{"id":285,"date":"2013-07-29T12:29:34","date_gmt":"2013-07-29T12:29:34","guid":{"rendered":"http:\/\/www.studyorgo.com\/blog\/?p=285"},"modified":"2014-02-21T03:22:02","modified_gmt":"2014-02-21T03:22:02","slug":"energy-diagram-module-series-part-two-gibbs-free-energy-and-spontaneity","status":"publish","type":"post","link":"https:\/\/www.studyorgo.com\/blog\/energy-diagram-module-series-part-two-gibbs-free-energy-and-spontaneity\/","title":{"rendered":"Energy Diagram Module Series- Part Two: Gibbs Free Energy and Spontaneity"},"content":{"rendered":"<p>This is part 2 of a four part series in the Energy Diagram Module. Stay tuned for the other parts!<\/p>\n<p>To see part 1 <a href=\"http:\/\/www.studyorgo.com\/blog\/energy-diagram-module-series-part-one-introduction\/?utm_source=rss&amp;utm_medium=rss&amp;utm_campaign=energy-diagram-module-series-part-one-introduction\">click here<\/a>.<\/p>\n<p>In order to talk about energy of the reaction, a few key concepts are needed.<\/p>\n<ul>\n<li>If the products have less potential energy than the reactants, the reaction will release a net amount of energy (an <b>exothermic reaction<\/b>).<\/li>\n<li>If the products are higher in energy than the reactions, the reaction will consume a net amount of energy (<b>an endothermic reaction<\/b>).<\/li>\n<\/ul>\n<p>But, in order to predict how well a reaction will progress, or how <b>spontaneous<\/b> the reaction will be, enthalpy is insufficient to make this estimation.\u00a0 Therefore, we rely on the thermodynamic calculation of <b>Gibbs Free Energy (<i>\u0394<\/i><\/b><b><i><\/i><\/b><b><i>G<sup>0<\/sup><\/i><\/b><b>)<\/b> which is represented by the equation;<\/p>\n<p align=\"center\"><i>\u0394<\/i><i>G<sup>0 <\/sup>= <i>\u0394<\/i><\/i><i><\/i><i>H<sup>0<\/sup> &#8211; T<i>\u0394<\/i><\/i><i><\/i><i>S<\/i><\/p>\n<p style=\"text-align: left;\" align=\"center\">The components of Gibbs Free Energy are:<\/p>\n<ul>\n<li><b>Enthalpy<\/b>, <i>\u0394<\/i><b><\/b><b>H<sup>0<\/sup> &#8211;<\/b> The heat consumed or released by the reaction. <i><\/i><\/li>\n<li><b>Temperature, T<\/b> \u2013 Temperature of the system.<i><\/i><\/li>\n<li><b>Entropy, <i>\u0394<\/i><\/b><b><\/b><b>S<\/b>.\u00a0 The change in the degree of disorder.<i><\/i><\/li>\n<\/ul>\n<p>The sign of the reaction indicates the release (negative) or absorption (positive) of heat during the reaction. Therefore, if\u00a0<i>\u0394<\/i><i><\/i><i><span style=\"text-decoration: underline;\">G<sup>0<\/sup><\/span><\/i><span style=\"text-decoration: underline;\"> is negative the reaction is <b>always<\/b> spontaneous<\/span>.\u00a0 Similarly, <span style=\"text-decoration: underline;\">if <i>\u0394<\/i><\/span><i><\/i><i><span style=\"text-decoration: underline;\">G<sup>0<\/sup><\/span><\/i><span style=\"text-decoration: underline;\"> is positive the reaction is <b>never<\/b> spontaneous<\/span>.\u00a0 <span style=\"text-decoration: underline;\">When <i>\u0394<\/i><\/span><i><\/i><i><span style=\"text-decoration: underline;\">G<sup>0<\/sup><\/span><\/i><span style=\"text-decoration: underline;\"> equals zero, the reaction is at<b> equilibrium<\/b>.<\/span> In accounting for these additional thermodynamic properties using Gibbs Free Energy, different terms are used.<\/p>\n<ul>\n<li>If a reaction has a negative \u00a0<i>\u0394<\/i><b><i><\/i><\/b><b><i>G<sup>0 <\/sup><\/i><\/b>, it is therefore spontaneous and is said to be <b>exergonic<\/b>.<i><\/i><\/li>\n<li>Conversely, a reaction that has a\u00a0 positive \u00a0<i>\u0394<\/i><b><i><\/i><\/b><b><i>G<sup>0 <\/sup><\/i><\/b>\u00a0is not spontaneous is considered <b>endergonic<\/b>. <i><\/i><\/li>\n<\/ul>\n<p>&nbsp;<\/p>\n<p><a href=\"http:\/\/www.studyorgo.com\/blog\/wp-content\/uploads\/2013\/07\/StudyOrgo.com-Graph-21.png\"><img loading=\"lazy\" decoding=\"async\" class=\"alignnone size-medium wp-image-281 aligncenter\" alt=\"StudyOrgo.com - Graph 2\" src=\"http:\/\/www.studyorgo.com\/blog\/wp-content\/uploads\/2013\/07\/StudyOrgo.com-Graph-21-300x152.png\" width=\"300\" height=\"152\" srcset=\"https:\/\/www.studyorgo.com\/blog\/wp-content\/uploads\/2013\/07\/StudyOrgo.com-Graph-21-300x152.png 300w, https:\/\/www.studyorgo.com\/blog\/wp-content\/uploads\/2013\/07\/StudyOrgo.com-Graph-21.png 1024w\" sizes=\"auto, (max-width: 300px) 100vw, 300px\" \/><\/a><\/p>\n<p style=\"text-align: center;\">Graph 2<\/p>\n<p style=\"text-align: center;\">(click on photo to enlarge)<\/p>\n<p>With the addition of the temperature variable in the Gibbs Free Energy equation, it is easy to see that an endergonic reactions can be driven forward simply by increasing the temperature of the reaction so that the term T<i>\u0394<\/i>S is more negative than <i>\u0394<\/i>H<sup>0<\/sup>, thus making <i>\u0394<\/i>G<sup>0<\/sup> negative and making the reaction spontaneous!<i><span style=\"text-decoration: underline;\"><br \/>\n<\/span><\/i><\/p>\n","protected":false},"excerpt":{"rendered":"<p>This is part 2 of a four part series in the Energy Diagram Module. Stay tuned for the other parts! To see part 1 click here. In order to talk about energy of the reaction, a few key concepts are needed. If the products have less potential energy than the reactants, the reaction will release [&hellip;]<\/p>\n","protected":false},"author":2,"featured_media":0,"comment_status":"open","ping_status":"open","sticky":false,"template":"","format":"standard","meta":{"_monsterinsights_skip_tracking":false,"_monsterinsights_sitenote_active":false,"_monsterinsights_sitenote_note":"","_monsterinsights_sitenote_category":0,"footnotes":""},"categories":[1],"tags":[],"class_list":["post-285","post","type-post","status-publish","format-standard","hentry","category-organic-chemistry"],"aioseo_notices":[],"_links":{"self":[{"href":"https:\/\/www.studyorgo.com\/blog\/wp-json\/wp\/v2\/posts\/285","targetHints":{"allow":["GET"]}}],"collection":[{"href":"https:\/\/www.studyorgo.com\/blog\/wp-json\/wp\/v2\/posts"}],"about":[{"href":"https:\/\/www.studyorgo.com\/blog\/wp-json\/wp\/v2\/types\/post"}],"author":[{"embeddable":true,"href":"https:\/\/www.studyorgo.com\/blog\/wp-json\/wp\/v2\/users\/2"}],"replies":[{"embeddable":true,"href":"https:\/\/www.studyorgo.com\/blog\/wp-json\/wp\/v2\/comments?post=285"}],"version-history":[{"count":4,"href":"https:\/\/www.studyorgo.com\/blog\/wp-json\/wp\/v2\/posts\/285\/revisions"}],"predecessor-version":[{"id":296,"href":"https:\/\/www.studyorgo.com\/blog\/wp-json\/wp\/v2\/posts\/285\/revisions\/296"}],"wp:attachment":[{"href":"https:\/\/www.studyorgo.com\/blog\/wp-json\/wp\/v2\/media?parent=285"}],"wp:term":[{"taxonomy":"category","embeddable":true,"href":"https:\/\/www.studyorgo.com\/blog\/wp-json\/wp\/v2\/categories?post=285"},{"taxonomy":"post_tag","embeddable":true,"href":"https:\/\/www.studyorgo.com\/blog\/wp-json\/wp\/v2\/tags?post=285"}],"curies":[{"name":"wp","href":"https:\/\/api.w.org\/{rel}","templated":true}]}}