{"id":1346,"date":"2017-10-12T15:18:06","date_gmt":"2017-10-12T15:18:06","guid":{"rendered":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/?post_type=chapter&#038;p=1346"},"modified":"2018-10-05T18:42:17","modified_gmt":"2018-10-05T18:42:17","slug":"the-predicted-stabilities-of-resonance-contributors","status":"publish","type":"chapter","link":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/chapter\/the-predicted-stabilities-of-resonance-contributors\/","title":{"raw":"The Predicted Stabilities of Resonance Contributors","rendered":"The Predicted Stabilities of Resonance Contributors"},"content":{"raw":"<div class=\"elm-header\">\r\n<div class=\"elm-header-custom\">\u00a0Resonance is a mental exercise and method within the <a title=\"Theoretical Chemistry\/Chemical Bonding\/Valence Bond Theory\" href=\"https:\/\/chem.libretexts.org\/Core\/Physical_and_Theoretical_Chemistry\/Chemical_Bonding\/Valence_Bond_Theory\" rel=\"internal\">Valence Bond Theory<\/a> of bonding that describes the delocalization of electrons within molecules. It compares and contrasts two or more possible\u00a0<a title=\"Theoretical Chemistry\/Chemical Bonding\/Lewis Theory of Bonding\/Lewis Structures\" rel=\"broken\">Lewis structures<\/a> that can represent a particular molecule. Resonance structures\u00a0are used\u00a0when one\u00a0Lewis structure for a single molecule\u00a0cannot\u00a0fully describe the\u00a0bonding that takes place\u00a0between\u00a0neighboring atoms relative to the empirical data for the actual bond lengths between those atoms. The net sum of valid resonance structures is defined as a resonance\u00a0hybrid, which represents the overall\u00a0delocalization\u00a0of electrons within the molecule. A molecule that has several resonance structures is more stable than one with fewer. Some resonance structures are more favorable than others.<\/div>\r\n<\/div>\r\n<div id=\"elm-main-content\" class=\"elm-content-container\">\r\n<div id=\"s2006\">\r\n<div id=\"section_1\">\r\n<h3 id=\"Introduction-2006\">Introduction<\/h3>\r\nElectrons have no fixed position in atoms, compounds and molecules (see image below) but have probabilities of being found in certain spaces (orbitals). Resonance forms illustrate areas of higher probabilities (electron densities). This is like holding your hat in either your right hand or your left. The term Resonance is applied when \u00a0there are two or more possibilities available. Resonance structures do not change the relative positions of the atoms like your arms in the metaphor. The skeleton of the <a title=\"Lewis Structures\" rel=\"broken\">Lewis Structure<\/a> remains the same, only the electron locations change.\r\n<div class=\"textbox\">\r\n<p class=\"boxtitle\"><strong>\"PICK THE CORRECT ARROW FOR THE JOB\"<\/strong><\/p>\r\nMost arrows in chemistry cannot be used interchangeably and care must be given to selecting the correct arrow for the job.\r\n<ul>\r\n \t<li><span style=\"font-style: normal;font-weight: normal;line-height: normal;font-size: 16px;text-indent: 0px;text-align: left;letter-spacing: normal;float: none;direction: ltr;max-width: none;max-height: none;min-width: 0px;min-height: 0px;border: 0px;padding: 0px;margin: 0px\">\u2194\u2194<\/span>: A double headed arrow on both ends of the arrow between Lewis structures is used to show their inter-connectivity<\/li>\r\n \t<li><span style=\"font-style: normal;font-weight: normal;line-height: normal;font-size: 16px;text-indent: 0px;text-align: left;letter-spacing: normal;float: none;direction: ltr;max-width: none;max-height: none;min-width: 0px;min-height: 0px;border: 0px;padding: 0px;margin: 0px\">\u21cc\u21cc<\/span>:\u00a0Double harpoons are used to designate equilibria<\/li>\r\n \t<li><span style=\"font-style: normal;font-weight: normal;line-height: normal;font-size: 16px;text-indent: 0px;text-align: left;letter-spacing: normal;float: none;direction: ltr;max-width: none;max-height: none;min-width: 0px;min-height: 0px;border: 0px;padding: 0px;margin: 0px\">\u21c0\u21c0<\/span>: A single harpoon on one end indicate the movement of\u00a0<strong>one\u00a0<\/strong>electron<\/li>\r\n \t<li><span style=\"font-style: normal;font-weight: normal;line-height: normal;font-size: 16px;text-indent: 0px;text-align: left;letter-spacing: normal;float: none;direction: ltr;max-width: none;max-height: none;min-width: 0px;min-height: 0px;border: 0px;padding: 0px;margin: 0px\">\u2192\u2192<\/span>:\u00a0A double headed arrow on one end is used to indicate the movement of\u00a0<strong>two\u00a0<\/strong>electrons<\/li>\r\n<\/ul>\r\n<\/div>\r\n<div class=\"textbox examples\">\r\n<h3>Example 1: Ozone<\/h3>\r\n<div>\r\n<div>\r\n<div id=\"example\">\r\n\r\n\u00a0Consider ozone (O<sub>3<\/sub>)\r\n<h3><strong>SOLUTION<\/strong><\/h3>\r\n<img class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145245\/ozone.jpg\" alt=\"ozone.jpg\" width=\"218\" height=\"125\" \/>\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0 <img class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145247\/Ozone-animation.gif\" alt=\"Ozone-animation.gif\" width=\"135px\" height=\"80px\" \/>\r\n\r\nAn animation of how one can do a resonance with ozone by moving electrons\r\n\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<div id=\"section_2\">\r\n<h3 id=\"Delocalization_and_Resonance_Structures_Rules-2006\">Delocalization\u00a0and\u00a0Resonance Structures Rules<\/h3>\r\nIn resonance structures, the electrons are able to move to help stabilize the molecule. This movement of the electrons is called <a title=\"Delocalization of Electrons\" rel=\"broken\">delocalization<\/a>.\r\n<ol>\r\n \t<li style=\"list-style-type: none\">\r\n<ol>\r\n \t<li>Resonance structures should have the same number of electrons, do not add or subtract any electrons. (check the number of electrons by simply counting them).<\/li>\r\n \t<li>All resonance structures must follow the rules of writing <a class=\"internal\" title=\"Organic Chemistry\/Fundamentals\/Lewis Structures\" href=\"https:\/\/chem.libretexts.org\/Core\/Organic_Chemistry\/Fundamentals\/Lewis_Structures\" rel=\"internal\">Lewis Structures<\/a>.<\/li>\r\n \t<li>The <a class=\"internal mt-disabled\" title=\"Physical Chemistry\/Quantum Mechanics\/Virtual: Atomic Orbitals\/Hybrid Orbitals\" rel=\"broken\">hybridization <\/a>of the structure must stay the same.<\/li>\r\n \t<li>The skeleton of the structure can not be changed (only the electrons move).<\/li>\r\n \t<li>Resonance structures must also have the same amount of lone pairs.<\/li>\r\n<\/ol>\r\n<\/li>\r\n<\/ol>\r\n<\/div>\r\n<div id=\"section_3\">\r\n<h3 id=\"Formal_Charge-2006\">Formal Charge<\/h3>\r\nEven though the structures look the same, the <a title=\"Formal Charges\" rel=\"broken\">formal charg<\/a>e (FC)\u00a0may not be. Formal charges are charges that are assigned to a specific atom in a molecule. If computed correctly, the overall formal charge of the molecule should be the same as the oxidation charge of the molecule (the charge when you write out the empirical and molecular formula). We want to choose the resonance structure with the least formal charges that add up to zero or the charge of the overall molecule. The equation for finding Formal Charge is:\r\n\r\n<strong>Formal Charge = (number of valence electrons in free orbital) - (number of lone-pair electrons) - ( $$ \\frac{1}{2} $$ number bond pair electrons)<\/strong>\r\n\r\nThe formal charge has to equal the molecule's overall charge,e.g., the $$CNS^-$$ has an overall charge of -1, so the Lewis structure's formal charge has to equal -1.\r\n<div>\r\n<div class=\"textbox examples\">\r\n<h3 class=\"boxtitle\">Example 2: Thiocyanate Ion<\/h3>\r\nConsider the thiocyanate ($$CNS^-$$) ion.\r\n<h3><strong>SOLUTION<\/strong><\/h3>\r\n1. Find the Lewis Structure of the molecule. (Remember the Lewis Structure rules.)\r\n\r\n<img class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145247\/CNS_lewis_structure.jpg\" alt=\"CNS lewis structure.jpg\" width=\"235px\" height=\"67px\" \/>\r\n\r\n2. Resonance:\u00a0All elements want an octet, and we can do that in multiple ways by moving the terminal atom's electrons around (bonds too).\r\n\r\n<img class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145249\/CNS_resonance.jpg\" alt=\"CNS resonance.jpg\" width=\"228px\" height=\"181px\" \/>\r\n\r\n3. Assign Formal Charges\r\n\r\nFormal Charge = (number of valence electrons in free orbital) - (number of lone-pair electrons) - ( $$ \\frac{1}{2} $$ number bond pair electrons)\r\n\r\nRemember to\u00a0determine the number of valence electron each atom has before assigning Formal Charges\r\n\r\nC = 4 valence e<sup>-<\/sup>, N = 5 valence e<sup>-<\/sup>, S = 6 valence e<sup>-<\/sup>, also add an extra electron for the (-1) charge. The total of valence electrons is 16.\r\n\r\n<img class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145252\/CNS_FC.jpg\" alt=\"CNS FC.jpg\" width=\"306\" height=\"238\" \/>\r\n\r\n4. Find the most ideal resonance structure. (Note: It is the one with the least formal charges that adds up to zero or to the molecule's overall charge.)\r\n\r\n<img class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145254\/CNS_FC1.jpg\" alt=\"CNS FC1.jpg\" width=\"257\" height=\"217\" \/>\u00a0\u00a0<img class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145256\/CSN_fc_2.jpg\" alt=\"CSN fc 2.jpg\" width=\"309\" height=\"197\" \/>\r\n\r\n5. Now we have to look at\u00a0electronegativity\u00a0for the \"Correct\" Lewis structure.\r\n\r\nThe most\u00a0<a title=\"Physical Chemistry\/Physical Properties of Matter\/Atomic and Molecular Properties\/Electronegativity\" href=\"https:\/\/chem.libretexts.org\/Core\/Physical_and_Theoretical_Chemistry\/Physical_Properties_of_Matter\/Atomic_and_Molecular_Properties\/Electronegativity\" rel=\"internal\">electronegative<\/a>\u00a0atom usually has the negative formal charge, while the least electronegative atom usually has the positive formal charges.\r\n\r\n<img class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145257\/CNS_best.jpg\" alt=\"CNS best.jpg\" width=\"295\" height=\"189\" \/>\r\n\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<div id=\"section_4\">\r\n<h3 id=\"Resonance_Hybrids-2006\">Resonance Hybrids<\/h3>\r\nResonance Structures are a representation of a <em>Resonance Hybrid<\/em>, which is the combination of all resonance structures. The resonance structure with the Formal Charge closest to zero is the most accepted structure, however, the correct Lewis structure is actually a combination of all the resonance structures and is not solely describe as one.\r\n<ol>\r\n \t<li style=\"list-style-type: none\">\r\n<ol>\r\n \t<li>Draw the Lewis Structure &amp; Resonance for the molecule (using solid lines for bonds).<\/li>\r\n \t<li>Where there <strong>can<\/strong> be a double or triple bond, draw a dotted line (-----) for a bond.<\/li>\r\n \t<li>Draw only the lone pairs found in all resonance structures, do not include the lone pairs that are not on all of the resonance structures.<\/li>\r\n<\/ol>\r\n<\/li>\r\n<\/ol>\r\n<div id=\"note\">\r\n<p class=\"boxtitle\">Note: The correct Lewis structure is actually a combination of all the resonance structures and hence is not solely described as one.<\/p>\r\n\r\n<\/div>\r\n<div>\r\n<div>\r\n<div>\r\n<div class=\"textbox examples\">\r\n<div id=\"section_4\">\r\n<div>\r\n<div>\r\n<div>\r\n<h3 class=\"boxtitle\">Example 3: Carbonate Ion<\/h3>\r\nConsider the carbonate ion:\u00a0CO<sub>3<\/sub><sup>2<\/sup><sup>-<\/sup>\r\n<h3><strong>SOLUTION<\/strong><\/h3>\r\n<img class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145259\/CO3-hybrid.jpg\" alt=\"CO3-hybrid.jpg\" width=\"203\" height=\"179\" \/>\r\n\r\nStep 1: Draw the Lewis Structure &amp; Resonance.\r\n\r\n<img class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145301\/CO3-resonance.jpg\" alt=\"CO3-resonance.jpg\" width=\"437\" height=\"144\" \/>\r\n\r\nStep 2: Combine the resonance structures by adding (dotted) bonds where other resonance bonds can be formed.\r\n\r\n<img class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145304\/CO3-step-2.jpg\" alt=\"CO3-step-2.jpg\" width=\"394\" height=\"271\" \/>\r\n\r\nStep 3: Add only the lone pairs found on\u00a0<strong>ALL<\/strong>\u00a0resonance structures.\r\n\r\n<img class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145306\/CO3-step-3.jpg\" alt=\"CO3-step-3.jpg\" width=\"413\" height=\"282\" \/>\r\n\r\nThe bottom is the finished resonance hybrid for\u00a0<strong>CO<sub>3<\/sub><sup>2-<\/sup><\/strong>.\r\n\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 id=\"section_5\">\r\n<h3 id=\"Rules_for_estimating_stability_of_resonance_structures-2006\">Rules for estimating stability of resonance structures<\/h3>\r\n<ol>\r\n \t<li style=\"list-style-type: none\">\r\n<ol>\r\n \t<li>The <strong>greater the number of covalent bonds<\/strong>, the greater the stability since more atoms will have complete octets<\/li>\r\n \t<li>The structure with the<strong>\u00a0least<\/strong>\u00a0<strong>number of formal charges<\/strong>\u00a0is more stable<\/li>\r\n \t<li>The structure with the<strong> least<\/strong> <strong>separation of formal charge<\/strong> is more stable<\/li>\r\n \t<li>A structure with a <strong>negative charge on the more electronegative atom<\/strong> will be more stable<\/li>\r\n \t<li><strong>Positive charges on the least electronegative atom <\/strong>(most electropositive) is more stable<\/li>\r\n \t<li>Resonance forms that are equivalent have no difference in stability and contribute equally (eg. benzene)<\/li>\r\n<\/ol>\r\n<\/li>\r\n<\/ol>\r\n<div>\r\n<div>\r\n<div class=\"textbox examples\">\r\n<h3 class=\"boxtitle\">Example 4: Benzene and\u00a0Aminophenol<\/h3>\r\nBenzene is an extremely stable molecule and it is accounted for its geometry and molecular orbital interaction, but most importantly it's due to its resonance structures. The delocalized electrons in the benzene ring make the molecule very stable and with its characteristics of a nucleophile, it will react with a strong electrophile only and after the first reactivity, the substituted benzene will depend on its resonance to direct the next position for the reaction to add a second substituent.\r\n\r\n<img class=\"internal aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145307\/3DExamples_of_Resonance_1_22._Benzene1.png\" alt=\"\" width=\"247px\" height=\"76px\" \/>\r\n\r\nAminophenol is a very stable molecule that is present in most biological systems, mainly in proteins. By studies of NMR spectroscopy and X-Ray crystallography it is confirmed that the stability of the amide is due to resonance which through molecular orbital interaction creates almost a double bond between the Nitrogen and the carbon.\r\n\r\n<img class=\"internal aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145309\/3DExamples_of_Resonance_1_22_._aminophenol.png\" alt=\"\" width=\"520.994px\" height=\"82.9972px\" \/>\r\n\r\n<\/div>\r\n<\/div>\r\n<div>\r\n<div>\r\n<div class=\"textbox examples\">\r\n<h3 class=\"boxtitle\">Example 5: Multiple Resonance of other Molecules<\/h3>\r\nMolecules with multiple\u00a0resonance forms\r\n\r\n<img class=\"internal aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145311\/Examples_of_Resonance_2_3.png\" alt=\"\" width=\"593.991px\" height=\"226.989px\" \/>\r\n\r\nSome structural resonance conformations are the major contributor or the dominant forms that the molecule exists. For example, if we look at the above rules for estimating the stability of a molecule, we see that for the third molecule the first and second forms are the major contributors for the overall stability of the molecule. The nitrogen is more electronegative than carbon so, it can handle the negative charge more than carbon. A carbon with a negative charge is the least favorable conformation for the molecule to exist, so the last resonance form contributes very little for the stability of the Ion.<a title=\"Examples of Resonance_3_3.png\" href=\"https:\/\/chem.libretexts.org\/@api\/deki\/files\/12250\/Examples_of_Resonance_3_3.png?revision=1\" rel=\"internal\"><img class=\"internal aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145313\/Examples_of_Resonance_3_3.png\" alt=\"\" width=\"554\" height=\"150\" \/><\/a>\r\n\r\nThe Hybrid Resonance forms show the different Lewis structures with the electron been\u00a0delocalized. This is very important for the reactivity of\u00a0chlorobenzene\u00a0because in the presence of an\u00a0electrophile\u00a0it will react and the formation of another bond will be\u00a0directed\u00a0and determine by resonance. The long pair of electrons\u00a0delocalized\u00a0in the aromatic substituted ring is where it can potentially form a new bond with an\u00a0electrophile, as it is shown there are three possible places that reactivity can take place, the first to react will take place at the\u00a0<em>para\u00a0<\/em>position with respect to the\u00a0chloro\u00a0substituent\u00a0and then to either\u00a0<em>ortho\u00a0<\/em>position.\r\n\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<div id=\"section_6\">\r\n<h3><\/h3>\r\n<h3 id=\"References-2006\">References<\/h3>\r\n<ol>\r\n \t<li style=\"list-style-type: none\">\r\n<ol>\r\n \t<li>Petrucci, Ralph H., et al. <em>General Chemistry: Principles and Modern Applications<\/em>. New Jersey: Pearson Prentice Hall, 2007.<\/li>\r\n \t<li>Ahmad, Wan-Yaacob and Zakaria, Mat B. \"Drawing Lewis Structures from Lewis Symbols: A Direct Electron Pairing Approach.\" Journal of Chemical Education: Journal 77.3.<\/li>\r\n<\/ol>\r\n<\/li>\r\n<\/ol>\r\n<div class=\"textbox exercises\">\r\n<h3>Exercises<\/h3>\r\n<div id=\"section_7\">\r\n<ol>\r\n \t<li style=\"list-style-type: none\">\r\n<ol>\r\n \t<li>True or False, The picture below is a resonance structure?<\/li>\r\n<\/ol>\r\n<\/li>\r\n<\/ol>\r\n<img class=\"internal aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145314\/problem1-truefalse.jpg\" alt=\"problem1-truefalse.jpg\" width=\"237.997px\" height=\"105px\" \/>\r\n<ol start=\"2\">\r\n \t<li style=\"list-style-type: none\">\r\n<ol start=\"2\">\r\n \t<li>Draw the Lewis Dot Structure for <strong>SO<\/strong><sub><strong>4<\/strong><\/sub><sup><strong>2<\/strong><\/sup><sup><strong>-<\/strong><\/sup> and all possible resonance structures. Which of the following resonance structure is not favored among the Lewis Structures? Explain why. Assign Formal Charges.<\/li>\r\n \t<li>Draw the Lewis Dot Structure for <strong>CH<sub>3<\/sub>COO<\/strong><sup><strong>-<\/strong> <\/sup>and all possible resonance structures. Assign Formal Charges. Choose the most favorable Lewis Structure.<\/li>\r\n \t<li>Draw the Lewis Dot Structure for <strong>H<\/strong><strong>PO<\/strong><strong><sub>3<\/sub><sup><sub>2<\/sub><\/sup><\/strong><strong><sup><sub>-<\/sub><\/sup><\/strong> and all possible resonance structures. Assign Formal Charges.<\/li>\r\n \t<li>Draw the Lewis Dot Structure for\u00a0<strong>CHO<sub>2<\/sub><sup>1<\/sup><sup>-<\/sup><\/strong> and all possible resonance structures. Assign Formal Charges.<\/li>\r\n \t<li>Draw the\u00a0Resonance Hybrid Structure for <strong>P<\/strong><strong>O<\/strong><sub><strong>4<\/strong><\/sub><sup><strong>3<\/strong><\/sup><sup><strong>-<\/strong><\/sup>.<\/li>\r\n \t<li>Draw the\u00a0Resonance Hybrid Structure for <strong>N<\/strong><strong>O<\/strong><sub><strong>3<\/strong><\/sub><sup><strong>-<\/strong><\/sup>.<\/li>\r\n<\/ol>\r\n<\/li>\r\n<\/ol>\r\n<div id=\"section_8\">\r\n<h4 id=\"Problems_.232-2006\"><strong>Problems #2<\/strong><\/h4>\r\n<a title=\"resonance_problems_1_4.png\" href=\"https:\/\/chem.libretexts.org\/@api\/deki\/files\/12253\/resonance_problems_1_4.png?revision=1\" rel=\"internal\"><img class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145316\/resonance_problems_1_4.png\" alt=\"resonance_problems_1_4.png\" width=\"720px\" height=\"433px\" \/><\/a>\r\n\r\n<a title=\"resonance_problems_4_6.png\" href=\"https:\/\/chem.libretexts.org\/@api\/deki\/files\/12254\/resonance_problems_4_6.png?revision=1\" rel=\"internal\"><img class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145318\/resonance_problems_4_6.png\" alt=\"resonance_problems_4_6.png\" width=\"720px\" height=\"331px\" \/><\/a>\r\n\r\n<\/div>\r\n<div id=\"section_9\">\r\n<h3 id=\"Answers-2006\">Answers<\/h3>\r\n[reveal-answer q=\"630305\"]Show Answer[\/reveal-answer]\r\n[hidden-answer a=\"630305\"]\r\n\r\n1. False, because the electrons were not moved around, only the atoms (this violates the Resonance Structure Rules).\r\n\r\n2. Below are the all Lewis dot structure with formal charges (in red) for Sulfate (<strong>SO<sub>4<\/sub><sup>2<\/sup><sup>-<\/sup><\/strong>). There isn't a most favorable resonance of the Sulfate ion because they are all identical in charge and there is no change in Electronegativity between the Oxygen atoms.\r\n\r\n<img class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145321\/Sulfate-resonance.jpg\" alt=\"Sulfate-resonance.jpg\" width=\"486\" height=\"284\" \/>\r\n\r\n3. Below is the resonance for <strong>CH<sub>3<\/sub>COO<\/strong><sup>-<\/sup>, formal charges are displayed in red.\u00a0The Lewis Structure with the most formal charges is not desirable, because we want the Lewis Structure with the least formal charge.\r\n\r\n<img class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145324\/CHCOO-answerr.jpg\" alt=\"CHCOO-answerr.jpg\" width=\"417\" height=\"338\" \/>\r\n\r\n4. The resonance for <strong>HPO<sub>3<\/sub><sup>2<\/sup><sup>-<\/sup><\/strong>, and the formal charges (in red).<a title=\"Phosphite-resonance.jpg\" href=\"https:\/\/chem.libretexts.org\/@api\/deki\/files\/10306\/Phosphite-resonance.jpg?revision=1\" rel=\"internal\"><img class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145325\/Phosphite-resonance.jpg\" alt=\"Phosphite-resonance.jpg\" width=\"720px\" height=\"122px\" \/><\/a>\r\n\r\n5. The resonance for <strong>CHO<sub>2<\/sub><sup>1<\/sup><sup>-<\/sup><\/strong>, and the formal charges (in red).\r\n\r\n<img class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145327\/CHO2-answer.jpg\" alt=\"CHO2-answer.jpg\" width=\"390\" height=\"159\" \/>\r\n\r\n6. The resonance hybrid for <strong>PO<\/strong><sub><strong>4<\/strong><\/sub><sup><strong>3<\/strong><\/sup><sup><strong>-<\/strong><\/sup>, hybrid bonds are in red.\r\n\r\n<img class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145329\/PO4.jpg\" alt=\"PO4.jpg\" width=\"250\" height=\"213\" \/>\r\n\r\n7. The resonance hybrid for <strong>NO<\/strong><sub><strong>3<\/strong><\/sub><sup><strong>-<\/strong><\/sup>, hybrid bonds are in red.\r\n\r\n<img class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145331\/NO3.jpg\" alt=\"NO3.jpg\" width=\"311px\" height=\"278px\" \/>[\/hidden-answer]\r\n\r\n<\/div>\r\n<div id=\"section_10\">\r\n\r\n&nbsp;\r\n\r\n<strong><strong>Problems #2<\/strong><\/strong>\r\n\r\n[reveal-answer q=\"837981\"]Show Answer[\/reveal-answer]\r\n[hidden-answer a=\"837981\"]\r\n\r\n<img class=\"internal default aligncenter\" src=\"https:\/\/chem.libretexts.org\/@api\/deki\/files\/12256\/resonance_solutions_5_6.png?revision=1&amp;size=bestfit&amp;width=720&amp;height=334\" alt=\"resonance_solutions_5_6.png\" width=\"720px\" height=\"334px\" \/>[\/hidden-answer]\r\n\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<div id=\"section_11\">\r\n<h3 id=\"Contributors-2006\">Contributors<\/h3>\r\n<ul>\r\n \t<li style=\"list-style-type: none\">\r\n<ul>\r\n \t<li>Sharon Wei (UCD), Liza Chu (UCD)<\/li>\r\n<\/ul>\r\n<\/li>\r\n<\/ul>\r\n<\/div>\r\n<\/div>\r\n<\/div>","rendered":"<div class=\"elm-header\">\n<div class=\"elm-header-custom\">\u00a0Resonance is a mental exercise and method within the <a title=\"Theoretical Chemistry\/Chemical Bonding\/Valence Bond Theory\" href=\"https:\/\/chem.libretexts.org\/Core\/Physical_and_Theoretical_Chemistry\/Chemical_Bonding\/Valence_Bond_Theory\" rel=\"internal\">Valence Bond Theory<\/a> of bonding that describes the delocalization of electrons within molecules. It compares and contrasts two or more possible\u00a0<a title=\"Theoretical Chemistry\/Chemical Bonding\/Lewis Theory of Bonding\/Lewis Structures\" rel=\"broken\">Lewis structures<\/a> that can represent a particular molecule. Resonance structures\u00a0are used\u00a0when one\u00a0Lewis structure for a single molecule\u00a0cannot\u00a0fully describe the\u00a0bonding that takes place\u00a0between\u00a0neighboring atoms relative to the empirical data for the actual bond lengths between those atoms. The net sum of valid resonance structures is defined as a resonance\u00a0hybrid, which represents the overall\u00a0delocalization\u00a0of electrons within the molecule. A molecule that has several resonance structures is more stable than one with fewer. Some resonance structures are more favorable than others.<\/div>\n<\/div>\n<div id=\"elm-main-content\" class=\"elm-content-container\">\n<div id=\"s2006\">\n<div id=\"section_1\">\n<h3 id=\"Introduction-2006\">Introduction<\/h3>\n<p>Electrons have no fixed position in atoms, compounds and molecules (see image below) but have probabilities of being found in certain spaces (orbitals). Resonance forms illustrate areas of higher probabilities (electron densities). This is like holding your hat in either your right hand or your left. The term Resonance is applied when \u00a0there are two or more possibilities available. Resonance structures do not change the relative positions of the atoms like your arms in the metaphor. The skeleton of the <a title=\"Lewis Structures\" rel=\"broken\">Lewis Structure<\/a> remains the same, only the electron locations change.<\/p>\n<div class=\"textbox\">\n<p class=\"boxtitle\"><strong>&#8220;PICK THE CORRECT ARROW FOR THE JOB&#8221;<\/strong><\/p>\n<p>Most arrows in chemistry cannot be used interchangeably and care must be given to selecting the correct arrow for the job.<\/p>\n<ul>\n<li><span style=\"font-style: normal;font-weight: normal;line-height: normal;font-size: 16px;text-indent: 0px;text-align: left;letter-spacing: normal;float: none;direction: ltr;max-width: none;max-height: none;min-width: 0px;min-height: 0px;border: 0px;padding: 0px;margin: 0px\">\u2194\u2194<\/span>: A double headed arrow on both ends of the arrow between Lewis structures is used to show their inter-connectivity<\/li>\n<li><span style=\"font-style: normal;font-weight: normal;line-height: normal;font-size: 16px;text-indent: 0px;text-align: left;letter-spacing: normal;float: none;direction: ltr;max-width: none;max-height: none;min-width: 0px;min-height: 0px;border: 0px;padding: 0px;margin: 0px\">\u21cc\u21cc<\/span>:\u00a0Double harpoons are used to designate equilibria<\/li>\n<li><span style=\"font-style: normal;font-weight: normal;line-height: normal;font-size: 16px;text-indent: 0px;text-align: left;letter-spacing: normal;float: none;direction: ltr;max-width: none;max-height: none;min-width: 0px;min-height: 0px;border: 0px;padding: 0px;margin: 0px\">\u21c0\u21c0<\/span>: A single harpoon on one end indicate the movement of\u00a0<strong>one\u00a0<\/strong>electron<\/li>\n<li><span style=\"font-style: normal;font-weight: normal;line-height: normal;font-size: 16px;text-indent: 0px;text-align: left;letter-spacing: normal;float: none;direction: ltr;max-width: none;max-height: none;min-width: 0px;min-height: 0px;border: 0px;padding: 0px;margin: 0px\">\u2192\u2192<\/span>:\u00a0A double headed arrow on one end is used to indicate the movement of\u00a0<strong>two\u00a0<\/strong>electrons<\/li>\n<\/ul>\n<\/div>\n<div class=\"textbox examples\">\n<h3>Example 1: Ozone<\/h3>\n<div>\n<div>\n<div id=\"example\">\n<p>\u00a0Consider ozone (O<sub>3<\/sub>)<\/p>\n<h3><strong>SOLUTION<\/strong><\/h3>\n<p><img loading=\"lazy\" decoding=\"async\" class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145245\/ozone.jpg\" alt=\"ozone.jpg\" width=\"218\" height=\"125\" \/>\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0 <img decoding=\"async\" class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145247\/Ozone-animation.gif\" alt=\"Ozone-animation.gif\" width=\"135px\" height=\"80px\" \/><\/p>\n<p>An animation of how one can do a resonance with ozone by moving electrons<\/p>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<div id=\"section_2\">\n<h3 id=\"Delocalization_and_Resonance_Structures_Rules-2006\">Delocalization\u00a0and\u00a0Resonance Structures Rules<\/h3>\n<p>In resonance structures, the electrons are able to move to help stabilize the molecule. This movement of the electrons is called <a title=\"Delocalization of Electrons\" rel=\"broken\">delocalization<\/a>.<\/p>\n<ol>\n<li style=\"list-style-type: none\">\n<ol>\n<li>Resonance structures should have the same number of electrons, do not add or subtract any electrons. (check the number of electrons by simply counting them).<\/li>\n<li>All resonance structures must follow the rules of writing <a class=\"internal\" title=\"Organic Chemistry\/Fundamentals\/Lewis Structures\" href=\"https:\/\/chem.libretexts.org\/Core\/Organic_Chemistry\/Fundamentals\/Lewis_Structures\" rel=\"internal\">Lewis Structures<\/a>.<\/li>\n<li>The <a class=\"internal mt-disabled\" title=\"Physical Chemistry\/Quantum Mechanics\/Virtual: Atomic Orbitals\/Hybrid Orbitals\" rel=\"broken\">hybridization <\/a>of the structure must stay the same.<\/li>\n<li>The skeleton of the structure can not be changed (only the electrons move).<\/li>\n<li>Resonance structures must also have the same amount of lone pairs.<\/li>\n<\/ol>\n<\/li>\n<\/ol>\n<\/div>\n<div id=\"section_3\">\n<h3 id=\"Formal_Charge-2006\">Formal Charge<\/h3>\n<p>Even though the structures look the same, the <a title=\"Formal Charges\" rel=\"broken\">formal charg<\/a>e (FC)\u00a0may not be. Formal charges are charges that are assigned to a specific atom in a molecule. If computed correctly, the overall formal charge of the molecule should be the same as the oxidation charge of the molecule (the charge when you write out the empirical and molecular formula). We want to choose the resonance structure with the least formal charges that add up to zero or the charge of the overall molecule. The equation for finding Formal Charge is:<\/p>\n<p><strong>Formal Charge = (number of valence electrons in free orbital) &#8211; (number of lone-pair electrons) &#8211; ( $$ \\frac{1}{2} $$ number bond pair electrons)<\/strong><\/p>\n<p>The formal charge has to equal the molecule&#8217;s overall charge,e.g., the $$CNS^-$$ has an overall charge of -1, so the Lewis structure&#8217;s formal charge has to equal -1.<\/p>\n<div>\n<div class=\"textbox examples\">\n<h3 class=\"boxtitle\">Example 2: Thiocyanate Ion<\/h3>\n<p>Consider the thiocyanate ($$CNS^-$$) ion.<\/p>\n<h3><strong>SOLUTION<\/strong><\/h3>\n<p>1. Find the Lewis Structure of the molecule. (Remember the Lewis Structure rules.)<\/p>\n<p><img decoding=\"async\" class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145247\/CNS_lewis_structure.jpg\" alt=\"CNS lewis structure.jpg\" width=\"235px\" height=\"67px\" \/><\/p>\n<p>2. Resonance:\u00a0All elements want an octet, and we can do that in multiple ways by moving the terminal atom&#8217;s electrons around (bonds too).<\/p>\n<p><img decoding=\"async\" class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145249\/CNS_resonance.jpg\" alt=\"CNS resonance.jpg\" width=\"228px\" height=\"181px\" \/><\/p>\n<p>3. Assign Formal Charges<\/p>\n<p>Formal Charge = (number of valence electrons in free orbital) &#8211; (number of lone-pair electrons) &#8211; ( $$ \\frac{1}{2} $$ number bond pair electrons)<\/p>\n<p>Remember to\u00a0determine the number of valence electron each atom has before assigning Formal Charges<\/p>\n<p>C = 4 valence e<sup>&#8211;<\/sup>, N = 5 valence e<sup>&#8211;<\/sup>, S = 6 valence e<sup>&#8211;<\/sup>, also add an extra electron for the (-1) charge. The total of valence electrons is 16.<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145252\/CNS_FC.jpg\" alt=\"CNS FC.jpg\" width=\"306\" height=\"238\" \/><\/p>\n<p>4. Find the most ideal resonance structure. (Note: It is the one with the least formal charges that adds up to zero or to the molecule&#8217;s overall charge.)<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145254\/CNS_FC1.jpg\" alt=\"CNS FC1.jpg\" width=\"257\" height=\"217\" \/>\u00a0\u00a0<img loading=\"lazy\" decoding=\"async\" class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145256\/CSN_fc_2.jpg\" alt=\"CSN fc 2.jpg\" width=\"309\" height=\"197\" \/><\/p>\n<p>5. Now we have to look at\u00a0electronegativity\u00a0for the &#8220;Correct&#8221; Lewis structure.<\/p>\n<p>The most\u00a0<a title=\"Physical Chemistry\/Physical Properties of Matter\/Atomic and Molecular Properties\/Electronegativity\" href=\"https:\/\/chem.libretexts.org\/Core\/Physical_and_Theoretical_Chemistry\/Physical_Properties_of_Matter\/Atomic_and_Molecular_Properties\/Electronegativity\" rel=\"internal\">electronegative<\/a>\u00a0atom usually has the negative formal charge, while the least electronegative atom usually has the positive formal charges.<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145257\/CNS_best.jpg\" alt=\"CNS best.jpg\" width=\"295\" height=\"189\" \/><\/p>\n<\/div>\n<\/div>\n<\/div>\n<div id=\"section_4\">\n<h3 id=\"Resonance_Hybrids-2006\">Resonance Hybrids<\/h3>\n<p>Resonance Structures are a representation of a <em>Resonance Hybrid<\/em>, which is the combination of all resonance structures. The resonance structure with the Formal Charge closest to zero is the most accepted structure, however, the correct Lewis structure is actually a combination of all the resonance structures and is not solely describe as one.<\/p>\n<ol>\n<li style=\"list-style-type: none\">\n<ol>\n<li>Draw the Lewis Structure &amp; Resonance for the molecule (using solid lines for bonds).<\/li>\n<li>Where there <strong>can<\/strong> be a double or triple bond, draw a dotted line (&#8212;&#8211;) for a bond.<\/li>\n<li>Draw only the lone pairs found in all resonance structures, do not include the lone pairs that are not on all of the resonance structures.<\/li>\n<\/ol>\n<\/li>\n<\/ol>\n<div id=\"note\">\n<p class=\"boxtitle\">Note: The correct Lewis structure is actually a combination of all the resonance structures and hence is not solely described as one.<\/p>\n<\/div>\n<div>\n<div>\n<div>\n<div class=\"textbox examples\">\n<div id=\"section_4\">\n<div>\n<div>\n<div>\n<h3 class=\"boxtitle\">Example 3: Carbonate Ion<\/h3>\n<p>Consider the carbonate ion:\u00a0CO<sub>3<\/sub><sup>2<\/sup><sup>&#8211;<\/sup><\/p>\n<h3><strong>SOLUTION<\/strong><\/h3>\n<p><img loading=\"lazy\" decoding=\"async\" class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145259\/CO3-hybrid.jpg\" alt=\"CO3-hybrid.jpg\" width=\"203\" height=\"179\" \/><\/p>\n<p>Step 1: Draw the Lewis Structure &amp; Resonance.<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145301\/CO3-resonance.jpg\" alt=\"CO3-resonance.jpg\" width=\"437\" height=\"144\" \/><\/p>\n<p>Step 2: Combine the resonance structures by adding (dotted) bonds where other resonance bonds can be formed.<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145304\/CO3-step-2.jpg\" alt=\"CO3-step-2.jpg\" width=\"394\" height=\"271\" \/><\/p>\n<p>Step 3: Add only the lone pairs found on\u00a0<strong>ALL<\/strong>\u00a0resonance structures.<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145306\/CO3-step-3.jpg\" alt=\"CO3-step-3.jpg\" width=\"413\" height=\"282\" \/><\/p>\n<p>The bottom is the finished resonance hybrid for\u00a0<strong>CO<sub>3<\/sub><sup>2-<\/sup><\/strong>.<\/p>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<div id=\"section_5\">\n<h3 id=\"Rules_for_estimating_stability_of_resonance_structures-2006\">Rules for estimating stability of resonance structures<\/h3>\n<ol>\n<li style=\"list-style-type: none\">\n<ol>\n<li>The <strong>greater the number of covalent bonds<\/strong>, the greater the stability since more atoms will have complete octets<\/li>\n<li>The structure with the<strong>\u00a0least<\/strong>\u00a0<strong>number of formal charges<\/strong>\u00a0is more stable<\/li>\n<li>The structure with the<strong> least<\/strong> <strong>separation of formal charge<\/strong> is more stable<\/li>\n<li>A structure with a <strong>negative charge on the more electronegative atom<\/strong> will be more stable<\/li>\n<li><strong>Positive charges on the least electronegative atom <\/strong>(most electropositive) is more stable<\/li>\n<li>Resonance forms that are equivalent have no difference in stability and contribute equally (eg. benzene)<\/li>\n<\/ol>\n<\/li>\n<\/ol>\n<div>\n<div>\n<div class=\"textbox examples\">\n<h3 class=\"boxtitle\">Example 4: Benzene and\u00a0Aminophenol<\/h3>\n<p>Benzene is an extremely stable molecule and it is accounted for its geometry and molecular orbital interaction, but most importantly it&#8217;s due to its resonance structures. The delocalized electrons in the benzene ring make the molecule very stable and with its characteristics of a nucleophile, it will react with a strong electrophile only and after the first reactivity, the substituted benzene will depend on its resonance to direct the next position for the reaction to add a second substituent.<\/p>\n<p><img decoding=\"async\" class=\"internal aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145307\/3DExamples_of_Resonance_1_22._Benzene1.png\" alt=\"\" width=\"247px\" height=\"76px\" \/><\/p>\n<p>Aminophenol is a very stable molecule that is present in most biological systems, mainly in proteins. By studies of NMR spectroscopy and X-Ray crystallography it is confirmed that the stability of the amide is due to resonance which through molecular orbital interaction creates almost a double bond between the Nitrogen and the carbon.<\/p>\n<p><img decoding=\"async\" class=\"internal aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145309\/3DExamples_of_Resonance_1_22_._aminophenol.png\" alt=\"\" width=\"520.994px\" height=\"82.9972px\" \/><\/p>\n<\/div>\n<\/div>\n<div>\n<div>\n<div class=\"textbox examples\">\n<h3 class=\"boxtitle\">Example 5: Multiple Resonance of other Molecules<\/h3>\n<p>Molecules with multiple\u00a0resonance forms<\/p>\n<p><img decoding=\"async\" class=\"internal aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145311\/Examples_of_Resonance_2_3.png\" alt=\"\" width=\"593.991px\" height=\"226.989px\" \/><\/p>\n<p>Some structural resonance conformations are the major contributor or the dominant forms that the molecule exists. For example, if we look at the above rules for estimating the stability of a molecule, we see that for the third molecule the first and second forms are the major contributors for the overall stability of the molecule. The nitrogen is more electronegative than carbon so, it can handle the negative charge more than carbon. A carbon with a negative charge is the least favorable conformation for the molecule to exist, so the last resonance form contributes very little for the stability of the Ion.<a title=\"Examples of Resonance_3_3.png\" href=\"https:\/\/chem.libretexts.org\/@api\/deki\/files\/12250\/Examples_of_Resonance_3_3.png?revision=1\" rel=\"internal\"><img loading=\"lazy\" decoding=\"async\" class=\"internal aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145313\/Examples_of_Resonance_3_3.png\" alt=\"\" width=\"554\" height=\"150\" \/><\/a><\/p>\n<p>The Hybrid Resonance forms show the different Lewis structures with the electron been\u00a0delocalized. This is very important for the reactivity of\u00a0chlorobenzene\u00a0because in the presence of an\u00a0electrophile\u00a0it will react and the formation of another bond will be\u00a0directed\u00a0and determine by resonance. The long pair of electrons\u00a0delocalized\u00a0in the aromatic substituted ring is where it can potentially form a new bond with an\u00a0electrophile, as it is shown there are three possible places that reactivity can take place, the first to react will take place at the\u00a0<em>para\u00a0<\/em>position with respect to the\u00a0chloro\u00a0substituent\u00a0and then to either\u00a0<em>ortho\u00a0<\/em>position.<\/p>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<div id=\"section_6\">\n<h3><\/h3>\n<h3 id=\"References-2006\">References<\/h3>\n<ol>\n<li style=\"list-style-type: none\">\n<ol>\n<li>Petrucci, Ralph H., et al. <em>General Chemistry: Principles and Modern Applications<\/em>. New Jersey: Pearson Prentice Hall, 2007.<\/li>\n<li>Ahmad, Wan-Yaacob and Zakaria, Mat B. &#8220;Drawing Lewis Structures from Lewis Symbols: A Direct Electron Pairing Approach.&#8221; Journal of Chemical Education: Journal 77.3.<\/li>\n<\/ol>\n<\/li>\n<\/ol>\n<div class=\"textbox exercises\">\n<h3>Exercises<\/h3>\n<div id=\"section_7\">\n<ol>\n<li style=\"list-style-type: none\">\n<ol>\n<li>True or False, The picture below is a resonance structure?<\/li>\n<\/ol>\n<\/li>\n<\/ol>\n<p><img decoding=\"async\" class=\"internal aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145314\/problem1-truefalse.jpg\" alt=\"problem1-truefalse.jpg\" width=\"237.997px\" height=\"105px\" \/><\/p>\n<ol start=\"2\">\n<li style=\"list-style-type: none\">\n<ol start=\"2\">\n<li>Draw the Lewis Dot Structure for <strong>SO<\/strong><sub><strong>4<\/strong><\/sub><sup><strong>2<\/strong><\/sup><sup><strong>&#8211;<\/strong><\/sup> and all possible resonance structures. Which of the following resonance structure is not favored among the Lewis Structures? Explain why. Assign Formal Charges.<\/li>\n<li>Draw the Lewis Dot Structure for <strong>CH<sub>3<\/sub>COO<\/strong><sup><strong>&#8211;<\/strong> <\/sup>and all possible resonance structures. Assign Formal Charges. Choose the most favorable Lewis Structure.<\/li>\n<li>Draw the Lewis Dot Structure for <strong>H<\/strong><strong>PO<\/strong><strong><sub>3<\/sub><sup><sub>2<\/sub><\/sup><\/strong><strong><sup><sub>&#8211;<\/sub><\/sup><\/strong> and all possible resonance structures. Assign Formal Charges.<\/li>\n<li>Draw the Lewis Dot Structure for\u00a0<strong>CHO<sub>2<\/sub><sup>1<\/sup><sup>&#8211;<\/sup><\/strong> and all possible resonance structures. Assign Formal Charges.<\/li>\n<li>Draw the\u00a0Resonance Hybrid Structure for <strong>P<\/strong><strong>O<\/strong><sub><strong>4<\/strong><\/sub><sup><strong>3<\/strong><\/sup><sup><strong>&#8211;<\/strong><\/sup>.<\/li>\n<li>Draw the\u00a0Resonance Hybrid Structure for <strong>N<\/strong><strong>O<\/strong><sub><strong>3<\/strong><\/sub><sup><strong>&#8211;<\/strong><\/sup>.<\/li>\n<\/ol>\n<\/li>\n<\/ol>\n<div id=\"section_8\">\n<h4 id=\"Problems_.232-2006\"><strong>Problems #2<\/strong><\/h4>\n<p><a title=\"resonance_problems_1_4.png\" href=\"https:\/\/chem.libretexts.org\/@api\/deki\/files\/12253\/resonance_problems_1_4.png?revision=1\" rel=\"internal\"><img decoding=\"async\" class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145316\/resonance_problems_1_4.png\" alt=\"resonance_problems_1_4.png\" width=\"720px\" height=\"433px\" \/><\/a><\/p>\n<p><a title=\"resonance_problems_4_6.png\" href=\"https:\/\/chem.libretexts.org\/@api\/deki\/files\/12254\/resonance_problems_4_6.png?revision=1\" rel=\"internal\"><img decoding=\"async\" class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145318\/resonance_problems_4_6.png\" alt=\"resonance_problems_4_6.png\" width=\"720px\" height=\"331px\" \/><\/a><\/p>\n<\/div>\n<div id=\"section_9\">\n<h3 id=\"Answers-2006\">Answers<\/h3>\n<div class=\"qa-wrapper\" style=\"display: block\"><span class=\"show-answer collapsed\" style=\"cursor: pointer\" data-target=\"q630305\">Show Answer<\/span><\/p>\n<div id=\"q630305\" class=\"hidden-answer\" style=\"display: none\">\n<p>1. False, because the electrons were not moved around, only the atoms (this violates the Resonance Structure Rules).<\/p>\n<p>2. Below are the all Lewis dot structure with formal charges (in red) for Sulfate (<strong>SO<sub>4<\/sub><sup>2<\/sup><sup>&#8211;<\/sup><\/strong>). There isn&#8217;t a most favorable resonance of the Sulfate ion because they are all identical in charge and there is no change in Electronegativity between the Oxygen atoms.<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145321\/Sulfate-resonance.jpg\" alt=\"Sulfate-resonance.jpg\" width=\"486\" height=\"284\" \/><\/p>\n<p>3. Below is the resonance for <strong>CH<sub>3<\/sub>COO<\/strong><sup>&#8211;<\/sup>, formal charges are displayed in red.\u00a0The Lewis Structure with the most formal charges is not desirable, because we want the Lewis Structure with the least formal charge.<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145324\/CHCOO-answerr.jpg\" alt=\"CHCOO-answerr.jpg\" width=\"417\" height=\"338\" \/><\/p>\n<p>4. The resonance for <strong>HPO<sub>3<\/sub><sup>2<\/sup><sup>&#8211;<\/sup><\/strong>, and the formal charges (in red).<a title=\"Phosphite-resonance.jpg\" href=\"https:\/\/chem.libretexts.org\/@api\/deki\/files\/10306\/Phosphite-resonance.jpg?revision=1\" rel=\"internal\"><img decoding=\"async\" class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145325\/Phosphite-resonance.jpg\" alt=\"Phosphite-resonance.jpg\" width=\"720px\" height=\"122px\" \/><\/a><\/p>\n<p>5. The resonance for <strong>CHO<sub>2<\/sub><sup>1<\/sup><sup>&#8211;<\/sup><\/strong>, and the formal charges (in red).<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145327\/CHO2-answer.jpg\" alt=\"CHO2-answer.jpg\" width=\"390\" height=\"159\" \/><\/p>\n<p>6. The resonance hybrid for <strong>PO<\/strong><sub><strong>4<\/strong><\/sub><sup><strong>3<\/strong><\/sup><sup><strong>&#8211;<\/strong><\/sup>, hybrid bonds are in red.<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145329\/PO4.jpg\" alt=\"PO4.jpg\" width=\"250\" height=\"213\" \/><\/p>\n<p>7. The resonance hybrid for <strong>NO<\/strong><sub><strong>3<\/strong><\/sub><sup><strong>&#8211;<\/strong><\/sup>, hybrid bonds are in red.<\/p>\n<p><img decoding=\"async\" class=\"internal default aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/1518\/2017\/10\/05145331\/NO3.jpg\" alt=\"NO3.jpg\" width=\"311px\" height=\"278px\" \/><\/div>\n<\/div>\n<\/div>\n<div id=\"section_10\">\n<p>&nbsp;<\/p>\n<p><strong><strong>Problems #2<\/strong><\/strong><\/p>\n<div class=\"qa-wrapper\" style=\"display: block\"><span class=\"show-answer collapsed\" style=\"cursor: pointer\" data-target=\"q837981\">Show Answer<\/span><\/p>\n<div id=\"q837981\" class=\"hidden-answer\" style=\"display: none\">\n<p><img decoding=\"async\" class=\"internal default aligncenter\" src=\"https:\/\/chem.libretexts.org\/@api\/deki\/files\/12256\/resonance_solutions_5_6.png?revision=1&amp;size=bestfit&amp;width=720&amp;height=334\" alt=\"resonance_solutions_5_6.png\" width=\"720px\" height=\"334px\" \/><\/div>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<div id=\"section_11\">\n<h3 id=\"Contributors-2006\">Contributors<\/h3>\n<ul>\n<li style=\"list-style-type: none\">\n<ul>\n<li>Sharon Wei (UCD), Liza Chu (UCD)<\/li>\n<\/ul>\n<\/li>\n<\/ul>\n<\/div>\n<\/div>\n<\/div>\n","protected":false},"author":44985,"menu_order":5,"template":"","meta":{"_candela_citation":"[]","CANDELA_OUTCOMES_GUID":"","pb_show_title":"on","pb_short_title":"","pb_subtitle":"","pb_authors":[],"pb_section_license":""},"chapter-type":[],"contributor":[],"license":[],"class_list":["post-1346","chapter","type-chapter","status-publish","hentry"],"part":26,"_links":{"self":[{"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/pressbooks\/v2\/chapters\/1346","targetHints":{"allow":["GET"]}}],"collection":[{"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/pressbooks\/v2\/chapters"}],"about":[{"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/wp\/v2\/types\/chapter"}],"author":[{"embeddable":true,"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/wp\/v2\/users\/44985"}],"version-history":[{"count":5,"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/pressbooks\/v2\/chapters\/1346\/revisions"}],"predecessor-version":[{"id":2330,"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/pressbooks\/v2\/chapters\/1346\/revisions\/2330"}],"part":[{"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/pressbooks\/v2\/parts\/26"}],"metadata":[{"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/pressbooks\/v2\/chapters\/1346\/metadata\/"}],"wp:attachment":[{"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/wp\/v2\/media?parent=1346"}],"wp:term":[{"taxonomy":"chapter-type","embeddable":true,"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/pressbooks\/v2\/chapter-type?post=1346"},{"taxonomy":"contributor","embeddable":true,"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/wp\/v2\/contributor?post=1346"},{"taxonomy":"license","embeddable":true,"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/wp-json\/wp\/v2\/license?post=1346"}],"curies":[{"name":"wp","href":"https:\/\/api.w.org\/{rel}","templated":true}]}}