{"id":293,"date":"2017-10-04T15:12:49","date_gmt":"2017-10-04T15:12:49","guid":{"rendered":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/?post_type=chapter&#038;p=293"},"modified":"2017-11-29T15:20:35","modified_gmt":"2017-11-29T15:20:35","slug":"hybridization-structure-of-acetylene","status":"publish","type":"chapter","link":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/chapter\/hybridization-structure-of-acetylene\/","title":{"raw":"Hybridization: Structure of Acetylene","rendered":"Hybridization: Structure of Acetylene"},"content":{"raw":"<div class=\"elm-header\">\r\n<h2 class=\"elm-header-custom\">sp Hybrid Orbitals and the Structure of Acetylene<\/h2>\r\n<\/div>\r\n<div id=\"elm-main-content\" class=\"elm-content-container\">\r\n<div>\r\n<div id=\"skills\">\r\n<div class=\"textbox learning-objectives\">\r\n<h3>Objectives<\/h3>\r\nAfter completing this section, you should be able to\r\n<ol>\r\n \t<li>use the concept of <em>sp<\/em> hybridization to account for the formation of carbon-carbon triple bonds, and describe a carbon-carbon triple bond as consisting of one \u03c3 bond and two \u03c0 bonds.<\/li>\r\n \t<li>list the approximate bond lengths associated with typical carbon-carbon single bonds, double bonds and triple bonds. [You may need to review Sections 1.7 and 1.8.]<\/li>\r\n \t<li>list the approximate bond angles associated with <em>sp<\/em><sup>3<\/sup>-, <em>sp<\/em><sup>2<\/sup>- and sp\u2011hybridized carbon atoms, and hence, predict the bond angles to be expected in given organic compounds. [If necessary, review Sections 1.6, 1.7 and 1.8.]<\/li>\r\n \t<li>account for the differences in bond length, bond strength and bond angles found in compounds containing <em>sp<\/em><sup>3<\/sup>-, <em>sp<\/em><sup>2<\/sup>- and sp\u2011hybridized carbon atoms, such as ethane, ethylene and acetylene.<\/li>\r\n<\/ol>\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>\r\n\r\nMake certain that you can define, and use in context, the key term below.\r\n<ul>\r\n \t<li><em>sp<\/em> hybrid<\/li>\r\n<\/ul>\r\n<\/div>\r\n<div id=\"note\">\r\n<p class=\"boxtitle\">Study Notes<\/p>\r\nThe bond angles associated with <em>sp<\/em><sup>3<\/sup>-, <em>sp<\/em><sup>2<\/sup>- and sp\u2011hybridized carbon atoms are approximately 109.5, 120 and 180\u00b0, respectively.\r\n\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<div id=\"section_1\">\r\n<h3 class=\"editable\">Bonding in acetylene<\/h3>\r\nFinally, the hybrid orbital concept applies well to triple-bonded groups, such as alkynes and nitriles. Consider, for example, the structure of ethyne (common\u00a0 name acetylene), the simplest alkyne.\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\/04151225\/image217.png\" alt=\"image210.png\" width=\"152px\" height=\"78px\" \/>\r\n\r\nThis molecule is linear: all four atoms lie in a straight line. The carbon-carbon triple bond is only 1.20\u00c5 long. In the hybrid orbital picture of acetylene, both carbons are <strong>sp-hybridized<\/strong>. In an sp-hybridized carbon, the 2<em>s<\/em> orbital combines with the 2<em>p<\/em><sub>x<\/sub> orbital to form two sp hybrid orbitals that are oriented at an angle of 180\u00b0with respect to each other (eg. along the x axis). The 2<em>p<\/em><sub>y<\/sub> and 2<em>p<\/em><sub>z<\/sub> orbitals remain unhybridized, and are oriented perpendicularly along the y and z axes, respectively.\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\/04151227\/image219.png\" alt=\"image212.png\" width=\"600px\" height=\"130px\" \/>\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\/04151229\/image221.png\" alt=\"image214.png\" width=\"238px\" height=\"162px\" \/>\r\n\r\nThe C-C sigma bond, then, is formed by the overlap of one sp orbital from each of the carbons, while the two C-H sigma bonds are formed by the overlap of the second sp orbital on each carbon with a 1<em>s<\/em> orbital on a hydrogen.\u00a0 Each carbon atom still has two half-filled 2<em>p<\/em><sub>y<\/sub> and 2<em>p<\/em><sub>z<\/sub> orbitals, which are perpendicular both to each other and to the line formed by the sigma bonds.\u00a0 These two perpendicular pairs of <em>p<\/em> orbitals form two pi bonds between the carbons, resulting in a triple bond overall (one sigma bond plus two pi bonds).\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\/04151231\/image223.png\" alt=\"image216.png\" width=\"447px\" height=\"123px\" \/>\r\n\r\nThe hybrid orbital concept nicely explains another experimental observation: single bonds adjacent to double and triple bonds are progressively shorter and stronger than \u2018normal\u2019 single bonds, such as the one in a simple alkane.\u00a0 The carbon-carbon bond in ethane (structure A below) results from the overlap of two sp<sup>3<\/sup> orbitals.\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\/04151233\/image225.png\" alt=\"image218.png\" width=\"597px\" height=\"182px\" \/>\r\n\r\nIn alkene B, however, the carbon-carbon single bond is the result of overlap between an sp<sup>2<\/sup> orbital and an sp<sup>3<\/sup> orbital, while in alkyne C the carbon-carbon single bond is the result of overlap between an sp orbital and an sp<sup>3<\/sup> orbital.\u00a0 These are all single bonds, but the bond in molecule C is shorter and stronger than the one in B, which is in turn shorter and stronger than the one in A.\r\n\r\nThe explanation here is relatively straightforward.\u00a0 An sp orbital is composed of one <em>s<\/em> orbital and one <em>p<\/em> orbital, and thus it has 50%\u00a0 <em>s<\/em> character and 50% <em>p<\/em> character.\u00a0 sp<sup>2<\/sup> orbitals, by comparison, have 33% <em>s<\/em> character and 67% <em>p<\/em> character, while sp<sup>3<\/sup> orbitals have 25% <em>s<\/em> character and 75% <em>p<\/em> character.\u00a0 Because of their spherical shape, 2<em>s<\/em> orbitals are smaller, and hold electrons closer and \u2018tighter\u2019 to the nucleus, compared to 2<em>p<\/em> orbitals.\u00a0 Consequently, bonds involving sp + sp<sup>3<\/sup> overlap (as in alkyne C)\u00a0 are\u00a0 shorter and stronger than bonds involving sp<sup>2<\/sup> + sp<sup>3<\/sup> overlap (as in alkene B).\u00a0 Bonds involving sp<sup>3<\/sup>-sp<sup>3<\/sup>overlap (as in alkane A)\u00a0 are the longest and weakest of the group, because of the 75% \u2018<em>p<\/em>\u2019 character of the hybrids.\r\n\r\n<\/div>\r\n<div id=\"section_2\">\r\n<h3 class=\"editable\">Comparison of C-C bonds Ethane, Ethylene, and Acetylene<\/h3>\r\n<table style=\"border-spacing: 1px\" border=\"1\">\r\n<tbody>\r\n<tr>\r\n<td><strong>Molecule<\/strong><\/td>\r\n<td><strong>Bond<\/strong><\/td>\r\n<td><strong>Bond Strength (kJ\/mol)<\/strong><\/td>\r\n<td><strong>Bond Length (pm)<\/strong><\/td>\r\n<\/tr>\r\n<tr>\r\n<td><strong>Ethane, CH<sub>3<\/sub>CH<sub>3<\/sub><\/strong><\/td>\r\n<td><strong>(<em>sp<sup>3<\/sup><\/em>) C-C (<em>sp<sup>3<\/sup><\/em>)<\/strong><\/td>\r\n<td><strong>376<\/strong><\/td>\r\n<td><strong>154<\/strong><\/td>\r\n<\/tr>\r\n<tr>\r\n<td><strong>Ethylene, H<sub>2<\/sub>C=CH2<\/strong><\/td>\r\n<td><strong>(<em>sp<sup>2<\/sup><\/em>) C=C (<em>sp<sup>2<\/sup><\/em>)<\/strong><\/td>\r\n<td><strong>728<\/strong><\/td>\r\n<td><strong>134<\/strong><\/td>\r\n<\/tr>\r\n<tr>\r\n<td>Acetylene, <img src=\"https:\/\/chem.libretexts.org\/LibreTexts\/Athabasca_University\/Chemistry_350%3A_Organic_Chemistry_I\/Chapter_1%3A_Structure_and_Bonding\/denied:data:image\/png;base64,iVBORw0KGgoAAAANSUhEUgAAAW4AAABQCAIAAAC2xwixAAAEfUlEQVR4nO3dMU7zQBCG4dwhN8gBuAQ99V9xDiqugbgDV6DhAPRIdEi0SKlI5b+IFCEnGY+9Y++34\/c5QLLZnflsr43ZdABQbFN7AAAyIEoABCBKAAQgSgAEIEoABCBKAAQgSgAEIEoABCBKAAQgSgAEIEoABCBKAAQgSgAEIEpmtIGM2rWQH1M8l9q9g77aFZHc6ub36fHeqLbnl9eoL1qsQ+AUtbJHixVSKybO78f727VJvP338Hs4jP3A76\/Pm+324gfudnc\/+\/20cQ5+eOyv+GvU12EBJau5fCG11WJd+igxxulREigl34s5TFvHWoXUSoudpI2S\/f7nbrcbu+QXlZyshgwAhSYvX8VC0m+xnpxRUngMOXf\/+DT2FyGBioUk3mLnEkZJ+PKX\/C60q24hKbfYRdmi5PXleezS+nFush7VC0m2xa5JFSUzHUb+env\/GD9baIxCIWm2mCFPlPi3xy6uovMoxGVOeiKFJNhitjxRYj8ytNlsttubz6\/vyb\/rZIVPH62KSCEJtpgtSZQMPjvk3+YYPChxYpKYTiGptdigJFFiH0nGbpcO1hM7JlnpFJJaiw3KECX2gk2bJvuKl1s5KUkVklSLeWSIEnu1pp1BXDs7ZaMkMalCkmoxj+aj5HD4ffh3GzuYo2NhsTOyEmqFpNNiTs1Hib25xUkEnNQKSafFnJqPEvuuG\/ujcFIrJJ0Wc2o+Sozr25AJwkqoFZJOizk1HyXG3Tu2OeCnVkg6LeYUHyXhjN9pb5VxyxZOgoUk0mJ+RMm8FpslGAaXSbCQRFrMr+0osXfd60bJYvMDJ2OxBAtJpMX8iJK5LDY\/cDIWS7CQRFrMjyiZy2LzAydjsQQLSaTF\/IiSuSw2P3AyFkuwkERazK\/tKBHcLftrsSnCIHulBAtJpMX8MkcJz5XASbCQRFrMj0fUgK7TKySdFnNqPkrUnndGo9QKSafFnJqPkvn+Cus4JP62eCXUCkmnxZyajxL7zVclG2a9wxSZkptaIem0mFPzUWJvmE2eI+NjeXFBSmqFpNNiTs1HSTf0at9pnT\/fMQqypApJqsU8MkRJ+Nt97QPUhhOTpKQKSarFPDJEyeCCjT2JsF8XzI2hrKQKSarFPDJEiT2esaMa\/EeNXN0kplNIai02KEmUdBH\/n3HwoOT5ELROpJAEW8yWJ0oG\/xPayfkFqmftj7glnJ5IIQm2mC1PlNijCsEuyUooFJJmixlSRUnnuECdjEubValeSLItdk22KOlmKwJuAK9N3UJSbrGLEkaJPbxpyJF1qlhI4i12LmeU2B84SuF1TfkAUG7y8lUsJP0W60kbJUeF56gl92tKvhdzmLyUVQqplRY7SR4lg6MNXPuesd+IuRUu6MKF1FaLdSuJkhP7bcCx93qvVxeqiVrcBQqpuRYLm1z0BBQ+QtWuiOSY37nUbhz01a6I5JhfAAGIEgABiBIAAYgSAAGIEgABiBIAAYgSAAGIEgABiBIAAYgSAAGIEgABiBIAAYgSAAGIEgABiBIAAYgSAAH+A5v5siz9fpA7AAAAAElFTkSuQmCC#fixme\" alt=\"\" width=\"78\" height=\"17\" \/><\/td>\r\n<td><strong>(<em>sp<\/em>) <\/strong><img src=\"https:\/\/chem.libretexts.org\/LibreTexts\/Athabasca_University\/Chemistry_350%3A_Organic_Chemistry_I\/Chapter_1%3A_Structure_and_Bonding\/denied:data:image\/png;base64,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#fixme\" alt=\"\" width=\"92\" height=\"27\" \/><strong> (<em>sp<\/em>)<\/strong><\/td>\r\n<td><strong>965<\/strong><\/td>\r\n<td><strong>120<\/strong><\/td>\r\n<\/tr>\r\n<\/tbody>\r\n<\/table>\r\nNotice that as the bond order increases the bond length decreases and the bond strength increases.\r\n\r\n<\/div>\r\n<div id=\"section_3\">\r\n<div id=\"s61688\">\r\n<div id=\"section_30\">\r\n<div class=\"textbox examples\">\r\n<h3>Example<\/h3>\r\n<div id=\"section_3\">\r\n<div id=\"s61688\">\r\n<div id=\"section_30\">\r\n\r\n<span>1-Cyclohexyne is a very strained molecule. By looking at the molecule explain why there is such a intermolecular strain using the knowledge of hybridization and bond angles.<\/span>\r\n<p><img class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/04151236\/1.9.png\" alt=\"\" width=\"94\" height=\"100\" \/><\/p>\r\n\r\n<\/div>\r\n<div id=\"section_31\">\r\n\r\n[reveal-answer q=\"47173\"]Show Answer[\/reveal-answer]\r\n[hidden-answer a=\"47173\"]\r\n\r\n<span><span>The alkyne is a sp hybridized orbital. By looking at a sp orbital, we can see that the bond angle is 180\u00b0, but in cyclohexane the regular angles would be 109.5\u00b0. Therefore the molecule would be strained to force the 180\u00b0 to be a 109\u00b0.\u00a0 \u00a0[\/hidden-answer]<\/span><\/span>\r\n\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<div id=\"section_4\">\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<h2 class=\"elm-header-custom\">sp Hybrid Orbitals and the Structure of Acetylene<\/h2>\n<\/div>\n<div id=\"elm-main-content\" class=\"elm-content-container\">\n<div>\n<div id=\"skills\">\n<div class=\"textbox learning-objectives\">\n<h3>Objectives<\/h3>\n<p>After completing this section, you should be able to<\/p>\n<ol>\n<li>use the concept of <em>sp<\/em> hybridization to account for the formation of carbon-carbon triple bonds, and describe a carbon-carbon triple bond as consisting of one \u03c3 bond and two \u03c0 bonds.<\/li>\n<li>list the approximate bond lengths associated with typical carbon-carbon single bonds, double bonds and triple bonds. [You may need to review Sections 1.7 and 1.8.]<\/li>\n<li>list the approximate bond angles associated with <em>sp<\/em><sup>3<\/sup>-, <em>sp<\/em><sup>2<\/sup>&#8211; and sp\u2011hybridized carbon atoms, and hence, predict the bond angles to be expected in given organic compounds. [If necessary, review Sections 1.6, 1.7 and 1.8.]<\/li>\n<li>account for the differences in bond length, bond strength and bond angles found in compounds containing <em>sp<\/em><sup>3<\/sup>-, <em>sp<\/em><sup>2<\/sup>&#8211; and sp\u2011hybridized carbon atoms, such as ethane, ethylene and acetylene.<\/li>\n<\/ol>\n<\/div>\n<\/div>\n<div>\n<div class=\"textbox key-takeaways\">\n<h3>Key Terms<\/h3>\n<div>\n<p>Make certain that you can define, and use in context, the key term below.<\/p>\n<ul>\n<li><em>sp<\/em> hybrid<\/li>\n<\/ul>\n<\/div>\n<div id=\"note\">\n<p class=\"boxtitle\">Study Notes<\/p>\n<p>The bond angles associated with <em>sp<\/em><sup>3<\/sup>-, <em>sp<\/em><sup>2<\/sup>&#8211; and sp\u2011hybridized carbon atoms are approximately 109.5, 120 and 180\u00b0, respectively.<\/p>\n<\/div>\n<\/div>\n<\/div>\n<div id=\"section_1\">\n<h3 class=\"editable\">Bonding in acetylene<\/h3>\n<p>Finally, the hybrid orbital concept applies well to triple-bonded groups, such as alkynes and nitriles. Consider, for example, the structure of ethyne (common\u00a0 name acetylene), the simplest alkyne.<\/p>\n<p><img decoding=\"async\" class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/04151225\/image217.png\" alt=\"image210.png\" width=\"152px\" height=\"78px\" \/><\/p>\n<p>This molecule is linear: all four atoms lie in a straight line. The carbon-carbon triple bond is only 1.20\u00c5 long. In the hybrid orbital picture of acetylene, both carbons are <strong>sp-hybridized<\/strong>. In an sp-hybridized carbon, the 2<em>s<\/em> orbital combines with the 2<em>p<\/em><sub>x<\/sub> orbital to form two sp hybrid orbitals that are oriented at an angle of 180\u00b0with respect to each other (eg. along the x axis). The 2<em>p<\/em><sub>y<\/sub> and 2<em>p<\/em><sub>z<\/sub> orbitals remain unhybridized, and are oriented perpendicularly along the y and z axes, respectively.<\/p>\n<p><img decoding=\"async\" class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/04151227\/image219.png\" alt=\"image212.png\" width=\"600px\" height=\"130px\" \/><\/p>\n<p><img decoding=\"async\" class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/04151229\/image221.png\" alt=\"image214.png\" width=\"238px\" height=\"162px\" \/><\/p>\n<p>The C-C sigma bond, then, is formed by the overlap of one sp orbital from each of the carbons, while the two C-H sigma bonds are formed by the overlap of the second sp orbital on each carbon with a 1<em>s<\/em> orbital on a hydrogen.\u00a0 Each carbon atom still has two half-filled 2<em>p<\/em><sub>y<\/sub> and 2<em>p<\/em><sub>z<\/sub> orbitals, which are perpendicular both to each other and to the line formed by the sigma bonds.\u00a0 These two perpendicular pairs of <em>p<\/em> orbitals form two pi bonds between the carbons, resulting in a triple bond overall (one sigma bond plus two pi bonds).<\/p>\n<p><img decoding=\"async\" class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/04151231\/image223.png\" alt=\"image216.png\" width=\"447px\" height=\"123px\" \/><\/p>\n<p>The hybrid orbital concept nicely explains another experimental observation: single bonds adjacent to double and triple bonds are progressively shorter and stronger than \u2018normal\u2019 single bonds, such as the one in a simple alkane.\u00a0 The carbon-carbon bond in ethane (structure A below) results from the overlap of two sp<sup>3<\/sup> orbitals.<\/p>\n<p><img decoding=\"async\" class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/04151233\/image225.png\" alt=\"image218.png\" width=\"597px\" height=\"182px\" \/><\/p>\n<p>In alkene B, however, the carbon-carbon single bond is the result of overlap between an sp<sup>2<\/sup> orbital and an sp<sup>3<\/sup> orbital, while in alkyne C the carbon-carbon single bond is the result of overlap between an sp orbital and an sp<sup>3<\/sup> orbital.\u00a0 These are all single bonds, but the bond in molecule C is shorter and stronger than the one in B, which is in turn shorter and stronger than the one in A.<\/p>\n<p>The explanation here is relatively straightforward.\u00a0 An sp orbital is composed of one <em>s<\/em> orbital and one <em>p<\/em> orbital, and thus it has 50%\u00a0 <em>s<\/em> character and 50% <em>p<\/em> character.\u00a0 sp<sup>2<\/sup> orbitals, by comparison, have 33% <em>s<\/em> character and 67% <em>p<\/em> character, while sp<sup>3<\/sup> orbitals have 25% <em>s<\/em> character and 75% <em>p<\/em> character.\u00a0 Because of their spherical shape, 2<em>s<\/em> orbitals are smaller, and hold electrons closer and \u2018tighter\u2019 to the nucleus, compared to 2<em>p<\/em> orbitals.\u00a0 Consequently, bonds involving sp + sp<sup>3<\/sup> overlap (as in alkyne C)\u00a0 are\u00a0 shorter and stronger than bonds involving sp<sup>2<\/sup> + sp<sup>3<\/sup> overlap (as in alkene B).\u00a0 Bonds involving sp<sup>3<\/sup>-sp<sup>3<\/sup>overlap (as in alkane A)\u00a0 are the longest and weakest of the group, because of the 75% \u2018<em>p<\/em>\u2019 character of the hybrids.<\/p>\n<\/div>\n<div id=\"section_2\">\n<h3 class=\"editable\">Comparison of C-C bonds Ethane, Ethylene, and Acetylene<\/h3>\n<table style=\"border-spacing: 1px\">\n<tbody>\n<tr>\n<td><strong>Molecule<\/strong><\/td>\n<td><strong>Bond<\/strong><\/td>\n<td><strong>Bond Strength (kJ\/mol)<\/strong><\/td>\n<td><strong>Bond Length (pm)<\/strong><\/td>\n<\/tr>\n<tr>\n<td><strong>Ethane, CH<sub>3<\/sub>CH<sub>3<\/sub><\/strong><\/td>\n<td><strong>(<em>sp<sup>3<\/sup><\/em>) C-C (<em>sp<sup>3<\/sup><\/em>)<\/strong><\/td>\n<td><strong>376<\/strong><\/td>\n<td><strong>154<\/strong><\/td>\n<\/tr>\n<tr>\n<td><strong>Ethylene, H<sub>2<\/sub>C=CH2<\/strong><\/td>\n<td><strong>(<em>sp<sup>2<\/sup><\/em>) C=C (<em>sp<sup>2<\/sup><\/em>)<\/strong><\/td>\n<td><strong>728<\/strong><\/td>\n<td><strong>134<\/strong><\/td>\n<\/tr>\n<tr>\n<td>Acetylene, <img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/chem.libretexts.org\/LibreTexts\/Athabasca_University\/Chemistry_350%3A_Organic_Chemistry_I\/Chapter_1%3A_Structure_and_Bonding\/denied:data:image\/png;base64,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#fixme\" alt=\"\" width=\"78\" height=\"17\" \/><\/td>\n<td><strong>(<em>sp<\/em>) <\/strong><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/chem.libretexts.org\/LibreTexts\/Athabasca_University\/Chemistry_350%3A_Organic_Chemistry_I\/Chapter_1%3A_Structure_and_Bonding\/denied:data:image\/png;base64,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#fixme\" alt=\"\" width=\"92\" height=\"27\" \/><strong> (<em>sp<\/em>)<\/strong><\/td>\n<td><strong>965<\/strong><\/td>\n<td><strong>120<\/strong><\/td>\n<\/tr>\n<\/tbody>\n<\/table>\n<p>Notice that as the bond order increases the bond length decreases and the bond strength increases.<\/p>\n<\/div>\n<div id=\"section_3\">\n<div id=\"s61688\">\n<div id=\"section_30\">\n<div class=\"textbox examples\">\n<h3>Example<\/h3>\n<div id=\"section_3\">\n<div id=\"s61688\">\n<div id=\"section_30\">\n<p><span>1-Cyclohexyne is a very strained molecule. By looking at the molecule explain why there is such a intermolecular strain using the knowledge of hybridization and bond angles.<\/span><\/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\/04151236\/1.9.png\" alt=\"\" width=\"94\" height=\"100\" \/><\/p>\n<\/div>\n<div id=\"section_31\">\n<div class=\"qa-wrapper\" style=\"display: block\"><span class=\"show-answer collapsed\" style=\"cursor: pointer\" data-target=\"q47173\">Show Answer<\/span><\/p>\n<div id=\"q47173\" class=\"hidden-answer\" style=\"display: none\">\n<p><span><span>The alkyne is a sp hybridized orbital. By looking at a sp orbital, we can see that the bond angle is 180\u00b0, but in cyclohexane the regular angles would be 109.5\u00b0. Therefore the molecule would be strained to force the 180\u00b0 to be a 109\u00b0.\u00a0 \u00a0<\/div>\n<\/div>\n<p><\/span><\/span><\/p>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<div id=\"section_4\">\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":311,"menu_order":12,"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-293","chapter","type-chapter","status-publish","hentry"],"part":76,"_links":{"self":[{"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/pressbooks\/v2\/chapters\/293","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\/311"}],"version-history":[{"count":6,"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/pressbooks\/v2\/chapters\/293\/revisions"}],"predecessor-version":[{"id":2208,"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/pressbooks\/v2\/chapters\/293\/revisions\/2208"}],"part":[{"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/pressbooks\/v2\/parts\/76"}],"metadata":[{"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/pressbooks\/v2\/chapters\/293\/metadata\/"}],"wp:attachment":[{"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/wp\/v2\/media?parent=293"}],"wp:term":[{"taxonomy":"chapter-type","embeddable":true,"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/pressbooks\/v2\/chapter-type?post=293"},{"taxonomy":"contributor","embeddable":true,"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/wp\/v2\/contributor?post=293"},{"taxonomy":"license","embeddable":true,"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/wp\/v2\/license?post=293"}],"curies":[{"name":"wp","href":"https:\/\/api.w.org\/{rel}","templated":true}]}}