{"id":267,"date":"2017-10-04T15:05:23","date_gmt":"2017-10-04T15:05:23","guid":{"rendered":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/?post_type=chapter&#038;p=267"},"modified":"2017-10-18T19:03:49","modified_gmt":"2017-10-18T19:03:49","slug":"hybridization-structure-of-methane","status":"publish","type":"chapter","link":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/chapter\/hybridization-structure-of-methane\/","title":{"raw":"Hybridization: Structure of Methane","rendered":"Hybridization: Structure of Methane"},"content":{"raw":"<div class=\"elm-header\">\r\n<div class=\"elm-header-custom\">\r\n\r\n\r\n<div class=\"textbox learning-objectives\">\r\n<h3>Objective<\/h3>\r\n<div class=\"elm-header\">\r\n<div class=\"elm-header-custom\">Objective<\/div>\r\n<\/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 describe the structure of methane in terms of the <em>sp<\/em><sup>3<\/sup> hybridization of the central carbon atom.\r\n\r\n<\/div>\r\n<div><\/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>\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 terms below.\r\n<ul>\r\n \t<li>bond angle<\/li>\r\n \t<li>hybridization<\/li>\r\n \t<li><em>sp<\/em><sup>3<\/sup> hybrid<\/li>\r\n<\/ul>\r\n<\/div>\r\n<\/div>\r\n<div id=\"note\">\r\n<p class=\"boxtitle\">Study Notes<\/p>\r\nThe tetrahedral shape is a very important one in organic chemistry, as it is the basic shape of all compounds in which a carbon atom is bonded to four other atoms. Note that the tetrahedral bond angle of H\u2212C\u2212H is 109.5\u00b0.\r\n\r\n<\/div>\r\n<div id=\"section_1\">\r\n<h3 class=\"editable\">Bonding in Methane, CH<sub>4<\/sub><\/h3>\r\nWe are starting with methane because it is the simplest case which illustrates the sort of processes involved. You will remember that the dots-and-crossed picture of methane looks like this.\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\/04150504\/ch4dandc.gif\" alt=\"\" width=\"131px\" height=\"130px\" \/>\r\n\r\nThere is a serious mismatch between this structure and the modern electronic structure of carbon, 1s<sup>2<\/sup>2s<sup>2<\/sup>2p<sub>x<\/sub><sup>1<\/sup>2p<sub>y<\/sub><sup>1<\/sup>. The modern structure shows that there are only 2 unpaired electrons to share with hydrogens, instead of the 4 which the simple view requires.\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\/04150506\/cground.gif\" alt=\"\" width=\"144px\" height=\"68px\" \/>\r\n\r\nYou can see this more readily using the electrons-in-boxes notation. Only the 2-level electrons are shown. The 1s<sup>2<\/sup> electrons are too deep inside the atom to be involved in bonding. The only electrons directly available for sharing are the 2p electrons. Why then isn't methane CH<sub>2<\/sub>?\r\n<div id=\"section_11\">\r\n<div id=\"section_15\">\r\n<div id=\"section_2\">\r\n<h3 class=\"editable\">Promotion of an electron<\/h3>\r\nWhen bonds are formed, energy is released and the system becomes more stable. If carbon forms 4 bonds rather than 2, twice as much energy is released and so the resulting molecule becomes even more stable. There is only a small energy gap between the 2s and 2p orbitals, and so it pays the carbon to provide a small amount of energy to promote an electron from the 2s to the empty 2p to give 4 unpaired electrons. The extra energy released when the bonds form more than compensates for the initial input.\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\/04150507\/cpromote.gif\" alt=\"\" width=\"144px\" height=\"212px\" \/>\r\n\r\nThe carbon atom is now said to be in an excited state. Now that we've got 4 unpaired electrons ready for bonding, another problem arises. In methane all the carbon-hydrogen bonds are identical, but our electrons are in two different kinds of orbitals. You aren't going to get four identical bonds unless you start from four identical orbitals.\r\n\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<div id=\"section_12\">\r\n<div id=\"section_16\">\r\n<div id=\"section_3\">\r\n<h3 class=\"editable\">Hybridization<\/h3>\r\nThe electrons rearrange themselves again in a process called hybridization. This reorganizes the electrons into four identical hybrid orbitals called sp<sup>3<\/sup> hybrids (because they are made from one s orbital and three p orbitals). You should read \"sp<sup>3<\/sup>\" as \"s p three\" - not as \"s p cubed\".\r\n\r\n<img class=\"size-medium wp-image-2013 aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/18190304\/hybridization-300x143.jpg\" alt=\"\" width=\"300\" height=\"143\" \/>\r\n\r\nsp<sup>3<\/sup> hybrid orbitals look a bit like half a p orbital, and they arrange themselves in space so that they are as far apart as possible. You can picture the nucleus as being at the center of a tetrahedron (a triangularly based pyramid) with the orbitals pointing to the corners. For clarity, the nucleus is drawn far larger than it really is.\r\n<div id=\"section_13\">\r\n<div id=\"section_17\">\r\n<div id=\"section_4\">\r\n<h4 class=\"editable\">What happens when the bonds are formed?<\/h4>\r\nRemember that hydrogen's electron is in a 1s orbital - a spherically symmetric region of space surrounding the nucleus where there is some fixed chance (say 95%) of finding the electron. When a covalent bond is formed, the atomic orbitals (the orbitals in the individual atoms) merge to produce a new molecular orbital which contains the electron pair which creates the bond.\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\/04150511\/ch4orbitals.gif\" alt=\"\" width=\"338px\" height=\"213px\" \/>\r\n\r\nFour molecular orbitals are formed, looking rather like the original sp<sup>3<\/sup> hybrids, but with a hydrogen nucleus embedded in each lobe. Each orbital holds the 2 electrons that we've previously drawn as a dot and a cross. The principles involved - promotion of electrons if necessary, then hybridization, followed by the formation of molecular orbitals - can be applied to any covalently-bound molecule.\r\n\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>\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<\/div>\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>Objective<\/h3>\n<div class=\"elm-header\">\n<div class=\"elm-header-custom\">Objective<\/div>\n<\/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 describe the structure of methane in terms of the <em>sp<\/em><sup>3<\/sup> hybridization of the central carbon atom.<\/p>\n<\/div>\n<div><\/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>\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 terms below.<\/p>\n<ul>\n<li>bond angle<\/li>\n<li>hybridization<\/li>\n<li><em>sp<\/em><sup>3<\/sup> hybrid<\/li>\n<\/ul>\n<\/div>\n<\/div>\n<div id=\"note\">\n<p class=\"boxtitle\">Study Notes<\/p>\n<p>The tetrahedral shape is a very important one in organic chemistry, as it is the basic shape of all compounds in which a carbon atom is bonded to four other atoms. Note that the tetrahedral bond angle of H\u2212C\u2212H is 109.5\u00b0.<\/p>\n<\/div>\n<div id=\"section_1\">\n<h3 class=\"editable\">Bonding in Methane, CH<sub>4<\/sub><\/h3>\n<p>We are starting with methane because it is the simplest case which illustrates the sort of processes involved. You will remember that the dots-and-crossed picture of methane looks like this.<\/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\/04150504\/ch4dandc.gif\" alt=\"\" width=\"131px\" height=\"130px\" \/><\/p>\n<p>There is a serious mismatch between this structure and the modern electronic structure of carbon, 1s<sup>2<\/sup>2s<sup>2<\/sup>2p<sub>x<\/sub><sup>1<\/sup>2p<sub>y<\/sub><sup>1<\/sup>. The modern structure shows that there are only 2 unpaired electrons to share with hydrogens, instead of the 4 which the simple view requires.<\/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\/04150506\/cground.gif\" alt=\"\" width=\"144px\" height=\"68px\" \/><\/p>\n<p>You can see this more readily using the electrons-in-boxes notation. Only the 2-level electrons are shown. The 1s<sup>2<\/sup> electrons are too deep inside the atom to be involved in bonding. The only electrons directly available for sharing are the 2p electrons. Why then isn&#8217;t methane CH<sub>2<\/sub>?<\/p>\n<div id=\"section_11\">\n<div id=\"section_15\">\n<div id=\"section_2\">\n<h3 class=\"editable\">Promotion of an electron<\/h3>\n<p>When bonds are formed, energy is released and the system becomes more stable. If carbon forms 4 bonds rather than 2, twice as much energy is released and so the resulting molecule becomes even more stable. There is only a small energy gap between the 2s and 2p orbitals, and so it pays the carbon to provide a small amount of energy to promote an electron from the 2s to the empty 2p to give 4 unpaired electrons. The extra energy released when the bonds form more than compensates for the initial input.<\/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\/04150507\/cpromote.gif\" alt=\"\" width=\"144px\" height=\"212px\" \/><\/p>\n<p>The carbon atom is now said to be in an excited state. Now that we&#8217;ve got 4 unpaired electrons ready for bonding, another problem arises. In methane all the carbon-hydrogen bonds are identical, but our electrons are in two different kinds of orbitals. You aren&#8217;t going to get four identical bonds unless you start from four identical orbitals.<\/p>\n<\/div>\n<\/div>\n<\/div>\n<div id=\"section_12\">\n<div id=\"section_16\">\n<div id=\"section_3\">\n<h3 class=\"editable\">Hybridization<\/h3>\n<p>The electrons rearrange themselves again in a process called hybridization. This reorganizes the electrons into four identical hybrid orbitals called sp<sup>3<\/sup> hybrids (because they are made from one s orbital and three p orbitals). You should read &#8220;sp<sup>3<\/sup>&#8221; as &#8220;s p three&#8221; &#8211; not as &#8220;s p cubed&#8221;.<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" class=\"size-medium wp-image-2013 aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/18190304\/hybridization-300x143.jpg\" alt=\"\" width=\"300\" height=\"143\" \/><\/p>\n<p>sp<sup>3<\/sup> hybrid orbitals look a bit like half a p orbital, and they arrange themselves in space so that they are as far apart as possible. You can picture the nucleus as being at the center of a tetrahedron (a triangularly based pyramid) with the orbitals pointing to the corners. For clarity, the nucleus is drawn far larger than it really is.<\/p>\n<div id=\"section_13\">\n<div id=\"section_17\">\n<div id=\"section_4\">\n<h4 class=\"editable\">What happens when the bonds are formed?<\/h4>\n<p>Remember that hydrogen&#8217;s electron is in a 1s orbital &#8211; a spherically symmetric region of space surrounding the nucleus where there is some fixed chance (say 95%) of finding the electron. When a covalent bond is formed, the atomic orbitals (the orbitals in the individual atoms) merge to produce a new molecular orbital which contains the electron pair which creates the bond.<\/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\/04150511\/ch4orbitals.gif\" alt=\"\" width=\"338px\" height=\"213px\" \/><\/p>\n<p>Four molecular orbitals are formed, looking rather like the original sp<sup>3<\/sup> hybrids, but with a hydrogen nucleus embedded in each lobe. Each orbital holds the 2 electrons that we&#8217;ve previously drawn as a dot and a cross. The principles involved &#8211; promotion of electrons if necessary, then hybridization, followed by the formation of molecular orbitals &#8211; can be applied to any covalently-bound molecule.<\/p>\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<\/div>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n","protected":false},"author":311,"menu_order":9,"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-267","chapter","type-chapter","status-publish","hentry"],"part":76,"_links":{"self":[{"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/pressbooks\/v2\/chapters\/267","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":4,"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/pressbooks\/v2\/chapters\/267\/revisions"}],"predecessor-version":[{"id":2014,"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/pressbooks\/v2\/chapters\/267\/revisions\/2014"}],"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\/267\/metadata\/"}],"wp:attachment":[{"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/wp\/v2\/media?parent=267"}],"wp:term":[{"taxonomy":"chapter-type","embeddable":true,"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/pressbooks\/v2\/chapter-type?post=267"},{"taxonomy":"contributor","embeddable":true,"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/wp\/v2\/contributor?post=267"},{"taxonomy":"license","embeddable":true,"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/wp\/v2\/license?post=267"}],"curies":[{"name":"wp","href":"https:\/\/api.w.org\/{rel}","templated":true}]}}