\" class=\"internal aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05155019\/electrns.gif\" alt=\"\" width=\"438px\" height=\"215px\" \/><\/p>\n<p>The energies noted above are sufficient to promote or excite a molecular electron to a higher energy orbital. Consequently, absorption spectroscopy carried out in this region is sometimes called &#8220;electronic spectroscopy&#8221;. A diagram showing the various kinds of electronic excitation that may occur in organic molecules is shown on the left. Of the six transitions outlined, only the two lowest energy ones (left-most, colored blue) are achieved by the energies available in the 200 to 800 nm spectrum. 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 <strong>excited state<\/strong>.<\/p>\n<p>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. An optical spectrometer records the wavelengths at which absorption occurs, together with the degree of absorption at each wavelength. 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. <strong>Absorbance<\/strong> usually ranges from 0 (no absorption) to 2 (99% absorption), and is precisely defined in context with spectrometer operation.<\/p>\n<\/div>\n<div id=\"section_3\">\n<h3 class=\"editable\">Electronic transitions<\/h3>\n<p>Let\u2019s take as our first example the simple case of molecular hydrogen, H<sub>2<\/sub>.\u00a0 As you may recall from section 2.1A, the molecular orbital picture for the hydrogen molecule consists of one bonding \u03c3 MO, and a higher energy antibonding \u03c3* MO.\u00a0 When the molecule is in the ground state, both electrons are paired in the lower-energy bonding orbital \u2013 this is the Highest Occupied Molecular Orbital (HOMO).\u00a0 The antibonding \u03c3* orbital, in turn, is the Lowest Unoccupied Molecular Orbital (LUMO).<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05155022\/image023.png\" alt=\"image024.png\" width=\"535\" height=\"212\" \/><\/p>\n<p>If the molecule is exposed to light of a wavelength with energy equal to <strong>\u0394<\/strong>E, 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 \u2013 in other words, from the \u03c3 to the \u03c3* orbital. This is referred to as a <strong>\u03c3 <\/strong><strong>&#8211; <\/strong><strong>\u03c3<\/strong><strong>* transition<\/strong>. <strong>\u0394<\/strong>E for this electronic transition is 258 kcal\/mol, corresponding to light with a wavelength of 111 nm.<\/p>\n<p>When a double-bonded molecule such as ethene (common name ethylene) absorbs light, it undergoes a <strong>\u03c0<\/strong><strong> &#8211; <\/strong><strong>\u03c0<\/strong><strong>* transition.\u00a0 <\/strong>Because \u03c0- \u03c0* energy gaps are narrower than \u03c3 <strong>&#8211;<\/strong> \u03c3<strong>* <\/strong>gaps, ethene absorbs light at 165 nm &#8211; a longer wavelength than molecular hydrogen.<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05155024\/image025.png\" alt=\"image026.png\" width=\"597\" height=\"162\" \/><\/p>\n<p>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 \u2013 700 nm.\u00a0 Where UV-vis spectroscopy becomes useful to most organic and biological chemists is in the study of molecules with conjugated pi systems.\u00a0 In these groups, the energy gap for \u03c0 -\u03c0* transitions is smaller than for isolated double bonds, and thus the wavelength absorbed is longer.\u00a0 Molecules or parts of molecules that absorb light strongly in the UV-vis region are called <strong>chromophores<\/strong>.<\/p>\n<p>Next, we&#8217;ll consider the 1,3-butadiene molecule. From valence orbital theory alone we might expect that the C<sub>2<\/sub>-C<sub>3<\/sub> bond in this molecule, because it is a sigma bond, would be able to rotate freely.<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05155027\/fig2-2-3.png\" alt=\"fig2-2-3.png\" width=\"158\" height=\"139\" \/><\/p>\n<p>Experimentally, however, it is observed that there is a significant barrier to rotation about the C<sub>2<\/sub>-C<sub>3<\/sub> bond, and that the entire molecule is planar.\u00a0 In addition, the C<sub>2<\/sub>-C<sub>3<\/sub> bond is 148 pm long,\u00a0 shorter than a typical carbon-carbon single bond (about 154 pm), though\u00a0 longer than a typical double bond (about 134 pm).<\/p>\n<p>Molecular orbital theory accounts for these observations with the concept of <strong>delocalized \u03c0 bonds<\/strong>.\u00a0 In this picture, the four <em>p<\/em> atomic orbitals combine mathematically to form four pi molecular orbitals of increasing energy.\u00a0 Two of these &#8211; the bonding pi orbitals &#8211; are lower in energy than the <em>p<\/em> atomic orbitals from which they are formed, while two\u00a0 &#8211; the antibonding pi orbitals &#8211; are higher in energy.<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05155030\/fig2-2-4.png\" alt=\"fig2-2-4.png\" width=\"511\" height=\"440\" \/><\/p>\n<p>The lowest energy molecular orbital, pi<sub>1<\/sub>, has only constructive interaction and zero nodes. Higher in energy, but still lower than the isolated <em>p<\/em> orbitals, the pi<sub>2 <\/sub>orbital has one node but two constructive interactions &#8211; thus it is still a bonding orbital overall. Looking at the two antibonding orbitals,\u00a0 pi<sub>3<\/sub>* has two nodes and one constructive interaction, while pi<sub>4<\/sub>* has three nodes and zero constructive interactions.<\/p>\n<p>By the <em>aufbau<\/em> principle, the four electrons from the isolated 2<em>p<\/em><sub>z<\/sub> atomic orbitals are placed in the bonding pi<sub>1<\/sub> and pi<sub>2<\/sub> MO\u2019s.\u00a0 Because pi<sub>1<\/sub> includes constructive interaction between C<sub>2<\/sub> and C<sub>3<\/sub>, there is a degree, in the 1,3-butadiene molecule, of pi-bonding interaction between these two carbons, which accounts for its shorter length and the barrier to rotation. The valence bond picture of 1,3-butadiene shows the two pi bonds as being isolated from one another, with each pair of pi electrons \u2018stuck\u2019 in its own pi bond.\u00a0 However, molecular orbital theory predicts (accurately) that the four pi electrons are to some extent delocalized, or \u2018spread out\u2019, over the whole pi system.<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05155033\/fig2-2-5.png\" alt=\"fig2-2-5.png\" width=\"408\" height=\"164\" \/><\/p>\n<p><a class=\"external\" title=\"http:\/\/wps.prenhall.com\/wps\/media\/objects\/340\/348272\/Instructor_Resources\/Chapter_15\/Text_Images\/FG15_02.JPG\" href=\"http:\/\/wps.prenhall.com\/wps\/media\/objects\/340\/348272\/Instructor_Resources\/Chapter_15\/Text_Images\/FG15_02.JPG\" target=\"_blank\" rel=\"external nofollow noopener\"><strong>space-filling view<\/strong><\/a><\/p>\n<p>1,3-butadiene is the simplest example of a system of <strong>conjugated pi<\/strong><strong> bonds<\/strong>.\u00a0 To be considered conjugated, two or more pi bonds must be separated by only one single bond \u2013 in other words, there cannot be an intervening <em>sp<sup>3<\/sup><\/em>-hybridized carbon, because this would break up the overlapping system of parallel <em>p<\/em> orbitals.\u00a0 In the compound below, for example,\u00a0 the C<sub>1<\/sub>-C<sub>2<\/sub> and C<sub>3<\/sub>-C<sub>4<\/sub> double bonds are conjugated, while the C<sub>6<\/sub>-C<sub>7<\/sub> double bond is <strong>isolated<\/strong> from the other two pi bonds by <em>sp<sup>3<\/sup><\/em>-hybridized C<sub>5<\/sub>.<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05155035\/fig2-2-6.png\" alt=\"fig2-2-6.png\" width=\"256\" height=\"143\" \/><\/p>\n<p>A very important concept to keep in mind is that <em>there is an inherent thermodynamic stability associated with conjugation.<\/em> This stability can be measured experimentally by comparing the\u00a0 <strong>heat of hydrogenation<\/strong> of two different dienes.\u00a0 (Hydrogenation is a reaction type that we will learn much more about in chapter 15: essentially, it is the process of adding a hydrogen molecule &#8211; two protons and two electrons &#8211; to a p bond). When the two <em>conjugated<\/em> double bonds of 1,3-pentadiene are &#8216;hydrogenated&#8217; to produce pentane, about 225 kJ is released per mole of pentane formed.\u00a0\u00a0 Compare that to the approximately 250 kJ\/mol released when the two <em>isolated<\/em> double bonds in 1,4-pentadiene are hydrogenated, also forming pentane.<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05155038\/fig2-2-7.png\" alt=\"fig2-2-7.png\" width=\"361\" height=\"224\" \/><\/p>\n<p>The conjugated diene is lower in energy: in other words, it is more stable.\u00a0 In general, conjugated pi bonds are more stable than isolated pi bonds.<\/p>\n<p>Conjugated pi systems can involve oxygen and nitrogen atoms as well as carbon.\u00a0 In the metabolism of fat molecules, some of the key reactions involve alkenes that are conjugated to carbonyl groups.<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05155040\/fig2-2-8.png\" alt=\"fig2-2-8.png\" width=\"113\" height=\"54\" \/><\/p>\n<p>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.\u00a0 Beta-carotene, with its system of 11 conjugated double bonds,\u00a0 absorbs light with wavelengths in the blue region of the visible spectrum while allowing other visible wavelengths \u2013 mainly those in the red-yellow region &#8211; to be transmitted. This is why carrots are orange.<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05155043\/image031.png\" alt=\"image032.png\" width=\"572\" height=\"136\" \/><\/p>\n<div>\n<div id=\"exercise\">\n<div class=\"textbox exercises\">\n<h3>Exercise 2.2.1<\/h3>\n<p>Identify all conjugated and isolated double bonds in the structures below. For each conjugated pi system, specify the number of overlapping <em>p<\/em> orbitals, and how many pi electrons are shared among them.<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05155045\/figE2-2-1.png\" alt=\"figE2-2-1.png\" width=\"346\" height=\"60\" \/><\/p>\n<\/div>\n<div class=\"textbox exercises\">\n<h3 class=\"boxtitle\">Exercise 2.2.2<\/h3>\n<p>Identify all isolated and conjugated pi bonds in lycopene, the red-colored compound in tomatoes.\u00a0 How many pi electrons are contained in the conjugated pi system?<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05155048\/figE2-2-2.png\" alt=\"figE2-2-2.png\" width=\"583\" height=\"87\" \/><\/p>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<div id=\"section_4\">\n<div class=\"textbox exercises\">\n<div id=\"section_4\">\n<h3 class=\"editable\">Exercises<\/h3>\n<div id=\"s61719\">\n<div id=\"section_26\">\n<p>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?<\/p>\n<h3>Solution<\/h3>\n<div class=\"qa-wrapper\" style=\"display: block\"><span class=\"show-answer collapsed\" style=\"cursor: pointer\" data-target=\"q18576\">Show Answer<\/span><\/p>\n<div id=\"q18576\" class=\"hidden-answer\" style=\"display: none\">\u00a0E = hc\/\u03bb E = (6.62 \u00d7 10<sup>-34<\/sup> Js)(3.00 \u00d7 10<sup>8<\/sup> m\/s)\/(3.00 \u00d7 10<sup>-7<\/sup>\u00a0m) E = 6.62 \u00d7 10<sup>-19<\/sup> J The range of 3.972 \u00d7 10<sup>-19<\/sup> to 6.62 \u00d7 10<sup>-19<\/sup>\u00a0joules. This energy range is greater in energy than the in NMR and IR.<\/div>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<div id=\"section_5\">\n<h3 class=\"editable\">Contributors<\/h3>\n<ul>\n<li><a class=\"external\" title=\"http:\/\/science.athabascau.ca\/staff-pages\/dietmark\" href=\"http:\/\/science.athabascau.ca\/staff-pages\/dietmark\" target=\"_blank\" rel=\"external nofollow noopener\">Dr. Dietmar Kennepohl<\/a> FCIC (Professor of Chemistry, <a class=\"external\" title=\"http:\/\/www.athabascau.ca\/\" href=\"http:\/\/www.athabascau.ca\/\" target=\"_blank\" rel=\"external nofollow noopener\">Athabasca University<\/a>)<\/li>\n<li>Prof. Steven Farmer (<a class=\"external\" title=\"http:\/\/www.sonoma.edu\" href=\"http:\/\/www.sonoma.edu\" target=\"_blank\" rel=\"external nofollow noopener\">Sonoma State University<\/a>)<\/li>\n<li><a title=\"Organic_Chemistry_With_a_Biological_Emphasis\" href=\"https:\/\/chem.libretexts.org\/Textbook_Maps\/Organic_Chemistry_Textbook_Maps\/Map%3A_Organic_Chemistry_with_a_Biological_Emphasis_(Soderberg)\" rel=\"internal\">Organic Chemistry With a Biological Emphasis <\/a>by\u00a0<a class=\"external\" title=\"http:\/\/facultypages.morris.umn.edu\/~soderbt\/\" href=\"http:\/\/facultypages.morris.umn.edu\/%7Esoderbt\/\" target=\"_blank\" rel=\"external nofollow noopener\">Tim Soderberg<\/a>\u00a0(University of Minnesota, Morris)<\/li>\n<\/ul>\n<\/div>\n<\/div>\n","protected":false},"author":44985,"menu_order":15,"template":"","meta":{"_candela_citation":"[]","CANDELA_OUTCOMES_GUID":"","pb_show_title":"on","pb_short_title":"","pb_subtitle":"","pb_authors":[],"pb_section_license":""},"chapter-type":[],"contributor":[],"license":[],"class_list":["post-1798","chapter","type-chapter","status-publish","hentry"],"part":29,"_links":{"self":[{"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/pressbooks\/v2\/chapters\/1798","targetHints":{"allow":["GET"]}}],"collection":[{"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/pressbooks\/v2\/chapters"}],"about":[{"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/wp\/v2\/types\/chapter"}],"author":[{"embeddable":true,"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/wp\/v2\/users\/44985"}],"version-history":[{"count":10,"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/pressbooks\/v2\/chapters\/1798\/revisions"}],"predecessor-version":[{"id":2362,"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/pressbooks\/v2\/chapters\/1798\/revisions\/2362"}],"part":[{"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/pressbooks\/v2\/parts\/29"}],"metadata":[{"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/pressbooks\/v2\/chapters\/1798\/metadata\/"}],"wp:attachment":[{"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/wp\/v2\/media?parent=1798"}],"wp:term":[{"taxonomy":"chapter-type","embeddable":true,"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/pressbooks\/v2\/chapter-type?post=1798"},{"taxonomy":"contributor","embeddable":true,"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/wp\/v2\/contributor?post=1798"},{"taxonomy":"license","embeddable":true,"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/wp\/v2\/license?post=1798"}],"curies":[{"name":"wp","href":"https:\/\/api.w.org\/{rel}","templated":true}]}}