{"id":243,"date":"2017-08-21T04:47:28","date_gmt":"2017-08-21T04:47:28","guid":{"rendered":"http:\/\/americanboard.org\/Subjects\/chemistry\/?page_id=243"},"modified":"2017-09-19T15:04:46","modified_gmt":"2017-09-19T15:04:46","slug":"lewis-structures","status":"publish","type":"page","link":"https:\/\/americanboard.org\/Subjects\/chemistry\/lewis-structures\/","title":{"rendered":"Lewis Structures"},"content":{"rendered":"<div class=\"twelve columns\" style=\"margin-top: 10%;\">\n<div class=\"advance\">\n<p><a class=\"button button-primary\" href=\"http:\/\/americanboard.org\/Subjects\/chemistry\/patterns-in-space\">\u2b05 Previous Lesson<\/a>\u00a0<a class=\"button\" href=\"http:\/\/americanboard.org\/Subjects\/chemistry\/chemical-naming-and-structure\">Workshop Index<\/a>\u00a0<a class=\"button button-primary\" href=\"http:\/\/americanboard.org\/Subjects\/chemistry\/vsepr\">Next Lesson \u27a1<\/a><\/p>\n<\/div>\n<p><!-- UPDATE NEXT\/PREVIOUS ABOVE --><\/p>\n<p><!-- CONTENT STARTS HERE --><\/p>\n<h1 id=\"title\">Chemical Naming and Structure: Lewis Structures<\/h1>\n<h4>Objective<\/h4>\n<p>We will review drawing Lewis structures for molecules and ions, including those that follow the Octet Rule and those that do not.<\/p>\n<h4>Previously we covered&#8230;<\/h4>\n<ul>\n<li>Ionic bonding occurs between a metal atom and a nonmetal atom, due to their large difference in electronegativity.<\/li>\n<li>Covalent bonding occurs between two identical nonmetal atoms, due to their identical electronegativities.<\/li>\n<li>Polar covalent bonding occurs between two different nonmetal atoms, due to their moderately different electronegativities.<\/li>\n<li>Nonmetals bond with each other covalently to form discrete particles called molecules.<\/li>\n<li>Metals bond with nonmetals to form oppositely-charged ions that alternate in a three-dimensional spatial pattern that is fully interconnected and indefinite.<\/li>\n<li>Very few substances (usually containing carbon or silicon) covalently bond in a three-dimensional, interconnected and indefinite array.<\/li>\n<li>Metal atoms bond with each other through a \u201csea of electrons\u201d\u2014electrons that freely float between metal atoms that are patterned in a three-dimensional network.<\/li>\n<\/ul>\n<section>\n<h3>Overview<\/h3>\n<p>Lewis structures are a useful tool in determining the arrangement of atoms and types of bonds in a molecule. A typical Lewis structure represents an atom\u2019s nucleus and its <abbr title=\"Electrons in all energy levels below the atom\u2019s highest occupied energy level\">inner electrons<\/abbr> with the element symbol, and the atom\u2019s <abbr title=\"Electrons in the highest occupied energy level of an atom\">valence electrons<\/abbr> with dots. It is helpful to remember that a Lewis structure is only a two-dimensional representation; therefore, it is not meant to show the three-dimensional shape of a molecule (although we will see how to infer that shape in a later lesson). It also behooves us to note that Lewis structures are not particularly useful for substances that do not form discrete molecules, such as ionic compounds, polymers or metals.<\/p>\n<h4>Steps for Drawing Lewis Structures for Molecules that Obey the Octet Rule<\/h4>\n<p>The following is a simple set of steps that we follow in drawing Lewis structures.<\/p>\n<ol>\n<li>Count the valence electrons available. Be sure to consider the effect of charge if the species is a <abbr title=\"A group of atoms with a positive or negative charge\">polyatomic ion<\/abbr>.<\/li>\n<li>Draw a skeleton structure of the molecule, using only <abbr title=\" A covalent bond in which two atoms share a single pair of electrons\"> single bonds<\/abbr>. This structure should usually be symmetrical.<\/li>\n<li>Distribute the electrons remaining from the initial count throughout the molecule as needed according to the Octet Rule. It is often helpful to begin this process with atoms on the perimeter of the molecule, saving the central atom for the end.<\/li>\n<li>If all electrons are used prior to each atom achieving an octet, it will be necessary to use <abbr title=\"A covalent bond in which two atoms share a single pair of electrons\">double bonds<\/abbr> or <abbr title=\"a covalent bond in which two atoms share three pairs of electrons\">triple bonds<\/abbr>.<\/li>\n<\/ol>\n<h4>Examples that Obey the Octet Rule<\/h4>\n<p>Let\u2019s start with a simple molecule whose complete Lewis structure we saw in a previous lesson: CCl<sub>4<\/sub>. Carbon has four valence electrons and chlorine has seven, so the valence electrons available are: 4 + (4\u00d77) = 32 electrons. In part (a) of the figure below, the skeleton structure is shown. Carbon has been placed in the center to make the molecule symmetrical, and there is a single bond from carbon to each chlorine atom. This uses up eight of the thirty-two electrons, so there are twenty-four electrons left. If we place six additional electrons (in pairs) around each chlorine atom, then all of our electrons have been used and each atom has eight valence electrons around it, completing an octet for each atom. This is shown in part (b).<\/p>\n<p><center><img decoding=\"async\" src=\"http:\/\/americanboard.org\/Subjects\/chemistry\/wp-content\/uploads\/sites\/3\/2017\/08\/patternsinspace2.CCl4_.gif\" \/><\/center>Now we\u2019ll try a somewhat more complex example: CO<sub>2<\/sub>. As before, carbon has four valence electrons and oxygen has six, so the valence electrons available are: 4 + (2\u00d76) = 16 electrons. Part (a) below shows the skeleton structure with carbon in the center for symmetry. Four electrons are used thus far, so twelve electrons are left. We could place six additional electrons around each oxygen, as in part (b), but that would leave the carbon electron-deficient; that is, the carbon would have less than eight electrons. Part (c) shows what would happen if we move a pair of electrons from each oxygen to the carbon. Note that carbon now has eight electrons, but each oxygen is now electron-deficient. What we need is to find a place to move the electrons where they will still belong to oxygen, but will also belong to carbon. Sharing an additional pair, and thus creating a double bond on both sides as shown in part (d) will accomplish this.<\/p>\n<table>\n<tbody>\n<tr>\n<td><img loading=\"lazy\" decoding=\"async\" src=\"http:\/\/americanboard.org\/Subjects\/chemistry\/wp-content\/uploads\/sites\/3\/2017\/08\/lewisstructures3.co2a.gif\" alt=\"Image showing part a stage of the Lewis structure for CO2\" width=\"121\" height=\"51\" align=\"absmiddle\" \/><\/td>\n<td><img loading=\"lazy\" decoding=\"async\" src=\"http:\/\/americanboard.org\/Subjects\/chemistry\/wp-content\/uploads\/sites\/3\/2017\/08\/lewisstructures3.co2b.gif\" alt=\"Image showing part b stage of the Lewis structure for CO2\" width=\"137\" height=\"51\" align=\"absmiddle\" \/><\/td>\n<td><img loading=\"lazy\" decoding=\"async\" src=\"http:\/\/americanboard.org\/Subjects\/chemistry\/wp-content\/uploads\/sites\/3\/2017\/08\/lewisstructures3.co2c.gif\" alt=\"Image showing part c stage of the Lewis structure for CO2\" width=\"137\" height=\"51\" align=\"absmiddle\" \/><\/td>\n<td><img loading=\"lazy\" decoding=\"async\" src=\"http:\/\/americanboard.org\/Subjects\/chemistry\/wp-content\/uploads\/sites\/3\/2017\/08\/lewisstructures3.co2d.gif\" alt=\"Image showing part d stage of the Lewis structure for CO2\" width=\"145\" height=\"51\" align=\"absmiddle\" \/><\/td>\n<\/tr>\n<\/tbody>\n<\/table>\n<p>It is time to look at a polyatomic ion: ClO<sub>3<\/sub><sup>1\u2212<\/sup>. Chlorine has seven valence electrons and each oxygen six, but we must also account for the particle\u2019s charge. A negative charge is created when an electron is added to a neutral species, so we need to add an additional electron for the 1\u2212 charge. Our electron count looks like this: 7 + (3\u00d76) + 1 =26 electrons. The skeleton structure has chlorine centrally located for symmetry, as shown in part (a) below. Placing six additional electrons around each oxygen provides an octet for them, as shown in part (b). At this point, we have used up twenty-four electrons, so there are two electrons left. Since the chlorine is still two electrons shy of eight, we can put the remaining electrons on the chlorine, creating a lone pair. <abbr title=\"A pair of electrons that is not shared between atoms in a bond, but belongs to one atom alone\">Lone pairs<\/abbr> will become important when we discuss shapes later. The only other thing we must do to our structure is show the charge; after all, the number of electrons is different for ClO<sub>3<\/sub> vs. ClO<sub>3<\/sub><sup>1\u2212<\/sup>. The final structure is shown in part (c).<\/p>\n<table>\n<tbody>\n<tr>\n<td><img loading=\"lazy\" decoding=\"async\" src=\"http:\/\/americanboard.org\/Subjects\/chemistry\/wp-content\/uploads\/sites\/3\/2017\/08\/lewisstructures3.clo2a.gif\" alt=\"Image showing stage a of the Lewis structure for ClO3\" width=\"116\" height=\"77\" \/><\/td>\n<td><img loading=\"lazy\" decoding=\"async\" src=\"http:\/\/americanboard.org\/Subjects\/chemistry\/wp-content\/uploads\/sites\/3\/2017\/08\/lewisstructures3.clo3b.gif\" alt=\"Image showing stage b of the Lewis structure for ClO3\" width=\"133\" height=\"96\" \/><\/td>\n<td><img loading=\"lazy\" decoding=\"async\" src=\"http:\/\/americanboard.org\/Subjects\/chemistry\/wp-content\/uploads\/sites\/3\/2017\/08\/lewisstructures3.clo3c.gif\" alt=\"Image showing stage c of the Lewis structure for ClO3\" width=\"180\" height=\"130\" \/><\/td>\n<\/tr>\n<\/tbody>\n<\/table>\n<h3>Resonance Structures<\/h3>\n<p>Sometimes, a single Lewis structure may be misleading. It may imply that one bond is different from another (single vs. double, for example) when in reality they are the same. This is a situation in which we use resonance structures<\/p>\n<p>Let\u2019s consider the carbonate ion: CO<sub>3<\/sub><sup>2\u2212<\/sup>. The electron count is as follows: 4 + (3\u00d76) + 2 = 24 electrons. Working through the steps from above should result in any one of the structures in the figure below. Each of these structures is considered valid, in that it uses exactly 24 electrons, obeys the Octet Rule for every atom and shows the appropriate charge. The problem with any single structure is that it implies that one of the carbon-oxygen bonds is different from the other two\u2014double vs. single. If this were correct, then experiments such as X-ray crystallography would show the double bond to be measurably shorter than the two single bonds. Instead, experimental results indicate that all three C-O bonds are of equal length, shorter than a single bond yet longer than a double bond. In other words, the \u201creal\u201d structure is not any one of those shown in the figure, but an average of them all.<\/p>\n<p><center><img decoding=\"async\" src=\"http:\/\/americanboard.org\/Subjects\/chemistry\/wp-content\/uploads\/sites\/3\/2017\/08\/lewisstructures4.lew_dot_CO3-2.gif\" \/><\/center><\/p>\n<h3>Exceptions to the Octet Rule<\/h3>\n<p>The Octet Rule is a marvelous rule of thumb, but it is far from set in stone. For example, hydrogen cannot have more than\u00a0two electrons due to the lack of a <em>p\u00a0<\/em>sublevel in its highest occupied energy level, <em>n\u00a0<\/em>= 1. Nonmetals in the second period (B, C, N, O and F) tend to\u00a0follow the Octet Rule rather rigidly, but nonmetals in the third period\u00a0and\u00a0beyond are capable of having more than eight valence electrons in a\u00a0stable\u00a0molecular structure. This is usually attributed to the presence of\u00a0empty <em>d<\/em> orbitals in the highest\u00a0occupied\u00a0energy level, which are capable of accepting electrons in a coordinated\u00a0<abbr title=\" A bond in which the shared pair of electrons is donated entirely from one atom to the other\">covalent bond<\/abbr>. In general,\u00a0extra electrons (those above eight) are found on the central atom.<\/p>\n<p>Let\u2019s\u00a0consider PCl<sub>5<\/sub>. Our electron count\u00a0is: 5 + (5\u00d77) = 40 electrons. We should recognize that the\u00a0Octet Rule is a\u00a0problem as soon as we draw the skeleton structure. If we connect each\u00a0Cl to the\u00a0P with a single bond (which would make the molecule symmetrical), then\u00a0P\u00a0already has ten electrons, not eight. This is shown in part (a) of the\u00a0figure\u00a0below. Arranging the remaining electrons around P finishes the\u00a0structure, which\u00a0is shown in part (b).<\/p>\n<p><center><img loading=\"lazy\" decoding=\"async\" src=\"http:\/\/americanboard.org\/Subjects\/chemistry\/wp-content\/uploads\/sites\/3\/2017\/08\/lewisstructures4.lewis_dot_PCl5.gif\" alt=\"Image showing 2 stages of the Lewis structure of PCl5\" width=\"486\" height=\"280\" \/><\/center><\/p>\n<section class=\"question\">\n<h4>Question<\/h4>\n<p>Which of the following describes the Lewis structure of the NO<sub>2<\/sub><sup>1\u2212<\/sup> ion?<\/p>\n<ol>\n<li>It obeys the Octet Rule with each O singly-bonded to the N.<\/li>\n<li>It obeys the Octet Rule with two resonance structures, each of which has one O doubly-bonded to N and one O singly-bonded.<\/li>\n<li>It obeys the Octet Rule with each O doubly-bonded to the N.<\/li>\n<li>It violates the Octet Rule with each O doubly-bonded to the N and a lone pair on the N<\/li>\n<\/ol>\n<p><a class=\"button button-primary q-answer\"> Reveal Answer <\/a><\/p>\n<p class=\"q-reveal\">Choice B is correct. The electron count is 5 + (2\u00d76) + 1 = 18 electrons. Choice A leaves the nitrogen electron-deficient with only six electrons. Choice C only uses 16 electrons. Choice D is incorrect because neither nitrogen nor oxygen have empty d orbitals so they must obey the Octet Rule.<\/p>\n<\/section>\n<p><!-- CONTENT ENDS HERE --><\/p>\n<p><!-- UPDATE NEXT\/PREVIOUS BELOW --><\/p>\n<div class=\"advance\">\n<p><a class=\"button button-primary\" href=\"http:\/\/americanboard.org\/Subjects\/chemistry\/patterns-in-space\">\u2b05 Previous Lesson<\/a>\u00a0<a class=\"button\" href=\"http:\/\/americanboard.org\/Subjects\/chemistry\/chemical-naming-and-structure\">Workshop Index<\/a>\u00a0<a class=\"button button-primary\" href=\"http:\/\/americanboard.org\/Subjects\/chemistry\/vsepr\">Next Lesson \u27a1<\/a><\/p>\n<\/div>\n<p><a class=\"backtotop\" href=\"#title\">Back to Top<\/a><\/p>\n<\/section>\n<\/div>\n","protected":false},"excerpt":{"rendered":"<p>\u2b05 Previous Lesson\u00a0Workshop Index\u00a0Next Lesson \u27a1 Chemical Naming and Structure: Lewis Structures Objective We will review drawing Lewis structures for molecules and ions, including those that follow the Octet Rule and those that do not. Previously we covered&#8230; Ionic bonding occurs between a metal atom and a nonmetal atom, due to their large difference in [&hellip;]<\/p>\n","protected":false},"author":1,"featured_media":0,"parent":0,"menu_order":0,"comment_status":"closed","ping_status":"closed","template":"","meta":{"footnotes":""},"class_list":["post-243","page","type-page","status-publish","hentry"],"_links":{"self":[{"href":"https:\/\/americanboard.org\/Subjects\/chemistry\/wp-json\/wp\/v2\/pages\/243","targetHints":{"allow":["GET"]}}],"collection":[{"href":"https:\/\/americanboard.org\/Subjects\/chemistry\/wp-json\/wp\/v2\/pages"}],"about":[{"href":"https:\/\/americanboard.org\/Subjects\/chemistry\/wp-json\/wp\/v2\/types\/page"}],"author":[{"embeddable":true,"href":"https:\/\/americanboard.org\/Subjects\/chemistry\/wp-json\/wp\/v2\/users\/1"}],"replies":[{"embeddable":true,"href":"https:\/\/americanboard.org\/Subjects\/chemistry\/wp-json\/wp\/v2\/comments?post=243"}],"version-history":[{"count":15,"href":"https:\/\/americanboard.org\/Subjects\/chemistry\/wp-json\/wp\/v2\/pages\/243\/revisions"}],"predecessor-version":[{"id":841,"href":"https:\/\/americanboard.org\/Subjects\/chemistry\/wp-json\/wp\/v2\/pages\/243\/revisions\/841"}],"wp:attachment":[{"href":"https:\/\/americanboard.org\/Subjects\/chemistry\/wp-json\/wp\/v2\/media?parent=243"}],"curies":[{"name":"wp","href":"https:\/\/api.w.org\/{rel}","templated":true}]}}