{"id":2484,"date":"2018-04-11T23:53:00","date_gmt":"2018-04-12T03:53:00","guid":{"rendered":"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/chapter\/covalent-bonds\/"},"modified":"2018-06-23T00:08:20","modified_gmt":"2018-06-23T04:08:20","slug":"covalent-bonds","status":"publish","type":"chapter","link":"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/chapter\/covalent-bonds\/","title":{"raw":"9.5 Covalent Bonds and Lewis Structures","rendered":"9.5 Covalent Bonds and Lewis Structures"},"content":{"raw":"<div class=\"section\" id=\"ball-ch09_s03\" lang=\"en\">\r\n<div class=\"learning_objectives editable block\" id=\"ball-ch09_s03_n01\">\r\n<div class=\"bcc-box bcc-highlight\">\r\n<h3>Learning Objectives<\/h3>\r\nBy the end of this section, you will be able to:\r\n<ul>\r\n \t<li>Define <em>covalent bond<\/em>.<\/li>\r\n \t<li>Illustrate covalent bond formation with Lewis electron dot diagrams.<\/li>\r\n \t<li>Draw Lewis structures depicting the bonding in simple molecules<\/li>\r\n<\/ul>\r\n<\/div>\r\n<\/div>\r\n<p id=\"ball-ch09_s03_p01\" class=\"para editable block\">Ionic bonding typically occurs when it is easy for one atom to lose one or more electrons and another atom to gain one or more electrons. However, some atoms won\u2019t give up or gain electrons easily. Yet they still participate in compound formation. How?<\/p>\r\n<p id=\"ball-ch09_s03_p02\" class=\"para editable block\">There is another mechanism for obtaining a complete valence shell: <em class=\"emphasis\">sharing<\/em> electrons. When electrons are shared between two atoms, they make a bond called a <span class=\"margin_term\"><a class=\"glossterm\">covalent bond<\/a><\/span>.<\/p>\r\n<p id=\"ball-ch09_s03_p03\" class=\"para editable block\">Let us illustrate a covalent bond by using H atoms, with the understanding that H atoms need only two electrons to fill the 1<em class=\"emphasis\">s<\/em> subshell. Each H atom starts with a single electron in its valence shell:<a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/H-H.png\"><img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/H-H-1.png\" alt=\"H-H\" width=\"400\" height=\"40\" class=\"wp-image-4428 aligncenter\" \/><\/a>The two H atoms can share their electrons:<\/p>\r\n\r\n<div class=\"informalfigure large block\">\r\n<p class=\"para editable block\"><a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/H-H-2.png\"><img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/H-H-2-1.png\" alt=\"H-H-2\" width=\"400\" height=\"40\" class=\"wp-image-4429 aligncenter\" \/><\/a>We can use circles to show that each H atom has two electrons around the nucleus, completely filling each atom\u2019s valence shell:<\/p>\r\n\r\n<div class=\"informalfigure large block\">\r\n<p class=\"para editable block\"><a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/H-H-3.png\"><img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/H-H-3-1.png\" alt=\"H-H-3\" width=\"400\" height=\"80\" class=\"wp-image-4430 aligncenter\" \/><\/a>Because each H atom has a filled valence shell, this bond is stable, and we have made a diatomic hydrogen molecule. (This explains why hydrogen is one of the diatomic elements.) For simplicity\u2019s sake, it is not unusual to represent the covalent bond with a dash, instead of with two dots:<\/p>\r\n\r\n<div class=\"informalfigure large block\">\r\n<p style=\"text-align: center\"><span class=\"informalequation block\"><span class=\"mathphrase\">H\u2013H<\/span><\/span><\/p>\r\n<p id=\"ball-ch09_s03_p07\" class=\"para editable block\">Because two atoms are sharing one pair of electrons, this covalent bond is called a <span class=\"margin_term\"><a class=\"glossterm\">single bond<\/a><\/span>.<\/p>\r\n<p id=\"ball-ch09_s03_p08\" class=\"para editable block\">As another example, consider fluorine. F atoms have seven electrons in their valence shell:<a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/F-F.png\"><img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/F-F-1.png\" alt=\"F-F\" width=\"400\" height=\"40\" class=\"wp-image-4431 aligncenter\" \/><\/a>These two atoms can do the same thing that the H atoms did; they share their unpaired electrons to make a covalent bond.<\/p>\r\n\r\n<div class=\"informalfigure large block\">\r\n<p class=\"para editable block\"><a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/F-F-2.png\"><img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/F-F-2-1.png\" alt=\"F-F-2\" width=\"400\" height=\"40\" class=\"wp-image-4432 aligncenter\" \/><\/a>Note that each F atom has a complete octet around it now:<\/p>\r\n\r\n<div class=\"informalfigure large block\">\r\n<p class=\"para editable block\"><a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/F-F-3.png\"><img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/F-F-3-1.png\" alt=\"F-F-3\" width=\"400\" height=\"80\" class=\"wp-image-4433 aligncenter\" \/><\/a>We can also write this using a dash to represent the shared electron pair:<\/p>\r\n\r\n<div class=\"informalfigure large block\">\r\n<p class=\"para editable block\"><a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/F-F-4.png\"><img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/F-F-4-1.png\" alt=\"F-F-4\" width=\"400\" height=\"40\" class=\"wp-image-4434 aligncenter\" \/><\/a>There are two different types of electrons in the fluorine diatomic molecule. The <span class=\"margin_term\"><a class=\"glossterm\">bonding electron pair<\/a><\/span>\u00a0makes the covalent bond. Each F atom has three other pairs of electrons that do not participate in the bonding; they are called <span class=\"margin_term\"><a class=\"glossterm\">lone electron pairs<\/a><\/span>. Each F atom has one bonding pair and three lone pairs of electrons.<\/p>\r\n\r\n<div class=\"informalfigure large block\">\r\n<p id=\"ball-ch09_s03_p13\" class=\"para editable block\">Covalent bonds can be made between different elements as well. One example is HF. Each atom starts out with an odd number of electrons in its valence shell:<a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/H-F.png\"><img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/H-F-1.png\" alt=\"H-F\" width=\"400\" height=\"40\" class=\"wp-image-4437 aligncenter\" \/><\/a>The two atoms can share their unpaired electrons to make a covalent bond:<\/p>\r\n\r\n<div class=\"informalfigure large block\">\r\n<p class=\"para editable block\"><a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/H-F-2.png\"><img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/H-F-2-1.png\" alt=\"H-F-2\" width=\"400\" height=\"40\" class=\"wp-image-4438 aligncenter\" \/><\/a>We note that the H atom has a full valence shell with two electrons, while the F atom has a complete octet of electrons.<\/p>\r\n\r\n<div class=\"informalfigure large block\">\r\n<div class=\"textbox shaded\">\r\n<h3 class=\"title\">Example 1<\/h3>\r\n<p id=\"ball-ch09_s03_p16\" class=\"para\">Use Lewis electron dot diagrams to illustrate the covalent bond formation in HBr.<\/p>\r\n&nbsp;\r\n<p class=\"simpara\"><strong>Solution<\/strong><\/p>\r\n<p id=\"ball-ch09_s03_p17\" class=\"para\">HBr is very similar to HF, except that it has Br instead of F. The atoms are as follows:<a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/H-Br.png\"><img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/H-Br-1.png\" alt=\"H-Br\" width=\"400\" height=\"40\" class=\"alignnone wp-image-4439 aligncenter\" \/><\/a><span style=\"font-size: 1em\">The two atoms can share their unpaired electron:<\/span><span style=\"font-size: 1em\"><\/span><\/p>\r\n\r\n<div class=\"informalfigure large\">\r\n<p class=\"para\"><a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/H-Br-2.png\"><img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/H-Br-2-1.png\" alt=\"H-Br-2\" width=\"400\" height=\"40\" class=\"alignnone wp-image-4440 aligncenter\" \/><\/a><\/p>\r\n\r\n<div class=\"informalfigure large\">\r\n<p class=\"simpara\"><strong><em class=\"emphasis bolditalic\">Test Yourself<\/em><\/strong><\/p>\r\n<p id=\"ball-ch09_s03_p19\" class=\"para\">Use Lewis electron dot diagrams to illustrate the covalent bond formation in Cl<sub class=\"subscript\">2<\/sub>.<\/p>\r\n&nbsp;\r\n<p class=\"simpara\"><strong><em class=\"emphasis\">Answer<\/em><\/strong><a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/Cl-Cl.png\"><img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/Cl-Cl-1.png\" alt=\"Cl-Cl\" width=\"400\" height=\"40\" class=\"alignnone wp-image-4441 aligncenter\" \/><\/a><\/p>\r\n\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<p id=\"ball-ch09_s03_p20\" class=\"para editable block\">More than two atoms can participate in covalent bonding, although any given covalent bond will be between two atoms only. Consider H and O atoms:<a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/H-O.png\"><img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/H-O-1.png\" alt=\"H-O\" width=\"400\" height=\"40\" class=\"wp-image-4443 aligncenter\" \/><\/a>The H and O atoms can share an electron to form a covalent bond:<\/p>\r\n\r\n<div class=\"informalfigure large block\">\r\n<p class=\"para editable block\"><a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/H-O-2.png\"><img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/H-O-2-1.png\" alt=\"H-O-2\" width=\"400\" height=\"40\" class=\"wp-image-4444 aligncenter\" \/><\/a>The H atom has a complete valence shell. However, the O atom has only seven electrons around it, which is not a complete octet. We fix this by including a second H atom, whose single electron will make a second covalent bond with the O atom:<\/p>\r\n\r\n<div class=\"informalfigure large block\">\r\n<p class=\"para editable block\"><a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/H-O-3.png\"><img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/H-O-3-1.png\" alt=\"H-O-3\" width=\"400\" height=\"60\" class=\"wp-image-4445 aligncenter\" \/><\/a>(It does not matter on what side the second H atom is positioned.) Now the O atom has a complete octet around it, and each H atom has two electrons, filling its valence shell. This is how a water molecule, H<sub class=\"subscript\">2<\/sub>O, is made.<\/p>\r\n\r\n<div class=\"informalfigure large block\">\r\n<div class=\"textbox shaded\">\r\n<h3 class=\"title\">Example 2<\/h3>\r\n<p id=\"ball-ch09_s03_p24\" class=\"para\">Use a Lewis electron dot diagram to show the covalent bonding in NH<sub class=\"subscript\">3<\/sub>.<\/p>\r\n&nbsp;\r\n<p class=\"simpara\"><strong>Solution<\/strong><\/p>\r\n<p id=\"ball-ch09_s03_p25\" class=\"para\">The N atom has the following Lewis electron dot diagram:<a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/N.png\"><img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/N-1.png\" alt=\"N\" width=\"400\" height=\"40\" class=\"wp-image-4446 aligncenter\" \/><\/a><span style=\"font-size: 1em\">It has three unpaired electrons, each of which can make a covalent bond by sharing electrons with an H atom. The electron dot diagram of NH<\/span><sub class=\"subscript\">3<\/sub><span style=\"font-size: 1em\"> is as follows:<\/span><span style=\"font-size: 1em\"><\/span><\/p>\r\n\r\n<div class=\"informalfigure large\">\r\n<p class=\"para\"><a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/N-H.png\"><img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/N-H-1.png\" alt=\"N-H\" width=\"400\" height=\"60\" class=\"wp-image-4447 aligncenter\" \/><\/a><\/p>\r\n\r\n<div class=\"informalfigure large\">\r\n<p class=\"simpara\"><strong><em class=\"emphasis bolditalic\">Test Yourself<\/em><\/strong><\/p>\r\n<p id=\"ball-ch09_s03_p27\" class=\"para\">Use a Lewis electron dot diagram to show the covalent bonding in PCl<sub class=\"subscript\">3<\/sub>.<\/p>\r\n&nbsp;\r\n<p class=\"simpara\"><strong><em class=\"emphasis\">Answer<\/em><\/strong><a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/Cl-P.png\"><img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/Cl-P-1.png\" alt=\"Cl-P\" width=\"400\" height=\"60\" class=\"wp-image-4448 aligncenter\" \/><\/a><\/p>\r\n\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<section id=\"fs-idm44772400\">\r\n<h2>Lewis Structures<\/h2>\r\n<p id=\"fs-idp28276192\">We also use Lewis symbols to indicate the formation of covalent bonds, which are shown in <strong>Lewis structures<\/strong>, drawings that describe the bonding in molecules and polyatomic ions. For example, when two chlorine atoms form a chlorine molecule, they share one pair of electrons:<\/p>\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_C12dot_img-2.jpg\" alt=\"A Lewis dot diagram shows a reaction. Two chlorine symbols, each surrounded by seven dots are separated by a plus sign. The dots on the first atom are all black and the dots on the second atom are all read. The phrase, \u201cChlorine atoms\u201d is written below. A right-facing arrow points to two chlorine symbols, each with six dots surrounding their outer edges and a shared pair of dots in between. One of the shared dots is black and one is red. The phrase, \u201cChlorine molecule\u201d is written below.\" class=\"aligncenter\" width=\"324\" height=\"104\" \/>\r\n<p id=\"fs-idm75288528\">The Lewis structure indicates that each Cl atom has three pairs of electrons that are not used in bonding (called <strong>lone pairs<\/strong>) and one shared pair of electrons (written between the atoms). A dash (or line) is sometimes used to indicate a shared pair of electrons:<\/p>\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Cl2dash_img-2.jpg\" alt=\"Two Lewis structures are shown. The left-hand structure shows two H atoms connected by a single bond. The right-hand structure shows two C l atoms connected by a single bond and each surrounded by six dots.\" class=\"aligncenter\" width=\"224\" height=\"54\" \/>\r\n<p id=\"fs-idm97531888\">A single shared pair of electrons is called a <strong>single bond<\/strong>. Each Cl atom interacts with eight valence electrons: the six in the lone pairs and the two in the single bond.<\/p>\r\n\r\n<section id=\"fs-idp113967376\">\r\n<h2>The Octet Rule<\/h2>\r\n<p id=\"fs-idm16266560\">The other halogen molecules (F<sub>2<\/sub>, Br<sub>2<\/sub>, I<sub>2<\/sub>, and At<sub>2<\/sub>) form bonds like those in the chlorine molecule: one single bond between atoms and three lone pairs of electrons per atom. This allows each halogen atom to have a noble gas electron configuration. The tendency of main group atoms to form enough bonds to obtain eight valence electrons is known as the <strong>octet rule<\/strong>.<\/p>\r\n<p id=\"fs-idm45369344\">The number of bonds that an atom can form can often be predicted from the number of electrons needed to reach an octet (eight valence electrons); this is especially true of the nonmetals of the second period of the periodic table (C, N, O, and F). For example, each atom of a group 14 element has four electrons in its outermost shell and therefore requires four more electrons to reach an octet. These four electrons can be gained by forming four covalent bonds, as illustrated here for carbon in CCl<sub>4<\/sub> (carbon tetrachloride) and silicon in SiH<sub>4<\/sub> (silane). Because hydrogen only needs two electrons to fill its valence shell, it is an exception to the octet rule. The transition elements and inner transition elements also do not follow the octet rule:<\/p>\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_XY4struc_img-2.jpg\" alt=\"Two sets of Lewis dot structures are shown. The left structures depict five C l symbols in a cross shape with eight dots around each, the word \u201cor\u201d and the same five C l symbols, connected by four single bonds in a cross shape. The name \u201cCarbon tetrachloride\u201d is written below the structure. The right hand structures show a S i symbol, surrounded by eight dots and four H symbols in a cross shape. The word \u201cor\u201d separates this from an S i symbol with four single bonds connecting the four H symbols in a cross shape. The name \u201cSilane\u201d is written below these diagrams.\" class=\"aligncenter\" width=\"631\" height=\"187\" \/>\r\n<p id=\"fs-idm5630800\">Group 15 elements such as nitrogen have five valence electrons in the atomic Lewis symbol: one lone pair and three unpaired electrons. To obtain an octet, these atoms form three covalent bonds, as in NH<sub>3<\/sub> (ammonia). Oxygen and other atoms in group 16 obtain an octet by forming two covalent bonds:<\/p>\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_numbonds_img-2.jpg\" alt=\"Three Lewis structures labeled, \u201cAmmonia,\u201d \u201cWater,\u201d and \u201cHydrogen fluoride\u201d are shown. The left structure shows a nitrogen atom with a lone pair of electrons and single bonded to three hydrogen atoms. The middle structure shows an oxygen atom with two lone pairs of electrons and two singly-bonded hydrogen atoms. The right structure shows a hydrogen atom single bonded to a fluorine atom that has three lone pairs of electrons.\" class=\"aligncenter\" width=\"454\" height=\"158\" \/>\r\n\r\n<\/section><section id=\"fs-idm81855664\">\r\n<h2>Double and Triple Bonds<\/h2>\r\n<p id=\"fs-idp27538208\">As previously mentioned, when a pair of atoms shares one pair of electrons, we call this a single bond. However, a pair of atoms may need to share more than one pair of electrons in order to achieve the requisite octet. A <strong>double bond<\/strong> forms when two pairs of electrons are shared between a pair of atoms, as between the carbon and oxygen atoms in CH<sub>2<\/sub>O (formaldehyde) and between the two carbon atoms in C<sub>2<\/sub>H<sub>4<\/sub> (ethylene):<\/p>\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_DoubleBond_img-2.jpg\" alt=\"Two pairs of Lewis structures are shown. The left pair of structures shows a carbon atom forming single bonds to two hydrogen atoms. There are four electrons between the C atom and an O atom. The O atom also has two pairs of dots. The word \u201cor\u201d separates this structure from the same diagram, except this time there is a double bond between the C atom and O atom. The name, \u201cFormaldehyde\u201d is written below these structures. A right-facing arrow leads to two more structures. The left shows two C atoms with four dots in between them and each forming single bonds to two H atoms. The word \u201cor\u201d lies to the left of the second structure, which is the same except that the C atoms form double bonds with one another. The name, \u201cEthylene\u201d is written below these structures.\" width=\"655\" height=\"151\" class=\"aligncenter\" \/>\r\n<p id=\"fs-idp15965904\">A <strong>triple bond<\/strong> forms when three electron pairs are shared by a pair of atoms, as in carbon monoxide (CO) and the cyanide ion (CN<sup>\u2013<\/sup>):<\/p>\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_COCN_img-2.jpg\" alt=\"Two pairs of Lewis structures are shown and connected by a right-facing arrow. The left pair of structures show a C atom and an O atom with six dots in between them and a lone pair on each. The word \u201cor\u201d and the same structure with a triple bond in between the C atom and O atom also are shown. The name \u201cCarbon monoxide\u201d is written below this structure. The right pair of structures show a C atom and an N atom with six dots in between them and a lone pair on each. The word \u201cor\u201d and the same structure with a triple bond in between the C atom and N atom also are shown. The name \u201cCyanide ion\u201d is written below this structure.\" class=\"aligncenter\" width=\"654\" height=\"87\" \/>\r\n\r\n<\/section><\/section><section id=\"fs-idm53492672\">\r\n<h2>Writing Lewis Structures with the Octet Rule<\/h2>\r\n<p id=\"fs-idm12599184\">For very simple molecules and molecular ions, we can write the Lewis structures by merely pairing up the unpaired electrons on the constituent atoms. See these examples:<\/p>\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Unprelec_img-2.jpg\" alt=\"Three reactions are shown with Lewis dot diagrams. The first shows a hydrogen with one red dot, a plus sign and a bromine with seven dots, one of which is red, connected by a right-facing arrow to a hydrogen and bromine with a pair of red dots in between them. There are also three lone pairs on the bromine. The second reaction shows a hydrogen with a coefficient of two and one red dot, a plus sign, and a sulfur atom with six dots, two of which are red, connected by a right facing arrow to two hydrogen atoms and one sulfur atom. There are two red dots in between the two hydrogen atoms and the sulfur atom. Both pairs of these dots are red. The sulfur atom also has two lone pairs of dots. The third reaction shows two nitrogen atoms each with five dots, three of which are red, separated by a plus sign, and connected by a right-facing arrow to two nitrogen atoms with six red electron dots in between one another. Each nitrogen atom also has one lone pair of electrons.\" width=\"542\" height=\"227\" class=\"aligncenter\" \/>\r\n<p id=\"fs-idm49249888\">For more complicated molecules and molecular ions, it is helpful to follow the step-by-step procedure outlined here:<\/p>\r\n\r\n<ol id=\"fs-idp55317056\">\r\n \t<li>Determine the total number of valence (outer shell) electrons. For cations, subtract one electron for each positive charge. For anions, add one electron for each negative charge.<\/li>\r\n \t<li>Draw a skeleton structure of the molecule or ion, arranging the atoms around a central atom. (Generally, the least electronegative element should be placed in the center.) Connect each atom to the central atom with a single bond (one electron pair).<\/li>\r\n \t<li>Distribute the remaining electrons as lone pairs on the terminal atoms (except hydrogen), completing an octet around each atom.<\/li>\r\n \t<li>Place all remaining electrons on the central atom.<\/li>\r\n \t<li>Rearrange the electrons of the outer atoms to make multiple bonds with the central atom in order to obtain octets wherever possible.<\/li>\r\n<\/ol>\r\n<p id=\"fs-idm19702384\">Let us determine the Lewis structures of SiH<sub>4<\/sub>, CHO<sub>2<\/sub>\u2212, NO<sup>+<\/sup>, and OF<sub>2<\/sub> as examples in following this procedure:<\/p>\r\n\r\n<ol id=\"fs-idm8107808\">\r\n \t<li>Determine the total number of valence (outer shell) electrons in the molecule or ion.\r\n<ul id=\"fs-idp47568320\">\r\n \t<li>For a molecule, we add the number of valence electrons on each atom in the molecule:\r\n<div class=\"equation\" id=\"fs-idm31186656\" style=\"text-align: center\">$latex \\begin{array}{r r l} \\text{SiH}_4 &amp; &amp; \\\\[1em] &amp; \\text{Si: 4 valence electrons\/atom} \\times 1 \\;\\text{atom} &amp; = 4 \\\\[1em] \\rule[-0.5ex]{21em}{0.1ex}\\hspace{-21em} + &amp; \\text{H: 1 valence electron\/atom} \\times 4 \\;\\text{atoms} &amp; = 4 \\\\[1em] &amp; &amp; = 8 \\;\\text{valence electrons} \\end{array}$<\/div><\/li>\r\n \t<li>For a <em>negative ion<\/em>, such as CHO<sub>2<\/sub><sup>\u2212<\/sup>, we add the number of valence electrons on the atoms to the number of negative charges on the ion (one electron is gained for each single negative charge):\r\n<div class=\"equation\" id=\"fs-idm69212352\" style=\"text-align: center\">$latex \\begin{array}{r r l} {\\text{CHO}_2}^{-} &amp; &amp; \\\\[1em] &amp; \\text{C: 4 valence electrons\/atom} \\times 1 \\;\\text{atom} &amp; = 4 \\\\[1em] &amp; \\text{H: 1 valence electron\/atom} \\times 1 \\;\\text{atom} &amp; = 1 \\\\[1em] &amp; \\text{O: 6 valence electrons\/atom} \\times 2 \\;\\text{atoms} &amp; = 12 \\\\[1em] \\rule[-0.5ex]{21.5em}{0.1ex}\\hspace{-21.5em} + &amp; 1\\;\\text{additional electron} &amp; = 1 \\\\[1em] &amp; &amp; = 18 \\;\\text{valence electrons} \\end{array}$<\/div><\/li>\r\n \t<li>For a <em>positive ion<\/em>, such as NO<sup>+<\/sup>, we add the number of valence electrons on the atoms in the ion and then subtract the number of positive charges on the ion (one electron is lost for each single positive charge) from the total number of valence electrons:\r\n<div class=\"equation\" id=\"fs-idm16450944\" style=\"text-align: center\">$latex \\begin{array}{r r l} \\text{NO}^{+} &amp; &amp; \\\\[1em] &amp; \\text{N: 5 valence electrons\/atom} \\times 1 \\;\\text{atom} &amp; = 5 \\\\[1em] &amp; \\text{O: 6 valence electrons\/atom} \\times 1 \\;\\text{atom} &amp; = 6 \\\\[1em] \\rule[-0.5ex]{21em}{0.1ex}\\hspace{-21em} + &amp; -1 \\;\\text{electron (positive charge)} &amp; = -1 \\\\[1em] &amp; &amp; = 10 \\;\\text{valence electrons} \\end{array}$<\/div><\/li>\r\n \t<li>Since OF<sub>2<\/sub> is a neutral molecule, we simply add the number of valence electrons:\r\n<div class=\"equation\" id=\"fs-idm18474448\" style=\"text-align: center\">$latex \\begin{array}{r r l} \\text{OF}_{2} &amp; &amp; \\\\[1em] &amp; \\text{O: 6 valence electrons\/atom} \\times 1 \\;\\text{atom} &amp; = 6 \\\\[1em] \\rule[-0.5ex]{21em}{0.1ex}\\hspace{-21em} + &amp; \\text{F: 7 valence electrons\/atom} \\times 2 \\;\\text{atoms} &amp; = 14 \\\\[1em] &amp; &amp; = 20 \\;\\text{valence electrons} \\end{array}$<\/div><\/li>\r\n<\/ul>\r\n<\/li>\r\n \t<li>Draw a skeleton structure of the molecule or ion, arranging the atoms around a central atom and connecting each atom to the central atom with a single (one electron pair) bond. (Note that we denote ions with brackets around the structure, indicating the charge outside the brackets:)<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Singlebond_img-2.jpg\" alt=\"Four Lewis diagrams are shown. The first shows one silicon single boned to four hydrogen atoms. The second shows a carbon which forms a single bond with an oxygen and a hydrogen and a double bond with a second oxygen. This structure is surrounded by brackets and has a superscripted negative sign near the upper right corner. The third structure shows a nitrogen single bonded to an oxygen and surrounded by brackets with a superscripted plus sign in the upper right corner. The last structure shows two fluorine atoms single bonded to a central oxygen.\" width=\"649\" height=\"104\" class=\"aligncenter\" \/>When several arrangements of atoms are possible, as for CHO<sub>2<\/sub><sup>\u2212<\/sup>, we must use experimental evidence to choose the correct one. In general, the less electronegative elements are more likely to be central atoms. In CHO<sub>2<\/sub><sup>\u2212<\/sup>, the less electronegative carbon atom occupies the central position with the oxygen and hydrogen atoms surrounding it. Other examples include P in POCl<sub>3<\/sub>, S in SO<sub>2<\/sub>, and Cl in ClO<sub>4<\/sub><sup>\u2212<\/sup>. An exception is that hydrogen is almost never a central atom. As the most electronegative element, fluorine also cannot be a central atom.<\/li>\r\n \t<li>Distribute the remaining electrons as lone pairs on the terminal atoms (except hydrogen) to complete their valence shells with an octet of electrons.\r\n<ul id=\"fs-idp233104\">\r\n \t<li>There are no remaining electrons on SiH<sub>4<\/sub>, so it is unchanged:<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_SiH4_img-2.jpg\" alt=\"Four Lewis structures are shown. The first shows one silicon single boned to four hydrogen atoms. The second shows a carbon single bonded to two oxygen atoms that each have three lone pairs and single bonded to a hydrogen. This structure is surrounded by brackets and has a superscripted negative sign near the upper right corner. The third structure shows a nitrogen single bonded to an oxygen, each with three lone pairs of electrons. This structure is surrounded by brackets with a superscripted plus sign in the upper right corner. The last structure shows two fluorine atoms, each with three lone pairs of electrons, single bonded to a central oxygen.\" width=\"606\" height=\"93\" class=\"aligncenter\" \/><\/li>\r\n<\/ul>\r\n<\/li>\r\n \t<li>Place all remaining electrons on the central atom.\r\n<ul id=\"fs-idp28108576\">\r\n \t<li>For SiH<sub>4<\/sub>, CHO<sub>2<\/sub><sup>\u2212<\/sup>, and NO<sup>+<\/sup>, there are no remaining electrons; we already placed all of the electrons determined in Step 1.<\/li>\r\n \t<li>For OF<sub>2<\/sub>, we had 16 electrons remaining in Step 3, and we placed 12, leaving 4 to be placed on the central atom:<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_oxydiflor2_img-2.jpg\" alt=\"A Lewis structure shows two fluorine atoms, each with three lone pairs of electrons, single bonded to a central oxygen which has two lone pairs of electrons.\" width=\"131\" height=\"48\" class=\"aligncenter\" \/><\/li>\r\n<\/ul>\r\n<\/li>\r\n \t<li>Rearrange the electrons of the outer atoms to make multiple bonds with the central atom in order to obtain octets wherever possible.\r\n<ul id=\"fs-idm72332608\">\r\n \t<li>SiH<sub>4<\/sub>: Si already has an octet, so nothing needs to be done.<\/li>\r\n \t<li>CHO<sub>2<\/sub><sup>\u2212<\/sup>: We have distributed the valence electrons as lone pairs on the oxygen atoms, but the carbon atom lacks an octet:<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_formate2_img-2.jpg\" alt=\"Two Lewis diagrams are shown with the word \u201cgives\u201d in between them. The left diagram, surrounded by brackets and with a superscripted negative sign, shows a carbon atom single bonded to two oxygen atoms, each with three lone pairs of electrons. The carbon atom also forms a single bond with a hydrogen atom. A curved arrow points from a lone pair on one of the oxygen atoms to the carbon atom. The right diagram, surrounded by brackets and with a superscripted negative sign, shows a carbon atom single bonded to an oxygen atom with three lone pairs of electrons, double bonded to an oxygen atom with two lone pairs of electrons, and single bonded to a hydrogen atom.\" width=\"571\" height=\"123\" class=\"aligncenter\" \/><\/li>\r\n \t<li>NO<sup>+<\/sup>: For this ion, we added eight valence electrons, but neither atom has an octet. We cannot add any more electrons since we have already used the total that we found in Step 1, so we must move electrons to form a multiple bond:<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_nitrosoni2_img-2.jpg\" alt=\"Two Lewis diagrams are shown with the word \u201cgives\u201d in between them. The left diagram, surrounded by brackets and with a superscripted positive sign, shows a nitrogen atom single bonded to an oxygen atom, each with two lone pairs of electrons. The right diagram, surrounded by brackets and with a superscripted positive sign, shows a nitrogen atom double bonded to an oxygen atom. The nitrogen atom has two lone pairs of electrons and the oxygen atom has one.\" width=\"547\" height=\"64\" class=\"aligncenter\" \/>This still does not produce an octet, so we must move another pair, forming a triple bond:<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_nitrosoni3_img-2.jpg\" alt=\"A Lewis structure shows a nitrogen atom with one lone pair of electrons triple bonded to an oxygen with a lone pair of electrons. The structure is surrounded by brackets and has a superscripted positive sign.\" \/><\/li>\r\n \t<li>In OF<sub>2<\/sub>, each atom has an octet as drawn, so nothing changes.<\/li>\r\n<\/ul>\r\n<\/li>\r\n<\/ol>\r\n<p id=\"ball-ch09_s03_p49\" class=\"para editable block\">Polyatomic ions are bonded together with covalent bonds, as seen in the example of CHO<sub>2<\/sub>\u2212. \u00a0Because they are ions, however, they participate in ionic bonding with other ions. So both major types of bonding can occur at the same time.<\/p>\r\n\r\n<div class=\"textbox shaded\" id=\"fs-idm36798944\">\r\n<h3>Example 3<\/h3>\r\n<p id=\"fs-idm67097232\">NASA\u2019s Cassini-Huygens mission detected a large cloud of toxic hydrogen cyanide (HCN) on Titan, one of Saturn\u2019s moons. Titan also contains ethane (H<sub>3<\/sub>CCH<sub>3<\/sub>), acetylene (HCCH), and ammonia (NH<sub>3<\/sub>). What are the Lewis structures of these molecules?<\/p>\r\n&nbsp;\r\n<p id=\"fs-idp90944512\"><strong>Solution<\/strong><\/p>\r\n\r\n<ol id=\"fs-idm4698672\" class=\"stepwise\">\r\n \t<li><em>Calculate the number of valence electrons.<\/em>HCN: (1 \u00d7 1) + (4 \u00d7 1) + (5 \u00d7 1) = 10H<sub>3<\/sub>CCH<sub>3<\/sub>: (1 \u00d7 3) + (2 \u00d7 4) + (1 \u00d7 3) = 14HCCH: (1 \u00d7 1) + (2 \u00d7 4) + (1 \u00d7 1) = 10NH<sub>3<\/sub>: (5 \u00d7 1) + (3 \u00d7 1) = 8<\/li>\r\n \t<li><em>Draw a skeleton and connect the atoms with single bonds.<\/em> Remember that H is never a central atom:<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_ex070301_1_img-2.jpg\" alt=\"Four Lewis structures are shown. The first structure shows a carbon atom single bonded to a hydrogen atom and a nitrogen atom. The second structure shows two carbon atoms single bonded to one another. Each is single bonded to three hydrogen atoms. The third structure shows two carbon atoms single bonded to one another and each single bonded to one hydrogen atom. The fourth structure shows a nitrogen atom single bonded to three hydrogen atoms.\" \/><\/li>\r\n \t<li><em>Where needed, distribute electrons to the terminal atoms:<\/em><img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_ex070301_2_img-2.jpg\" alt=\"Four Lewis structures are shown. The first structure shows a carbon atom single bonded to a hydrogen atom and a nitrogen atom, which has three lone pairs of electrons. The second structure shows two carbon atoms single bonded to one another. Each is single bonded to three hydrogen atoms. The third structure shows two carbon atoms single bonded to one another and each single bonded to one hydrogen atom. The fourth structure shows a nitrogen atom single bonded to three hydrogen atoms.\" \/>HCN: six electrons placed on NH<sub>3<\/sub>CCH<sub>3<\/sub>: no electrons remainHCCH: no terminal atoms capable of accepting electronsNH<sub>3<\/sub>: no terminal atoms capable of accepting electrons<\/li>\r\n \t<li><em>Where needed, place remaining electrons on the central atom:<\/em><img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_ex070301_3_img-2.jpg\" alt=\"Four Lewis structures are shown. The first structure shows a carbon atom single bonded to a hydrogen atom and a nitrogen atom, which has three lone pairs of electrons. The second structure shows two carbon atoms single bonded to one another. Each is single bonded to three hydrogen atoms. The third structure shows two carbon atoms, each with a lone pair of electrons, single bonded to one another and each single bonded to one hydrogen atom. The fourth structure shows a nitrogen atom with a lone pair of electrons single bonded to three hydrogen atoms.\" \/>HCN: no electrons remainH<sub>3<\/sub>CCH<sub>3<\/sub>: no electrons remainHCCH: four electrons placed on carbonNH<sub>3<\/sub>: two electrons placed on nitrogen<\/li>\r\n \t<li><em>Where needed, rearrange electrons to form multiple bonds in order to obtain an octet on each atom:<\/em>HCN: form two more C\u2013N bondsH<sub>3<\/sub>CCH<sub>3<\/sub>: all atoms have the correct number of electronsHCCH: form a triple bond between the two carbon atomsNH<sub>3<\/sub>: all atoms have the correct number of electrons<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_ex070301_4_img-2.jpg\" alt=\"Four Lewis structures are shown. The first structure shows a carbon atom single bonded to a hydrogen atom and a nitrogen atom, which has three lone pairs of electrons. Two curved arrows point from the nitrogen to the carbon. Below this structure is the word \u201cgives\u201d and below that is the same structure, but this time there is a triple bond between the carbon and nitrogen. The second structure shows two carbons single bonded to one another and each single bonded to three hydrogen atoms. The third structure shows two carbon atoms, each with a lone pair of electrons, single bonded to one another and each single bonded to one hydrogen atom. Two curved arrows point from the carbon atoms to the space in between the two. Below this structure is the word \u201cgives\u201d and the same structure, but this time with a triple bond between the two carbons. The fourth structure shows a nitrogen atom with a lone pair of electrons single bonded to three hydrogen atoms.\" \/><\/li>\r\n<\/ol>\r\n&nbsp;\r\n<p id=\"fs-idm58848848\"><em><strong>Test yourself<\/strong><\/em>\r\nBoth carbon monoxide, CO, and carbon dioxide, CO<sub>2<\/sub>, are products of the combustion of fossil fuels. Both of these gases also cause problems: CO is toxic and CO<sub>2<\/sub> has been implicated in global climate change. What are the Lewis structures of these two molecules?<\/p>\r\n&nbsp;\r\n\r\n<em><strong>Answers<\/strong><\/em>\r\n\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_COCO2_img-2.jpg\" alt=\"Two Lewis structures are shown. The left shows a carbon triple bonded to an oxygen, each with a lone electron pair. The right structure shows a carbon double bonded to an oxygen on each side. Each oxygen has two lone pairs of electrons.\" \/>\r\n\r\n<\/div>\r\n<div class=\"textbox shaded\">\r\n<h3 class=\"title\">Example 4<\/h3>\r\n<p id=\"ball-ch09_s03_p42\" class=\"para\">What is the proper Lewis electron dot diagram for CO<sub class=\"subscript\">2<\/sub>?<\/p>\r\n&nbsp;\r\n<p class=\"simpara\"><strong>Solution<\/strong><\/p>\r\n<p id=\"ball-ch09_s03_p43\" class=\"para\">The central atom is a C atom, with O atoms as surrounding atoms. We have a total of 4 +\u00a06 +\u00a06 = 16 valence electrons. Following the rules for Lewis electron dot diagrams for compounds gives us<a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/C-O.png\"><img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/C-O-1.png\" alt=\"C-O\" width=\"400\" height=\"40\" class=\"wp-image-4457 aligncenter\" \/><\/a><span style=\"font-size: 1em\">The O atoms have complete octets around them, but the C atom has only four electrons around it. The way to solve this dilemma is to make a double bond between carbon and <\/span><em class=\"emphasis\" style=\"font-size: 1em\">each<\/em><span style=\"font-size: 1em\"> O atom:<\/span><span style=\"font-size: 1em\"><\/span><\/p>\r\n\r\n<div class=\"informalfigure large\">\r\n<p class=\"para\"><a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/C-O-2.png\"><img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/C-O-2-1.png\" alt=\"C-O-2\" width=\"400\" height=\"40\" class=\"wp-image-4458 aligncenter\" \/><\/a><span style=\"font-size: 1em\">Each O atom still has eight electrons around it, but now the C atom also has a complete octet. This is an acceptable Lewis electron dot diagram for CO<\/span><sub class=\"subscript\">2<\/sub><span style=\"font-size: 1em\">.<\/span><\/p>\r\n\r\n<div class=\"informalfigure large\">\r\n\r\n&nbsp;\r\n<p class=\"simpara\"><strong><em class=\"emphasis bolditalic\">Test Yourself<\/em><\/strong><\/p>\r\n<p id=\"ball-ch09_s03_p46\" class=\"para\">What is the proper Lewis electron dot diagram for carbonyl sulfide (COS)?<\/p>\r\n&nbsp;\r\n<p class=\"simpara\"><strong><em class=\"emphasis\">Answer<\/em><\/strong><a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/C-S-O.png\"><img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/C-S-O-1.png\" alt=\"C-S-O\" width=\"400\" height=\"40\" class=\"wp-image-4459 aligncenter\" \/><\/a><\/p>\r\n\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<div id=\"fs-idp177244128\" class=\"textbox shaded\">\r\n<h3 class=\"title\">Fullerene Chemistry<\/h3>\r\n<p id=\"fs-idm18160736\">Carbon soot has been known to man since prehistoric times, but it was not until fairly recently that the molecular structure of the main component of soot was discovered. In 1996, the Nobel Prize in Chemistry was awarded to Richard <strong class=\"no-emphasis\">Smalley<\/strong> (<a href=\"#CNX_Chem_07_03_Smalley\" class=\"autogenerated-content\">Figure 1<\/a>), Robert Curl, and Harold Kroto for their work in discovering a new form of carbon, the C<sub>60<\/sub> buckminsterfullerene molecule (<a href=\"https:\/\/opentextbc.ca\/chemistry\/chapter\/introduction-8\/#CNX_Chem_07_00_Bucky\" class=\"autogenerated-content\">Figure 1 in Chapter 8 Introduction<\/a>). An entire class of compounds, including spheres and tubes of various shapes, were discovered based on C<sub>60.<\/sub> This type of molecule, called a fullerene, shows promise in a variety of applications. Because of their size and shape, fullerenes can encapsulate other molecules, so they have shown potential in various applications from hydrogen storage to targeted drug delivery systems. They also possess unique electronic and optical properties that have been put to good use in solar powered devices and chemical sensors.<\/p>\r\n\r\n<figure id=\"CNX_Chem_07_03_Smalley\">\r\n\r\n[caption id=\"\" align=\"aligncenter\" width=\"650\"]<a href=\"https:\/\/opentextbc.ca\/chemistry\/wp-content\/uploads\/sites\/150\/2016\/05\/CNX_Chem_07_03_Smalley.jpg\"><img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Smalley-2.jpg\" alt=\"A photo of Richard Smalley is shown.\" width=\"650\" height=\"497\" \/><\/a> <strong>Figure 1.<\/strong> Richard Smalley (1943\u20132005), a professor of physics, chemistry, and astronomy at Rice University, was one of the leading advocates for fullerene chemistry. Upon his death in 2005, the US Senate honored him as the \u201cFather of Nanotechnology.\u201d (credit: United States Department of Energy)[\/caption]<\/figure>\r\n<\/div>\r\n<\/section><section id=\"fs-idm61779936\">\r\n<h2>Exceptions to the Octet Rule<\/h2>\r\n<p id=\"fs-idm1862800\">Many covalent molecules have central atoms that do not have eight electrons in their Lewis structures. These molecules fall into three categories:<\/p>\r\n\r\n<ul id=\"fs-idm41738208\">\r\n \t<li>Odd-electron molecules have an odd number of valence electrons, and therefore have an unpaired electron.<\/li>\r\n \t<li>Electron-deficient molecules have a central atom that has fewer electrons than needed for a noble gas configuration.<\/li>\r\n \t<li>Hypervalent molecules have a central atom that has more electrons than needed for a noble gas configuration.<\/li>\r\n<\/ul>\r\n<section id=\"fs-idm33391760\">Examples of these will be covered later chemistry courses.\r\n<div class=\"textbox shaded\">\r\n<h3 class=\"title\">Food and Drink App: Vitamins and Minerals<\/h3>\r\n<p id=\"ball-ch09_s03_p50\" class=\"para\">Vitamins are nutrients that our bodies need in small amounts but cannot synthesize; therefore, they must be obtained from the diet. The word <em class=\"emphasis\">vitamin<\/em> comes from \u201cvital amine\u201d because it was once thought that all these compounds had an amine group (NH<sub class=\"subscript\">2<\/sub>) in it. This is not actually true, but the name stuck anyway.<\/p>\r\n<p id=\"ball-ch09_s03_p51\" class=\"para\">All vitamins are covalently bonded molecules. Most of them are commonly named with a letter, although all of them also have formal chemical names. Thus vitamin A is also called retinol, vitamin C is called ascorbic acid, and vitamin E is called tocopherol. There is no single vitamin B; there is a group of substances called the <em class=\"emphasis\">B complex vitamins<\/em> that are all water soluble and participate in cell metabolism. If a diet is lacking in a vitamin, diseases such as scurvy or rickets develop. Luckily, all vitamins are available as supplements, so any dietary deficiency in a vitamin can be easily corrected.<\/p>\r\n<p id=\"ball-ch09_s03_p52\" class=\"para\">A mineral is any chemical element other than carbon, hydrogen, oxygen, or nitrogen that is needed by the body. Minerals that the body needs in quantity include sodium, potassium, magnesium, calcium, phosphorus, sulfur, and chlorine. Essential minerals that the body needs in tiny quantities (so-called <em class=\"emphasis\">trace elements<\/em>) include manganese, iron, cobalt, nickel, copper, zinc, molybdenum, selenium, and iodine. Minerals are also obtained from the diet. Interestingly, most minerals are consumed in ionic form, rather than as elements or from covalent molecules. Like vitamins, most minerals are available in pill form, so any deficiency can be compensated for by taking supplements.<\/p>\r\n<p class=\"para\"><a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/Nutrition-Facts.png\"><img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/Nutrition-Facts-1.png\" alt=\"Nutrition-Facts\" width=\"280\" height=\"567\" class=\"alignnone wp-image-4461\" \/><\/a><\/p>\r\n<strong>Figure 2.<\/strong> Vitamin and Mineral Supplements\r\n<div class=\"informalfigure medium\" id=\"ball-ch09_s03_f01\">\r\n<p class=\"para\">Every entry down through pantothenic acid is a vitamin, and everything from calcium and below is a mineral.<\/p>\r\n\r\n<\/div>\r\n<\/div>\r\n<\/section><\/section><section id=\"fs-idm103622960\" class=\"summary\">\r\n<h2>Key Concepts and Summary<\/h2>\r\n<p id=\"fs-idm18966656\">Valence electronic structures can be visualized by drawing Lewis symbols (for atoms and monatomic ions) and Lewis structures (for molecules and polyatomic ions). Lone pairs, unpaired electrons, and single, double, or triple bonds are used to indicate where the valence electrons are located around each atom in a Lewis structure. Most structures\u2014especially those containing second row elements\u2014obey the octet rule, in which every atom (except H) is surrounded by eight electrons. Exceptions to the octet rule occur for odd-electron molecules (free radicals), electron-deficient molecules, and hypervalent molecules.<\/p>\r\n\r\n<\/section><section id=\"fs-idm14528032\" class=\"exercises\">\r\n<div class=\"bcc-box bcc-info\">\r\n<h3>Exercises<\/h3>\r\n1. Write the Lewis symbols for each of the following ions:\r\n<p id=\"fs-idp177307680\">a) As<sup>3\u2013 \u00a0 \u00a0 \u00a0 \u00a0\u00a0<\/sup>b) I<sup>\u2013 \u00a0 \u00a0 \u00a0 \u00a0\u00a0<\/sup>c) Be<sup>2+ \u00a0 \u00a0 \u00a0 \u00a0\u00a0<\/sup>d) O<sup>2\u2013 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0<\/sup>e) Ga<sup>3+ \u00a0 \u00a0 \u00a0 \u00a0 \u00a0<\/sup>f) Li<sup>+ \u00a0 \u00a0 \u00a0 \u00a0\u00a0<\/sup>g) N<sup>3\u2013<\/sup><\/p>\r\n2. Write the Lewis symbols of the ions in each of the following ionic compounds and the Lewis symbols of the atom from which they are formed:\r\n<p id=\"fs-idm81853760\">a) MgS \u00a0 \u00a0 \u00a0 \u00a0\u00a0b) Al<sub>2<\/sub>O<sub>3 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0<\/sub>c) GaCl<sub>3 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0<\/sub>d) K<sub>2<\/sub>O \u00a0 \u00a0 \u00a0 \u00a0e) Li<sub>3<\/sub>N \u00a0 \u00a0 \u00a0 \u00a0\u00a0f) KF<\/p>\r\n3. Write the Lewis structure for the diatomic molecule P<sub>2<\/sub>, an unstable form of phosphorus found in high-temperature phosphorus vapor.\r\n\r\n4. Write Lewis structures for the following:\r\n<p id=\"fs-idm45128544\">a) O<sub>2 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0\u00a0<\/sub>b) H<sub>2<\/sub>CO \u00a0 \u00a0 \u00a0 \u00a0\u00a0c) AsF<sub>3 \u00a0 \u00a0 \u00a0 \u00a0\u00a0<\/sub>d) ClNO \u00a0 \u00a0 \u00a0 \u00a0\u00a0e) SiCl<sub>4<\/sub><\/p>\r\n<p id=\"fs-idm5797408\">f) H<sub>3<\/sub>O<sup>+ \u00a0 \u00a0 \u00a0 \u00a0<\/sup>g) NH<sub>4<\/sub><sup>+ \u00a0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0<\/sup>h) BF<sub>4<\/sub><sup>\u2212 \u00a0 \u00a0 \u00a0 \u00a0\u00a0<\/sup>i) HCCH \u00a0 \u00a0 \u00a0 \u00a0\u00a0j) ClCN \u00a0 \u00a0 \u00a0 \u00a0 \u00a0k) C<sub>2<\/sub><sup>2+<\/sup><\/p>\r\n5. Write Lewis structures for the following:\r\n<p id=\"fs-idm17992976\">a) \u00a0SeCl<sub>3<\/sub><sup>+ \u00a0 \u00a0 \u00a0 \u00a0 \u00a0<\/sup>b) Cl<sub>2<\/sub>BBCl<sub>2<\/sub> (contains a B\u2013B bond)<\/p>\r\n6. Correct the following statement: \u201cThe bonds in solid PbCl<sub>2<\/sub> are ionic; the bond in a HCl molecule is covalent. Thus, all of the valence electrons in PbCl<sub>2<\/sub> are located on the Cl<sup>\u2013<\/sup> ions, and all of the valence electrons in a HCl molecule are shared between the H and Cl atoms.\u201d\r\n\r\n7. Methanol, H<sub>3<\/sub>COH, is used as the fuel in some race cars. Ethanol, C<sub>2<\/sub>H<sub>5<\/sub>OH, is used extensively as motor fuel in Brazil. Both methanol and ethanol produce CO<sub>2<\/sub> and H<sub>2<\/sub>O when they burn. Write the chemical equations for these combustion reactions using Lewis structures instead of chemical formulas.\r\n\r\n8. Carbon tetrachloride was formerly used in fire extinguishers for electrical fires. It is no longer used for this purpose because of the formation of the toxic gas phosgene, Cl<sub>2<\/sub>CO. Write the Lewis structures for carbon tetrachloride and phosgene.\r\n\r\n9. The arrangement of atoms in several biologically important molecules is given here. Complete the Lewis structures of these molecules by adding multiple bonds and lone pairs. Do not add any more atoms.\r\n<p id=\"fs-idm80239808\">a) the amino acid serine:<\/p>\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Impbioa_img-2.jpg\" alt=\"A Lewis structure is shown. A nitrogen atom is single bonded to two hydrogen atoms and a carbon atom. The carbon atom is single bonded to a hydrogen atom and two other carbon atoms. One of these carbon atoms is single bonded to two hydrogen atoms and an oxygen atom. The oxygen atom is bonded to a hydrogen atom. The other carbon atom is single bonded to two oxygen atoms, one of which is bonded to a hydrogen atom.\" width=\"228\" height=\"181\" class=\"\" \/>\r\n<p id=\"fs-idm147749792\">b) urea:<\/p>\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Impbiob_img-2.jpg\" alt=\"A Lewis structure is shown. A nitrogen atom is single bonded to two hydrogen atoms and a carbon atom. The carbon atom is single bonded to an oxygen atom and another nitrogen atom. That nitrogen atom is then single bonded to two hydrogen atoms.\" width=\"196\" height=\"69\" class=\"\" \/>\r\n<p id=\"fs-idp28905856\">c) pyruvic acid:<\/p>\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Impbioc_img-2.jpg\" alt=\"A Lewis structure is shown. A carbon atom is single bonded to three hydrogen atoms and another carbon atom. The second carbon atom is single bonded to an oxygen atom and a third carbon atom. This carbon is then single bonded to two oxygen atoms, one of which is single bonded to a hydrogen atom.\" width=\"237\" height=\"111\" class=\"\" \/>\r\n<p id=\"fs-idm104403040\">d) uracil:<\/p>\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Impbiod_img-2.jpg\" alt=\"A Lewis hexagonal ring structure is shown. From the top of the ring (moving clockwise), three carbon atoms, one nitrogen atom, a carbon atom, and a nitrogen atom are single bonded to each another. The top carbon atom is single bonded to an oxygen atom. The second and third carbons and the nitrogen atom are each single bonded to a hydrogen atom. The next carbon atom is single bonded to an oxygen atom, and the last nitrogen atom is single bonded to a hydrogen atom.\" width=\"172\" height=\"196\" class=\"\" \/>\r\n<p id=\"fs-idp106603552\">e) carbonic acid:<\/p>\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Impbioe_img-2.jpg\" alt=\"A Lewis structure is shown. A carbon atom is single bonded to three oxygen atoms. Two of those oxygen atoms are each single bonded to a hydrogen atom.\" width=\"193\" height=\"68\" class=\"\" \/>\r\n\r\n10. A compound with a molar mass of about 42 g\/mol contains 85.7% carbon and 14.3% hydrogen by mass. Write the Lewis structure for a molecule of the compound.\r\n\r\n11. How are single, double, and triple bonds similar? How do they differ?\r\n\r\n<span style=\"font-size: 1em\">12. How many electrons will be in the valence shell of H atoms when it makes a covalent bond?<\/span>\r\n\r\n13. What is the Lewis electron dot diagram of I<sub class=\"subscript\">2<\/sub>? Circle the electrons around each atom to verify that each valence shell is filled.\r\n\r\n14. What is the Lewis electron dot diagram of NCl<sub class=\"subscript\">3<\/sub>? Circle the electrons around each atom to verify that each valence shell is filled.\r\n\r\n15. Draw the Lewis electron dot diagram for each substance. \u00a0\u00a0<span style=\"font-size: 1em\">a) \u00a0SF<\/span><sub class=\"subscript\">2 \u00a0 \u00a0 \u00a0 \u00a0\u00a0<\/sub><span style=\"font-size: 1em\">b) \u00a0BH<\/span><sub class=\"subscript\">4<\/sub><sup class=\"superscript\">\u2212<\/sup>\r\n\r\n<span style=\"font-size: 1em\">16. Draw the Lewis electron dot diagram for each substance. \u00a0\u00a0<\/span><span style=\"font-size: 1em\">a) \u00a0GeH<\/span><sub class=\"subscript\">4 \u00a0 \u00a0\u00a0<\/sub><span style=\"font-size: 1em\">b) \u00a0ClF<\/span>\r\n\r\n<span style=\"font-size: 1em\">17. Draw the Lewis electron dot diagram for each substance. Double or triple bonds may be needed. \u00a0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0\u00a0<\/span>\r\n\r\n<span style=\"font-size: 1em\">a) \u00a0SiO<\/span><sub class=\"subscript\">2 \u00a0 \u00a0<\/sub><span style=\"font-size: 1em\">b) \u00a0C<\/span><sub class=\"subscript\">2<\/sub><span style=\"font-size: 1em\">H<\/span><sub class=\"subscript\">4<\/sub><span style=\"font-size: 1em\"> (assume two central atoms)<\/span>\r\n\r\n<span style=\"font-size: 1em\">18. Draw the Lewis electron dot diagram for each substance. Double or triple bonds may be needed.<\/span>\r\n<div class=\"question\">\r\n\r\na) \u00a0CS<sub class=\"subscript\">2 \u00a0 \u00a0 \u00a0\u00a0<\/sub>b) \u00a0NH<sub class=\"subscript\">2<\/sub>CONH<sub class=\"subscript\">2<\/sub> (assume that the N and C atoms are the central atoms)\r\n\r\n<\/div>\r\n&nbsp;\r\n\r\n<strong>Answers<\/strong>\r\n<p id=\"fs-idm9743232\">1. a) eight electrons:\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Question1a_img-2.jpg\" alt=\"A Lewis dot diagram shows the symbol for arsenic, A s, surrounded by eight dots and a superscripted three negative sign.\" width=\"272\" height=\"51\" class=\"\" \/><\/p>\r\nb) eight electrons:\r\n\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Question1b_img-2.jpg\" alt=\"A Lewis dot diagram shows the symbol for iodine, I, surrounded by eight dots and a superscripted negative sign.\" width=\"272\" height=\"51\" class=\"\" \/>\r\n\r\nc) no electrons: \u00a0\u00a0Be<sup>2+<\/sup>\r\n\r\nd) eight electrons:\r\n\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Question1d_img-2.jpg\" alt=\"A Lewis dot diagram shows the symbol for oxygen, O, surrounded by eight dots and a superscripted two negative sign.\" width=\"272\" height=\"51\" class=\"\" \/>\r\n\r\ne) no electrons: \u00a0Ga<sup>3+<\/sup>\r\n\r\nf) no electrons:\u00a0Li<sup>+<\/sup>\r\n\r\ng) eight electrons:\r\n\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Question1g_img-2.jpg\" alt=\"A Lewis dot diagram shows the symbol for nitrogen, N, surrounded by eight dots and a superscripted three negative sign.\" width=\"266\" height=\"50\" class=\"\" \/>\r\n<p id=\"fs-idp56026944\">2. a)<\/p>\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Exercise3a_img-2.jpg\" alt=\"Two Lewis structures are shown. The left shows the symbol M g with a superscripted two positive sign while the right shows the symbol S surrounded by eight dots and a superscripted two negative sign.\" width=\"261\" height=\"49\" class=\"\" \/>\r\n\r\nb)\r\n\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Exercise3b_img-2.jpg\" alt=\"Two Lewis structures are shown. The left shows the symbol A l with a superscripted three positive sign while the right shows the symbol O surrounded by eight dots and a superscripted two negative sign.\" width=\"261\" height=\"49\" class=\"\" \/>\r\n\r\nc)\r\n\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Exercise3c_img-2.jpg\" alt=\"Two Lewis structures are shown. The left shows the symbol G a with a superscripted three positive sign while the right shows the symbol C l surrounded by eight dots and a superscripted negative sign.\" width=\"256\" height=\"48\" class=\"\" \/>\r\n\r\nd)\r\n\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Exercise3d_img-2.jpg\" alt=\"Two Lewis structures are shown. The left shows the symbol K with a superscripted positive sign while the right shows the symbol O surrounded by eight dots and a superscripted two negative sign.\" width=\"256\" height=\"48\" class=\"\" \/>\r\n\r\ne)\r\n\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Exercise3e_img-2.jpg\" alt=\"Two Lewis structures are shown. The left shows the symbol L i with a superscripted positive sign while the right shows the symbol N surrounded by eight dots and a superscripted three negative sign.\" width=\"250\" height=\"47\" class=\"\" \/>\r\n\r\nf)\r\n\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Exercise3f_img-2.jpg\" alt=\"Two Lewis structures are shown. The left shows the symbol K with a superscripted positive sign while the right shows the symbol F surrounded by eight dots and a superscripted negative sign.\" width=\"250\" height=\"47\" class=\"\" \/>\r\n\r\n3.\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Question5_img-2.jpg\" alt=\"A Lewis diagram shows two phosphorus atoms triple bonded together each with one lone electron pair.\" width=\"244\" height=\"27\" class=\"\" \/>\r\n<p id=\"fs-idm41104416\">4. a)\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Question7a_img-2.jpg\" alt=\"A Lewis structure shows two oxygen atoms double bonded together, and each has two lone pairs of electrons.\" width=\"97\" height=\"39\" class=\"\" \/><\/p>\r\nIn this case, the Lewis structure is inadequate to depict the fact that experimental studies have shown two unpaired electrons in each oxygen molecule<em>.<\/em>\r\n\r\nb)\r\n\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Question7b_img-2.jpg\" alt=\"A Lewis structure shows a carbon atom that is single bonded to two hydrogen atoms and double bonded to an oxygen atom. The oxygen atom has two lone pairs of electrons.\" width=\"117\" height=\"84\" class=\"\" \/>\r\n\r\nc)\r\n\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Question7c_img-2.jpg\" alt=\"A Lewis structure shows an arsenic atom single bonded to three fluorine atoms. Each fluorine atom has a lone pair of electrons.\" width=\"126\" height=\"83\" class=\"\" \/>\r\n\r\nd)\r\n\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Question7d_img-2.jpg\" alt=\"A Lewis structure shows a nitrogen atom with a lone pair of electrons single bonded to a chlorine atom that has three lone pairs of electrons. The nitrogen is also double bonded to an oxygen which has two lone pairs of electrons.\" width=\"273\" height=\"47\" class=\"\" \/>\r\n\r\ne)\r\n\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Question7e_img-2.jpg\" alt=\"A Lewis structure shows a silicon atom that is single bonded to four chlorine atoms. Each chlorine atom has three lone pairs of electrons.\" width=\"141\" height=\"139\" class=\"\" \/>\r\n\r\nf)\r\n\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Question7f_img-2.jpg\" alt=\"A Lewis structure shows an oxygen atom with a lone pair of electrons single bonded to three hydrogen atoms. The structure is surrounded by brackets with a superscripted positive sign.\" width=\"270\" height=\"104\" class=\"\" \/>\r\n\r\ng)\r\n\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Question7g_img-2.jpg\" alt=\"A Lewis structure shows a nitrogen atom single bonded to four hydrogen atoms. The structure is surrounded by brackets with a superscripted positive sign.\" width=\"254\" height=\"129\" class=\"\" \/>\r\n\r\nh)\r\n\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Question7h_img-2.jpg\" alt=\"A Lewis structure shows a boron atom single bonded to four fluorine atoms. Each fluorine atom has three lone pairs of electrons. The structure is surrounded by brackets with a superscripted negative sign.\" width=\"250\" height=\"127\" class=\"\" \/>\r\n\r\ni)\r\n\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Question7i_img-2.jpg\" alt=\"A Lewis structure shows two carbon atoms that are triple bonded together. Each carbon is also single bonded to a hydrogen atom.\" width=\"159\" height=\"29\" class=\"\" \/>\r\n\r\nj)\r\n\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Question7j_img-2.jpg\" alt=\"A Lewis structure shows a carbon atom that is triple bonded to a nitrogen atom that has one lone pair of electrons. The carbon is also single bonded to a chlorine atom that has three lone pairs of electrons.\" width=\"138\" height=\"49\" class=\"\" \/>\r\n\r\nk)\r\n\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Question7k_img-2.jpg\" alt=\"A Lewis structure shows two carbon atoms joined with a triple bond. A superscripted 2 positive sign lies to the right of the second carbon.\" width=\"86\" height=\"35\" class=\"\" \/>\r\n<p id=\"fs-idm123049856\">5. a) SeCl<sub>3<\/sub><sup>+<\/sup>:<\/p>\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Question9c_img-2.jpg\" alt=\"A Lewis structure shows a selenium atom with one lone pair of electrons single bonded to three chlorine atoms each with three lone pairs of electrons. The whole structure is surrounded by brackets.\" width=\"265\" height=\"102\" class=\"\" \/>\r\n\r\nb) Cl<sub>2<\/sub>BBCl<sub>2<\/sub>:\r\n\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Question9d_img-2.jpg\" alt=\"A Lewis structure shows two boron atoms that are single bonded together. Each is also single bonded to two chlorine atoms that both have three lone pairs of electrons.\" width=\"146\" height=\"109\" class=\"\" \/>\r\n<p id=\"fs-idm7862384\">6. Two valence electrons per Pb atom are transferred to Cl atoms; the resulting Pb<sup>2+<\/sup> ion has a 6<em>s<\/em><sup>2<\/sup> valence shell configuration. Two of the valence electrons in the HCl molecule are shared, and the other six are located on the Cl atom as lone pairs of electrons.<\/p>\r\n7.\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Question18_img-2.jpg\" alt=\"Two reactions are shown using Lewis structures. The top reaction shows a carbon atom, single bonded to three hydrogen atoms and single bonded to an oxygen atom with two lone pairs of electrons. The oxygen atom is also bonded to a hydrogen atom. This is followed by a plus sign and the number one point five, followed by two oxygen atoms bonded together with a double bond and each with two lone pairs of electrons. A right-facing arrow leads to a carbon atom that is double bonded to two oxygen atoms, each of which has two lone pairs of electrons. This structure is followed by a plus sign, a number two, and a structure made up of an oxygen with two lone pairs of electrons single bonded to two hydrogen atoms. The bottom reaction shows a carbon atom, single bonded to three hydrogen atoms and single bonded to another carbon atom. The second carbon atom is single bonded to two hydrogen atoms and one oxygen atom with two lone pairs of electrons. The oxygen atom is also bonded to a hydrogen atom. This is followed by a plus sign and the number three, followed by two oxygen atoms bonded together with a double bond. Each oxygen atom has two lone pairs of electrons. A right-facing arrow leads to a number two and a carbon atom that is double bonded to two oxygen atoms, each of which has two lone pairs of electrons. This structure is followed by a plus sign, a number three, and a structure made up of an oxygen with two lone pairs of electrons single bonded to two hydrogen atoms.\" \/>\r\n\r\n8.\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Carbtet_img-2.jpg\" alt=\"Two Lewis structures are shown. The left depicts a carbon atom single bonded to four chlorine atoms, each with three lone pairs of electrons. The right shows a carbon atom double bonded to an oxygen atom that has two lone pairs of electrons. The carbon atom is also single bonded to two chlorine atoms, each of which has three lone pairs of electrons.\" width=\"273\" height=\"134\" class=\"\" \/>\r\n<p id=\"fs-idm17951840\">9. a)\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Impbioansa_img-2.jpg\" alt=\"A Lewis structure is shown. A nitrogen atom is single bonded to two hydrogen atoms and a carbon atom. The carbon atom is single bonded to a hydrogen atom and two other carbon atoms. One of these carbon atoms is single bonded to two hydrogen atoms and an oxygen atom. The oxygen atom is bonded to a hydrogen atom. The other carbon is single bonded to two oxygen atoms, one of which is bonded to a hydrogen atom. The oxygen atoms have two lone pairs of electron dots, and the nitrogen atom has one lone pair of electron dots.\" width=\"232\" height=\"194\" class=\"\" \/><\/p>\r\nb)\r\n\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Impbioansb_img-2.jpg\" alt=\"A Lewis structure is shown. A nitrogen atom is single bonded to two hydrogen atoms and a carbon atom. The carbon atom is single bonded to an oxygen atom and one nitrogen atom. That nitrogen atom is then single bonded to two hydrogen atoms. The oxygen atom has two lone pairs of electron dots, and the nitrogen atoms have one lone pair of electron dots each.\" width=\"203\" height=\"94\" class=\"\" \/>\r\n\r\nc)\r\n\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Impbioansc_img-2.jpg\" alt=\"A Lewis structure is shown. A carbon atom is single bonded to three hydrogen atoms and a carbon atom. The carbon atom is single bonded to an oxygen atom and a third carbon atom. This carbon is then single bonded to two oxygen atoms, one of which is single bonded to a hydrogen atom. Each oxygen atom has two lone pairs of electron dots.\" width=\"243\" height=\"122\" class=\"\" \/>\r\n\r\nd)\r\n\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Impbioansd_img-2.jpg\" alt=\"A Lewis hexagonal ring structure is shown. From the top of the ring, three carbon atoms, one nitrogen atom, a carbon atom and a nitrogen atom are single bonded to one another. The top carbon is single bonded to an oxygen, the second and third carbons and the nitrogen atom are each single bonded to a hydrogen atom. The next carbon is single bonded to an oxygen atom and the last nitrogen is single bonded to a hydrogen atom. The oxygen atoms have two lone pairs of electron dots, and the nitrogen atoms have one lone pair of electron dots.\" width=\"187\" height=\"210\" class=\"\" \/>\r\n\r\ne)\r\n\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Impbioanse_img-2.jpg\" alt=\"A Lewis structure is shown. A carbon atom is single bonded to three oxygen atoms. Two of those oxygen atoms are each single bonded to a hydrogen atom. Each oxygen atom has two lone pairs of electron dots.\" width=\"205\" height=\"95\" class=\"\" \/>\r\n\r\n10.\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Exercise25_img-2.jpg\" alt=\"A Lewis structure is shown. A carbon atom is single bonded to three hydrogen atoms and another carbon atom. The second carbon atom is double bonded to another carbon atom and single bonded to a hydrogen atom. The last carbon is single bonded to two hydrogen atoms.\" width=\"269\" height=\"119\" class=\"\" \/>\r\n<p id=\"fs-idp205940864\">11. Each bond includes a sharing of electrons between atoms. Two electrons are shared in a single bond; four electrons are shared in a double bond; and six electrons are shared in a triple bond.<\/p>\r\n12. two\r\n\r\n13.\r\n\r\n<strong><a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/I-I.png\"><img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/I-I-1.png\" alt=\"I-I\" width=\"343\" height=\"84\" class=\"alignnone wp-image-4462\" \/><\/a><\/strong>\r\n\r\n14.\r\n\r\n<strong><a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/N-Cl.png\"><img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/N-Cl-1.png\" alt=\"N-Cl\" width=\"343\" height=\"102\" class=\"alignnone wp-image-4463\" \/><\/a><\/strong>\r\n\r\n15.\r\n<p style=\"padding-left: 30px\">a) \u00a0\u00a0<a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/S-F.png\"><img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/S-F-1.png\" alt=\"S-F\" width=\"339\" height=\"101\" class=\"alignnone wp-image-4464\" \/><\/a><\/p>\r\n<p style=\"padding-left: 30px\">b) \u00a0\u00a0<a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/H-B.png\"><img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/H-B-1.png\" alt=\"H-B\" width=\"333\" height=\"135\" class=\"alignnone wp-image-4465\" \/><\/a><\/p>\r\n16.\r\n<p style=\"padding-left: 30px\">a) \u00a0\u00a0<a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/H-Ge.png\"><img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/H-Ge-1.png\" alt=\"H-Ge\" width=\"331\" height=\"134\" class=\"alignnone wp-image-4467\" \/><\/a><\/p>\r\n<p style=\"padding-left: 30px\">b) \u00a0\u00a0<a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/Cl-F.png\"><img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/Cl-F-1.png\" alt=\"Cl-F\" width=\"343\" height=\"84\" class=\"alignnone wp-image-4468\" \/><\/a><\/p>\r\n17.\r\n<p style=\"padding-left: 30px\">a) \u00a0\u00a0<a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/Si-O.png\"><img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/Si-O-1.png\" alt=\"Si-O\" width=\"327\" height=\"80\" class=\"alignnone wp-image-4469\" \/><\/a><\/p>\r\n<p style=\"padding-left: 30px\">b) \u00a0\u00a0<a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/C-H.png\"><img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/C-H-1.png\" alt=\"C-H\" width=\"286\" height=\"116\" class=\"alignnone wp-image-4470\" \/><\/a><\/p>\r\n18.\r\n<p style=\"padding-left: 30px\">a) \u00a0\u00a0<a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/S-C.png\"><img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/S-C-1.png\" alt=\"S-C\" width=\"327\" height=\"80\" class=\"alignnone wp-image-4471\" \/><\/a><\/p>\r\n<p style=\"padding-left: 30px\">b) \u00a0\u00a0<a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/C-N-H-O.png\"><img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/C-N-H-O-1.png\" alt=\"C-N-H-O\" width=\"300\" height=\"154\" class=\"alignnone wp-image-4472\" \/><\/a><\/p>\r\n\r\n<\/div>\r\n<\/section>\r\n<div>\r\n<h2>Glossary<\/h2>\r\n<strong>double bond:\u00a0<\/strong>covalent bond in which two pairs of electrons are shared between two atoms\r\n\r\n<strong>free radical:\u00a0<\/strong>molecule that contains an odd number of electrons\r\n\r\n<strong>hypervalent molecule:\u00a0<\/strong>molecule containing at least one main group element that has more than eight electrons in its valence shell\r\n\r\n<strong>Lewis structure:\u00a0<\/strong>diagram showing lone pairs and bonding pairs of electrons in a molecule or an ion\r\n\r\n<strong>Lewis symbol:\u00a0<\/strong>symbol for an element or monatomic ion that uses a dot to represent each valence electron in the element or ion\r\n\r\n<strong>lone pair:\u00a0<\/strong>two (a pair of) valence electrons that are not used to form a covalent bond\r\n\r\n<strong>octet rule:\u00a0<\/strong>guideline that states main group atoms will form structures in which eight valence electrons interact with each nucleus, counting bonding electrons as interacting with both atoms connected by the bond\r\n\r\n<strong>single bond:\u00a0<\/strong>bond in which a single pair of electrons is shared between two atoms\r\n\r\n<strong>triple bond:\u00a0<\/strong>bond in which three pairs of electrons are shared between two atoms\r\n<dl id=\"fs-idm68093504\" class=\"definition\">\r\n \t<dt><a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/C-N-H-O.png\"><\/a><\/dt>\r\n<\/dl>\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>\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<\/div>","rendered":"<div class=\"section\" id=\"ball-ch09_s03\" lang=\"en\">\n<div class=\"learning_objectives editable block\" id=\"ball-ch09_s03_n01\">\n<div class=\"bcc-box bcc-highlight\">\n<h3>Learning Objectives<\/h3>\n<p>By the end of this section, you will be able to:<\/p>\n<ul>\n<li>Define <em>covalent bond<\/em>.<\/li>\n<li>Illustrate covalent bond formation with Lewis electron dot diagrams.<\/li>\n<li>Draw Lewis structures depicting the bonding in simple molecules<\/li>\n<\/ul>\n<\/div>\n<\/div>\n<p id=\"ball-ch09_s03_p01\" class=\"para editable block\">Ionic bonding typically occurs when it is easy for one atom to lose one or more electrons and another atom to gain one or more electrons. However, some atoms won\u2019t give up or gain electrons easily. Yet they still participate in compound formation. How?<\/p>\n<p id=\"ball-ch09_s03_p02\" class=\"para editable block\">There is another mechanism for obtaining a complete valence shell: <em class=\"emphasis\">sharing<\/em> electrons. When electrons are shared between two atoms, they make a bond called a <span class=\"margin_term\"><a class=\"glossterm\">covalent bond<\/a><\/span>.<\/p>\n<p id=\"ball-ch09_s03_p03\" class=\"para editable block\">Let us illustrate a covalent bond by using H atoms, with the understanding that H atoms need only two electrons to fill the 1<em class=\"emphasis\">s<\/em> subshell. Each H atom starts with a single electron in its valence shell:<a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/H-H.png\"><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/H-H-1.png\" alt=\"H-H\" width=\"400\" height=\"40\" class=\"wp-image-4428 aligncenter\" \/><\/a>The two H atoms can share their electrons:<\/p>\n<div class=\"informalfigure large block\">\n<p class=\"para editable block\"><a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/H-H-2.png\"><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/H-H-2-1.png\" alt=\"H-H-2\" width=\"400\" height=\"40\" class=\"wp-image-4429 aligncenter\" \/><\/a>We can use circles to show that each H atom has two electrons around the nucleus, completely filling each atom\u2019s valence shell:<\/p>\n<div class=\"informalfigure large block\">\n<p class=\"para editable block\"><a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/H-H-3.png\"><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/H-H-3-1.png\" alt=\"H-H-3\" width=\"400\" height=\"80\" class=\"wp-image-4430 aligncenter\" \/><\/a>Because each H atom has a filled valence shell, this bond is stable, and we have made a diatomic hydrogen molecule. (This explains why hydrogen is one of the diatomic elements.) For simplicity\u2019s sake, it is not unusual to represent the covalent bond with a dash, instead of with two dots:<\/p>\n<div class=\"informalfigure large block\">\n<p style=\"text-align: center\"><span class=\"informalequation block\"><span class=\"mathphrase\">H\u2013H<\/span><\/span><\/p>\n<p id=\"ball-ch09_s03_p07\" class=\"para editable block\">Because two atoms are sharing one pair of electrons, this covalent bond is called a <span class=\"margin_term\"><a class=\"glossterm\">single bond<\/a><\/span>.<\/p>\n<p id=\"ball-ch09_s03_p08\" class=\"para editable block\">As another example, consider fluorine. F atoms have seven electrons in their valence shell:<a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/F-F.png\"><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/F-F-1.png\" alt=\"F-F\" width=\"400\" height=\"40\" class=\"wp-image-4431 aligncenter\" \/><\/a>These two atoms can do the same thing that the H atoms did; they share their unpaired electrons to make a covalent bond.<\/p>\n<div class=\"informalfigure large block\">\n<p class=\"para editable block\"><a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/F-F-2.png\"><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/F-F-2-1.png\" alt=\"F-F-2\" width=\"400\" height=\"40\" class=\"wp-image-4432 aligncenter\" \/><\/a>Note that each F atom has a complete octet around it now:<\/p>\n<div class=\"informalfigure large block\">\n<p class=\"para editable block\"><a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/F-F-3.png\"><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/F-F-3-1.png\" alt=\"F-F-3\" width=\"400\" height=\"80\" class=\"wp-image-4433 aligncenter\" \/><\/a>We can also write this using a dash to represent the shared electron pair:<\/p>\n<div class=\"informalfigure large block\">\n<p class=\"para editable block\"><a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/F-F-4.png\"><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/F-F-4-1.png\" alt=\"F-F-4\" width=\"400\" height=\"40\" class=\"wp-image-4434 aligncenter\" \/><\/a>There are two different types of electrons in the fluorine diatomic molecule. The <span class=\"margin_term\"><a class=\"glossterm\">bonding electron pair<\/a><\/span>\u00a0makes the covalent bond. Each F atom has three other pairs of electrons that do not participate in the bonding; they are called <span class=\"margin_term\"><a class=\"glossterm\">lone electron pairs<\/a><\/span>. Each F atom has one bonding pair and three lone pairs of electrons.<\/p>\n<div class=\"informalfigure large block\">\n<p id=\"ball-ch09_s03_p13\" class=\"para editable block\">Covalent bonds can be made between different elements as well. One example is HF. Each atom starts out with an odd number of electrons in its valence shell:<a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/H-F.png\"><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/H-F-1.png\" alt=\"H-F\" width=\"400\" height=\"40\" class=\"wp-image-4437 aligncenter\" \/><\/a>The two atoms can share their unpaired electrons to make a covalent bond:<\/p>\n<div class=\"informalfigure large block\">\n<p class=\"para editable block\"><a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/H-F-2.png\"><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/H-F-2-1.png\" alt=\"H-F-2\" width=\"400\" height=\"40\" class=\"wp-image-4438 aligncenter\" \/><\/a>We note that the H atom has a full valence shell with two electrons, while the F atom has a complete octet of electrons.<\/p>\n<div class=\"informalfigure large block\">\n<div class=\"textbox shaded\">\n<h3 class=\"title\">Example 1<\/h3>\n<p id=\"ball-ch09_s03_p16\" class=\"para\">Use Lewis electron dot diagrams to illustrate the covalent bond formation in HBr.<\/p>\n<p>&nbsp;<\/p>\n<p class=\"simpara\"><strong>Solution<\/strong><\/p>\n<p id=\"ball-ch09_s03_p17\" class=\"para\">HBr is very similar to HF, except that it has Br instead of F. The atoms are as follows:<a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/H-Br.png\"><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/H-Br-1.png\" alt=\"H-Br\" width=\"400\" height=\"40\" class=\"alignnone wp-image-4439 aligncenter\" \/><\/a><span style=\"font-size: 1em\">The two atoms can share their unpaired electron:<\/span><span style=\"font-size: 1em\"><\/span><\/p>\n<div class=\"informalfigure large\">\n<p class=\"para\"><a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/H-Br-2.png\"><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/H-Br-2-1.png\" alt=\"H-Br-2\" width=\"400\" height=\"40\" class=\"alignnone wp-image-4440 aligncenter\" \/><\/a><\/p>\n<div class=\"informalfigure large\">\n<p class=\"simpara\"><strong><em class=\"emphasis bolditalic\">Test Yourself<\/em><\/strong><\/p>\n<p id=\"ball-ch09_s03_p19\" class=\"para\">Use Lewis electron dot diagrams to illustrate the covalent bond formation in Cl<sub class=\"subscript\">2<\/sub>.<\/p>\n<p>&nbsp;<\/p>\n<p class=\"simpara\"><strong><em class=\"emphasis\">Answer<\/em><\/strong><a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/Cl-Cl.png\"><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/Cl-Cl-1.png\" alt=\"Cl-Cl\" width=\"400\" height=\"40\" class=\"alignnone wp-image-4441 aligncenter\" \/><\/a><\/p>\n<\/div>\n<\/div>\n<\/div>\n<p id=\"ball-ch09_s03_p20\" class=\"para editable block\">More than two atoms can participate in covalent bonding, although any given covalent bond will be between two atoms only. Consider H and O atoms:<a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/H-O.png\"><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/H-O-1.png\" alt=\"H-O\" width=\"400\" height=\"40\" class=\"wp-image-4443 aligncenter\" \/><\/a>The H and O atoms can share an electron to form a covalent bond:<\/p>\n<div class=\"informalfigure large block\">\n<p class=\"para editable block\"><a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/H-O-2.png\"><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/H-O-2-1.png\" alt=\"H-O-2\" width=\"400\" height=\"40\" class=\"wp-image-4444 aligncenter\" \/><\/a>The H atom has a complete valence shell. However, the O atom has only seven electrons around it, which is not a complete octet. We fix this by including a second H atom, whose single electron will make a second covalent bond with the O atom:<\/p>\n<div class=\"informalfigure large block\">\n<p class=\"para editable block\"><a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/H-O-3.png\"><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/H-O-3-1.png\" alt=\"H-O-3\" width=\"400\" height=\"60\" class=\"wp-image-4445 aligncenter\" \/><\/a>(It does not matter on what side the second H atom is positioned.) Now the O atom has a complete octet around it, and each H atom has two electrons, filling its valence shell. This is how a water molecule, H<sub class=\"subscript\">2<\/sub>O, is made.<\/p>\n<div class=\"informalfigure large block\">\n<div class=\"textbox shaded\">\n<h3 class=\"title\">Example 2<\/h3>\n<p id=\"ball-ch09_s03_p24\" class=\"para\">Use a Lewis electron dot diagram to show the covalent bonding in NH<sub class=\"subscript\">3<\/sub>.<\/p>\n<p>&nbsp;<\/p>\n<p class=\"simpara\"><strong>Solution<\/strong><\/p>\n<p id=\"ball-ch09_s03_p25\" class=\"para\">The N atom has the following Lewis electron dot diagram:<a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/N.png\"><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/N-1.png\" alt=\"N\" width=\"400\" height=\"40\" class=\"wp-image-4446 aligncenter\" \/><\/a><span style=\"font-size: 1em\">It has three unpaired electrons, each of which can make a covalent bond by sharing electrons with an H atom. The electron dot diagram of NH<\/span><sub class=\"subscript\">3<\/sub><span style=\"font-size: 1em\"> is as follows:<\/span><span style=\"font-size: 1em\"><\/span><\/p>\n<div class=\"informalfigure large\">\n<p class=\"para\"><a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/N-H.png\"><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/N-H-1.png\" alt=\"N-H\" width=\"400\" height=\"60\" class=\"wp-image-4447 aligncenter\" \/><\/a><\/p>\n<div class=\"informalfigure large\">\n<p class=\"simpara\"><strong><em class=\"emphasis bolditalic\">Test Yourself<\/em><\/strong><\/p>\n<p id=\"ball-ch09_s03_p27\" class=\"para\">Use a Lewis electron dot diagram to show the covalent bonding in PCl<sub class=\"subscript\">3<\/sub>.<\/p>\n<p>&nbsp;<\/p>\n<p class=\"simpara\"><strong><em class=\"emphasis\">Answer<\/em><\/strong><a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/Cl-P.png\"><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/Cl-P-1.png\" alt=\"Cl-P\" width=\"400\" height=\"60\" class=\"wp-image-4448 aligncenter\" \/><\/a><\/p>\n<\/div>\n<\/div>\n<\/div>\n<section id=\"fs-idm44772400\">\n<h2>Lewis Structures<\/h2>\n<p id=\"fs-idp28276192\">We also use Lewis symbols to indicate the formation of covalent bonds, which are shown in <strong>Lewis structures<\/strong>, drawings that describe the bonding in molecules and polyatomic ions. For example, when two chlorine atoms form a chlorine molecule, they share one pair of electrons:<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_C12dot_img-2.jpg\" alt=\"A Lewis dot diagram shows a reaction. Two chlorine symbols, each surrounded by seven dots are separated by a plus sign. The dots on the first atom are all black and the dots on the second atom are all read. The phrase, \u201cChlorine atoms\u201d is written below. A right-facing arrow points to two chlorine symbols, each with six dots surrounding their outer edges and a shared pair of dots in between. One of the shared dots is black and one is red. The phrase, \u201cChlorine molecule\u201d is written below.\" class=\"aligncenter\" width=\"324\" height=\"104\" \/><\/p>\n<p id=\"fs-idm75288528\">The Lewis structure indicates that each Cl atom has three pairs of electrons that are not used in bonding (called <strong>lone pairs<\/strong>) and one shared pair of electrons (written between the atoms). A dash (or line) is sometimes used to indicate a shared pair of electrons:<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Cl2dash_img-2.jpg\" alt=\"Two Lewis structures are shown. The left-hand structure shows two H atoms connected by a single bond. The right-hand structure shows two C l atoms connected by a single bond and each surrounded by six dots.\" class=\"aligncenter\" width=\"224\" height=\"54\" \/><\/p>\n<p id=\"fs-idm97531888\">A single shared pair of electrons is called a <strong>single bond<\/strong>. Each Cl atom interacts with eight valence electrons: the six in the lone pairs and the two in the single bond.<\/p>\n<section id=\"fs-idp113967376\">\n<h2>The Octet Rule<\/h2>\n<p id=\"fs-idm16266560\">The other halogen molecules (F<sub>2<\/sub>, Br<sub>2<\/sub>, I<sub>2<\/sub>, and At<sub>2<\/sub>) form bonds like those in the chlorine molecule: one single bond between atoms and three lone pairs of electrons per atom. This allows each halogen atom to have a noble gas electron configuration. The tendency of main group atoms to form enough bonds to obtain eight valence electrons is known as the <strong>octet rule<\/strong>.<\/p>\n<p id=\"fs-idm45369344\">The number of bonds that an atom can form can often be predicted from the number of electrons needed to reach an octet (eight valence electrons); this is especially true of the nonmetals of the second period of the periodic table (C, N, O, and F). For example, each atom of a group 14 element has four electrons in its outermost shell and therefore requires four more electrons to reach an octet. These four electrons can be gained by forming four covalent bonds, as illustrated here for carbon in CCl<sub>4<\/sub> (carbon tetrachloride) and silicon in SiH<sub>4<\/sub> (silane). Because hydrogen only needs two electrons to fill its valence shell, it is an exception to the octet rule. The transition elements and inner transition elements also do not follow the octet rule:<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_XY4struc_img-2.jpg\" alt=\"Two sets of Lewis dot structures are shown. The left structures depict five C l symbols in a cross shape with eight dots around each, the word \u201cor\u201d and the same five C l symbols, connected by four single bonds in a cross shape. The name \u201cCarbon tetrachloride\u201d is written below the structure. The right hand structures show a S i symbol, surrounded by eight dots and four H symbols in a cross shape. The word \u201cor\u201d separates this from an S i symbol with four single bonds connecting the four H symbols in a cross shape. The name \u201cSilane\u201d is written below these diagrams.\" class=\"aligncenter\" width=\"631\" height=\"187\" \/><\/p>\n<p id=\"fs-idm5630800\">Group 15 elements such as nitrogen have five valence electrons in the atomic Lewis symbol: one lone pair and three unpaired electrons. To obtain an octet, these atoms form three covalent bonds, as in NH<sub>3<\/sub> (ammonia). Oxygen and other atoms in group 16 obtain an octet by forming two covalent bonds:<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_numbonds_img-2.jpg\" alt=\"Three Lewis structures labeled, \u201cAmmonia,\u201d \u201cWater,\u201d and \u201cHydrogen fluoride\u201d are shown. The left structure shows a nitrogen atom with a lone pair of electrons and single bonded to three hydrogen atoms. The middle structure shows an oxygen atom with two lone pairs of electrons and two singly-bonded hydrogen atoms. The right structure shows a hydrogen atom single bonded to a fluorine atom that has three lone pairs of electrons.\" class=\"aligncenter\" width=\"454\" height=\"158\" \/><\/p>\n<\/section>\n<section id=\"fs-idm81855664\">\n<h2>Double and Triple Bonds<\/h2>\n<p id=\"fs-idp27538208\">As previously mentioned, when a pair of atoms shares one pair of electrons, we call this a single bond. However, a pair of atoms may need to share more than one pair of electrons in order to achieve the requisite octet. A <strong>double bond<\/strong> forms when two pairs of electrons are shared between a pair of atoms, as between the carbon and oxygen atoms in CH<sub>2<\/sub>O (formaldehyde) and between the two carbon atoms in C<sub>2<\/sub>H<sub>4<\/sub> (ethylene):<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_DoubleBond_img-2.jpg\" alt=\"Two pairs of Lewis structures are shown. The left pair of structures shows a carbon atom forming single bonds to two hydrogen atoms. There are four electrons between the C atom and an O atom. The O atom also has two pairs of dots. The word \u201cor\u201d separates this structure from the same diagram, except this time there is a double bond between the C atom and O atom. The name, \u201cFormaldehyde\u201d is written below these structures. A right-facing arrow leads to two more structures. The left shows two C atoms with four dots in between them and each forming single bonds to two H atoms. The word \u201cor\u201d lies to the left of the second structure, which is the same except that the C atoms form double bonds with one another. The name, \u201cEthylene\u201d is written below these structures.\" width=\"655\" height=\"151\" class=\"aligncenter\" \/><\/p>\n<p id=\"fs-idp15965904\">A <strong>triple bond<\/strong> forms when three electron pairs are shared by a pair of atoms, as in carbon monoxide (CO) and the cyanide ion (CN<sup>\u2013<\/sup>):<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_COCN_img-2.jpg\" alt=\"Two pairs of Lewis structures are shown and connected by a right-facing arrow. The left pair of structures show a C atom and an O atom with six dots in between them and a lone pair on each. The word \u201cor\u201d and the same structure with a triple bond in between the C atom and O atom also are shown. The name \u201cCarbon monoxide\u201d is written below this structure. The right pair of structures show a C atom and an N atom with six dots in between them and a lone pair on each. The word \u201cor\u201d and the same structure with a triple bond in between the C atom and N atom also are shown. The name \u201cCyanide ion\u201d is written below this structure.\" class=\"aligncenter\" width=\"654\" height=\"87\" \/><\/p>\n<\/section>\n<\/section>\n<section id=\"fs-idm53492672\">\n<h2>Writing Lewis Structures with the Octet Rule<\/h2>\n<p id=\"fs-idm12599184\">For very simple molecules and molecular ions, we can write the Lewis structures by merely pairing up the unpaired electrons on the constituent atoms. See these examples:<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Unprelec_img-2.jpg\" alt=\"Three reactions are shown with Lewis dot diagrams. The first shows a hydrogen with one red dot, a plus sign and a bromine with seven dots, one of which is red, connected by a right-facing arrow to a hydrogen and bromine with a pair of red dots in between them. There are also three lone pairs on the bromine. The second reaction shows a hydrogen with a coefficient of two and one red dot, a plus sign, and a sulfur atom with six dots, two of which are red, connected by a right facing arrow to two hydrogen atoms and one sulfur atom. There are two red dots in between the two hydrogen atoms and the sulfur atom. Both pairs of these dots are red. The sulfur atom also has two lone pairs of dots. The third reaction shows two nitrogen atoms each with five dots, three of which are red, separated by a plus sign, and connected by a right-facing arrow to two nitrogen atoms with six red electron dots in between one another. Each nitrogen atom also has one lone pair of electrons.\" width=\"542\" height=\"227\" class=\"aligncenter\" \/><\/p>\n<p id=\"fs-idm49249888\">For more complicated molecules and molecular ions, it is helpful to follow the step-by-step procedure outlined here:<\/p>\n<ol id=\"fs-idp55317056\">\n<li>Determine the total number of valence (outer shell) electrons. For cations, subtract one electron for each positive charge. For anions, add one electron for each negative charge.<\/li>\n<li>Draw a skeleton structure of the molecule or ion, arranging the atoms around a central atom. (Generally, the least electronegative element should be placed in the center.) Connect each atom to the central atom with a single bond (one electron pair).<\/li>\n<li>Distribute the remaining electrons as lone pairs on the terminal atoms (except hydrogen), completing an octet around each atom.<\/li>\n<li>Place all remaining electrons on the central atom.<\/li>\n<li>Rearrange the electrons of the outer atoms to make multiple bonds with the central atom in order to obtain octets wherever possible.<\/li>\n<\/ol>\n<p id=\"fs-idm19702384\">Let us determine the Lewis structures of SiH<sub>4<\/sub>, CHO<sub>2<\/sub>\u2212, NO<sup>+<\/sup>, and OF<sub>2<\/sub> as examples in following this procedure:<\/p>\n<ol id=\"fs-idm8107808\">\n<li>Determine the total number of valence (outer shell) electrons in the molecule or ion.\n<ul id=\"fs-idp47568320\">\n<li>For a molecule, we add the number of valence electrons on each atom in the molecule:\n<div class=\"equation\" id=\"fs-idm31186656\" style=\"text-align: center\">[latex]\\begin{array}{r r l} \\text{SiH}_4 & & \\\\[1em] & \\text{Si: 4 valence electrons\/atom} \\times 1 \\;\\text{atom} & = 4 \\\\[1em] \\rule[-0.5ex]{21em}{0.1ex}\\hspace{-21em} + & \\text{H: 1 valence electron\/atom} \\times 4 \\;\\text{atoms} & = 4 \\\\[1em] & & = 8 \\;\\text{valence electrons} \\end{array}[\/latex]<\/div>\n<\/li>\n<li>For a <em>negative ion<\/em>, such as CHO<sub>2<\/sub><sup>\u2212<\/sup>, we add the number of valence electrons on the atoms to the number of negative charges on the ion (one electron is gained for each single negative charge):\n<div class=\"equation\" id=\"fs-idm69212352\" style=\"text-align: center\">[latex]\\begin{array}{r r l} {\\text{CHO}_2}^{-} & & \\\\[1em] & \\text{C: 4 valence electrons\/atom} \\times 1 \\;\\text{atom} & = 4 \\\\[1em] & \\text{H: 1 valence electron\/atom} \\times 1 \\;\\text{atom} & = 1 \\\\[1em] & \\text{O: 6 valence electrons\/atom} \\times 2 \\;\\text{atoms} & = 12 \\\\[1em] \\rule[-0.5ex]{21.5em}{0.1ex}\\hspace{-21.5em} + & 1\\;\\text{additional electron} & = 1 \\\\[1em] & & = 18 \\;\\text{valence electrons} \\end{array}[\/latex]<\/div>\n<\/li>\n<li>For a <em>positive ion<\/em>, such as NO<sup>+<\/sup>, we add the number of valence electrons on the atoms in the ion and then subtract the number of positive charges on the ion (one electron is lost for each single positive charge) from the total number of valence electrons:\n<div class=\"equation\" id=\"fs-idm16450944\" style=\"text-align: center\">[latex]\\begin{array}{r r l} \\text{NO}^{+} & & \\\\[1em] & \\text{N: 5 valence electrons\/atom} \\times 1 \\;\\text{atom} & = 5 \\\\[1em] & \\text{O: 6 valence electrons\/atom} \\times 1 \\;\\text{atom} & = 6 \\\\[1em] \\rule[-0.5ex]{21em}{0.1ex}\\hspace{-21em} + & -1 \\;\\text{electron (positive charge)} & = -1 \\\\[1em] & & = 10 \\;\\text{valence electrons} \\end{array}[\/latex]<\/div>\n<\/li>\n<li>Since OF<sub>2<\/sub> is a neutral molecule, we simply add the number of valence electrons:\n<div class=\"equation\" id=\"fs-idm18474448\" style=\"text-align: center\">[latex]\\begin{array}{r r l} \\text{OF}_{2} & & \\\\[1em] & \\text{O: 6 valence electrons\/atom} \\times 1 \\;\\text{atom} & = 6 \\\\[1em] \\rule[-0.5ex]{21em}{0.1ex}\\hspace{-21em} + & \\text{F: 7 valence electrons\/atom} \\times 2 \\;\\text{atoms} & = 14 \\\\[1em] & & = 20 \\;\\text{valence electrons} \\end{array}[\/latex]<\/div>\n<\/li>\n<\/ul>\n<\/li>\n<li>Draw a skeleton structure of the molecule or ion, arranging the atoms around a central atom and connecting each atom to the central atom with a single (one electron pair) bond. (Note that we denote ions with brackets around the structure, indicating the charge outside the brackets:)<img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Singlebond_img-2.jpg\" alt=\"Four Lewis diagrams are shown. The first shows one silicon single boned to four hydrogen atoms. The second shows a carbon which forms a single bond with an oxygen and a hydrogen and a double bond with a second oxygen. This structure is surrounded by brackets and has a superscripted negative sign near the upper right corner. The third structure shows a nitrogen single bonded to an oxygen and surrounded by brackets with a superscripted plus sign in the upper right corner. The last structure shows two fluorine atoms single bonded to a central oxygen.\" width=\"649\" height=\"104\" class=\"aligncenter\" \/>When several arrangements of atoms are possible, as for CHO<sub>2<\/sub><sup>\u2212<\/sup>, we must use experimental evidence to choose the correct one. In general, the less electronegative elements are more likely to be central atoms. In CHO<sub>2<\/sub><sup>\u2212<\/sup>, the less electronegative carbon atom occupies the central position with the oxygen and hydrogen atoms surrounding it. Other examples include P in POCl<sub>3<\/sub>, S in SO<sub>2<\/sub>, and Cl in ClO<sub>4<\/sub><sup>\u2212<\/sup>. An exception is that hydrogen is almost never a central atom. As the most electronegative element, fluorine also cannot be a central atom.<\/li>\n<li>Distribute the remaining electrons as lone pairs on the terminal atoms (except hydrogen) to complete their valence shells with an octet of electrons.\n<ul id=\"fs-idp233104\">\n<li>There are no remaining electrons on SiH<sub>4<\/sub>, so it is unchanged:<img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_SiH4_img-2.jpg\" alt=\"Four Lewis structures are shown. The first shows one silicon single boned to four hydrogen atoms. The second shows a carbon single bonded to two oxygen atoms that each have three lone pairs and single bonded to a hydrogen. This structure is surrounded by brackets and has a superscripted negative sign near the upper right corner. The third structure shows a nitrogen single bonded to an oxygen, each with three lone pairs of electrons. This structure is surrounded by brackets with a superscripted plus sign in the upper right corner. The last structure shows two fluorine atoms, each with three lone pairs of electrons, single bonded to a central oxygen.\" width=\"606\" height=\"93\" class=\"aligncenter\" \/><\/li>\n<\/ul>\n<\/li>\n<li>Place all remaining electrons on the central atom.\n<ul id=\"fs-idp28108576\">\n<li>For SiH<sub>4<\/sub>, CHO<sub>2<\/sub><sup>\u2212<\/sup>, and NO<sup>+<\/sup>, there are no remaining electrons; we already placed all of the electrons determined in Step 1.<\/li>\n<li>For OF<sub>2<\/sub>, we had 16 electrons remaining in Step 3, and we placed 12, leaving 4 to be placed on the central atom:<img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_oxydiflor2_img-2.jpg\" alt=\"A Lewis structure shows two fluorine atoms, each with three lone pairs of electrons, single bonded to a central oxygen which has two lone pairs of electrons.\" width=\"131\" height=\"48\" class=\"aligncenter\" \/><\/li>\n<\/ul>\n<\/li>\n<li>Rearrange the electrons of the outer atoms to make multiple bonds with the central atom in order to obtain octets wherever possible.\n<ul id=\"fs-idm72332608\">\n<li>SiH<sub>4<\/sub>: Si already has an octet, so nothing needs to be done.<\/li>\n<li>CHO<sub>2<\/sub><sup>\u2212<\/sup>: We have distributed the valence electrons as lone pairs on the oxygen atoms, but the carbon atom lacks an octet:<img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_formate2_img-2.jpg\" alt=\"Two Lewis diagrams are shown with the word \u201cgives\u201d in between them. The left diagram, surrounded by brackets and with a superscripted negative sign, shows a carbon atom single bonded to two oxygen atoms, each with three lone pairs of electrons. The carbon atom also forms a single bond with a hydrogen atom. A curved arrow points from a lone pair on one of the oxygen atoms to the carbon atom. The right diagram, surrounded by brackets and with a superscripted negative sign, shows a carbon atom single bonded to an oxygen atom with three lone pairs of electrons, double bonded to an oxygen atom with two lone pairs of electrons, and single bonded to a hydrogen atom.\" width=\"571\" height=\"123\" class=\"aligncenter\" \/><\/li>\n<li>NO<sup>+<\/sup>: For this ion, we added eight valence electrons, but neither atom has an octet. We cannot add any more electrons since we have already used the total that we found in Step 1, so we must move electrons to form a multiple bond:<img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_nitrosoni2_img-2.jpg\" alt=\"Two Lewis diagrams are shown with the word \u201cgives\u201d in between them. The left diagram, surrounded by brackets and with a superscripted positive sign, shows a nitrogen atom single bonded to an oxygen atom, each with two lone pairs of electrons. The right diagram, surrounded by brackets and with a superscripted positive sign, shows a nitrogen atom double bonded to an oxygen atom. The nitrogen atom has two lone pairs of electrons and the oxygen atom has one.\" width=\"547\" height=\"64\" class=\"aligncenter\" \/>This still does not produce an octet, so we must move another pair, forming a triple bond:<img decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_nitrosoni3_img-2.jpg\" alt=\"A Lewis structure shows a nitrogen atom with one lone pair of electrons triple bonded to an oxygen with a lone pair of electrons. The structure is surrounded by brackets and has a superscripted positive sign.\" \/><\/li>\n<li>In OF<sub>2<\/sub>, each atom has an octet as drawn, so nothing changes.<\/li>\n<\/ul>\n<\/li>\n<\/ol>\n<p id=\"ball-ch09_s03_p49\" class=\"para editable block\">Polyatomic ions are bonded together with covalent bonds, as seen in the example of CHO<sub>2<\/sub>\u2212. \u00a0Because they are ions, however, they participate in ionic bonding with other ions. So both major types of bonding can occur at the same time.<\/p>\n<div class=\"textbox shaded\" id=\"fs-idm36798944\">\n<h3>Example 3<\/h3>\n<p id=\"fs-idm67097232\">NASA\u2019s Cassini-Huygens mission detected a large cloud of toxic hydrogen cyanide (HCN) on Titan, one of Saturn\u2019s moons. Titan also contains ethane (H<sub>3<\/sub>CCH<sub>3<\/sub>), acetylene (HCCH), and ammonia (NH<sub>3<\/sub>). What are the Lewis structures of these molecules?<\/p>\n<p>&nbsp;<\/p>\n<p id=\"fs-idp90944512\"><strong>Solution<\/strong><\/p>\n<ol id=\"fs-idm4698672\" class=\"stepwise\">\n<li><em>Calculate the number of valence electrons.<\/em>HCN: (1 \u00d7 1) + (4 \u00d7 1) + (5 \u00d7 1) = 10H<sub>3<\/sub>CCH<sub>3<\/sub>: (1 \u00d7 3) + (2 \u00d7 4) + (1 \u00d7 3) = 14HCCH: (1 \u00d7 1) + (2 \u00d7 4) + (1 \u00d7 1) = 10NH<sub>3<\/sub>: (5 \u00d7 1) + (3 \u00d7 1) = 8<\/li>\n<li><em>Draw a skeleton and connect the atoms with single bonds.<\/em> Remember that H is never a central atom:<img decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_ex070301_1_img-2.jpg\" alt=\"Four Lewis structures are shown. The first structure shows a carbon atom single bonded to a hydrogen atom and a nitrogen atom. The second structure shows two carbon atoms single bonded to one another. Each is single bonded to three hydrogen atoms. The third structure shows two carbon atoms single bonded to one another and each single bonded to one hydrogen atom. The fourth structure shows a nitrogen atom single bonded to three hydrogen atoms.\" \/><\/li>\n<li><em>Where needed, distribute electrons to the terminal atoms:<\/em><img decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_ex070301_2_img-2.jpg\" alt=\"Four Lewis structures are shown. The first structure shows a carbon atom single bonded to a hydrogen atom and a nitrogen atom, which has three lone pairs of electrons. The second structure shows two carbon atoms single bonded to one another. Each is single bonded to three hydrogen atoms. The third structure shows two carbon atoms single bonded to one another and each single bonded to one hydrogen atom. The fourth structure shows a nitrogen atom single bonded to three hydrogen atoms.\" \/>HCN: six electrons placed on NH<sub>3<\/sub>CCH<sub>3<\/sub>: no electrons remainHCCH: no terminal atoms capable of accepting electronsNH<sub>3<\/sub>: no terminal atoms capable of accepting electrons<\/li>\n<li><em>Where needed, place remaining electrons on the central atom:<\/em><img decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_ex070301_3_img-2.jpg\" alt=\"Four Lewis structures are shown. The first structure shows a carbon atom single bonded to a hydrogen atom and a nitrogen atom, which has three lone pairs of electrons. The second structure shows two carbon atoms single bonded to one another. Each is single bonded to three hydrogen atoms. The third structure shows two carbon atoms, each with a lone pair of electrons, single bonded to one another and each single bonded to one hydrogen atom. The fourth structure shows a nitrogen atom with a lone pair of electrons single bonded to three hydrogen atoms.\" \/>HCN: no electrons remainH<sub>3<\/sub>CCH<sub>3<\/sub>: no electrons remainHCCH: four electrons placed on carbonNH<sub>3<\/sub>: two electrons placed on nitrogen<\/li>\n<li><em>Where needed, rearrange electrons to form multiple bonds in order to obtain an octet on each atom:<\/em>HCN: form two more C\u2013N bondsH<sub>3<\/sub>CCH<sub>3<\/sub>: all atoms have the correct number of electronsHCCH: form a triple bond between the two carbon atomsNH<sub>3<\/sub>: all atoms have the correct number of electrons<img decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_ex070301_4_img-2.jpg\" alt=\"Four Lewis structures are shown. The first structure shows a carbon atom single bonded to a hydrogen atom and a nitrogen atom, which has three lone pairs of electrons. Two curved arrows point from the nitrogen to the carbon. Below this structure is the word \u201cgives\u201d and below that is the same structure, but this time there is a triple bond between the carbon and nitrogen. The second structure shows two carbons single bonded to one another and each single bonded to three hydrogen atoms. The third structure shows two carbon atoms, each with a lone pair of electrons, single bonded to one another and each single bonded to one hydrogen atom. Two curved arrows point from the carbon atoms to the space in between the two. Below this structure is the word \u201cgives\u201d and the same structure, but this time with a triple bond between the two carbons. The fourth structure shows a nitrogen atom with a lone pair of electrons single bonded to three hydrogen atoms.\" \/><\/li>\n<\/ol>\n<p>&nbsp;<\/p>\n<p id=\"fs-idm58848848\"><em><strong>Test yourself<\/strong><\/em><br \/>\nBoth carbon monoxide, CO, and carbon dioxide, CO<sub>2<\/sub>, are products of the combustion of fossil fuels. Both of these gases also cause problems: CO is toxic and CO<sub>2<\/sub> has been implicated in global climate change. What are the Lewis structures of these two molecules?<\/p>\n<p>&nbsp;<\/p>\n<p><em><strong>Answers<\/strong><\/em><\/p>\n<p><img decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_COCO2_img-2.jpg\" alt=\"Two Lewis structures are shown. The left shows a carbon triple bonded to an oxygen, each with a lone electron pair. The right structure shows a carbon double bonded to an oxygen on each side. Each oxygen has two lone pairs of electrons.\" \/><\/p>\n<\/div>\n<div class=\"textbox shaded\">\n<h3 class=\"title\">Example 4<\/h3>\n<p id=\"ball-ch09_s03_p42\" class=\"para\">What is the proper Lewis electron dot diagram for CO<sub class=\"subscript\">2<\/sub>?<\/p>\n<p>&nbsp;<\/p>\n<p class=\"simpara\"><strong>Solution<\/strong><\/p>\n<p id=\"ball-ch09_s03_p43\" class=\"para\">The central atom is a C atom, with O atoms as surrounding atoms. We have a total of 4 +\u00a06 +\u00a06 = 16 valence electrons. Following the rules for Lewis electron dot diagrams for compounds gives us<a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/C-O.png\"><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/C-O-1.png\" alt=\"C-O\" width=\"400\" height=\"40\" class=\"wp-image-4457 aligncenter\" \/><\/a><span style=\"font-size: 1em\">The O atoms have complete octets around them, but the C atom has only four electrons around it. The way to solve this dilemma is to make a double bond between carbon and <\/span><em class=\"emphasis\" style=\"font-size: 1em\">each<\/em><span style=\"font-size: 1em\"> O atom:<\/span><span style=\"font-size: 1em\"><\/span><\/p>\n<div class=\"informalfigure large\">\n<p class=\"para\"><a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/C-O-2.png\"><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/C-O-2-1.png\" alt=\"C-O-2\" width=\"400\" height=\"40\" class=\"wp-image-4458 aligncenter\" \/><\/a><span style=\"font-size: 1em\">Each O atom still has eight electrons around it, but now the C atom also has a complete octet. This is an acceptable Lewis electron dot diagram for CO<\/span><sub class=\"subscript\">2<\/sub><span style=\"font-size: 1em\">.<\/span><\/p>\n<div class=\"informalfigure large\">\n<p>&nbsp;<\/p>\n<p class=\"simpara\"><strong><em class=\"emphasis bolditalic\">Test Yourself<\/em><\/strong><\/p>\n<p id=\"ball-ch09_s03_p46\" class=\"para\">What is the proper Lewis electron dot diagram for carbonyl sulfide (COS)?<\/p>\n<p>&nbsp;<\/p>\n<p class=\"simpara\"><strong><em class=\"emphasis\">Answer<\/em><\/strong><a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/C-S-O.png\"><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/C-S-O-1.png\" alt=\"C-S-O\" width=\"400\" height=\"40\" class=\"wp-image-4459 aligncenter\" \/><\/a><\/p>\n<\/div>\n<\/div>\n<\/div>\n<div id=\"fs-idp177244128\" class=\"textbox shaded\">\n<h3 class=\"title\">Fullerene Chemistry<\/h3>\n<p id=\"fs-idm18160736\">Carbon soot has been known to man since prehistoric times, but it was not until fairly recently that the molecular structure of the main component of soot was discovered. In 1996, the Nobel Prize in Chemistry was awarded to Richard <strong class=\"no-emphasis\">Smalley<\/strong> (<a href=\"#CNX_Chem_07_03_Smalley\" class=\"autogenerated-content\">Figure 1<\/a>), Robert Curl, and Harold Kroto for their work in discovering a new form of carbon, the C<sub>60<\/sub> buckminsterfullerene molecule (<a href=\"https:\/\/opentextbc.ca\/chemistry\/chapter\/introduction-8\/#CNX_Chem_07_00_Bucky\" class=\"autogenerated-content\">Figure 1 in Chapter 8 Introduction<\/a>). An entire class of compounds, including spheres and tubes of various shapes, were discovered based on C<sub>60.<\/sub> This type of molecule, called a fullerene, shows promise in a variety of applications. Because of their size and shape, fullerenes can encapsulate other molecules, so they have shown potential in various applications from hydrogen storage to targeted drug delivery systems. They also possess unique electronic and optical properties that have been put to good use in solar powered devices and chemical sensors.<\/p>\n<figure id=\"CNX_Chem_07_03_Smalley\">\n<figure style=\"width: 650px\" class=\"wp-caption aligncenter\"><a href=\"https:\/\/opentextbc.ca\/chemistry\/wp-content\/uploads\/sites\/150\/2016\/05\/CNX_Chem_07_03_Smalley.jpg\"><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Smalley-2.jpg\" alt=\"A photo of Richard Smalley is shown.\" width=\"650\" height=\"497\" \/><\/a><figcaption class=\"wp-caption-text\"><strong>Figure 1.<\/strong> Richard Smalley (1943\u20132005), a professor of physics, chemistry, and astronomy at Rice University, was one of the leading advocates for fullerene chemistry. Upon his death in 2005, the US Senate honored him as the \u201cFather of Nanotechnology.\u201d (credit: United States Department of Energy)<\/figcaption><\/figure>\n<\/figure>\n<\/div>\n<\/section>\n<section id=\"fs-idm61779936\">\n<h2>Exceptions to the Octet Rule<\/h2>\n<p id=\"fs-idm1862800\">Many covalent molecules have central atoms that do not have eight electrons in their Lewis structures. These molecules fall into three categories:<\/p>\n<ul id=\"fs-idm41738208\">\n<li>Odd-electron molecules have an odd number of valence electrons, and therefore have an unpaired electron.<\/li>\n<li>Electron-deficient molecules have a central atom that has fewer electrons than needed for a noble gas configuration.<\/li>\n<li>Hypervalent molecules have a central atom that has more electrons than needed for a noble gas configuration.<\/li>\n<\/ul>\n<section id=\"fs-idm33391760\">Examples of these will be covered later chemistry courses.<\/p>\n<div class=\"textbox shaded\">\n<h3 class=\"title\">Food and Drink App: Vitamins and Minerals<\/h3>\n<p id=\"ball-ch09_s03_p50\" class=\"para\">Vitamins are nutrients that our bodies need in small amounts but cannot synthesize; therefore, they must be obtained from the diet. The word <em class=\"emphasis\">vitamin<\/em> comes from \u201cvital amine\u201d because it was once thought that all these compounds had an amine group (NH<sub class=\"subscript\">2<\/sub>) in it. This is not actually true, but the name stuck anyway.<\/p>\n<p id=\"ball-ch09_s03_p51\" class=\"para\">All vitamins are covalently bonded molecules. Most of them are commonly named with a letter, although all of them also have formal chemical names. Thus vitamin A is also called retinol, vitamin C is called ascorbic acid, and vitamin E is called tocopherol. There is no single vitamin B; there is a group of substances called the <em class=\"emphasis\">B complex vitamins<\/em> that are all water soluble and participate in cell metabolism. If a diet is lacking in a vitamin, diseases such as scurvy or rickets develop. Luckily, all vitamins are available as supplements, so any dietary deficiency in a vitamin can be easily corrected.<\/p>\n<p id=\"ball-ch09_s03_p52\" class=\"para\">A mineral is any chemical element other than carbon, hydrogen, oxygen, or nitrogen that is needed by the body. Minerals that the body needs in quantity include sodium, potassium, magnesium, calcium, phosphorus, sulfur, and chlorine. Essential minerals that the body needs in tiny quantities (so-called <em class=\"emphasis\">trace elements<\/em>) include manganese, iron, cobalt, nickel, copper, zinc, molybdenum, selenium, and iodine. Minerals are also obtained from the diet. Interestingly, most minerals are consumed in ionic form, rather than as elements or from covalent molecules. Like vitamins, most minerals are available in pill form, so any deficiency can be compensated for by taking supplements.<\/p>\n<p class=\"para\"><a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/Nutrition-Facts.png\"><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/Nutrition-Facts-1.png\" alt=\"Nutrition-Facts\" width=\"280\" height=\"567\" class=\"alignnone wp-image-4461\" \/><\/a><\/p>\n<p><strong>Figure 2.<\/strong> Vitamin and Mineral Supplements<\/p>\n<div class=\"informalfigure medium\" id=\"ball-ch09_s03_f01\">\n<p class=\"para\">Every entry down through pantothenic acid is a vitamin, and everything from calcium and below is a mineral.<\/p>\n<\/div>\n<\/div>\n<\/section>\n<\/section>\n<section id=\"fs-idm103622960\" class=\"summary\">\n<h2>Key Concepts and Summary<\/h2>\n<p id=\"fs-idm18966656\">Valence electronic structures can be visualized by drawing Lewis symbols (for atoms and monatomic ions) and Lewis structures (for molecules and polyatomic ions). Lone pairs, unpaired electrons, and single, double, or triple bonds are used to indicate where the valence electrons are located around each atom in a Lewis structure. Most structures\u2014especially those containing second row elements\u2014obey the octet rule, in which every atom (except H) is surrounded by eight electrons. Exceptions to the octet rule occur for odd-electron molecules (free radicals), electron-deficient molecules, and hypervalent molecules.<\/p>\n<\/section>\n<section id=\"fs-idm14528032\" class=\"exercises\">\n<div class=\"bcc-box bcc-info\">\n<h3>Exercises<\/h3>\n<p>1. Write the Lewis symbols for each of the following ions:<\/p>\n<p id=\"fs-idp177307680\">a) As<sup>3\u2013 \u00a0 \u00a0 \u00a0 \u00a0\u00a0<\/sup>b) I<sup>\u2013 \u00a0 \u00a0 \u00a0 \u00a0\u00a0<\/sup>c) Be<sup>2+ \u00a0 \u00a0 \u00a0 \u00a0\u00a0<\/sup>d) O<sup>2\u2013 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0<\/sup>e) Ga<sup>3+ \u00a0 \u00a0 \u00a0 \u00a0 \u00a0<\/sup>f) Li<sup>+ \u00a0 \u00a0 \u00a0 \u00a0\u00a0<\/sup>g) N<sup>3\u2013<\/sup><\/p>\n<p>2. Write the Lewis symbols of the ions in each of the following ionic compounds and the Lewis symbols of the atom from which they are formed:<\/p>\n<p id=\"fs-idm81853760\">a) MgS \u00a0 \u00a0 \u00a0 \u00a0\u00a0b) Al<sub>2<\/sub>O<sub>3 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0<\/sub>c) GaCl<sub>3 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0<\/sub>d) K<sub>2<\/sub>O \u00a0 \u00a0 \u00a0 \u00a0e) Li<sub>3<\/sub>N \u00a0 \u00a0 \u00a0 \u00a0\u00a0f) KF<\/p>\n<p>3. Write the Lewis structure for the diatomic molecule P<sub>2<\/sub>, an unstable form of phosphorus found in high-temperature phosphorus vapor.<\/p>\n<p>4. Write Lewis structures for the following:<\/p>\n<p id=\"fs-idm45128544\">a) O<sub>2 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0\u00a0<\/sub>b) H<sub>2<\/sub>CO \u00a0 \u00a0 \u00a0 \u00a0\u00a0c) AsF<sub>3 \u00a0 \u00a0 \u00a0 \u00a0\u00a0<\/sub>d) ClNO \u00a0 \u00a0 \u00a0 \u00a0\u00a0e) SiCl<sub>4<\/sub><\/p>\n<p id=\"fs-idm5797408\">f) H<sub>3<\/sub>O<sup>+ \u00a0 \u00a0 \u00a0 \u00a0<\/sup>g) NH<sub>4<\/sub><sup>+ \u00a0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0<\/sup>h) BF<sub>4<\/sub><sup>\u2212 \u00a0 \u00a0 \u00a0 \u00a0\u00a0<\/sup>i) HCCH \u00a0 \u00a0 \u00a0 \u00a0\u00a0j) ClCN \u00a0 \u00a0 \u00a0 \u00a0 \u00a0k) C<sub>2<\/sub><sup>2+<\/sup><\/p>\n<p>5. Write Lewis structures for the following:<\/p>\n<p id=\"fs-idm17992976\">a) \u00a0SeCl<sub>3<\/sub><sup>+ \u00a0 \u00a0 \u00a0 \u00a0 \u00a0<\/sup>b) Cl<sub>2<\/sub>BBCl<sub>2<\/sub> (contains a B\u2013B bond)<\/p>\n<p>6. Correct the following statement: \u201cThe bonds in solid PbCl<sub>2<\/sub> are ionic; the bond in a HCl molecule is covalent. Thus, all of the valence electrons in PbCl<sub>2<\/sub> are located on the Cl<sup>\u2013<\/sup> ions, and all of the valence electrons in a HCl molecule are shared between the H and Cl atoms.\u201d<\/p>\n<p>7. Methanol, H<sub>3<\/sub>COH, is used as the fuel in some race cars. Ethanol, C<sub>2<\/sub>H<sub>5<\/sub>OH, is used extensively as motor fuel in Brazil. Both methanol and ethanol produce CO<sub>2<\/sub> and H<sub>2<\/sub>O when they burn. Write the chemical equations for these combustion reactions using Lewis structures instead of chemical formulas.<\/p>\n<p>8. Carbon tetrachloride was formerly used in fire extinguishers for electrical fires. It is no longer used for this purpose because of the formation of the toxic gas phosgene, Cl<sub>2<\/sub>CO. Write the Lewis structures for carbon tetrachloride and phosgene.<\/p>\n<p>9. The arrangement of atoms in several biologically important molecules is given here. Complete the Lewis structures of these molecules by adding multiple bonds and lone pairs. Do not add any more atoms.<\/p>\n<p id=\"fs-idm80239808\">a) the amino acid serine:<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Impbioa_img-2.jpg\" alt=\"A Lewis structure is shown. A nitrogen atom is single bonded to two hydrogen atoms and a carbon atom. The carbon atom is single bonded to a hydrogen atom and two other carbon atoms. One of these carbon atoms is single bonded to two hydrogen atoms and an oxygen atom. The oxygen atom is bonded to a hydrogen atom. The other carbon atom is single bonded to two oxygen atoms, one of which is bonded to a hydrogen atom.\" width=\"228\" height=\"181\" class=\"\" \/><\/p>\n<p id=\"fs-idm147749792\">b) urea:<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Impbiob_img-2.jpg\" alt=\"A Lewis structure is shown. A nitrogen atom is single bonded to two hydrogen atoms and a carbon atom. The carbon atom is single bonded to an oxygen atom and another nitrogen atom. That nitrogen atom is then single bonded to two hydrogen atoms.\" width=\"196\" height=\"69\" class=\"\" \/><\/p>\n<p id=\"fs-idp28905856\">c) pyruvic acid:<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Impbioc_img-2.jpg\" alt=\"A Lewis structure is shown. A carbon atom is single bonded to three hydrogen atoms and another carbon atom. The second carbon atom is single bonded to an oxygen atom and a third carbon atom. This carbon is then single bonded to two oxygen atoms, one of which is single bonded to a hydrogen atom.\" width=\"237\" height=\"111\" class=\"\" \/><\/p>\n<p id=\"fs-idm104403040\">d) uracil:<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Impbiod_img-2.jpg\" alt=\"A Lewis hexagonal ring structure is shown. From the top of the ring (moving clockwise), three carbon atoms, one nitrogen atom, a carbon atom, and a nitrogen atom are single bonded to each another. The top carbon atom is single bonded to an oxygen atom. The second and third carbons and the nitrogen atom are each single bonded to a hydrogen atom. The next carbon atom is single bonded to an oxygen atom, and the last nitrogen atom is single bonded to a hydrogen atom.\" width=\"172\" height=\"196\" class=\"\" \/><\/p>\n<p id=\"fs-idp106603552\">e) carbonic acid:<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Impbioe_img-2.jpg\" alt=\"A Lewis structure is shown. A carbon atom is single bonded to three oxygen atoms. Two of those oxygen atoms are each single bonded to a hydrogen atom.\" width=\"193\" height=\"68\" class=\"\" \/><\/p>\n<p>10. A compound with a molar mass of about 42 g\/mol contains 85.7% carbon and 14.3% hydrogen by mass. Write the Lewis structure for a molecule of the compound.<\/p>\n<p>11. How are single, double, and triple bonds similar? How do they differ?<\/p>\n<p><span style=\"font-size: 1em\">12. How many electrons will be in the valence shell of H atoms when it makes a covalent bond?<\/span><\/p>\n<p>13. What is the Lewis electron dot diagram of I<sub class=\"subscript\">2<\/sub>? Circle the electrons around each atom to verify that each valence shell is filled.<\/p>\n<p>14. What is the Lewis electron dot diagram of NCl<sub class=\"subscript\">3<\/sub>? Circle the electrons around each atom to verify that each valence shell is filled.<\/p>\n<p>15. Draw the Lewis electron dot diagram for each substance. \u00a0\u00a0<span style=\"font-size: 1em\">a) \u00a0SF<\/span><sub class=\"subscript\">2 \u00a0 \u00a0 \u00a0 \u00a0\u00a0<\/sub><span style=\"font-size: 1em\">b) \u00a0BH<\/span><sub class=\"subscript\">4<\/sub><sup class=\"superscript\">\u2212<\/sup><\/p>\n<p><span style=\"font-size: 1em\">16. Draw the Lewis electron dot diagram for each substance. \u00a0\u00a0<\/span><span style=\"font-size: 1em\">a) \u00a0GeH<\/span><sub class=\"subscript\">4 \u00a0 \u00a0\u00a0<\/sub><span style=\"font-size: 1em\">b) \u00a0ClF<\/span><\/p>\n<p><span style=\"font-size: 1em\">17. Draw the Lewis electron dot diagram for each substance. Double or triple bonds may be needed. \u00a0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0\u00a0<\/span><\/p>\n<p><span style=\"font-size: 1em\">a) \u00a0SiO<\/span><sub class=\"subscript\">2 \u00a0 \u00a0<\/sub><span style=\"font-size: 1em\">b) \u00a0C<\/span><sub class=\"subscript\">2<\/sub><span style=\"font-size: 1em\">H<\/span><sub class=\"subscript\">4<\/sub><span style=\"font-size: 1em\"> (assume two central atoms)<\/span><\/p>\n<p><span style=\"font-size: 1em\">18. Draw the Lewis electron dot diagram for each substance. Double or triple bonds may be needed.<\/span><\/p>\n<div class=\"question\">\n<p>a) \u00a0CS<sub class=\"subscript\">2 \u00a0 \u00a0 \u00a0\u00a0<\/sub>b) \u00a0NH<sub class=\"subscript\">2<\/sub>CONH<sub class=\"subscript\">2<\/sub> (assume that the N and C atoms are the central atoms)<\/p>\n<\/div>\n<p>&nbsp;<\/p>\n<p><strong>Answers<\/strong><\/p>\n<p id=\"fs-idm9743232\">1. a) eight electrons:<br \/>\n<img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Question1a_img-2.jpg\" alt=\"A Lewis dot diagram shows the symbol for arsenic, A s, surrounded by eight dots and a superscripted three negative sign.\" width=\"272\" height=\"51\" class=\"\" \/><\/p>\n<p>b) eight electrons:<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Question1b_img-2.jpg\" alt=\"A Lewis dot diagram shows the symbol for iodine, I, surrounded by eight dots and a superscripted negative sign.\" width=\"272\" height=\"51\" class=\"\" \/><\/p>\n<p>c) no electrons: \u00a0\u00a0Be<sup>2+<\/sup><\/p>\n<p>d) eight electrons:<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Question1d_img-2.jpg\" alt=\"A Lewis dot diagram shows the symbol for oxygen, O, surrounded by eight dots and a superscripted two negative sign.\" width=\"272\" height=\"51\" class=\"\" \/><\/p>\n<p>e) no electrons: \u00a0Ga<sup>3+<\/sup><\/p>\n<p>f) no electrons:\u00a0Li<sup>+<\/sup><\/p>\n<p>g) eight electrons:<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Question1g_img-2.jpg\" alt=\"A Lewis dot diagram shows the symbol for nitrogen, N, surrounded by eight dots and a superscripted three negative sign.\" width=\"266\" height=\"50\" class=\"\" \/><\/p>\n<p id=\"fs-idp56026944\">2. a)<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Exercise3a_img-2.jpg\" alt=\"Two Lewis structures are shown. The left shows the symbol M g with a superscripted two positive sign while the right shows the symbol S surrounded by eight dots and a superscripted two negative sign.\" width=\"261\" height=\"49\" class=\"\" \/><\/p>\n<p>b)<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Exercise3b_img-2.jpg\" alt=\"Two Lewis structures are shown. The left shows the symbol A l with a superscripted three positive sign while the right shows the symbol O surrounded by eight dots and a superscripted two negative sign.\" width=\"261\" height=\"49\" class=\"\" \/><\/p>\n<p>c)<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Exercise3c_img-2.jpg\" alt=\"Two Lewis structures are shown. The left shows the symbol G a with a superscripted three positive sign while the right shows the symbol C l surrounded by eight dots and a superscripted negative sign.\" width=\"256\" height=\"48\" class=\"\" \/><\/p>\n<p>d)<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Exercise3d_img-2.jpg\" alt=\"Two Lewis structures are shown. The left shows the symbol K with a superscripted positive sign while the right shows the symbol O surrounded by eight dots and a superscripted two negative sign.\" width=\"256\" height=\"48\" class=\"\" \/><\/p>\n<p>e)<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Exercise3e_img-2.jpg\" alt=\"Two Lewis structures are shown. The left shows the symbol L i with a superscripted positive sign while the right shows the symbol N surrounded by eight dots and a superscripted three negative sign.\" width=\"250\" height=\"47\" class=\"\" \/><\/p>\n<p>f)<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Exercise3f_img-2.jpg\" alt=\"Two Lewis structures are shown. The left shows the symbol K with a superscripted positive sign while the right shows the symbol F surrounded by eight dots and a superscripted negative sign.\" width=\"250\" height=\"47\" class=\"\" \/><\/p>\n<p>3.<br \/>\n<img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Question5_img-2.jpg\" alt=\"A Lewis diagram shows two phosphorus atoms triple bonded together each with one lone electron pair.\" width=\"244\" height=\"27\" class=\"\" \/><\/p>\n<p id=\"fs-idm41104416\">4. a)<br \/>\n<img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Question7a_img-2.jpg\" alt=\"A Lewis structure shows two oxygen atoms double bonded together, and each has two lone pairs of electrons.\" width=\"97\" height=\"39\" class=\"\" \/><\/p>\n<p>In this case, the Lewis structure is inadequate to depict the fact that experimental studies have shown two unpaired electrons in each oxygen molecule<em>.<\/em><\/p>\n<p>b)<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Question7b_img-2.jpg\" alt=\"A Lewis structure shows a carbon atom that is single bonded to two hydrogen atoms and double bonded to an oxygen atom. The oxygen atom has two lone pairs of electrons.\" width=\"117\" height=\"84\" class=\"\" \/><\/p>\n<p>c)<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Question7c_img-2.jpg\" alt=\"A Lewis structure shows an arsenic atom single bonded to three fluorine atoms. Each fluorine atom has a lone pair of electrons.\" width=\"126\" height=\"83\" class=\"\" \/><\/p>\n<p>d)<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Question7d_img-2.jpg\" alt=\"A Lewis structure shows a nitrogen atom with a lone pair of electrons single bonded to a chlorine atom that has three lone pairs of electrons. The nitrogen is also double bonded to an oxygen which has two lone pairs of electrons.\" width=\"273\" height=\"47\" class=\"\" \/><\/p>\n<p>e)<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Question7e_img-2.jpg\" alt=\"A Lewis structure shows a silicon atom that is single bonded to four chlorine atoms. Each chlorine atom has three lone pairs of electrons.\" width=\"141\" height=\"139\" class=\"\" \/><\/p>\n<p>f)<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Question7f_img-2.jpg\" alt=\"A Lewis structure shows an oxygen atom with a lone pair of electrons single bonded to three hydrogen atoms. The structure is surrounded by brackets with a superscripted positive sign.\" width=\"270\" height=\"104\" class=\"\" \/><\/p>\n<p>g)<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Question7g_img-2.jpg\" alt=\"A Lewis structure shows a nitrogen atom single bonded to four hydrogen atoms. The structure is surrounded by brackets with a superscripted positive sign.\" width=\"254\" height=\"129\" class=\"\" \/><\/p>\n<p>h)<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Question7h_img-2.jpg\" alt=\"A Lewis structure shows a boron atom single bonded to four fluorine atoms. Each fluorine atom has three lone pairs of electrons. The structure is surrounded by brackets with a superscripted negative sign.\" width=\"250\" height=\"127\" class=\"\" \/><\/p>\n<p>i)<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Question7i_img-2.jpg\" alt=\"A Lewis structure shows two carbon atoms that are triple bonded together. Each carbon is also single bonded to a hydrogen atom.\" width=\"159\" height=\"29\" class=\"\" \/><\/p>\n<p>j)<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Question7j_img-2.jpg\" alt=\"A Lewis structure shows a carbon atom that is triple bonded to a nitrogen atom that has one lone pair of electrons. The carbon is also single bonded to a chlorine atom that has three lone pairs of electrons.\" width=\"138\" height=\"49\" class=\"\" \/><\/p>\n<p>k)<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Question7k_img-2.jpg\" alt=\"A Lewis structure shows two carbon atoms joined with a triple bond. A superscripted 2 positive sign lies to the right of the second carbon.\" width=\"86\" height=\"35\" class=\"\" \/><\/p>\n<p id=\"fs-idm123049856\">5. a) SeCl<sub>3<\/sub><sup>+<\/sup>:<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Question9c_img-2.jpg\" alt=\"A Lewis structure shows a selenium atom with one lone pair of electrons single bonded to three chlorine atoms each with three lone pairs of electrons. The whole structure is surrounded by brackets.\" width=\"265\" height=\"102\" class=\"\" \/><\/p>\n<p>b) Cl<sub>2<\/sub>BBCl<sub>2<\/sub>:<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Question9d_img-2.jpg\" alt=\"A Lewis structure shows two boron atoms that are single bonded together. Each is also single bonded to two chlorine atoms that both have three lone pairs of electrons.\" width=\"146\" height=\"109\" class=\"\" \/><\/p>\n<p id=\"fs-idm7862384\">6. Two valence electrons per Pb atom are transferred to Cl atoms; the resulting Pb<sup>2+<\/sup> ion has a 6<em>s<\/em><sup>2<\/sup> valence shell configuration. Two of the valence electrons in the HCl molecule are shared, and the other six are located on the Cl atom as lone pairs of electrons.<\/p>\n<p>7.<br \/>\n<img decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Question18_img-2.jpg\" alt=\"Two reactions are shown using Lewis structures. The top reaction shows a carbon atom, single bonded to three hydrogen atoms and single bonded to an oxygen atom with two lone pairs of electrons. The oxygen atom is also bonded to a hydrogen atom. This is followed by a plus sign and the number one point five, followed by two oxygen atoms bonded together with a double bond and each with two lone pairs of electrons. A right-facing arrow leads to a carbon atom that is double bonded to two oxygen atoms, each of which has two lone pairs of electrons. This structure is followed by a plus sign, a number two, and a structure made up of an oxygen with two lone pairs of electrons single bonded to two hydrogen atoms. The bottom reaction shows a carbon atom, single bonded to three hydrogen atoms and single bonded to another carbon atom. The second carbon atom is single bonded to two hydrogen atoms and one oxygen atom with two lone pairs of electrons. The oxygen atom is also bonded to a hydrogen atom. This is followed by a plus sign and the number three, followed by two oxygen atoms bonded together with a double bond. Each oxygen atom has two lone pairs of electrons. A right-facing arrow leads to a number two and a carbon atom that is double bonded to two oxygen atoms, each of which has two lone pairs of electrons. This structure is followed by a plus sign, a number three, and a structure made up of an oxygen with two lone pairs of electrons single bonded to two hydrogen atoms.\" \/><\/p>\n<p>8.<br \/>\n<img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Carbtet_img-2.jpg\" alt=\"Two Lewis structures are shown. The left depicts a carbon atom single bonded to four chlorine atoms, each with three lone pairs of electrons. The right shows a carbon atom double bonded to an oxygen atom that has two lone pairs of electrons. The carbon atom is also single bonded to two chlorine atoms, each of which has three lone pairs of electrons.\" width=\"273\" height=\"134\" class=\"\" \/><\/p>\n<p id=\"fs-idm17951840\">9. a)<br \/>\n<img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Impbioansa_img-2.jpg\" alt=\"A Lewis structure is shown. A nitrogen atom is single bonded to two hydrogen atoms and a carbon atom. The carbon atom is single bonded to a hydrogen atom and two other carbon atoms. One of these carbon atoms is single bonded to two hydrogen atoms and an oxygen atom. The oxygen atom is bonded to a hydrogen atom. The other carbon is single bonded to two oxygen atoms, one of which is bonded to a hydrogen atom. The oxygen atoms have two lone pairs of electron dots, and the nitrogen atom has one lone pair of electron dots.\" width=\"232\" height=\"194\" class=\"\" \/><\/p>\n<p>b)<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Impbioansb_img-2.jpg\" alt=\"A Lewis structure is shown. A nitrogen atom is single bonded to two hydrogen atoms and a carbon atom. The carbon atom is single bonded to an oxygen atom and one nitrogen atom. That nitrogen atom is then single bonded to two hydrogen atoms. The oxygen atom has two lone pairs of electron dots, and the nitrogen atoms have one lone pair of electron dots each.\" width=\"203\" height=\"94\" class=\"\" \/><\/p>\n<p>c)<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Impbioansc_img-2.jpg\" alt=\"A Lewis structure is shown. A carbon atom is single bonded to three hydrogen atoms and a carbon atom. The carbon atom is single bonded to an oxygen atom and a third carbon atom. This carbon is then single bonded to two oxygen atoms, one of which is single bonded to a hydrogen atom. Each oxygen atom has two lone pairs of electron dots.\" width=\"243\" height=\"122\" class=\"\" \/><\/p>\n<p>d)<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Impbioansd_img-2.jpg\" alt=\"A Lewis hexagonal ring structure is shown. From the top of the ring, three carbon atoms, one nitrogen atom, a carbon atom and a nitrogen atom are single bonded to one another. The top carbon is single bonded to an oxygen, the second and third carbons and the nitrogen atom are each single bonded to a hydrogen atom. The next carbon is single bonded to an oxygen atom and the last nitrogen is single bonded to a hydrogen atom. The oxygen atoms have two lone pairs of electron dots, and the nitrogen atoms have one lone pair of electron dots.\" width=\"187\" height=\"210\" class=\"\" \/><\/p>\n<p>e)<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Impbioanse_img-2.jpg\" alt=\"A Lewis structure is shown. A carbon atom is single bonded to three oxygen atoms. Two of those oxygen atoms are each single bonded to a hydrogen atom. Each oxygen atom has two lone pairs of electron dots.\" width=\"205\" height=\"95\" class=\"\" \/><\/p>\n<p>10.<br \/>\n<img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_03_Exercise25_img-2.jpg\" alt=\"A Lewis structure is shown. A carbon atom is single bonded to three hydrogen atoms and another carbon atom. The second carbon atom is double bonded to another carbon atom and single bonded to a hydrogen atom. The last carbon is single bonded to two hydrogen atoms.\" width=\"269\" height=\"119\" class=\"\" \/><\/p>\n<p id=\"fs-idp205940864\">11. Each bond includes a sharing of electrons between atoms. Two electrons are shared in a single bond; four electrons are shared in a double bond; and six electrons are shared in a triple bond.<\/p>\n<p>12. two<\/p>\n<p>13.<\/p>\n<p><strong><a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/I-I.png\"><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/I-I-1.png\" alt=\"I-I\" width=\"343\" height=\"84\" class=\"alignnone wp-image-4462\" \/><\/a><\/strong><\/p>\n<p>14.<\/p>\n<p><strong><a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/N-Cl.png\"><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/N-Cl-1.png\" alt=\"N-Cl\" width=\"343\" height=\"102\" class=\"alignnone wp-image-4463\" \/><\/a><\/strong><\/p>\n<p>15.<\/p>\n<p style=\"padding-left: 30px\">a) \u00a0\u00a0<a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/S-F.png\"><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/S-F-1.png\" alt=\"S-F\" width=\"339\" height=\"101\" class=\"alignnone wp-image-4464\" \/><\/a><\/p>\n<p style=\"padding-left: 30px\">b) \u00a0\u00a0<a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/H-B.png\"><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/H-B-1.png\" alt=\"H-B\" width=\"333\" height=\"135\" class=\"alignnone wp-image-4465\" \/><\/a><\/p>\n<p>16.<\/p>\n<p style=\"padding-left: 30px\">a) \u00a0\u00a0<a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/H-Ge.png\"><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/H-Ge-1.png\" alt=\"H-Ge\" width=\"331\" height=\"134\" class=\"alignnone wp-image-4467\" \/><\/a><\/p>\n<p style=\"padding-left: 30px\">b) \u00a0\u00a0<a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/Cl-F.png\"><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/Cl-F-1.png\" alt=\"Cl-F\" width=\"343\" height=\"84\" class=\"alignnone wp-image-4468\" \/><\/a><\/p>\n<p>17.<\/p>\n<p style=\"padding-left: 30px\">a) \u00a0\u00a0<a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/Si-O.png\"><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/Si-O-1.png\" alt=\"Si-O\" width=\"327\" height=\"80\" class=\"alignnone wp-image-4469\" \/><\/a><\/p>\n<p style=\"padding-left: 30px\">b) \u00a0\u00a0<a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/C-H.png\"><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/C-H-1.png\" alt=\"C-H\" width=\"286\" height=\"116\" class=\"alignnone wp-image-4470\" \/><\/a><\/p>\n<p>18.<\/p>\n<p style=\"padding-left: 30px\">a) \u00a0\u00a0<a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/S-C.png\"><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/S-C-1.png\" alt=\"S-C\" width=\"327\" height=\"80\" class=\"alignnone wp-image-4471\" \/><\/a><\/p>\n<p style=\"padding-left: 30px\">b) \u00a0\u00a0<a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/C-N-H-O.png\"><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/C-N-H-O-1.png\" alt=\"C-N-H-O\" width=\"300\" height=\"154\" class=\"alignnone wp-image-4472\" \/><\/a><\/p>\n<\/div>\n<\/section>\n<div>\n<h2>Glossary<\/h2>\n<p><strong>double bond:\u00a0<\/strong>covalent bond in which two pairs of electrons are shared between two atoms<\/p>\n<p><strong>free radical:\u00a0<\/strong>molecule that contains an odd number of electrons<\/p>\n<p><strong>hypervalent molecule:\u00a0<\/strong>molecule containing at least one main group element that has more than eight electrons in its valence shell<\/p>\n<p><strong>Lewis structure:\u00a0<\/strong>diagram showing lone pairs and bonding pairs of electrons in a molecule or an ion<\/p>\n<p><strong>Lewis symbol:\u00a0<\/strong>symbol for an element or monatomic ion that uses a dot to represent each valence electron in the element or ion<\/p>\n<p><strong>lone pair:\u00a0<\/strong>two (a pair of) valence electrons that are not used to form a covalent bond<\/p>\n<p><strong>octet rule:\u00a0<\/strong>guideline that states main group atoms will form structures in which eight valence electrons interact with each nucleus, counting bonding electrons as interacting with both atoms connected by the bond<\/p>\n<p><strong>single bond:\u00a0<\/strong>bond in which a single pair of electrons is shared between two atoms<\/p>\n<p><strong>triple bond:\u00a0<\/strong>bond in which three pairs of electrons are shared between two atoms<\/p>\n<dl id=\"fs-idm68093504\" class=\"definition\">\n<dt><a href=\"http:\/\/opentextbc.ca\/introductorychemistry\/wp-content\/uploads\/sites\/17\/2014\/09\/C-N-H-O.png\"><\/a><\/dt>\n<\/dl>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n","protected":false},"author":330,"menu_order":6,"template":"","meta":{"pb_show_title":"on","pb_short_title":"9.5 Covalent Bonds and Lewis Structures","pb_subtitle":"","pb_authors":[],"pb_section_license":"cc-by-nc-sa"},"chapter-type":[],"contributor":[],"license":[54],"class_list":["post-2484","chapter","type-chapter","status-publish","hentry","license-cc-by-nc-sa"],"part":1538,"_links":{"self":[{"href":"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-json\/pressbooks\/v2\/chapters\/2484","targetHints":{"allow":["GET"]}}],"collection":[{"href":"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-json\/pressbooks\/v2\/chapters"}],"about":[{"href":"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-json\/wp\/v2\/types\/chapter"}],"author":[{"embeddable":true,"href":"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-json\/wp\/v2\/users\/330"}],"version-history":[{"count":13,"href":"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-json\/pressbooks\/v2\/chapters\/2484\/revisions"}],"predecessor-version":[{"id":4707,"href":"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-json\/pressbooks\/v2\/chapters\/2484\/revisions\/4707"}],"part":[{"href":"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-json\/pressbooks\/v2\/parts\/1538"}],"metadata":[{"href":"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-json\/pressbooks\/v2\/chapters\/2484\/metadata\/"}],"wp:attachment":[{"href":"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-json\/wp\/v2\/media?parent=2484"}],"wp:term":[{"taxonomy":"chapter-type","embeddable":true,"href":"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-json\/pressbooks\/v2\/chapter-type?post=2484"},{"taxonomy":"contributor","embeddable":true,"href":"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-json\/wp\/v2\/contributor?post=2484"},{"taxonomy":"license","embeddable":true,"href":"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-json\/wp\/v2\/license?post=2484"}],"curies":[{"name":"wp","href":"https:\/\/api.w.org\/{rel}","templated":true}]}}