{"id":1637,"date":"2017-10-12T13:15:43","date_gmt":"2017-10-12T13:15:43","guid":{"rendered":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/?post_type=chapter&#038;p=1637"},"modified":"2017-11-20T19:43:41","modified_gmt":"2017-11-20T19:43:41","slug":"infrared-spectroscopy","status":"publish","type":"chapter","link":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/chapter\/infrared-spectroscopy\/","title":{"raw":"Infrared Spectroscopy","rendered":"Infrared Spectroscopy"},"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 class=\"elm-header\"><\/div>\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 (by wavelength, wavenumber, or both) the region of the electromagnetic spectrum which is used in infrared (IR) spectroscopy.<\/li>\r\n \t<li>interconvert between wavelength and wavenumber.<\/li>\r\n \t<li>discuss, in general terms, the effect that the absorption of infrared radiation can have on a molecule.<\/li>\r\n<\/ol>\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<div class=\"textbox key-takeaways\">\r\n<h3>Key Terms<\/h3>\r\n<div id=\"elm-main-content\" class=\"elm-content-container\">\r\n<div>\r\n<div>\r\n\r\nMake certain that you can define, and use in context, the key terms below.\r\n<ul>\r\n \t<li>infrared spectrum<\/li>\r\n \t<li>wavenumber (reciprocal centimetres)<\/li>\r\n<\/ul>\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<div class=\"textbox\">\r\n<div class=\"elm-header\">\r\n<div class=\"elm-header-custom\">\r\n<h3>Study Notes<\/h3>\r\n<\/div>\r\n<\/div>\r\n<div id=\"elm-main-content\" class=\"elm-content-container\">\r\n<div>\r\n<div id=\"note\">\r\n\r\nNotice that the scale at the bottom of the infrared spectrum for 2-hexanone shown is calibrated in wavenumbers (cm<sup>\u22121<\/sup>). A wavenumber is the reciprocal of a wavelength (1\/\u03bb); thus, a wavenumber of 1600 cm<sup>\u22121<\/sup> corresponds to a wavelength of\r\n\r\n1 1600\u2009 cm \u22121 =6.25\u00d7 10 \u22124 cm\u00a0or\u00a06.25\u2009\u03bc\u2009m\r\n\r\nOrganic chemists find it more convenient to deal with wavenumbers rather than wavelengths when discussing infrared spectra.\r\n\r\nYou will obtain infrared spectra for a number of the compounds you will synthesize in the laboratory component of this course.\r\n\r\nThe inverted peaks observed in the spectra correspond to molecular stretching and bending vibrations that only occur at certain quantized frequencies. When infrared radiation matching these frequencies falls on the molecule, the molecule absorbs energy and becomes excited. Eventually the molecule returns to its original (ground) state, and the energy which was absorbed is released as heat.\r\n\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 id=\"section_1\">\r\n<h3 class=\"editable\">The electromagnetic spectrum<\/h3>\r\nElectromagnetic radiation, as you may recall from a previous chemistry or physics class, is composed of electrical and magnetic waves which oscillate on perpendicular planes. Visible light is electromagnetic radiation. So are the gamma rays that are emitted by spent nuclear fuel, the x-rays that a doctor uses to visualize your bones, the ultraviolet light that causes a painful sunburn when you forget to apply sun block, the infrared light that the army uses in night-vision goggles, the microwaves that you use to heat up your frozen burritos, and the radio-frequency waves that bring music to anybody who is old-fashioned enough to still listen to FM or AM radio.\r\n\r\nJust like ocean waves, electromagnetic waves travel in a defined direction. While the speed of ocean waves can vary, however, the speed of electromagnetic waves \u2013 commonly referred to as the speed of light \u2013 is essentially a constant, approximately 300 million meters per second. This is true whether we are talking about gamma radiation or visible light. Obviously, there is a big difference between these two types of waves \u2013 we are surrounded by the latter for more than half of our time on earth, whereas we hopefully never become exposed to the former to any significant degree.\u00a0 The different properties of the various types of electromagnetic radiation are due to differences in their wavelengths, and the corresponding differences in their energies: <em>shorter wavelengths correspond to higher energy.\u00a0 <\/em>\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\/05153552\/image001.png\" alt=\"image002.png\" width=\"391\" height=\"120\" \/>\r\n\r\nHigh-energy radiation (such as gamma- and x-rays) is composed of very short waves \u2013 as short as 10<sup>-16<\/sup> meter from crest to crest.\u00a0 Longer waves are far less energetic, and thus are less dangerous to living things.\u00a0 Visible light waves are in the range of 400 \u2013 700 nm (nanometers, or 10<sup>-9<\/sup> m), while radio waves can be several hundred meters in length.\r\n\r\nThe notion that electromagnetic radiation contains a quantifiable amount of energy can perhaps be better understood if we talk about light as a stream of <em>particles<\/em>, called <strong>photons<\/strong>, rather than as a wave. (Recall the concept known as \u2018wave-particle duality\u2019:\u00a0 at the quantum level, wave behavior and particle behavior become indistinguishable, and very small particles have an observable \u2018wavelength\u2019). If we describe light as a stream of photons, the energy of a particular wavelength can be expressed as:\r\n\r\nE =\u00a0 <em>hc<\/em>\/<span><strong>\u03bb<\/strong><\/span>\r\n\r\nwhere E is energy in kcal\/mol,\u00a0 <span><strong>\u03bb<\/strong><\/span> (the Greek letter <em>lambda<\/em>) is wavelength in meters, <em>c<\/em> is 3.00 x 10<sup>8<\/sup> m\/s (the speed of light), and <em>h<\/em> is 9.537 x 10<sup>-14<\/sup> kcal<span>\u2022<\/span>s<span>\u2022<\/span>mol<sup>-1<\/sup>, a number known as Planck\u2019s constant.\r\n\r\nBecause electromagnetic radiation travels at a constant speed, each wavelength corresponds to a given frequency, which is the number of times per second that a crest passes a given point. Longer waves have lower frequencies, and shorter waves have higher frequencies.\u00a0 Frequency is commonly reported in hertz (Hz),\u00a0 meaning \u2018cycles per second\u2019, or \u2018waves per second\u2019. The standard unit for frequency is s<sup>-1<\/sup>.\r\n\r\nWhen talking about electromagnetic waves, we can refer either to wavelength or to frequency - the two values are interconverted using the simple expression:\r\n\r\n<span><strong>\u03bb<\/strong><\/span><span><strong>\u03bd<\/strong><\/span> = <em>c<\/em>\r\n\r\nwhere <span><strong>\u03bd <\/strong><\/span>(the Greek letter \u2018<em>nu\u2019<\/em>) is frequency in s<sup>-1<\/sup>.\u00a0 Visible red light with a wavelength of 700 nm, for example, has a frequency of 4.29 x 10<sup>14<\/sup> Hz, and an energy of 40.9 kcal per mole of photons.\r\n\r\nThe full range of electromagnetic radiation wavelengths is referred to as the <strong>electromagnetic spectrum<\/strong>.\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\/05153554\/image003.png\" alt=\"image004.png\" width=\"749\" height=\"315\" \/>\r\n\r\nNotice in the figure above that visible light takes up just a narrow band of the full spectrum.\u00a0 White light from the sun or a light bulb is a mixture of all of the visible wavelengths.\u00a0 You see the visible region of the electromagnetic spectrum divided into its different wavelengths every time you see a rainbow: violet light has the shortest wavelength, and red light has the longest.\r\n<div>\r\n<div id=\"example\">\r\n<div class=\"textbox examples\">\r\n<h3>Example<\/h3>\r\n<div>\r\n<div id=\"example\">\r\n\r\nVisible light has a wavelength range of about 400-700 nm.\u00a0 What is the corresponding frequency range?\u00a0 What is the corresponding energy range, in kcal\/mol of photons?\r\n[reveal-answer q=\"218386\"]Show Answer[\/reveal-answer]\r\n[hidden-answer a=\"218386\"]\r\n\r\nsing\u00a0\u00a0<span><span>\u03bb<\/span><\/span><span><span>\u03bd<\/span><\/span>\u00a0=\u00a0<em>c<\/em>, we first rearrange to\u00a0<span><span>\u03bd<\/span><\/span>=\u00a0<em>c<\/em>\/<span><span>\u03bb<\/span><\/span>\u00a0 to solve for frequency.\r\n\r\nFor light with a wavelength of 400 nm, the frequency is 7.50 x 10<sup>14<\/sup>\u00a0Hz:\r\n\r\n<img class=\"internal default\" src=\"https:\/\/chem.libretexts.org\/@api\/deki\/files\/6448\/image241.png?revision=1\" alt=\"image242.png\" width=\"304\" height=\"55\" \/>\r\n\r\nIn the same way, we calculate that light with a wavelength of 700 nm has a frequency of 4.29 x 10<sup>14<\/sup>\u00a0Hz.\r\n\r\nTo calculate corresponding energies:\r\n\r\nUsing\u00a0<em>hc<\/em>\/<span><span>\u03bb<\/span><\/span>, we find for light at 400 nm:\r\n\r\n<img class=\"internal default\" src=\"https:\/\/chem.libretexts.org\/@api\/deki\/files\/6450\/image243.png?revision=1\" alt=\"image244.png\" width=\"421\" height=\"108\" \/>\r\n\r\nUsing the same equation, we find that light at 700 nm corresponds to 40.9 kcal\/mol.\r\n\r\n[\/hidden-answer]\r\n\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<p class=\"boxtitle\">Covalent bonds in organic molecules are not rigid sticks \u2013 rather, they behave more like springs.\u00a0 At room temperature, organic molecules are always in motion, as their bonds stretch, bend, and twist.\u00a0 These complex vibrations can be broken down mathematically into individual <strong>vibrational modes<\/strong>, a few of which are illustrated below.<\/p>\r\n\r\n<\/div>\r\n<\/div>\r\n<img class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05153558\/image007.png\" alt=\"image008.png\" width=\"457\" height=\"380\" \/>\r\n\r\nThe energy of molecular vibration is <em>quantized<\/em> rather than continuous, meaning that a molecule can only stretch and bend at certain 'allowed' frequencies.\u00a0 If a molecule is exposed to electromagnetic radiation that matches the frequency of one of its vibrational modes,\u00a0 it will in most cases absorb energy from the radiation and jump to a higher vibrational energy state - what this means is that the <em>amplitude<\/em> of the vibration will increase, but the vibrational <em>frequency<\/em> will remain the same.\u00a0 The difference in energy between the two vibrational states is equal to the energy associated with the wavelength of radiation that was absorbed.\u00a0 It turns out that it is the <em>infrared<\/em> region of the electromagnetic spectrum which contains frequencies corresponding to the vibrational frequencies of organic bonds.\r\n\r\nLet's take 2-hexanone as an example.\u00a0 Picture the carbonyl bond of the ketone group as a spring.\u00a0 This spring is constantly bouncing back and forth, stretching and compressing, pushing the carbon and oxygen atoms further apart and then pulling them together.\u00a0 This is the <strong>stretching mode<\/strong> of the carbonyl bond.\u00a0 In the space of one second, the spring 'bounces' back and forth 5.15 x 10<sup>13<\/sup> times - in other words,\u00a0 the ground-state frequency of carbonyl\u00a0 stretching for a the ketone group is about 5.15 x 10<sup>13<\/sup> Hz.\r\n\r\nIf 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<sup>-6<\/sup> m and an energy of 4.91 kcal\/mol.\u00a0 When the carbonyl bond absorbs this energy, it jumps up to an excited vibrational state.\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\/05153601\/image009.png\" alt=\"image010.png\" width=\"680\" height=\"310\" \/>\r\n\r\nThe value of <span><strong>\u0394<\/strong><\/span>E - the energy difference between the low energy (ground)\u00a0 and high energy (excited) vibrational states - is equal to 4.91 kcal\/mol, the same as the energy associated with the absorbed light frequency.\u00a0 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.\r\n\r\nWith an instrument called an infrared spectrophotometer, we can 'see' this vibrational transition.\u00a0 In the spectrophotometer, infrared light with frequencies ranging from about 10<sup>13<\/sup> to 10<sup>14<\/sup> Hz\u00a0 is passed though our sample of cyclohexane.\u00a0 Most frequencies pass right through the sample and are recorded by a detector on the other side.\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\/05153605\/image011.png\" alt=\"image012.png\" width=\"583\" height=\"197\" \/>\r\n\r\nOur 5.15 x 10<sup>13<\/sup> 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.\r\n\r\nThe vibrations of a 2-hexanone molecule are not, of course, limited to the simple stretching of the carbonyl bond.\u00a0 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.\r\n\r\nThe power of infrared spectroscopy arises from the observation that <em>different functional groups have different characteristic absorption frequencies<\/em>.\u00a0\u00a0 The carbonyl bond in a ketone, as we saw with our 2-hexanone example, typically absorbs in the range of\u00a0 5.11 -\u00a0 5.18 x 10<sup>13<\/sup> Hz, depending on the molecule.\u00a0\u00a0 The carbon-carbon triple bond of an alkyne, on the other hand, absorbs in the range\u00a0 6.30 - 6.80 x 10<sup>13<\/sup> Hz.\u00a0\u00a0 The technique is therefore very useful as a means of identifying which functional groups are present in a molecule of interest.\u00a0 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.\r\n\r\nSome bonds absorb infrared light more strongly than others, and some bonds do not absorb at all. <em>In order for a vibrational mode to absorb infrared light, it must result in a periodic change in the dipole moment of the molecule<\/em>.\u00a0 Such vibrations are said to be <strong>infrared active<\/strong>. In general, the greater the polarity of the bond, the stronger its IR absorption.\u00a0 The carbonyl bond is very polar, and absorbs very strongly.\u00a0 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.\r\n\r\nSome kinds of vibrations are <strong>infrared inactive<\/strong>.\u00a0 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 <em>do<\/em> absorb IR light).\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\/05153608\/image013.png\" alt=\"image014.png\" width=\"320\" height=\"112\" \/>\r\n\r\nNow, let's\u00a0 look at some actual output from IR spectroscopy experiments.\u00a0 Below is the IR spectrum for 2-hexanone.\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\/05153611\/image015.png\" alt=\"image016.png\" width=\"750\" height=\"445\" \/>\r\n\r\nThere are a number of things that need to be explained in order for you to understand what it is that we are looking at.\u00a0 On the horizontal axis we see IR wavelengths expressed in terms of a unit called <strong>wavenumber<\/strong> (cm<sup>-1<\/sup>), which tells us how many waves fit into one centimeter.\u00a0\u00a0 On the vertical axis we see \u2018<strong>% transmittance<\/strong>\u2019, which tells us how strongly light was absorbed at each frequency (100% transmittance means no absorption occurred at that frequency).\u00a0 The solid line traces the values of % transmittance for every wavelength \u2013 the \u2018peaks\u2019 (which are actually pointing down) show regions of strong absorption.\u00a0 For some reason, it is typical in IR spectroscopy to report wavenumber values rather than wavelength (in meters) or frequency (in Hz).\u00a0 The \u2018upside down\u2019 vertical axis, with absorbance peaks pointing down rather than up, is also a curious convention in IR spectroscopy.\u00a0 We wouldn\u2019t\u00a0 want to make things too easy for you!\r\n\r\n<\/div>\r\n<div id=\"section_2\">\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>William Reusch, Professor Emeritus (<a class=\"external\" title=\"http:\/\/www.msu.edu\/\" href=\"http:\/\/www.msu.edu\/\" target=\"_blank\" rel=\"external nofollow noopener\">Michigan State U.<\/a>), <a class=\"external\" title=\"http:\/\/www.cem.msu.edu\/~reusch\/VirtualText\/intro1.htm\" href=\"http:\/\/www.cem.msu.edu\/%7Ereusch\/VirtualText\/intro1.htm\" target=\"_blank\" rel=\"external nofollow noopener\">Virtual Textbook of\u00a0Organic\u00a0Chemistry<\/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>\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 class=\"elm-header\"><\/div>\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 (by wavelength, wavenumber, or both) the region of the electromagnetic spectrum which is used in infrared (IR) spectroscopy.<\/li>\n<li>interconvert between wavelength and wavenumber.<\/li>\n<li>discuss, in general terms, the effect that the absorption of infrared radiation can have on a molecule.<\/li>\n<\/ol>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<div class=\"textbox key-takeaways\">\n<h3>Key Terms<\/h3>\n<div id=\"elm-main-content\" class=\"elm-content-container\">\n<div>\n<div>\n<p>Make certain that you can define, and use in context, the key terms below.<\/p>\n<ul>\n<li>infrared spectrum<\/li>\n<li>wavenumber (reciprocal centimetres)<\/li>\n<\/ul>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<div class=\"textbox\">\n<div class=\"elm-header\">\n<div class=\"elm-header-custom\">\n<h3>Study Notes<\/h3>\n<\/div>\n<\/div>\n<div id=\"elm-main-content\" class=\"elm-content-container\">\n<div>\n<div id=\"note\">\n<p>Notice that the scale at the bottom of the infrared spectrum for 2-hexanone shown is calibrated in wavenumbers (cm<sup>\u22121<\/sup>). A wavenumber is the reciprocal of a wavelength (1\/\u03bb); thus, a wavenumber of 1600 cm<sup>\u22121<\/sup> corresponds to a wavelength of<\/p>\n<p>1 1600\u2009 cm \u22121 =6.25\u00d7 10 \u22124 cm\u00a0or\u00a06.25\u2009\u03bc\u2009m<\/p>\n<p>Organic chemists find it more convenient to deal with wavenumbers rather than wavelengths when discussing infrared spectra.<\/p>\n<p>You will obtain infrared spectra for a number of the compounds you will synthesize in the laboratory component of this course.<\/p>\n<p>The inverted peaks observed in the spectra correspond to molecular stretching and bending vibrations that only occur at certain quantized frequencies. When infrared radiation matching these frequencies falls on the molecule, the molecule absorbs energy and becomes excited. Eventually the molecule returns to its original (ground) state, and the energy which was absorbed is released as heat.<\/p>\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 id=\"section_1\">\n<h3 class=\"editable\">The electromagnetic spectrum<\/h3>\n<p>Electromagnetic radiation, as you may recall from a previous chemistry or physics class, is composed of electrical and magnetic waves which oscillate on perpendicular planes. Visible light is electromagnetic radiation. So are the gamma rays that are emitted by spent nuclear fuel, the x-rays that a doctor uses to visualize your bones, the ultraviolet light that causes a painful sunburn when you forget to apply sun block, the infrared light that the army uses in night-vision goggles, the microwaves that you use to heat up your frozen burritos, and the radio-frequency waves that bring music to anybody who is old-fashioned enough to still listen to FM or AM radio.<\/p>\n<p>Just like ocean waves, electromagnetic waves travel in a defined direction. While the speed of ocean waves can vary, however, the speed of electromagnetic waves \u2013 commonly referred to as the speed of light \u2013 is essentially a constant, approximately 300 million meters per second. This is true whether we are talking about gamma radiation or visible light. Obviously, there is a big difference between these two types of waves \u2013 we are surrounded by the latter for more than half of our time on earth, whereas we hopefully never become exposed to the former to any significant degree.\u00a0 The different properties of the various types of electromagnetic radiation are due to differences in their wavelengths, and the corresponding differences in their energies: <em>shorter wavelengths correspond to higher energy.\u00a0 <\/em><\/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\/05153552\/image001.png\" alt=\"image002.png\" width=\"391\" height=\"120\" \/><\/p>\n<p>High-energy radiation (such as gamma- and x-rays) is composed of very short waves \u2013 as short as 10<sup>-16<\/sup> meter from crest to crest.\u00a0 Longer waves are far less energetic, and thus are less dangerous to living things.\u00a0 Visible light waves are in the range of 400 \u2013 700 nm (nanometers, or 10<sup>-9<\/sup> m), while radio waves can be several hundred meters in length.<\/p>\n<p>The notion that electromagnetic radiation contains a quantifiable amount of energy can perhaps be better understood if we talk about light as a stream of <em>particles<\/em>, called <strong>photons<\/strong>, rather than as a wave. (Recall the concept known as \u2018wave-particle duality\u2019:\u00a0 at the quantum level, wave behavior and particle behavior become indistinguishable, and very small particles have an observable \u2018wavelength\u2019). If we describe light as a stream of photons, the energy of a particular wavelength can be expressed as:<\/p>\n<p>E =\u00a0 <em>hc<\/em>\/<span><strong>\u03bb<\/strong><\/span><\/p>\n<p>where E is energy in kcal\/mol,\u00a0 <span><strong>\u03bb<\/strong><\/span> (the Greek letter <em>lambda<\/em>) is wavelength in meters, <em>c<\/em> is 3.00 x 10<sup>8<\/sup> m\/s (the speed of light), and <em>h<\/em> is 9.537 x 10<sup>-14<\/sup> kcal<span>\u2022<\/span>s<span>\u2022<\/span>mol<sup>-1<\/sup>, a number known as Planck\u2019s constant.<\/p>\n<p>Because electromagnetic radiation travels at a constant speed, each wavelength corresponds to a given frequency, which is the number of times per second that a crest passes a given point. Longer waves have lower frequencies, and shorter waves have higher frequencies.\u00a0 Frequency is commonly reported in hertz (Hz),\u00a0 meaning \u2018cycles per second\u2019, or \u2018waves per second\u2019. The standard unit for frequency is s<sup>-1<\/sup>.<\/p>\n<p>When talking about electromagnetic waves, we can refer either to wavelength or to frequency &#8211; the two values are interconverted using the simple expression:<\/p>\n<p><span><strong>\u03bb<\/strong><\/span><span><strong>\u03bd<\/strong><\/span> = <em>c<\/em><\/p>\n<p>where <span><strong>\u03bd <\/strong><\/span>(the Greek letter \u2018<em>nu\u2019<\/em>) is frequency in s<sup>-1<\/sup>.\u00a0 Visible red light with a wavelength of 700 nm, for example, has a frequency of 4.29 x 10<sup>14<\/sup> Hz, and an energy of 40.9 kcal per mole of photons.<\/p>\n<p>The full range of electromagnetic radiation wavelengths is referred to as the <strong>electromagnetic spectrum<\/strong>.<\/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\/05153554\/image003.png\" alt=\"image004.png\" width=\"749\" height=\"315\" \/><\/p>\n<p>Notice in the figure above that visible light takes up just a narrow band of the full spectrum.\u00a0 White light from the sun or a light bulb is a mixture of all of the visible wavelengths.\u00a0 You see the visible region of the electromagnetic spectrum divided into its different wavelengths every time you see a rainbow: violet light has the shortest wavelength, and red light has the longest.<\/p>\n<div>\n<div id=\"example\">\n<div class=\"textbox examples\">\n<h3>Example<\/h3>\n<div>\n<div id=\"example\">\n<p>Visible light has a wavelength range of about 400-700 nm.\u00a0 What is the corresponding frequency range?\u00a0 What is the corresponding energy range, in kcal\/mol of photons?<\/p>\n<div class=\"qa-wrapper\" style=\"display: block\"><span class=\"show-answer collapsed\" style=\"cursor: pointer\" data-target=\"q218386\">Show Answer<\/span><\/p>\n<div id=\"q218386\" class=\"hidden-answer\" style=\"display: none\">\n<p>sing\u00a0\u00a0<span><span>\u03bb<\/span><\/span><span><span>\u03bd<\/span><\/span>\u00a0=\u00a0<em>c<\/em>, we first rearrange to\u00a0<span><span>\u03bd<\/span><\/span>=\u00a0<em>c<\/em>\/<span><span>\u03bb<\/span><\/span>\u00a0 to solve for frequency.<\/p>\n<p>For light with a wavelength of 400 nm, the frequency is 7.50 x 10<sup>14<\/sup>\u00a0Hz:<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" class=\"internal default\" src=\"https:\/\/chem.libretexts.org\/@api\/deki\/files\/6448\/image241.png?revision=1\" alt=\"image242.png\" width=\"304\" height=\"55\" \/><\/p>\n<p>In the same way, we calculate that light with a wavelength of 700 nm has a frequency of 4.29 x 10<sup>14<\/sup>\u00a0Hz.<\/p>\n<p>To calculate corresponding energies:<\/p>\n<p>Using\u00a0<em>hc<\/em>\/<span><span>\u03bb<\/span><\/span>, we find for light at 400 nm:<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" class=\"internal default\" src=\"https:\/\/chem.libretexts.org\/@api\/deki\/files\/6450\/image243.png?revision=1\" alt=\"image244.png\" width=\"421\" height=\"108\" \/><\/p>\n<p>Using the same equation, we find that light at 700 nm corresponds to 40.9 kcal\/mol.<\/p>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<p class=\"boxtitle\">Covalent bonds in organic molecules are not rigid sticks \u2013 rather, they behave more like springs.\u00a0 At room temperature, organic molecules are always in motion, as their bonds stretch, bend, and twist.\u00a0 These complex vibrations can be broken down mathematically into individual <strong>vibrational modes<\/strong>, a few of which are illustrated below.<\/p>\n<\/div>\n<\/div>\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\/05153558\/image007.png\" alt=\"image008.png\" width=\"457\" height=\"380\" \/><\/p>\n<p>The energy of molecular vibration is <em>quantized<\/em> rather than continuous, meaning that a molecule can only stretch and bend at certain &#8216;allowed&#8217; frequencies.\u00a0 If a molecule is exposed to electromagnetic radiation that matches the frequency of one of its vibrational modes,\u00a0 it will in most cases absorb energy from the radiation and jump to a higher vibrational energy state &#8211; what this means is that the <em>amplitude<\/em> of the vibration will increase, but the vibrational <em>frequency<\/em> will remain the same.\u00a0 The difference in energy between the two vibrational states is equal to the energy associated with the wavelength of radiation that was absorbed.\u00a0 It turns out that it is the <em>infrared<\/em> region of the electromagnetic spectrum which contains frequencies corresponding to the vibrational frequencies of organic bonds.<\/p>\n<p>Let&#8217;s take 2-hexanone as an example.\u00a0 Picture the carbonyl bond of the ketone group as a spring.\u00a0 This spring is constantly bouncing back and forth, stretching and compressing, pushing the carbon and oxygen atoms further apart and then pulling them together.\u00a0 This is the <strong>stretching mode<\/strong> of the carbonyl bond.\u00a0 In the space of one second, the spring &#8216;bounces&#8217; back and forth 5.15 x 10<sup>13<\/sup> times &#8211; in other words,\u00a0 the ground-state frequency of carbonyl\u00a0 stretching for a the ketone group is about 5.15 x 10<sup>13<\/sup> Hz.<\/p>\n<p>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<sup>-6<\/sup> m and an energy of 4.91 kcal\/mol.\u00a0 When the carbonyl bond absorbs this energy, it jumps up to an excited vibrational state.<\/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\/05153601\/image009.png\" alt=\"image010.png\" width=\"680\" height=\"310\" \/><\/p>\n<p>The value of <span><strong>\u0394<\/strong><\/span>E &#8211; the energy difference between the low energy (ground)\u00a0 and high energy (excited) vibrational states &#8211; is equal to 4.91 kcal\/mol, the same as the energy associated with the absorbed light frequency.\u00a0 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.<\/p>\n<p>With an instrument called an infrared spectrophotometer, we can &#8216;see&#8217; this vibrational transition.\u00a0 In the spectrophotometer, infrared light with frequencies ranging from about 10<sup>13<\/sup> to 10<sup>14<\/sup> Hz\u00a0 is passed though our sample of cyclohexane.\u00a0 Most frequencies pass right through the sample and are recorded by a detector on the other side.<\/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\/05153605\/image011.png\" alt=\"image012.png\" width=\"583\" height=\"197\" \/><\/p>\n<p>Our 5.15 x 10<sup>13<\/sup> 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.<\/p>\n<p>The vibrations of a 2-hexanone molecule are not, of course, limited to the simple stretching of the carbonyl bond.\u00a0 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.<\/p>\n<p>The power of infrared spectroscopy arises from the observation that <em>different functional groups have different characteristic absorption frequencies<\/em>.\u00a0\u00a0 The carbonyl bond in a ketone, as we saw with our 2-hexanone example, typically absorbs in the range of\u00a0 5.11 &#8211;\u00a0 5.18 x 10<sup>13<\/sup> Hz, depending on the molecule.\u00a0\u00a0 The carbon-carbon triple bond of an alkyne, on the other hand, absorbs in the range\u00a0 6.30 &#8211; 6.80 x 10<sup>13<\/sup> Hz.\u00a0\u00a0 The technique is therefore very useful as a means of identifying which functional groups are present in a molecule of interest.\u00a0 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.<\/p>\n<p>Some bonds absorb infrared light more strongly than others, and some bonds do not absorb at all. <em>In order for a vibrational mode to absorb infrared light, it must result in a periodic change in the dipole moment of the molecule<\/em>.\u00a0 Such vibrations are said to be <strong>infrared active<\/strong>. In general, the greater the polarity of the bond, the stronger its IR absorption.\u00a0 The carbonyl bond is very polar, and absorbs very strongly.\u00a0 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.<\/p>\n<p>Some kinds of vibrations are <strong>infrared inactive<\/strong>.\u00a0 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 <em>do<\/em> absorb IR light).<\/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\/05153608\/image013.png\" alt=\"image014.png\" width=\"320\" height=\"112\" \/><\/p>\n<p>Now, let&#8217;s\u00a0 look at some actual output from IR spectroscopy experiments.\u00a0 Below is the IR spectrum for 2-hexanone.<\/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\/05153611\/image015.png\" alt=\"image016.png\" width=\"750\" height=\"445\" \/><\/p>\n<p>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.\u00a0 On the horizontal axis we see IR wavelengths expressed in terms of a unit called <strong>wavenumber<\/strong> (cm<sup>-1<\/sup>), which tells us how many waves fit into one centimeter.\u00a0\u00a0 On the vertical axis we see \u2018<strong>% transmittance<\/strong>\u2019, which tells us how strongly light was absorbed at each frequency (100% transmittance means no absorption occurred at that frequency).\u00a0 The solid line traces the values of % transmittance for every wavelength \u2013 the \u2018peaks\u2019 (which are actually pointing down) show regions of strong absorption.\u00a0 For some reason, it is typical in IR spectroscopy to report wavenumber values rather than wavelength (in meters) or frequency (in Hz).\u00a0 The \u2018upside down\u2019 vertical axis, with absorbance peaks pointing down rather than up, is also a curious convention in IR spectroscopy.\u00a0 We wouldn\u2019t\u00a0 want to make things too easy for you!<\/p>\n<\/div>\n<div id=\"section_2\">\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>William Reusch, Professor Emeritus (<a class=\"external\" title=\"http:\/\/www.msu.edu\/\" href=\"http:\/\/www.msu.edu\/\" target=\"_blank\" rel=\"external nofollow noopener\">Michigan State U.<\/a>), <a class=\"external\" title=\"http:\/\/www.cem.msu.edu\/~reusch\/VirtualText\/intro1.htm\" href=\"http:\/\/www.cem.msu.edu\/%7Ereusch\/VirtualText\/intro1.htm\" target=\"_blank\" rel=\"external nofollow noopener\">Virtual Textbook of\u00a0Organic\u00a0Chemistry<\/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<\/div>\n","protected":false},"author":44985,"menu_order":2,"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-1637","chapter","type-chapter","status-publish","hentry"],"part":29,"_links":{"self":[{"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/pressbooks\/v2\/chapters\/1637","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":5,"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/pressbooks\/v2\/chapters\/1637\/revisions"}],"predecessor-version":[{"id":2140,"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/pressbooks\/v2\/chapters\/1637\/revisions\/2140"}],"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\/1637\/metadata\/"}],"wp:attachment":[{"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/wp\/v2\/media?parent=1637"}],"wp:term":[{"taxonomy":"chapter-type","embeddable":true,"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/pressbooks\/v2\/chapter-type?post=1637"},{"taxonomy":"contributor","embeddable":true,"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/wp\/v2\/contributor?post=1637"},{"taxonomy":"license","embeddable":true,"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/wp\/v2\/license?post=1637"}],"curies":[{"name":"wp","href":"https:\/\/api.w.org\/{rel}","templated":true}]}}