{"id":144040,"date":"2024-09-24T09:42:34","date_gmt":"2024-09-24T09:42:34","guid":{"rendered":"https:\/\/learnexams.com\/blog\/?p=144040"},"modified":"2024-09-24T09:42:37","modified_gmt":"2024-09-24T09:42:37","slug":"draw-the-lewis-structure-for-the-sf4-and-sf%e2%82%86-molecules-and-determine-if-each-molecule-is-polar-or-nonpolar","status":"publish","type":"post","link":"https:\/\/www.learnexams.com\/blog\/2024\/09\/24\/draw-the-lewis-structure-for-the-sf4-and-sf%e2%82%86-molecules-and-determine-if-each-molecule-is-polar-or-nonpolar\/","title":{"rendered":"Draw the Lewis structure for the SF4 and SF\u2086 molecules and determine if each molecule is polar or nonpolar."},"content":{"rendered":"\n<p>Draw the Lewis structure for the SF4 and SF\u2086 molecules and determine if each molecule is polar or nonpolar.<\/p>\n\n\n\n<p><strong><mark style=\"background-color:rgba(0, 0, 0, 0)\" class=\"has-inline-color has-ast-global-color-6-color\">The Correct answer and Explanation is:<\/mark><\/strong><\/p>\n\n\n\n<h3 class=\"wp-block-heading\">Lewis Structures of SF\u2084 and SF\u2086<\/h3>\n\n\n\n<p><strong>SF\u2084 (Sulfur Tetrafluoride):<\/strong><\/p>\n\n\n\n<ol class=\"wp-block-list\">\n<li><strong>Step 1: Count valence electrons.<\/strong><\/li>\n<\/ol>\n\n\n\n<ul class=\"wp-block-list\">\n<li>Sulfur (S) has 6 valence electrons.<\/li>\n\n\n\n<li>Fluorine (F) has 7 valence electrons. There are 4 fluorine atoms, so 4 \u00d7 7 = 28 electrons from fluorine.<\/li>\n\n\n\n<li>Total valence electrons = 6 + 28 = 34.<\/li>\n<\/ul>\n\n\n\n<ol class=\"wp-block-list\">\n<li><strong>Step 2: Place sulfur in the center and attach four fluorine atoms around it.<\/strong><\/li>\n<\/ol>\n\n\n\n<ul class=\"wp-block-list\">\n<li>Each sulfur-fluorine (S-F) bond contains 2 electrons, so 4 bonds use up 8 electrons (4 \u00d7 2 = 8).<\/li>\n<\/ul>\n\n\n\n<ol class=\"wp-block-list\">\n<li><strong>Step 3: Distribute the remaining electrons.<\/strong><\/li>\n<\/ol>\n\n\n\n<ul class=\"wp-block-list\">\n<li>34 total valence electrons minus 8 (for bonds) leaves 26 electrons. These are placed as lone pairs on the fluorine atoms (6 electrons on each fluorine).<\/li>\n<\/ul>\n\n\n\n<ol class=\"wp-block-list\">\n<li><strong>Step 4: Complete the sulfur\u2019s octet.<\/strong><\/li>\n<\/ol>\n\n\n\n<ul class=\"wp-block-list\">\n<li>Sulfur can have more than 8 electrons due to its expanded octet capability. There are 2 remaining electrons, which form a lone pair on sulfur. The structure is <strong>see-saw shaped<\/strong> with 4 bonds and 1 lone pair.<\/li>\n<\/ul>\n\n\n\n<p><strong>SF\u2086 (Sulfur Hexafluoride):<\/strong><\/p>\n\n\n\n<ol class=\"wp-block-list\">\n<li><strong>Step 1: Count valence electrons.<\/strong><\/li>\n<\/ol>\n\n\n\n<ul class=\"wp-block-list\">\n<li>Sulfur has 6 valence electrons.<\/li>\n\n\n\n<li>Fluorine has 7 valence electrons. There are 6 fluorine atoms, so 6 \u00d7 7 = 42 electrons from fluorine.<\/li>\n\n\n\n<li>Total valence electrons = 6 + 42 = 48.<\/li>\n<\/ul>\n\n\n\n<ol class=\"wp-block-list\">\n<li><strong>Step 2: Place sulfur in the center and attach six fluorine atoms around it.<\/strong><\/li>\n<\/ol>\n\n\n\n<ul class=\"wp-block-list\">\n<li>Each sulfur-fluorine bond uses 2 electrons, so 6 bonds use 12 electrons.<\/li>\n<\/ul>\n\n\n\n<ol class=\"wp-block-list\">\n<li><strong>Step 3: Distribute the remaining electrons.<\/strong><\/li>\n<\/ol>\n\n\n\n<ul class=\"wp-block-list\">\n<li>48 total valence electrons minus 12 (for bonds) leaves 36 electrons, distributed as lone pairs on the fluorine atoms (6 electrons on each fluorine). The structure is <strong>octahedral<\/strong> with 6 bonds and no lone pairs on sulfur.<\/li>\n<\/ul>\n\n\n\n<h3 class=\"wp-block-heading\">Polarity of SF\u2084 and SF\u2086<\/h3>\n\n\n\n<ul class=\"wp-block-list\">\n<li><strong>SF\u2084<\/strong> is <strong>polar<\/strong>. Its asymmetrical see-saw shape causes an uneven distribution of charge. The lone pair on sulfur causes a dipole moment, making SF\u2084 polar.<\/li>\n\n\n\n<li><strong>SF\u2086<\/strong> is <strong>nonpolar<\/strong>. It has a symmetrical octahedral shape. Even though sulfur forms polar S-F bonds, the symmetry cancels out the dipole moments, making the molecule nonpolar.<\/li>\n<\/ul>\n\n\n\n<h3 class=\"wp-block-heading\">Conclusion:<\/h3>\n\n\n\n<ul class=\"wp-block-list\">\n<li>SF\u2084 is <strong>polar<\/strong> due to its asymmetry and lone pair.<\/li>\n\n\n\n<li>SF\u2086 is <strong>nonpolar<\/strong> because of its symmetrical octahedral shape, which leads to an even distribution of charge.<\/li>\n<\/ul>\n","protected":false},"excerpt":{"rendered":"<p>Draw the Lewis structure for the SF4 and SF\u2086 molecules and determine if each molecule is polar or nonpolar. The Correct answer and Explanation is: Lewis Structures of SF\u2084 and SF\u2086 SF\u2084 (Sulfur Tetrafluoride): SF\u2086 (Sulfur Hexafluoride): Polarity of SF\u2084 and SF\u2086 Conclusion:<\/p>\n","protected":false},"author":1,"featured_media":0,"comment_status":"closed","ping_status":"closed","sticky":false,"template":"","format":"standard","meta":{"site-sidebar-layout":"default","site-content-layout":"","ast-site-content-layout":"default","site-content-style":"default","site-sidebar-style":"default","ast-global-header-display":"","ast-banner-title-visibility":"","ast-main-header-display":"","ast-hfb-above-header-display":"","ast-hfb-below-header-display":"","ast-hfb-mobile-header-display":"","site-post-title":"","ast-breadcrumbs-content":"","ast-featured-img":"","footer-sml-layout":"","ast-disable-related-posts":"","theme-transparent-header-meta":"","adv-header-id-meta":"","stick-header-meta":"","header-above-stick-meta":"","header-main-stick-meta":"","header-below-stick-meta":"","astra-migrate-meta-layouts":"default","ast-page-background-enabled":"default","ast-page-background-meta":{"desktop":{"background-color":"","background-image":"","background-repeat":"repeat","background-position":"center center","background-size":"auto","background-attachment":"scroll","background-type":"","background-media":"","overlay-type":"","overlay-color":"","overlay-opacity":"","overlay-gradient":""},"tablet":{"background-color":"","background-image":"","background-repeat":"repeat","background-position":"center center","background-size":"auto","background-attachment":"scroll","background-type":"","background-media":"","overlay-type":"","overlay-color":"","overlay-opacity":"","overlay-gradient":""},"mobile":{"background-color":"","background-image":"","background-repeat":"repeat","background-position":"center center","background-size":"auto","background-attachment":"scroll","background-type":"","background-media":"","overlay-type":"","overlay-color":"","overlay-opacity":"","overlay-gradient":""}},"ast-content-background-meta":{"desktop":{"background-color":"var(--ast-global-color-5)","background-image":"","background-repeat":"repeat","background-position":"center center","background-size":"auto","background-attachment":"scroll","background-type":"","background-media":"","overlay-type":"","overlay-color":"","overlay-opacity":"","overlay-gradient":""},"tablet":{"background-color":"var(--ast-global-color-5)","background-image":"","background-repeat":"repeat","background-position":"center center","background-size":"auto","background-attachment":"scroll","background-type":"","background-media":"","overlay-type":"","overlay-color":"","overlay-opacity":"","overlay-gradient":""},"mobile":{"background-color":"var(--ast-global-color-5)","background-image":"","background-repeat":"repeat","background-position":"center center","background-size":"auto","background-attachment":"scroll","background-type":"","background-media":"","overlay-type":"","overlay-color":"","overlay-opacity":"","overlay-gradient":""}},"footnotes":""},"categories":[25],"tags":[],"class_list":["post-144040","post","type-post","status-publish","format-standard","hentry","category-exams-certification"],"_links":{"self":[{"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/posts\/144040","targetHints":{"allow":["GET"]}}],"collection":[{"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/posts"}],"about":[{"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/types\/post"}],"author":[{"embeddable":true,"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/users\/1"}],"replies":[{"embeddable":true,"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/comments?post=144040"}],"version-history":[{"count":0,"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/posts\/144040\/revisions"}],"wp:attachment":[{"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/media?parent=144040"}],"wp:term":[{"taxonomy":"category","embeddable":true,"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/categories?post=144040"},{"taxonomy":"post_tag","embeddable":true,"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/tags?post=144040"}],"curies":[{"name":"wp","href":"https:\/\/api.w.org\/{rel}","templated":true}]}}