{"id":1663,"date":"2018-04-11T22:52:39","date_gmt":"2018-04-12T02:52:39","guid":{"rendered":"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/chapter\/7-4-formal-charges-and-resonance\/"},"modified":"2018-06-23T00:09:27","modified_gmt":"2018-06-23T04:09:27","slug":"7-4-formal-charges-and-resonance","status":"publish","type":"chapter","link":"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/chapter\/7-4-formal-charges-and-resonance\/","title":{"raw":"9.6 Formal Charges and Resonance","rendered":"9.6 Formal Charges and Resonance"},"content":{"raw":"<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>Compute formal charges for atoms in any Lewis structure<\/li>\r\n \t<li>Use formal charges to identify the most reasonable Lewis structure for a given molecule<\/li>\r\n \t<li>Explain the concept of resonance and draw Lewis structures representing resonance forms for a given molecule<\/li>\r\n<\/ul>\r\n<\/div>\r\n<p id=\"fs-idm49375088\">In the previous section, we discussed how to write Lewis structures for molecules and polyatomic ions. As we have seen, however, in some cases, there is seemingly more than one valid structure for a molecule. We can use the concept of formal charges to help us predict the most appropriate Lewis structure when more than one is reasonable.<\/p>\r\n\r\n<section id=\"fs-idp11710128\">\r\n<h2>Calculating Formal Charge<\/h2>\r\n<p id=\"fs-idm16217792\">The <strong>formal charge<\/strong> of an atom in a molecule is the <em>hypothetical<\/em> charge the atom would have if we could redistribute the electrons in the bonds evenly between the atoms. Another way of saying this is that formal charge results when we take the number of valence electrons of a neutral atom, subtract the nonbonding electrons, and then subtract the number of bonds connected to that atom in the Lewis structure.<\/p>\r\n<p id=\"fs-idp126842624\">Thus, we calculate formal charge as follows:<\/p>\r\n\r\n<div class=\"equation\" id=\"fs-idp90018832\" style=\"text-align: center\">$latex \\text{formal charge} = \\# \\;\\text{valence shell electrons (free atom)} \\; - \\;\\# \\;\\text{lone pair electrons}\\; - \\frac{1}{2} \\# \\;\\text{bonding electrons}$<\/div>\r\n<p id=\"fs-idp88204016\">We can double-check formal charge calculations by determining the sum of the formal charges for the whole structure. The sum of the formal charges of all atoms in a molecule must be zero; the sum of the formal charges in an ion should equal the charge of the ion.<\/p>\r\n<p id=\"fs-idp15200480\">We must remember that the formal charge calculated for an atom is not the <em>actual<\/em> charge of the atom in the molecule. Formal charge is only a useful bookkeeping procedure; it does not indicate the presence of actual charges.<\/p>\r\n\r\n<div class=\"textbox shaded\" id=\"fs-idm21031296\">\r\n<h3>Example 1<\/h3>\r\n<p id=\"fs-idp63607984\">Assign formal charges to each atom in the interhalogen ion ICl<sub>4<\/sub><sup>\u2212<\/sup>.<\/p>\r\n&nbsp;\r\n<p id=\"fs-idp251046608\"><strong>Solution<\/strong><\/p>\r\n\r\n<ol id=\"fs-idm30182496\" class=\"stepwise\">\r\n \t<li><em>We divide the bonding electron pairs equally for all I\u2013Cl bonds:\r\n<\/em><img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_ICL4_img-2.jpg\" alt=\"A Lewis structure is shown. An iodine atom with two lone pairs of electrons is single bonded to four chlorine atoms, each of which has three lone pairs of electrons. Brackets surround the structure and there is a superscripted negative sign.\" class=\"aligncenter\" width=\"186\" height=\"137\" \/><\/li>\r\n \t<li><em>We assign lone pairs of electrons to their atoms<\/em>. Each Cl atom now has seven electrons assigned to it, and the I atom has eight.<\/li>\r\n \t<li><em><em>Subtract this number from the number of valence electrons for the neutral atom:<\/em><\/em>I: 7 \u2013 8 = \u20131Cl: 7 \u2013 7 = 0The sum of the formal charges of all the atoms equals \u20131, which is identical to the charge of the ion (\u20131).<\/li>\r\n<\/ol>\r\n&nbsp;\r\n<p id=\"fs-idp65521520\"><em><strong>Test Yourself<\/strong><\/em>\r\nCalculate the formal charge for each atom in the carbon monoxide molecule:<\/p>\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_CO_img-2.jpg\" alt=\"A Lewis structure is shown. A carbon atom with one lone pair of electrons is triple bonded to an oxygen with one lone pair of electrons.\" class=\"aligncenter\" width=\"101\" height=\"33\" \/>\r\n\r\n&nbsp;\r\n\r\n<em><strong>Answer<\/strong><\/em>\r\n\r\nC \u22121, O +1\r\n\r\n<\/div>\r\n<div class=\"textbox shaded\" id=\"fs-idp59620704\">\r\n<h3>Example 2<\/h3>\r\n<p id=\"fs-idp26752688\">Assign formal charges to each atom in the interhalogen molecule BrCl<sub>3<\/sub>.<\/p>\r\n&nbsp;\r\n<p id=\"fs-idp34669824\"><strong>Solution<\/strong><\/p>\r\n\r\n<ol id=\"fs-idp8784496\" class=\"stepwise\">\r\n \t<li><em>Assign one of the electrons in each Br\u2013Cl bond to the Br atom and one to the Cl atom in that bond:\r\n<\/em><img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_BRCL3_img-2.jpg\" alt=\"A Lewis structure is shown. A bromine atom with two lone pairs of electrons is single bonded to three chlorine atoms, each of which has three lone pairs of electrons.\" class=\"aligncenter\" width=\"150\" height=\"99\" \/><\/li>\r\n \t<li><em>Assign the lone pairs to their atom.<\/em> Now each Cl atom has seven electrons and the Br atom has seven electrons.<\/li>\r\n \t<li><em>Subtract this number from the number of valence electrons for the neutral atom.<\/em> This gives the formal charge:Br: 7 \u2013 7 = 0Cl: 7 \u2013 7 = 0All atoms in BrCl<sub>3<\/sub> have a formal charge of zero, and the sum of the formal charges totals zero, as it must in a neutral molecule.<\/li>\r\n<\/ol>\r\n&nbsp;\r\n<p id=\"fs-idp245106544\"><em><strong>Test yourself<\/strong><\/em>\r\nDetermine the formal charge for each atom in NCl<sub>3<\/sub>.<\/p>\r\n&nbsp;\r\n\r\n<em><strong>Answer<\/strong><\/em>\r\n<p id=\"fs-idm22650544\">N: 0; all three Cl atoms: 0<\/p>\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_Ex070402_img-2.jpg\" alt=\"A Lewis structure is shown. A nitrogen atom with one lone pair of electrons is single bonded to three chlorine atoms, each of which has three lone pairs of electrons.\" class=\"aligncenter\" width=\"275\" height=\"99\" \/>\r\n\r\n<\/div>\r\n<\/section><section id=\"fs-idp27335568\">\r\n<h2>Using Formal Charge to Predict Molecular Structure<\/h2>\r\n<p id=\"fs-idp93581584\">The arrangement of atoms in a molecule or ion is called its <strong>molecular structure<\/strong>. In many cases, following the steps for writing Lewis structures may lead to more than one possible molecular structure\u2014different multiple bond and lone-pair electron placements or different arrangements of atoms, for instance. A few guidelines involving formal charge can be helpful in deciding which of the possible structures is most likely for a particular molecule or ion:<\/p>\r\n\r\n<ol id=\"fs-idp69636800\">\r\n \t<li>A molecular structure in which all formal charges are zero is preferable to one in which some formal charges are not zero.<\/li>\r\n \t<li>If the Lewis structure must have nonzero formal charges, the arrangement with the smallest nonzero formal charges is preferable.<\/li>\r\n \t<li>Lewis structures are preferable when adjacent formal charges are zero or of the opposite sign.<\/li>\r\n \t<li>When we must choose among several Lewis structures with similar distributions of formal charges, the structure with the negative formal charges on the more electronegative atoms is preferable.<\/li>\r\n<\/ol>\r\n<p id=\"fs-idp52450560\">To see how these guidelines apply, let us consider some possible structures for carbon dioxide, CO<sub>2<\/sub>. We know from our previous discussion that the less electronegative atom typically occupies the central position, but formal charges allow us to understand <em>why<\/em> this occurs. We can draw three possibilities for the structure: carbon in the center and double bonds, carbon in the center with a single and triple bond, and oxygen in the center with double bonds:<\/p>\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_CO2pos_img-2.jpg\" alt=\"Three Lewis structures are shown. The left and right structures show a carbon atom double bonded to two oxygen atoms, each of which has two lone pairs of electrons. The center structure shows a carbon atom that is triple bonded to an oxygen atom with one lone pair of electrons and single bonded to an oxygen atom with three lone pairs of electrons. The third structure shows an oxygen atom double bonded to another oxygen atom with to lone pairs of electrons. The first oxygen atom is also double bonded to a carbon atom with two lone pairs of electrons.\" \/>\r\n<p id=\"fs-idp58268848\">Comparing the three formal charges, we can definitively identify the structure on the left as preferable because it has only formal charges of zero (Guideline 1).<\/p>\r\n<p id=\"fs-idp198975168\">As another example, the thiocyanate ion, an ion formed from a carbon atom, a nitrogen atom, and a sulfur atom, could have three different molecular structures: CNS<sup>\u2013<\/sup>, NCS<sup>\u2013<\/sup>, or CSN<sup>\u2013<\/sup>. The formal charges present in each of these molecular structures can help us pick the most likely arrangement of atoms. Possible Lewis structures and the formal charges for each of the three possible structures for the thiocyanate ion are shown here:<\/p>\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_Thiocyan_img-2.jpg\" alt=\"Two rows of structures and numbers are shown. The top row is labeled, \u201cStructure\u201d and depicts three Lewis structures and the bottom row is labeled, \u201cFormal charge.\u201d The left structure shows a carbon atom double bonded to a nitrogen atom with two lone electron pairs on one side and double bonded to a sulfur atom with two lone electron pairs on the other. The structure is surrounded by brackets and has a superscripted negative sign. Below this structure are the numbers negative one, zero, and zero. The middle structure shows a carbon atom with two lone pairs of electrons double bonded to a nitrogen atom that is double bonded to a sulfur atom with two lone electron pairs. The structure is surrounded by brackets and has a superscripted negative sign. Below this structure are the numbers negative two, positive one, and zero. The right structure shows a carbon atom with two lone electron pairs double bonded to a sulfur atom that is double bonded to a nitrogen atom with two lone electron pairs. The structure is surrounded by brackets and has a superscripted negative sign. Below this structure are the numbers negative two, positive two, and one.\" \/>\r\n<p id=\"fs-idp142093872\">Note that the sum of the formal charges in each case is equal to the charge of the ion (\u20131). However, the first arrangement of atoms is preferred because it has the lowest number of atoms with nonzero formal charges (Guideline 2). Also, it places the least electronegative atom in the center, and the negative charge on the more electronegative element (Guideline 4).<\/p>\r\n\r\n<div class=\"textbox shaded\" id=\"fs-idp56478288\">\r\n<h3>Example 3<\/h3>\r\n<p id=\"fs-idm17761552\">Nitrous oxide, N<sub>2<\/sub>O, commonly known as laughing gas, is used as an anesthetic in minor surgeries, such as the routine extraction of wisdom teeth. Which is the likely structure for nitrous oxide?<\/p>\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_N2O_img-2.jpg\" alt=\"Two Lewis structures are shown with the word \u201cor\u201d in between them. The left structure depicts a nitrogen atom with two lone pairs of electrons double bonded to a nitrogen that is double bonded to an oxygen with two lone pairs of electrons. The right structure shows a nitrogen atom with two lone pairs of electrons double bonded to an oxygen atom that is double bonded to a nitrogen atom with two lone pairs of electrons.\" class=\"aligncenter\" \/>\r\n\r\n&nbsp;\r\n<p id=\"fs-idp200124576\"><strong>Solution<\/strong>\r\nDetermining formal charge yields the following:<\/p>\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_N2Ofc_img-2.jpg\" alt=\"Two Lewis structures are shown with the word \u201cor\u201d in between them. The left structure depicts a nitrogen atom with two lone pairs of electrons double bonded to a nitrogen atom that is double bonded to an oxygen atom with two lone pairs of electrons. The numbers negative one, positive one, and zero are written above this structure. The right structure shows a nitrogen atom with two lone pairs of electrons double bonded to an oxygen atom that is double bonded to a nitrogen atom with two lone pairs of electrons. The numbers negative one, positive two, and negative one are written above this structure.\" class=\"aligncenter\" \/>\r\n<p id=\"fs-idp52242480\">The structure with a terminal oxygen atom best satisfies the criteria for the most stable distribution of formal charge:<\/p>\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_NNO_img-2.jpg\" alt=\"A Lewis structure is shown. A nitrogen atom with two lone pairs of electrons is double bonded to a nitrogen atom that is double bonded to an oxygen atom with two lone pairs of electrons.\" class=\"aligncenter\" \/>\r\n<p id=\"fs-idp29868224\">The number of atoms with formal charges are minimized (Guideline 2), and there is no formal charge larger than one (Guideline 2). This is again consistent with the preference for having the less electronegative atom in the central position.<\/p>\r\n&nbsp;\r\n<p id=\"fs-idp67397856\"><em><strong>Test Yourself<\/strong><\/em>\r\nWhich is the most likely molecular structure for the nitrite (NO<sub>2<\/sub><sup>\u2212<\/sup>) ion?<\/p>\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_NO2ion_img-2.jpg\" alt=\"Two Lewis structures are shown with the word \u201cor\u201d written between them. The left structure shows a nitrogen atom with two lone pairs of electrons double bonded to an oxygen atom with one lone pair of electrons that is single bonded to an oxygen atom with three lone pairs of electrons. Brackets surround this structure and there is a superscripted negative sign. The right structure shows an oxygen atom with two lone pairs of electrons double bonded to a nitrogen atom with one lone pair of electrons that is single bonded to an oxygen with three lone pairs of electrons. Brackets surround this structure and there is a superscripted negative sign.\" class=\"aligncenter\" \/>\r\n\r\n&nbsp;\r\n\r\n<strong>Answer<\/strong>\r\n\r\nONO<sup>\u2013<\/sup>\r\n\r\n<\/div>\r\n<\/section><section id=\"fs-idp26392016\">\r\n<h2>Resonance<\/h2>\r\n<p id=\"fs-idp81027344\">You may have noticed that the nitrite anion in <a href=\"#fs-idp56478288\" class=\"autogenerated-content\">Example 3<\/a> can have two possible structures with the atoms in the same positions. The electrons involved in the N\u2013O double bond, however, are in different positions:<\/p>\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_NO2res_img-2.jpg\" alt=\"Two Lewis structures are shown. The left structure shows an oxygen atom with three lone pairs of electrons single bonded to a nitrogen atom with one lone pair of electrons that is double bonded to an oxygen with two lone pairs of electrons. Brackets surround this structure, and there is a superscripted negative sign. The right structure shows an oxygen atom with two lone pairs of electrons double bonded to a nitrogen atom with one lone pair of electrons that is single bonded to an oxygen atom with three lone pairs of electrons. Brackets surround this structure, and there is a superscripted negative sign.\" class=\"aligncenter\" \/>\r\n<p id=\"fs-idm7645504\">If nitrite ions do indeed contain a single and a double bond, we would expect for the two bond lengths to be different. A double bond between two atoms is shorter (and stronger) than a single bond between the same two atoms. Experiments show, however, that both N\u2013O bonds in NO<sub>2<\/sub><sup>\u2212<\/sup> have the same strength and length, and are identical in all other properties.<\/p>\r\n<p id=\"fs-idp203952864\">It is not possible to write a single Lewis structure for NO<sub>2<\/sub><sup>\u2212<\/sup> in which nitrogen has an octet and both bonds are equivalent. Instead, we use the concept of <strong>resonance<\/strong>: if two or more Lewis structures with the same arrangement of atoms can be written for a molecule or ion, the actual distribution of electrons is an <em>average<\/em> of that shown by the various Lewis structures. The actual distribution of electrons in each of the nitrogen-oxygen bonds in NO<sub>2<\/sub><sup>\u2212<\/sup> is the average of a double bond and a single bond. We call the individual Lewis structures <strong>resonance forms<\/strong>. The actual electronic structure of the molecule (the average of the resonance forms) is called a <strong>resonance hybrid<\/strong> of the individual resonance forms. A double-headed arrow between Lewis structures indicates that they are resonance forms. Thus, the electronic structure of the NO<sub>2<\/sub><sup>\u2212<\/sup> ion is shown as:<\/p>\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_NO2resarr_img-2.jpg\" alt=\"Two Lewis structures are shown with a double headed arrow drawn between them. The left structure shows an oxygen atom with two lone pairs of electrons double bonded to a nitrogen atom with one lone pair of electrons that is single bonded to an oxygen atom with three lone pairs of electrons. Brackets surround this structure, and there is a superscripted negative sign. The right structure shows an oxygen atom with three lone pairs of electrons single bonded to a nitrogen atom with one lone pair of electrons that is double bonded to an oxygen atom with two lone pairs of electrons. Brackets surround this structure, and there is a superscripted negative sign.\" class=\"aligncenter\" \/>\r\n<p id=\"fs-idm12853328\">We should remember that a molecule described as a resonance hybrid <em>never<\/em> possesses an electronic structure described by either resonance form. It does not fluctuate between resonance forms; rather, the actual electronic structure is <em>always<\/em> the average of that shown by all resonance forms. George Wheland, one of the pioneers of resonance theory, used a historical analogy to describe the relationship between resonance forms and resonance hybrids. A medieval traveler, having never before seen a rhinoceros, described it as a hybrid of a dragon and a unicorn because it had many properties in common with both. Just as a rhinoceros is neither a dragon sometimes nor a unicorn at other times, a resonance hybrid is neither of its resonance forms at any given time. Like a rhinoceros, it is a real entity that experimental evidence has shown to exist. It has some characteristics in common with its resonance forms, but the resonance forms themselves are convenient, imaginary images (like the unicorn and the dragon).<\/p>\r\n<p id=\"fs-idp245720144\">The carbonate anion, CO<sub>3<\/sub><sup>2\u2212<\/sup>, provides a second example of resonance:<\/p>\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_CO3res_img-2.jpg\" alt=\"Three Lewis structures are shown with double headed arrows in between. Each structure is surrounded by brackets, and each has a superscripted two negative sign. The left structure depicts a carbon atom bonded to three oxygen atoms. It is single bonded to two of these oxygen atoms, each of which has three lone pairs of electrons, and double bonded to the third, which has two lone pairs of electrons. The double bond is located between the lower left oxygen atom and the carbon atom. The central and right structures are the same as the first, but the position of the double bonded oxygen has moved to the lower right oxygen in the central structure and to the top oxygen in the right structure.\" class=\"aligncenter\" width=\"637\" height=\"155\" \/>\r\n<p id=\"fs-idp6151664\">One oxygen atom must have a double bond to carbon to complete the octet on the central atom. All oxygen atoms, however, are equivalent, and the double bond could form from any one of the three atoms. This gives rise to three resonance forms of the carbonate ion. Because we can write three identical resonance structures, we know that the actual arrangement of electrons in the carbonate ion is the average of the three structures. Again, experiments show that all three C\u2013O bonds are exactly the same.<\/p>\r\n\r\n<div id=\"fs-idp77544496\" class=\"textbox shaded\">\r\n\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/OSC_Interactive_200-11-2.png\" alt=\"\u00a0\" width=\"116\" height=\"72\" class=\"alignleft\" \/>\r\n\r\n&nbsp;\r\n<p id=\"fs-idp64654736\">The online <a href=\"http:\/\/openstaxcollege.org\/l\/16LewisMake\">Lewis Structure Make<\/a> includes many examples to practice drawing resonance structures.<\/p>\r\n\r\n<\/div>\r\n<\/section><section id=\"fs-idm30786384\" class=\"summary\">\r\n<h2>Key Concepts and Summary<\/h2>\r\n<p id=\"fs-idp91470432\">In a Lewis structure, formal charges can be assigned to each atom by treating each bond as if one-half of the electrons are assigned to each atom. These hypothetical formal charges are a guide to determining the most appropriate Lewis structure. A structure in which the formal charges are as close to zero as possible is preferred. Resonance occurs in cases where two or more Lewis structures with identical arrangements of atoms but different distributions of electrons can be written. The actual distribution of electrons (the resonance hybrid) is an average of the distribution indicated by the individual Lewis structures (the resonance forms).<\/p>\r\n\r\n<\/section><section id=\"fs-idp77377632\" class=\"key-equations\">\r\n<h2>Key Equations<\/h2>\r\n<ul id=\"fs-idp63226528\">\r\n \t<li>$latex \\text{formal charge} = \\# \\;\\text{valence shell electrons (free atom)} \\; - \\;\\# \\;\\text{lone pair electrons}\\; - \\frac{1}{2} \\# \\;\\text{bonding electrons}$<\/li>\r\n<\/ul>\r\n<\/section><section id=\"fs-idp27655616\" class=\"exercises\">\r\n<div class=\"bcc-box bcc-info\">\r\n<h3>Exercises<\/h3>\r\n1. Write resonance forms that describe the distribution of electrons in each of these molecules or ions.\r\n<p id=\"fs-idm12563008\">a) sulfur dioxide, SO<sub>2<\/sub><\/p>\r\n<p id=\"fs-idm34089184\">b) carbonate ion, CO<sub>3<\/sub><sup>2\u2212<\/sup><\/p>\r\n<p id=\"fs-idp28289264\">c) hydrogen carbonate ion, HCO<sub>3<\/sub><sup>\u2212<\/sup> (C is bonded to an OH group and two O atoms)<\/p>\r\n<p id=\"fs-idp4938192\">d) pyridine:<\/p>\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_SO2resd_img-2.jpg\" alt=\"A Lewis structure depicts a hexagonal ring composed of five carbon atoms and one nitrogen atom. Each carbon atom is single bonded to a hydrogen atom.\" width=\"168\" height=\"151\" class=\"\" \/>\r\n<p id=\"fs-idp27230032\">e) the allyl ion:<\/p>\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_SO2rese_img-2.jpg\" alt=\"A Lewis structure shows a carbon atom single bonded to two hydrogen atoms and a second carbon atom. The second carbon atom is single bonded to a hydrogen atom and a third carbon atom. The third carbon atom is single bonded to two hydrogen atoms. The whole structure is surrounded by brackets, and there is a superscripted negative sign.\" width=\"253\" height=\"99\" class=\"\" \/>\r\n\r\n2. Sodium nitrite, which has been used to preserve bacon and other meats, is an ionic compound. Write the resonance forms of the nitrite ion, NO<sub>2<\/sub><sup>\u2013<\/sup>.\r\n\r\n3. Write the Lewis structures for the following, and include resonance structures where appropriate. Indicate which has the strongest carbon-oxygen bond.\r\n<p id=\"fs-idp9268576\">a) CO<sub>2 \u00a0 \u00a0 \u00a0 \u00a0\u00a0<\/sub>b) CO<\/p>\r\n4. Determine the formal charge of each element in the following:\r\n<p id=\"fs-idp36610048\">a) HCl \u00a0 \u00a0 \u00a0 \u00a0\u00a0b) CF<sub>4 \u00a0 \u00a0 \u00a0 \u00a0\u00a0<\/sub>c) PCl<sub>3 \u00a0 \u00a0 \u00a0 \u00a0\u00a0<\/sub>d) PF<sub>5<\/sub><\/p>\r\n5. Calculate the formal charge of chlorine in the molecules Cl<sub>2<\/sub>, BeCl<sub>2<\/sub>, and ClF<sub>5<\/sub>.\r\n\r\n6. Draw all possible resonance structures for each of these compounds. Determine the formal charge on each atom in each of the resonance structures:\r\n<p id=\"fs-idp131483136\">a) O<sub>3 \u00a0 \u00a0 \u00a0 \u00a0\u00a0<\/sub>b) SO<sub>2 \u00a0 \u00a0 \u00a0 \u00a0\u00a0<\/sub>c) NO<sub>2<\/sub><sup>\u2212 \u00a0 \u00a0 \u00a0 \u00a0\u00a0<\/sup>d) NO<sub>3<\/sub><sup>\u2212<\/sup><\/p>\r\n7. Based on formal charge considerations, which of the following would likely be the correct arrangement of atoms in hypochlorous acid: HOCl or OClH?\r\n\r\n8. Draw the structure of hydroxylamine, H<sub>3<\/sub>NO, and assign formal charges; look up the structure. Is the actual structure consistent with the formal charges?\r\n\r\n9. Write the Lewis structure and chemical formula of the compound with a molar mass of about 70 g\/mol that contains 19.7% nitrogen and 80.3% fluorine by mass, and determine the formal charge of the atoms in this compound.\r\n\r\n10. Sulfuric acid is the industrial chemical produced in greatest quantity worldwide. About 90 billion pounds are produced each year in the United States alone. Write the Lewis structure for sulfuric acid, H<sub>2<\/sub>SO<sub>4<\/sub>, which has two oxygen atoms and two OH groups bonded to the sulfur.\r\n\r\n&nbsp;\r\n\r\n<strong>Answers<\/strong>\r\n<p id=\"fs-idp22600016\">1. a)\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_Ques2ansa_img-2.jpg\" alt=\"Two Lewis structures are shown with a double-headed arrow in between. The left structure shows a sulfur atom with a lone pair of electrons single bonded to the left to an oxygen atom with three lone pairs of electrons. The sulfur atom is also double bonded on the right to an oxygen atom with two lone pairs of electrons. The right structure depicts the same atoms, but this time the double bond is between the left oxygen and the sulfur atom. The lone pairs of electrons have also shifted to account for the change of bond types. The sulfur atom in the right structures, also has a third electron dot below it.\" width=\"430\" height=\"57\" class=\"\" \/><\/p>\r\nb)\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_Ques2ansb_img-2.jpg\" alt=\"Three Lewis structures are shown, with double-headed arrows in between, each surrounded by brackets and a superscripted two negative sign. The left structure depicts a carbon atom bonded to three oxygen atoms. It is single bonded to two of these oxygen atoms, each of which has three lone pairs of electrons, and double bonded to the third, which has two lone pairs of electrons. The double bond is located between the bottom oxygen and the carbon. The central and right structures are the same as the first, but the position of the double bonded oxygen has moved to the left oxygen in the right structure while the central structure only has single bonds. The lone pairs of electrons change to correspond with the bonds as well.\" \/>\r\n\r\nc)\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_Ques2ansc_img-2.jpg\" alt=\"Two Lewis structures are shown, with a double-headed arrow in between, each surrounded by brackets and a superscripted negative sign. The left structure depicts a carbon atom bonded to three oxygen atoms. It is single bonded to one of these oxygen atoms, which has three lone pairs of electrons, and double bonded to the other two, which have two lone pairs of electrons. One of the double bonded oxygen atoms also has a single bond to a hydrogen atom. The right structure is the same as the first, but there is only one double bonded oxygen. The oxygen with the single bonded hydrogen now has a single bond to the carbon atom. The lone pairs of electrons have also changed to correspond with the bonds.\" width=\"552\" height=\"130\" class=\"\" \/>\r\n\r\nd)\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_Ques2ansd_img-2.jpg\" alt=\"Two Lewis structures are shown with a double-headed arrow in between. The left structure depicts a hexagonal ring composed of five carbon atoms, each single bonded to a hydrogen atom, and one nitrogen atom that has a lone pair of electrons. The ring has alternating single and double bonds. The right structure is the same as the first, but each double bond has rotated to a new position.\" width=\"466\" height=\"169\" class=\"\" \/>\r\n\r\ne)\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_Ques2anse_img-2.jpg\" alt=\"Two Lewis structures are shown with a double-headed arrow in between. The left structure shows a carbon atom single bonded to two hydrogen atoms and a second carbon atom. The second carbon atom is single bonded to a hydrogen atom and double bonded to a third carbon atom. The third carbon atom is single bonded to two hydrogen atoms. The whole structure is surrounded by brackets and a superscripted negative sign. The right structure shows a carbon atom single bonded to two hydrogen atoms and double bonded to a second carbon atom. The second carbon atom is single bonded to a hydrogen atom and a third carbon atom. The third carbon atom is single bonded to two hydrogen atoms. The whole structure is surrounded by brackets and a superscripted negative sign.\" width=\"611\" height=\"81\" class=\"\" \/>\r\n\r\n2.\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_Ques11ans_img-2.jpg\" alt=\"Two pairs of Lewis structures are shown with a double-headed arrow in between each pair. The left structure of the first pair shows a nitrogen atom with one lone pair of electrons single bonded to an oxygen atom with three lone pairs of electrons. It is also double bonded to an oxygen with two lone pairs of electrons. The right image of this pair depicts the mirror image of the left. Both images are surrounded by brackets and a superscripted negative sign. They are labeled, \u201cFor N O subscript two superscript negative sign.\u201d The left structure of the second pair shows an oxygen atom with one lone pair of electrons single bonded to an oxygen atom with three lone pairs of electrons. It is also double bonded to an oxygen atom with two lone pairs of electrons. The right structure appears as a mirror image of the left. These structures are labeled, \u201cFor O subscript three.\u201d\" \/>\r\n<p id=\"fs-idp32184800\">3. a)\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_Ques13ansb_img-2.jpg\" alt=\"This structure shows a carbon atom double bonded to two oxygen atoms, each of which has two lone pairs of electrons.\" width=\"276\" height=\"51\" class=\"\" \/><\/p>\r\nb)\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_Ques13ansc_img-2.jpg\" alt=\"The right structure of this pair shows a carbon atom with one lone pair of electrons triple bonded to an oxygen with one lone pair of electrons.\" width=\"273\" height=\"26\" class=\"\" \/>\r\nCO has the strongest carbon-oxygen bond because there is a triple bond joining C and O. CO<sub>2<\/sub> has double bonds.\r\n<p id=\"fs-idm21325744\">4. a) H: 0, Cl: 0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0b) C: 0, F: 0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0 c) P: 0, Cl 0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0d) P: 0, F: 0<\/p>\r\n<p id=\"fs-idm369848352\">5. Cl in Cl<sub>2<\/sub>: 0; \u00a0 \u00a0 Cl in BeCl<sub>2<\/sub>: 0; \u00a0 \u00a0 Cl in ClF<sub>5<\/sub>: 0<\/p>\r\n<p id=\"fs-idp124490112\">6. a)\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_Exercis12a_img-2.jpg\" alt=\"Two Lewis structures are shown with a double-headed arrow in between. The left structure shows an oxygen atom with one lone pair of electrons single bonded to an oxygen atom with three lone pairs of electrons. It is also double bonded to an oxygen atom with two lone pairs of electrons. The symbols and numbers below this structure read, \u201c( 0 ), ( positive 1 ), ( negative 1 ).\u201d The phrase, \u201cFormal charge,\u201d and a right-facing arrow lie to the left of this structure. The right structure appears as a mirror image of the left and the symbols and numbers below this structure read, \u201c( negative 1 ), ( positive 1 ), ( 0 ).\u201d\" \/><\/p>\r\nb)\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_Exercis12b_img-2.jpg\" alt=\"Two Lewis structures are shown, with a double-headed arrow in between. The left structure shows a sulfur atom with one lone pair of electrons single bonded to an oxygen atom with three lone pairs of electrons. The sulfur atom also double bonded to an oxygen atom with two lone pairs of electrons. The symbols and numbers below this structure read, \u201c( negative 1 ), ( positive 1 ), ( 0 ).\u201d The right structure appears as a mirror image of the left and the symbols and numbers below this structure read, \u201c( 0 ), ( positive 1 ), ( negative 1 ).\u201d\" \/>\r\n\r\nc)\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_Exercis12c_img-2.jpg\" alt=\"[Two Lewis structures are shown, with brackets surrounding each with a superscripted negative sign and a double ended arrow in between. The left structure shows a nitrogen atom with one lone pair of electrons single bonded to an oxygen atom with three lone pairs of electrons and double bonded to an oxygen atom with two lone pairs of electrons. The symbols and numbers below this structure read \u201copen parenthesis, 0, close parenthesis, open parenthesis, 0, close parenthesis, open parenthesis, negative 1, close parenthesis. The right structure appears as a mirror image of the left and the symbols and numbers below this structure read \u201copen parenthesis, negative 1, close parenthesis, open parenthesis, 0, close parenthesis, open parenthesis, 0, close parenthesis.]\" \/>\r\n\r\nd)\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_Exercis12d_img-2.jpg\" alt=\"[Three Lewis structures are shown, with brackets surrounding each with a superscripted negative sign and a double ended arrow in between. The left structure shows a nitrogen atom single bonded to two oxygen atoms, each with three lone pairs of electrons and double bonded to an oxygen atom with two lone pairs of electrons. The single bonded oxygen atoms are labeled, from the top of the structure and going clockwise, \u201copen parenthesis, negative 1, close parenthesis, open parenthesis, positive 1, close parenthesis\u201d. The symbols and numbers below this structure read \u201copen parenthesis, 0, close parenthesis, open parenthesis, negative 1, close parenthesis. The middle structure shows a nitrogen atom single bonded to two oxygen atoms, each with three lone pairs of electrons, one of which is labeled \u201copen parenthesis, positive 1, close parenthesis\u201d and double bonded to an oxygen atom with two lone pairs of electrons labeled \u201copen parenthesis, 0, close parenthesis\u201d. The symbols and numbers below this structure read \u201copen parenthesis, negative 1, close parenthesis, open parenthesis, negative 1, close parenthesis. The right structure shows a nitrogen atom single bonded to two oxygen atoms, each with three lone pairs of electrons and double bonded to an oxygen atom with two lone pairs of electrons. One of the single bonded oxygen atoms is labeled, \u201copen parenthesis, negative 1, close parenthesis while the double bonded oxygen is labeled, \u201copen parenthesis, positive 1, close parenthesis\u201d. The symbols and numbers below this structure read \u201copen parenthesis, negative 1, close parenthesis\u201d and \u201copen parenthesis, 0, close parenthesis\u201d.]\" \/>\r\n<p id=\"fs-idm333542560\">7. HOCl<\/p>\r\n<p id=\"fs-idp6445072\">8. The structure that gives zero formal charges is consistent with the actual structure:<\/p>\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_OHamine_img-2.jpg\" alt=\"A Lewis structure shows a nitrogen atom with one lone pair of electrons single bonded to two hydrogen atoms and an oxygen atom which has two lone pairs of electrons. The oxygen atom is single bonded to a hydrogen atom.\" width=\"134\" height=\"101\" class=\"\" \/>\r\n<p id=\"fs-idp35441312\">9. NF<sub>3<\/sub><\/p>\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_NF3_img-2.jpg\" alt=\"A Lewis structure shows a nitrogen atom with one lone pair of electrons single bonded to three fluorine atoms, each with three lone pairs of electrons.\" width=\"82\" height=\"125\" class=\"\" \/>\r\n\r\n10.\r\n<img src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_H2SO4-2.jpg\" alt=\"A Lewis structure shows a hydrogen atom single bonded to an oxygen atom with two lone pairs of electrons. The oxygen atom is single bonded to a sulfur atom. The sulfur atom is double bonded to two oxygen atoms, each of which have three lone pairs of electrons, and single bonded to an oxygen atom with two lone pairs of electrons. This oxygen atom is single bonded to a hydrogen atom.\" width=\"171\" height=\"137\" class=\"\" \/>\r\n\r\n<\/div>\r\n<\/section>\r\n<div>\r\n<h2>Glossary<\/h2>\r\n<strong>formal charge:\u00a0<\/strong>charge that would result on an atom by taking the number of valence electrons on the neutral atom and subtracting the nonbonding electrons and the number of bonds (one-half of the bonding electrons)\r\n\r\n<strong>molecular structure:\u00a0<\/strong>arrangement of atoms in a molecule or ion\r\n\r\n<strong>resonance:\u00a0<\/strong>situation in which one Lewis structure is insufficient to describe the bonding in a molecule and the average of multiple structures is observed\r\n\r\n<strong>resonance forms:\u00a0<\/strong>two or more Lewis structures that have the same arrangement of atoms but different arrangements of electrons\r\n\r\n<strong>resonance hybrid:\u00a0<\/strong>average of the resonance forms shown by the individual Lewis structures\r\n\r\n<\/div>","rendered":"<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>Compute formal charges for atoms in any Lewis structure<\/li>\n<li>Use formal charges to identify the most reasonable Lewis structure for a given molecule<\/li>\n<li>Explain the concept of resonance and draw Lewis structures representing resonance forms for a given molecule<\/li>\n<\/ul>\n<\/div>\n<p id=\"fs-idm49375088\">In the previous section, we discussed how to write Lewis structures for molecules and polyatomic ions. As we have seen, however, in some cases, there is seemingly more than one valid structure for a molecule. We can use the concept of formal charges to help us predict the most appropriate Lewis structure when more than one is reasonable.<\/p>\n<section id=\"fs-idp11710128\">\n<h2>Calculating Formal Charge<\/h2>\n<p id=\"fs-idm16217792\">The <strong>formal charge<\/strong> of an atom in a molecule is the <em>hypothetical<\/em> charge the atom would have if we could redistribute the electrons in the bonds evenly between the atoms. Another way of saying this is that formal charge results when we take the number of valence electrons of a neutral atom, subtract the nonbonding electrons, and then subtract the number of bonds connected to that atom in the Lewis structure.<\/p>\n<p id=\"fs-idp126842624\">Thus, we calculate formal charge as follows:<\/p>\n<div class=\"equation\" id=\"fs-idp90018832\" style=\"text-align: center\">[latex]\\text{formal charge} = \\# \\;\\text{valence shell electrons (free atom)} \\; - \\;\\# \\;\\text{lone pair electrons}\\; - \\frac{1}{2} \\# \\;\\text{bonding electrons}[\/latex]<\/div>\n<p id=\"fs-idp88204016\">We can double-check formal charge calculations by determining the sum of the formal charges for the whole structure. The sum of the formal charges of all atoms in a molecule must be zero; the sum of the formal charges in an ion should equal the charge of the ion.<\/p>\n<p id=\"fs-idp15200480\">We must remember that the formal charge calculated for an atom is not the <em>actual<\/em> charge of the atom in the molecule. Formal charge is only a useful bookkeeping procedure; it does not indicate the presence of actual charges.<\/p>\n<div class=\"textbox shaded\" id=\"fs-idm21031296\">\n<h3>Example 1<\/h3>\n<p id=\"fs-idp63607984\">Assign formal charges to each atom in the interhalogen ion ICl<sub>4<\/sub><sup>\u2212<\/sup>.<\/p>\n<p>&nbsp;<\/p>\n<p id=\"fs-idp251046608\"><strong>Solution<\/strong><\/p>\n<ol id=\"fs-idm30182496\" class=\"stepwise\">\n<li><em>We divide the bonding electron pairs equally for all I\u2013Cl bonds:<br \/>\n<\/em><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_ICL4_img-2.jpg\" alt=\"A Lewis structure is shown. An iodine atom with two lone pairs of electrons is single bonded to four chlorine atoms, each of which has three lone pairs of electrons. Brackets surround the structure and there is a superscripted negative sign.\" class=\"aligncenter\" width=\"186\" height=\"137\" \/><\/li>\n<li><em>We assign lone pairs of electrons to their atoms<\/em>. Each Cl atom now has seven electrons assigned to it, and the I atom has eight.<\/li>\n<li><em><em>Subtract this number from the number of valence electrons for the neutral atom:<\/em><\/em>I: 7 \u2013 8 = \u20131Cl: 7 \u2013 7 = 0The sum of the formal charges of all the atoms equals \u20131, which is identical to the charge of the ion (\u20131).<\/li>\n<\/ol>\n<p>&nbsp;<\/p>\n<p id=\"fs-idp65521520\"><em><strong>Test Yourself<\/strong><\/em><br \/>\nCalculate the formal charge for each atom in the carbon monoxide molecule:<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_CO_img-2.jpg\" alt=\"A Lewis structure is shown. A carbon atom with one lone pair of electrons is triple bonded to an oxygen with one lone pair of electrons.\" class=\"aligncenter\" width=\"101\" height=\"33\" \/><\/p>\n<p>&nbsp;<\/p>\n<p><em><strong>Answer<\/strong><\/em><\/p>\n<p>C \u22121, O +1<\/p>\n<\/div>\n<div class=\"textbox shaded\" id=\"fs-idp59620704\">\n<h3>Example 2<\/h3>\n<p id=\"fs-idp26752688\">Assign formal charges to each atom in the interhalogen molecule BrCl<sub>3<\/sub>.<\/p>\n<p>&nbsp;<\/p>\n<p id=\"fs-idp34669824\"><strong>Solution<\/strong><\/p>\n<ol id=\"fs-idp8784496\" class=\"stepwise\">\n<li><em>Assign one of the electrons in each Br\u2013Cl bond to the Br atom and one to the Cl atom in that bond:<br \/>\n<\/em><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_BRCL3_img-2.jpg\" alt=\"A Lewis structure is shown. A bromine atom with two lone pairs of electrons is single bonded to three chlorine atoms, each of which has three lone pairs of electrons.\" class=\"aligncenter\" width=\"150\" height=\"99\" \/><\/li>\n<li><em>Assign the lone pairs to their atom.<\/em> Now each Cl atom has seven electrons and the Br atom has seven electrons.<\/li>\n<li><em>Subtract this number from the number of valence electrons for the neutral atom.<\/em> This gives the formal charge:Br: 7 \u2013 7 = 0Cl: 7 \u2013 7 = 0All atoms in BrCl<sub>3<\/sub> have a formal charge of zero, and the sum of the formal charges totals zero, as it must in a neutral molecule.<\/li>\n<\/ol>\n<p>&nbsp;<\/p>\n<p id=\"fs-idp245106544\"><em><strong>Test yourself<\/strong><\/em><br \/>\nDetermine the formal charge for each atom in NCl<sub>3<\/sub>.<\/p>\n<p>&nbsp;<\/p>\n<p><em><strong>Answer<\/strong><\/em><\/p>\n<p id=\"fs-idm22650544\">N: 0; all three Cl atoms: 0<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_Ex070402_img-2.jpg\" alt=\"A Lewis structure is shown. A nitrogen atom with one lone pair of electrons is single bonded to three chlorine atoms, each of which has three lone pairs of electrons.\" class=\"aligncenter\" width=\"275\" height=\"99\" \/><\/p>\n<\/div>\n<\/section>\n<section id=\"fs-idp27335568\">\n<h2>Using Formal Charge to Predict Molecular Structure<\/h2>\n<p id=\"fs-idp93581584\">The arrangement of atoms in a molecule or ion is called its <strong>molecular structure<\/strong>. In many cases, following the steps for writing Lewis structures may lead to more than one possible molecular structure\u2014different multiple bond and lone-pair electron placements or different arrangements of atoms, for instance. A few guidelines involving formal charge can be helpful in deciding which of the possible structures is most likely for a particular molecule or ion:<\/p>\n<ol id=\"fs-idp69636800\">\n<li>A molecular structure in which all formal charges are zero is preferable to one in which some formal charges are not zero.<\/li>\n<li>If the Lewis structure must have nonzero formal charges, the arrangement with the smallest nonzero formal charges is preferable.<\/li>\n<li>Lewis structures are preferable when adjacent formal charges are zero or of the opposite sign.<\/li>\n<li>When we must choose among several Lewis structures with similar distributions of formal charges, the structure with the negative formal charges on the more electronegative atoms is preferable.<\/li>\n<\/ol>\n<p id=\"fs-idp52450560\">To see how these guidelines apply, let us consider some possible structures for carbon dioxide, CO<sub>2<\/sub>. We know from our previous discussion that the less electronegative atom typically occupies the central position, but formal charges allow us to understand <em>why<\/em> this occurs. We can draw three possibilities for the structure: carbon in the center and double bonds, carbon in the center with a single and triple bond, and oxygen in the center with double bonds:<\/p>\n<p><img decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_CO2pos_img-2.jpg\" alt=\"Three Lewis structures are shown. The left and right structures show a carbon atom double bonded to two oxygen atoms, each of which has two lone pairs of electrons. The center structure shows a carbon atom that is triple bonded to an oxygen atom with one lone pair of electrons and single bonded to an oxygen atom with three lone pairs of electrons. The third structure shows an oxygen atom double bonded to another oxygen atom with to lone pairs of electrons. The first oxygen atom is also double bonded to a carbon atom with two lone pairs of electrons.\" \/><\/p>\n<p id=\"fs-idp58268848\">Comparing the three formal charges, we can definitively identify the structure on the left as preferable because it has only formal charges of zero (Guideline 1).<\/p>\n<p id=\"fs-idp198975168\">As another example, the thiocyanate ion, an ion formed from a carbon atom, a nitrogen atom, and a sulfur atom, could have three different molecular structures: CNS<sup>\u2013<\/sup>, NCS<sup>\u2013<\/sup>, or CSN<sup>\u2013<\/sup>. The formal charges present in each of these molecular structures can help us pick the most likely arrangement of atoms. Possible Lewis structures and the formal charges for each of the three possible structures for the thiocyanate ion are shown here:<\/p>\n<p><img decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_Thiocyan_img-2.jpg\" alt=\"Two rows of structures and numbers are shown. The top row is labeled, \u201cStructure\u201d and depicts three Lewis structures and the bottom row is labeled, \u201cFormal charge.\u201d The left structure shows a carbon atom double bonded to a nitrogen atom with two lone electron pairs on one side and double bonded to a sulfur atom with two lone electron pairs on the other. The structure is surrounded by brackets and has a superscripted negative sign. Below this structure are the numbers negative one, zero, and zero. The middle structure shows a carbon atom with two lone pairs of electrons double bonded to a nitrogen atom that is double bonded to a sulfur atom with two lone electron pairs. The structure is surrounded by brackets and has a superscripted negative sign. Below this structure are the numbers negative two, positive one, and zero. The right structure shows a carbon atom with two lone electron pairs double bonded to a sulfur atom that is double bonded to a nitrogen atom with two lone electron pairs. The structure is surrounded by brackets and has a superscripted negative sign. Below this structure are the numbers negative two, positive two, and one.\" \/><\/p>\n<p id=\"fs-idp142093872\">Note that the sum of the formal charges in each case is equal to the charge of the ion (\u20131). However, the first arrangement of atoms is preferred because it has the lowest number of atoms with nonzero formal charges (Guideline 2). Also, it places the least electronegative atom in the center, and the negative charge on the more electronegative element (Guideline 4).<\/p>\n<div class=\"textbox shaded\" id=\"fs-idp56478288\">\n<h3>Example 3<\/h3>\n<p id=\"fs-idm17761552\">Nitrous oxide, N<sub>2<\/sub>O, commonly known as laughing gas, is used as an anesthetic in minor surgeries, such as the routine extraction of wisdom teeth. Which is the likely structure for nitrous oxide?<\/p>\n<p><img decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_N2O_img-2.jpg\" alt=\"Two Lewis structures are shown with the word \u201cor\u201d in between them. The left structure depicts a nitrogen atom with two lone pairs of electrons double bonded to a nitrogen that is double bonded to an oxygen with two lone pairs of electrons. The right structure shows a nitrogen atom with two lone pairs of electrons double bonded to an oxygen atom that is double bonded to a nitrogen atom with two lone pairs of electrons.\" class=\"aligncenter\" \/><\/p>\n<p>&nbsp;<\/p>\n<p id=\"fs-idp200124576\"><strong>Solution<\/strong><br \/>\nDetermining formal charge yields the following:<\/p>\n<p><img decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_N2Ofc_img-2.jpg\" alt=\"Two Lewis structures are shown with the word \u201cor\u201d in between them. The left structure depicts a nitrogen atom with two lone pairs of electrons double bonded to a nitrogen atom that is double bonded to an oxygen atom with two lone pairs of electrons. The numbers negative one, positive one, and zero are written above this structure. The right structure shows a nitrogen atom with two lone pairs of electrons double bonded to an oxygen atom that is double bonded to a nitrogen atom with two lone pairs of electrons. The numbers negative one, positive two, and negative one are written above this structure.\" class=\"aligncenter\" \/><\/p>\n<p id=\"fs-idp52242480\">The structure with a terminal oxygen atom best satisfies the criteria for the most stable distribution of formal charge:<\/p>\n<p><img decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_NNO_img-2.jpg\" alt=\"A Lewis structure is shown. A nitrogen atom with two lone pairs of electrons is double bonded to a nitrogen atom that is double bonded to an oxygen atom with two lone pairs of electrons.\" class=\"aligncenter\" \/><\/p>\n<p id=\"fs-idp29868224\">The number of atoms with formal charges are minimized (Guideline 2), and there is no formal charge larger than one (Guideline 2). This is again consistent with the preference for having the less electronegative atom in the central position.<\/p>\n<p>&nbsp;<\/p>\n<p id=\"fs-idp67397856\"><em><strong>Test Yourself<\/strong><\/em><br \/>\nWhich is the most likely molecular structure for the nitrite (NO<sub>2<\/sub><sup>\u2212<\/sup>) ion?<\/p>\n<p><img decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_NO2ion_img-2.jpg\" alt=\"Two Lewis structures are shown with the word \u201cor\u201d written between them. The left structure shows a nitrogen atom with two lone pairs of electrons double bonded to an oxygen atom with one lone pair of electrons that is single bonded to an oxygen atom with three lone pairs of electrons. Brackets surround this structure and there is a superscripted negative sign. The right structure shows an oxygen atom with two lone pairs of electrons double bonded to a nitrogen atom with one lone pair of electrons that is single bonded to an oxygen with three lone pairs of electrons. Brackets surround this structure and there is a superscripted negative sign.\" class=\"aligncenter\" \/><\/p>\n<p>&nbsp;<\/p>\n<p><strong>Answer<\/strong><\/p>\n<p>ONO<sup>\u2013<\/sup><\/p>\n<\/div>\n<\/section>\n<section id=\"fs-idp26392016\">\n<h2>Resonance<\/h2>\n<p id=\"fs-idp81027344\">You may have noticed that the nitrite anion in <a href=\"#fs-idp56478288\" class=\"autogenerated-content\">Example 3<\/a> can have two possible structures with the atoms in the same positions. The electrons involved in the N\u2013O double bond, however, are in different positions:<\/p>\n<p><img decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_NO2res_img-2.jpg\" alt=\"Two Lewis structures are shown. The left structure shows an oxygen atom with three lone pairs of electrons single bonded to a nitrogen atom with one lone pair of electrons that is double bonded to an oxygen with two lone pairs of electrons. Brackets surround this structure, and there is a superscripted negative sign. The right structure shows an oxygen atom with two lone pairs of electrons double bonded to a nitrogen atom with one lone pair of electrons that is single bonded to an oxygen atom with three lone pairs of electrons. Brackets surround this structure, and there is a superscripted negative sign.\" class=\"aligncenter\" \/><\/p>\n<p id=\"fs-idm7645504\">If nitrite ions do indeed contain a single and a double bond, we would expect for the two bond lengths to be different. A double bond between two atoms is shorter (and stronger) than a single bond between the same two atoms. Experiments show, however, that both N\u2013O bonds in NO<sub>2<\/sub><sup>\u2212<\/sup> have the same strength and length, and are identical in all other properties.<\/p>\n<p id=\"fs-idp203952864\">It is not possible to write a single Lewis structure for NO<sub>2<\/sub><sup>\u2212<\/sup> in which nitrogen has an octet and both bonds are equivalent. Instead, we use the concept of <strong>resonance<\/strong>: if two or more Lewis structures with the same arrangement of atoms can be written for a molecule or ion, the actual distribution of electrons is an <em>average<\/em> of that shown by the various Lewis structures. The actual distribution of electrons in each of the nitrogen-oxygen bonds in NO<sub>2<\/sub><sup>\u2212<\/sup> is the average of a double bond and a single bond. We call the individual Lewis structures <strong>resonance forms<\/strong>. The actual electronic structure of the molecule (the average of the resonance forms) is called a <strong>resonance hybrid<\/strong> of the individual resonance forms. A double-headed arrow between Lewis structures indicates that they are resonance forms. Thus, the electronic structure of the NO<sub>2<\/sub><sup>\u2212<\/sup> ion is shown as:<\/p>\n<p><img decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_NO2resarr_img-2.jpg\" alt=\"Two Lewis structures are shown with a double headed arrow drawn between them. The left structure shows an oxygen atom with two lone pairs of electrons double bonded to a nitrogen atom with one lone pair of electrons that is single bonded to an oxygen atom with three lone pairs of electrons. Brackets surround this structure, and there is a superscripted negative sign. The right structure shows an oxygen atom with three lone pairs of electrons single bonded to a nitrogen atom with one lone pair of electrons that is double bonded to an oxygen atom with two lone pairs of electrons. Brackets surround this structure, and there is a superscripted negative sign.\" class=\"aligncenter\" \/><\/p>\n<p id=\"fs-idm12853328\">We should remember that a molecule described as a resonance hybrid <em>never<\/em> possesses an electronic structure described by either resonance form. It does not fluctuate between resonance forms; rather, the actual electronic structure is <em>always<\/em> the average of that shown by all resonance forms. George Wheland, one of the pioneers of resonance theory, used a historical analogy to describe the relationship between resonance forms and resonance hybrids. A medieval traveler, having never before seen a rhinoceros, described it as a hybrid of a dragon and a unicorn because it had many properties in common with both. Just as a rhinoceros is neither a dragon sometimes nor a unicorn at other times, a resonance hybrid is neither of its resonance forms at any given time. Like a rhinoceros, it is a real entity that experimental evidence has shown to exist. It has some characteristics in common with its resonance forms, but the resonance forms themselves are convenient, imaginary images (like the unicorn and the dragon).<\/p>\n<p id=\"fs-idp245720144\">The carbonate anion, CO<sub>3<\/sub><sup>2\u2212<\/sup>, provides a second example of resonance:<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_CO3res_img-2.jpg\" alt=\"Three Lewis structures are shown with double headed arrows in between. Each structure is surrounded by brackets, and each has a superscripted two negative sign. The left structure depicts a carbon atom bonded to three oxygen atoms. It is single bonded to two of these oxygen atoms, each of which has three lone pairs of electrons, and double bonded to the third, which has two lone pairs of electrons. The double bond is located between the lower left oxygen atom and the carbon atom. The central and right structures are the same as the first, but the position of the double bonded oxygen has moved to the lower right oxygen in the central structure and to the top oxygen in the right structure.\" class=\"aligncenter\" width=\"637\" height=\"155\" \/><\/p>\n<p id=\"fs-idp6151664\">One oxygen atom must have a double bond to carbon to complete the octet on the central atom. All oxygen atoms, however, are equivalent, and the double bond could form from any one of the three atoms. This gives rise to three resonance forms of the carbonate ion. Because we can write three identical resonance structures, we know that the actual arrangement of electrons in the carbonate ion is the average of the three structures. Again, experiments show that all three C\u2013O bonds are exactly the same.<\/p>\n<div id=\"fs-idp77544496\" class=\"textbox shaded\">\n<p><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/OSC_Interactive_200-11-2.png\" alt=\"\u00a0\" width=\"116\" height=\"72\" class=\"alignleft\" \/><\/p>\n<p>&nbsp;<\/p>\n<p id=\"fs-idp64654736\">The online <a href=\"http:\/\/openstaxcollege.org\/l\/16LewisMake\">Lewis Structure Make<\/a> includes many examples to practice drawing resonance structures.<\/p>\n<\/div>\n<\/section>\n<section id=\"fs-idm30786384\" class=\"summary\">\n<h2>Key Concepts and Summary<\/h2>\n<p id=\"fs-idp91470432\">In a Lewis structure, formal charges can be assigned to each atom by treating each bond as if one-half of the electrons are assigned to each atom. These hypothetical formal charges are a guide to determining the most appropriate Lewis structure. A structure in which the formal charges are as close to zero as possible is preferred. Resonance occurs in cases where two or more Lewis structures with identical arrangements of atoms but different distributions of electrons can be written. The actual distribution of electrons (the resonance hybrid) is an average of the distribution indicated by the individual Lewis structures (the resonance forms).<\/p>\n<\/section>\n<section id=\"fs-idp77377632\" class=\"key-equations\">\n<h2>Key Equations<\/h2>\n<ul id=\"fs-idp63226528\">\n<li>[latex]\\text{formal charge} = \\# \\;\\text{valence shell electrons (free atom)} \\; - \\;\\# \\;\\text{lone pair electrons}\\; - \\frac{1}{2} \\# \\;\\text{bonding electrons}[\/latex]<\/li>\n<\/ul>\n<\/section>\n<section id=\"fs-idp27655616\" class=\"exercises\">\n<div class=\"bcc-box bcc-info\">\n<h3>Exercises<\/h3>\n<p>1. Write resonance forms that describe the distribution of electrons in each of these molecules or ions.<\/p>\n<p id=\"fs-idm12563008\">a) sulfur dioxide, SO<sub>2<\/sub><\/p>\n<p id=\"fs-idm34089184\">b) carbonate ion, CO<sub>3<\/sub><sup>2\u2212<\/sup><\/p>\n<p id=\"fs-idp28289264\">c) hydrogen carbonate ion, HCO<sub>3<\/sub><sup>\u2212<\/sup> (C is bonded to an OH group and two O atoms)<\/p>\n<p id=\"fs-idp4938192\">d) pyridine:<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_SO2resd_img-2.jpg\" alt=\"A Lewis structure depicts a hexagonal ring composed of five carbon atoms and one nitrogen atom. Each carbon atom is single bonded to a hydrogen atom.\" width=\"168\" height=\"151\" class=\"\" \/><\/p>\n<p id=\"fs-idp27230032\">e) the allyl ion:<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_SO2rese_img-2.jpg\" alt=\"A Lewis structure shows a carbon atom single bonded to two hydrogen atoms and a second carbon atom. The second carbon atom is single bonded to a hydrogen atom and a third carbon atom. The third carbon atom is single bonded to two hydrogen atoms. The whole structure is surrounded by brackets, and there is a superscripted negative sign.\" width=\"253\" height=\"99\" class=\"\" \/><\/p>\n<p>2. Sodium nitrite, which has been used to preserve bacon and other meats, is an ionic compound. Write the resonance forms of the nitrite ion, NO<sub>2<\/sub><sup>\u2013<\/sup>.<\/p>\n<p>3. Write the Lewis structures for the following, and include resonance structures where appropriate. Indicate which has the strongest carbon-oxygen bond.<\/p>\n<p id=\"fs-idp9268576\">a) CO<sub>2 \u00a0 \u00a0 \u00a0 \u00a0\u00a0<\/sub>b) CO<\/p>\n<p>4. Determine the formal charge of each element in the following:<\/p>\n<p id=\"fs-idp36610048\">a) HCl \u00a0 \u00a0 \u00a0 \u00a0\u00a0b) CF<sub>4 \u00a0 \u00a0 \u00a0 \u00a0\u00a0<\/sub>c) PCl<sub>3 \u00a0 \u00a0 \u00a0 \u00a0\u00a0<\/sub>d) PF<sub>5<\/sub><\/p>\n<p>5. Calculate the formal charge of chlorine in the molecules Cl<sub>2<\/sub>, BeCl<sub>2<\/sub>, and ClF<sub>5<\/sub>.<\/p>\n<p>6. Draw all possible resonance structures for each of these compounds. Determine the formal charge on each atom in each of the resonance structures:<\/p>\n<p id=\"fs-idp131483136\">a) O<sub>3 \u00a0 \u00a0 \u00a0 \u00a0\u00a0<\/sub>b) SO<sub>2 \u00a0 \u00a0 \u00a0 \u00a0\u00a0<\/sub>c) NO<sub>2<\/sub><sup>\u2212 \u00a0 \u00a0 \u00a0 \u00a0\u00a0<\/sup>d) NO<sub>3<\/sub><sup>\u2212<\/sup><\/p>\n<p>7. Based on formal charge considerations, which of the following would likely be the correct arrangement of atoms in hypochlorous acid: HOCl or OClH?<\/p>\n<p>8. Draw the structure of hydroxylamine, H<sub>3<\/sub>NO, and assign formal charges; look up the structure. Is the actual structure consistent with the formal charges?<\/p>\n<p>9. Write the Lewis structure and chemical formula of the compound with a molar mass of about 70 g\/mol that contains 19.7% nitrogen and 80.3% fluorine by mass, and determine the formal charge of the atoms in this compound.<\/p>\n<p>10. Sulfuric acid is the industrial chemical produced in greatest quantity worldwide. About 90 billion pounds are produced each year in the United States alone. Write the Lewis structure for sulfuric acid, H<sub>2<\/sub>SO<sub>4<\/sub>, which has two oxygen atoms and two OH groups bonded to the sulfur.<\/p>\n<p>&nbsp;<\/p>\n<p><strong>Answers<\/strong><\/p>\n<p id=\"fs-idp22600016\">1. a)<br \/>\n<img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_Ques2ansa_img-2.jpg\" alt=\"Two Lewis structures are shown with a double-headed arrow in between. The left structure shows a sulfur atom with a lone pair of electrons single bonded to the left to an oxygen atom with three lone pairs of electrons. The sulfur atom is also double bonded on the right to an oxygen atom with two lone pairs of electrons. The right structure depicts the same atoms, but this time the double bond is between the left oxygen and the sulfur atom. The lone pairs of electrons have also shifted to account for the change of bond types. The sulfur atom in the right structures, also has a third electron dot below it.\" width=\"430\" height=\"57\" class=\"\" \/><\/p>\n<p>b)<br \/>\n<img decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_Ques2ansb_img-2.jpg\" alt=\"Three Lewis structures are shown, with double-headed arrows in between, each surrounded by brackets and a superscripted two negative sign. The left structure depicts a carbon atom bonded to three oxygen atoms. It is single bonded to two of these oxygen atoms, each of which has three lone pairs of electrons, and double bonded to the third, which has two lone pairs of electrons. The double bond is located between the bottom oxygen and the carbon. The central and right structures are the same as the first, but the position of the double bonded oxygen has moved to the left oxygen in the right structure while the central structure only has single bonds. The lone pairs of electrons change to correspond with the bonds as well.\" \/><\/p>\n<p>c)<br \/>\n<img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_Ques2ansc_img-2.jpg\" alt=\"Two Lewis structures are shown, with a double-headed arrow in between, each surrounded by brackets and a superscripted negative sign. The left structure depicts a carbon atom bonded to three oxygen atoms. It is single bonded to one of these oxygen atoms, which has three lone pairs of electrons, and double bonded to the other two, which have two lone pairs of electrons. One of the double bonded oxygen atoms also has a single bond to a hydrogen atom. The right structure is the same as the first, but there is only one double bonded oxygen. The oxygen with the single bonded hydrogen now has a single bond to the carbon atom. The lone pairs of electrons have also changed to correspond with the bonds.\" width=\"552\" height=\"130\" class=\"\" \/><\/p>\n<p>d)<br \/>\n<img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_Ques2ansd_img-2.jpg\" alt=\"Two Lewis structures are shown with a double-headed arrow in between. The left structure depicts a hexagonal ring composed of five carbon atoms, each single bonded to a hydrogen atom, and one nitrogen atom that has a lone pair of electrons. The ring has alternating single and double bonds. The right structure is the same as the first, but each double bond has rotated to a new position.\" width=\"466\" height=\"169\" class=\"\" \/><\/p>\n<p>e)<br \/>\n<img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_Ques2anse_img-2.jpg\" alt=\"Two Lewis structures are shown with a double-headed arrow in between. The left structure shows a carbon atom single bonded to two hydrogen atoms and a second carbon atom. The second carbon atom is single bonded to a hydrogen atom and double bonded to a third carbon atom. The third carbon atom is single bonded to two hydrogen atoms. The whole structure is surrounded by brackets and a superscripted negative sign. The right structure shows a carbon atom single bonded to two hydrogen atoms and double bonded to a second carbon atom. The second carbon atom is single bonded to a hydrogen atom and a third carbon atom. The third carbon atom is single bonded to two hydrogen atoms. The whole structure is surrounded by brackets and a superscripted negative sign.\" width=\"611\" height=\"81\" class=\"\" \/><\/p>\n<p>2.<br \/>\n<img decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_Ques11ans_img-2.jpg\" alt=\"Two pairs of Lewis structures are shown with a double-headed arrow in between each pair. The left structure of the first pair shows a nitrogen atom with one lone pair of electrons single bonded to an oxygen atom with three lone pairs of electrons. It is also double bonded to an oxygen with two lone pairs of electrons. The right image of this pair depicts the mirror image of the left. Both images are surrounded by brackets and a superscripted negative sign. They are labeled, \u201cFor N O subscript two superscript negative sign.\u201d The left structure of the second pair shows an oxygen atom with one lone pair of electrons single bonded to an oxygen atom with three lone pairs of electrons. It is also double bonded to an oxygen atom with two lone pairs of electrons. The right structure appears as a mirror image of the left. These structures are labeled, \u201cFor O subscript three.\u201d\" \/><\/p>\n<p id=\"fs-idp32184800\">3. a)<br \/>\n<img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_Ques13ansb_img-2.jpg\" alt=\"This structure shows a carbon atom double bonded to two oxygen atoms, each of which has two lone pairs of electrons.\" width=\"276\" height=\"51\" class=\"\" \/><\/p>\n<p>b)<br \/>\n<img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_Ques13ansc_img-2.jpg\" alt=\"The right structure of this pair shows a carbon atom with one lone pair of electrons triple bonded to an oxygen with one lone pair of electrons.\" width=\"273\" height=\"26\" class=\"\" \/><br \/>\nCO has the strongest carbon-oxygen bond because there is a triple bond joining C and O. CO<sub>2<\/sub> has double bonds.<\/p>\n<p id=\"fs-idm21325744\">4. a) H: 0, Cl: 0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0b) C: 0, F: 0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0 c) P: 0, Cl 0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0d) P: 0, F: 0<\/p>\n<p id=\"fs-idm369848352\">5. Cl in Cl<sub>2<\/sub>: 0; \u00a0 \u00a0 Cl in BeCl<sub>2<\/sub>: 0; \u00a0 \u00a0 Cl in ClF<sub>5<\/sub>: 0<\/p>\n<p id=\"fs-idp124490112\">6. a)<br \/>\n<img decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_Exercis12a_img-2.jpg\" alt=\"Two Lewis structures are shown with a double-headed arrow in between. The left structure shows an oxygen atom with one lone pair of electrons single bonded to an oxygen atom with three lone pairs of electrons. It is also double bonded to an oxygen atom with two lone pairs of electrons. The symbols and numbers below this structure read, \u201c( 0 ), ( positive 1 ), ( negative 1 ).\u201d The phrase, \u201cFormal charge,\u201d and a right-facing arrow lie to the left of this structure. The right structure appears as a mirror image of the left and the symbols and numbers below this structure read, \u201c( negative 1 ), ( positive 1 ), ( 0 ).\u201d\" \/><\/p>\n<p>b)<br \/>\n<img decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_Exercis12b_img-2.jpg\" alt=\"Two Lewis structures are shown, with a double-headed arrow in between. The left structure shows a sulfur atom with one lone pair of electrons single bonded to an oxygen atom with three lone pairs of electrons. The sulfur atom also double bonded to an oxygen atom with two lone pairs of electrons. The symbols and numbers below this structure read, \u201c( negative 1 ), ( positive 1 ), ( 0 ).\u201d The right structure appears as a mirror image of the left and the symbols and numbers below this structure read, \u201c( 0 ), ( positive 1 ), ( negative 1 ).\u201d\" \/><\/p>\n<p>c)<br \/>\n<img decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_Exercis12c_img-2.jpg\" alt=\"[Two Lewis structures are shown, with brackets surrounding each with a superscripted negative sign and a double ended arrow in between. The left structure shows a nitrogen atom with one lone pair of electrons single bonded to an oxygen atom with three lone pairs of electrons and double bonded to an oxygen atom with two lone pairs of electrons. The symbols and numbers below this structure read \u201copen parenthesis, 0, close parenthesis, open parenthesis, 0, close parenthesis, open parenthesis, negative 1, close parenthesis. The right structure appears as a mirror image of the left and the symbols and numbers below this structure read \u201copen parenthesis, negative 1, close parenthesis, open parenthesis, 0, close parenthesis, open parenthesis, 0, close parenthesis.]\" \/><\/p>\n<p>d)<br \/>\n<img decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_Exercis12d_img-2.jpg\" alt=\"[Three Lewis structures are shown, with brackets surrounding each with a superscripted negative sign and a double ended arrow in between. The left structure shows a nitrogen atom single bonded to two oxygen atoms, each with three lone pairs of electrons and double bonded to an oxygen atom with two lone pairs of electrons. The single bonded oxygen atoms are labeled, from the top of the structure and going clockwise, \u201copen parenthesis, negative 1, close parenthesis, open parenthesis, positive 1, close parenthesis\u201d. The symbols and numbers below this structure read \u201copen parenthesis, 0, close parenthesis, open parenthesis, negative 1, close parenthesis. The middle structure shows a nitrogen atom single bonded to two oxygen atoms, each with three lone pairs of electrons, one of which is labeled \u201copen parenthesis, positive 1, close parenthesis\u201d and double bonded to an oxygen atom with two lone pairs of electrons labeled \u201copen parenthesis, 0, close parenthesis\u201d. The symbols and numbers below this structure read \u201copen parenthesis, negative 1, close parenthesis, open parenthesis, negative 1, close parenthesis. The right structure shows a nitrogen atom single bonded to two oxygen atoms, each with three lone pairs of electrons and double bonded to an oxygen atom with two lone pairs of electrons. One of the single bonded oxygen atoms is labeled, \u201copen parenthesis, negative 1, close parenthesis while the double bonded oxygen is labeled, \u201copen parenthesis, positive 1, close parenthesis\u201d. The symbols and numbers below this structure read \u201copen parenthesis, negative 1, close parenthesis\u201d and \u201copen parenthesis, 0, close parenthesis\u201d.]\" \/><\/p>\n<p id=\"fs-idm333542560\">7. HOCl<\/p>\n<p id=\"fs-idp6445072\">8. The structure that gives zero formal charges is consistent with the actual structure:<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-content\/uploads\/sites\/387\/2018\/04\/CNX_Chem_07_04_OHamine_img-2.jpg\" alt=\"A Lewis structure shows a nitrogen atom with one lone pair of electrons single bonded to two hydrogen atoms and an oxygen atom which has two lone pairs of electrons. The oxygen atom is single bonded to a hydrogen atom.\" width=\"134\" height=\"101\" class=\"\" \/><\/p>\n<p id=\"fs-idp35441312\">9. NF<sub>3<\/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_04_NF3_img-2.jpg\" alt=\"A Lewis structure shows a nitrogen atom with one lone pair of electrons single bonded to three fluorine atoms, each with three lone pairs of electrons.\" width=\"82\" height=\"125\" 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_04_H2SO4-2.jpg\" alt=\"A Lewis structure shows a hydrogen atom single bonded to an oxygen atom with two lone pairs of electrons. The oxygen atom is single bonded to a sulfur atom. The sulfur atom is double bonded to two oxygen atoms, each of which have three lone pairs of electrons, and single bonded to an oxygen atom with two lone pairs of electrons. This oxygen atom is single bonded to a hydrogen atom.\" width=\"171\" height=\"137\" class=\"\" \/><\/p>\n<\/div>\n<\/section>\n<div>\n<h2>Glossary<\/h2>\n<p><strong>formal charge:\u00a0<\/strong>charge that would result on an atom by taking the number of valence electrons on the neutral atom and subtracting the nonbonding electrons and the number of bonds (one-half of the bonding electrons)<\/p>\n<p><strong>molecular structure:\u00a0<\/strong>arrangement of atoms in a molecule or ion<\/p>\n<p><strong>resonance:\u00a0<\/strong>situation in which one Lewis structure is insufficient to describe the bonding in a molecule and the average of multiple structures is observed<\/p>\n<p><strong>resonance forms:\u00a0<\/strong>two or more Lewis structures that have the same arrangement of atoms but different arrangements of electrons<\/p>\n<p><strong>resonance hybrid:\u00a0<\/strong>average of the resonance forms shown by the individual Lewis structures<\/p>\n<\/div>\n","protected":false},"author":330,"menu_order":7,"template":"","meta":{"pb_show_title":"on","pb_short_title":"9.6 Formal Charges and Resonance","pb_subtitle":"","pb_authors":[],"pb_section_license":"cc-by"},"chapter-type":[],"contributor":[],"license":[50],"class_list":["post-1663","chapter","type-chapter","status-publish","hentry","license-cc-by"],"part":1538,"_links":{"self":[{"href":"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-json\/pressbooks\/v2\/chapters\/1663","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":6,"href":"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-json\/pressbooks\/v2\/chapters\/1663\/revisions"}],"predecessor-version":[{"id":4708,"href":"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-json\/pressbooks\/v2\/chapters\/1663\/revisions\/4708"}],"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\/1663\/metadata\/"}],"wp:attachment":[{"href":"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-json\/wp\/v2\/media?parent=1663"}],"wp:term":[{"taxonomy":"chapter-type","embeddable":true,"href":"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-json\/pressbooks\/v2\/chapter-type?post=1663"},{"taxonomy":"contributor","embeddable":true,"href":"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-json\/wp\/v2\/contributor?post=1663"},{"taxonomy":"license","embeddable":true,"href":"https:\/\/pressbooks.bccampus.ca\/chem1114langaracollege\/wp-json\/wp\/v2\/license?post=1663"}],"curies":[{"name":"wp","href":"https:\/\/api.w.org\/{rel}","templated":true}]}}