{"id":132,"date":"2017-10-04T14:50:25","date_gmt":"2017-10-04T14:50:25","guid":{"rendered":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/?post_type=chapter&p=132"},"modified":"2018-09-28T19:27:43","modified_gmt":"2018-09-28T19:27:43","slug":"lewis-structures","status":"publish","type":"chapter","link":"https:\/\/courses.lumenlearning.com\/suny-mcc-organicchemistry\/chapter\/lewis-structures\/","title":{"raw":"Lewis Structures","rendered":"Lewis Structures"},"content":{"raw":"
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Using Lewis Dot Symbols to Describe Covalent Bonding<\/h2>\r\n<\/div>\r\n
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\r\n\r\nThis sharing of electrons allowing atoms to \"stick\" together is the basis of covalent bonding.\u00a0 There is some intermediate distant, generally a bit longer than 0.1 nm, or if you prefer 100 pm, at which the attractive forces significantly outweigh the repulsive forces and a bond will be formed if both atoms can achieve a completen s<\/em>2<\/sup>np<\/em>6<\/sup> configuration.\u00a0 It is this behavior that Lewis captured in his octet rule.\u00a0 The valence electron configurations of the constituent atoms of a covalent compound are important factors in determining its structure, stoichiometry, and properties. For example, chlorine, with seven valence electrons, is one electron short of an octet. If two chlorine atoms share their unpaired electrons by making a covalent bond and forming Cl2<\/sub>, they can each complete their valence shell:\r\n
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\"99cbb09f87af5a4b09dd56499e843e08.jpg\"<\/a><\/div>\r\n
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Each chlorine atom now has an octet. The electron pair being shared by the atoms is called a bonding <\/a>pair\u00a0<\/a>; the other three pairs of electrons on each chlorine atom are called lone <\/a>pairs<\/a>. Lone pairs are not involved in covalent bonding. If both electrons in a covalent bond come from the same atom, the bond is called a coordinate covalent <\/a>bond<\/a>.<\/p>\r\n

We can illustrate the formation of a water molecule from two hydrogen atoms and an oxygen atom using Lewis dot symbols:<\/p>\r\n\r\n

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\"254015a5c4a159aa4e955da701127624.jpg\"<\/a><\/div>\r\n
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The structure on the right is the Lewis electron structure<\/em>, or Lewis structure<\/em>, for H2<\/sub>O. With two bonding pairs and two lone pairs, the oxygen atom has now completed its octet. Moreover, by sharing a bonding pair with oxygen, each hydrogen atom now has a full valence shell of two electrons. Chemists usually indicate a bonding pair by a single line, as shown here for our two examples:<\/p>\r\n\r\n

\"67f67edbdb9eef2f11217fe1ff08c85b.jpg\"<\/a><\/div>\r\n
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The following procedure can be used to construct Lewis electron structures for more complex molecules and ions:<\/p>\r\n\r\n

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  1. Arrange the atoms to show specific connections.<\/strong> When there is a central atom, it is usually the least electronegative element in the compound. Chemists usually list this central atom first in the chemical formula (as in CCl4<\/sub> and CO3<\/sub>2\u2212<\/sup>, which both have C as the central atom), which is another clue to the compound\u2019s structure. Hydrogen and the halogens are almost always connected to only one other atom, so they are usually terminal<\/em> rather than central.<\/li>\r\n<\/ol>\r\n
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    Note:<\/strong><\/p>\r\n\r\n

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    The central atom is usually the least electronegative element in the molecule or ion; hydrogen and the halogens are usually terminal.<\/p>\r\n\r\n<\/div>\r\n<\/div>\r\n

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    1. Determine the total number of valence electrons in the molecule or ion<\/strong>. Add together the valence electrons from each atom. (Recall from Chapter 2 <\/a>\u00a0that the number of valence electrons is indicated by the position of the element in the periodic table.) If the species is a polyatomic ion, remember to add or subtract the number of electrons necessary to give the total charge on the ion. For CO3<\/sub>2\u2212<\/sup>, for example, we add two electrons to the total because of the \u22122 charge.<\/li>\r\n \t
    2. Place a bonding pair of electrons between each pair of adjacent atoms to give a single bond<\/strong>. In H2<\/sub>O, for example, there is a bonding pair of electrons between oxygen and each hydrogen.<\/li>\r\n \t
    3. Beginning with the terminal atoms, add enough electrons to each atom to give each atom an octet (two for hydrogen)<\/strong>. These electrons will usually be lone pairs.<\/li>\r\n \t
    4. If any electrons are left over, place them on the central atom<\/strong>. Some atoms are able to accommodate more than eight electrons.<\/li>\r\n \t
    5. If the central atom has fewer electrons than an octet, use lone pairs from terminal atoms to form multiple (double or triple) bonds to the central atom to achieve an octet<\/strong>. This will not change the number of electrons on the terminal atoms.<\/li>\r\n<\/ol>\r\n

      Now let\u2019s apply this procedure to some particular compounds, beginning with one we have already discussed.<\/p>\r\n\r\n

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      H2<\/sub>O<\/h2>\r\n

      1. Because H atoms are almost always terminal, the arrangement within the molecule must be HOH.<\/p>\r\n\r\n<\/div>\r\n<\/div>\r\n

      2. Each H atom (group 1) has 1 valence electron, and the O atom (group 16) has 6 valence electrons, for a total of 8 valence electrons<\/p>\r\n\r\n<\/div>\r\n3. Placing one bonding pair of electrons between the O atom and each H atom gives H:O:H, with 4 electrons left over.\r\n\r\n4. Each H atom has a full valence shell of 2 electrons.\r\n

      5. Adding the remaining 4 electrons to the oxygen (as two lone pairs) gives the following structure:<\/p>\r\n\r\n

      \"f6d9a24622184ca561822f70eb85e58b.jpg\"<\/a><\/div>\r\n

      This is the Lewis structure we drew earlier. Because it gives oxygen an octet and each hydrogen two electrons, we do not need to use step 6.<\/p>\r\n\r\n

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      OCl\u2212<\/sup><\/h2>\r\n<\/div>\r\n<\/div>\r\n

      1. With only two atoms in the molecule, there is no central atom.<\/p>\r\n\r\n<\/div>\r\n

      2. Oxygen (group 16) has 6 valence electrons, and chlorine (group 17) has 7 valence electrons; we must add one more for the negative charge on the ion, giving a total of 14 valence electrons.<\/p>\r\n

      3. Placing a bonding pair of electrons between O and Cl gives O:Cl, with 12 electrons left over.<\/p>\r\n

      4. If we place six electrons (as three lone pairs) on each atom, we obtain the following structure:<\/p>\r\n\r\n

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      \"7129216d30dabf7573f148e54f8dcc51.jpg\"<\/a><\/p>\r\n\r\n<\/div>\r\n

      Each atom now has an octet of electrons, so steps 5 and 6 are not needed. The Lewis electron structure is drawn within brackets as is customary for an ion, with the overall charge indicated outside the brackets, and the bonding pair of electrons is indicated by a solid line. OCl\u2212<\/sup> is the hypochlorite ion, the active ingredient in chlorine laundry bleach and swimming pool disinfectant.<\/p>\r\n\r\n

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      CH2<\/sub>O<\/h2>\r\n<\/div>\r\n<\/div>\r\n

      1. Because carbon is less electronegative than oxygen and hydrogen is normally terminal, C must be the central atom. One possible arrangement is as follows:<\/p>\r\n\r\n<\/div>\r\n

      \"ef5c576c0c84ded384007a33c963a2e1.jpg\"<\/a><\/div>\r\n

      2. Each hydrogen atom (group 1) has one valence electron, carbon (group 14) has 4 valence electrons, and oxygen (group 16) has 6 valence electrons, for a total of [(2)(1)\u00a0+\u00a04\u00a0+\u00a06]\u00a0=\u00a012 valence electrons.<\/p>\r\n

      3. Placing a bonding pair of electrons between each pair of bonded atoms gives the following:<\/p>\r\n\r\n

      \"c30b4f42355f76f67e81f90a19a8f0bd.jpg\"<\/a><\/div>\r\n

      Six electrons are used, and 6 are left over.<\/p>\r\n

      4. Adding all 6 remaining electrons to oxygen (as three lone pairs) gives the following:<\/p>\r\n\r\n

      \"c8ddcf2f0df54fc20f18d0e60408951c.jpg\"<\/a><\/div>\r\n

      Although oxygen now has an octet and each hydrogen has 2 electrons, carbon has only 6 electrons.<\/p>\r\n

      5. There are no electrons left to place on the central atom.<\/p>\r\n

      6. To give carbon an octet of electrons, we use one of the lone pairs of electrons on oxygen to form a carbon\u2013oxygen double bond:<\/p>\r\n\r\n

      \"2f734e95e4e4513e5106a1b61e604cc8.jpg\"<\/a><\/div>\r\n

      Both the oxygen and the carbon now have an octet of electrons, so this is an acceptable Lewis electron structure. The O has two bonding pairs and two lone pairs, and C has four bonding pairs. This is the structure of formaldehyde, which is used in embalming fluid.<\/p>\r\n

      An alternative structure can be drawn with one H bonded to O. Formal charges<\/em>, discussed later in this section, suggest that such a structure is less stable than that shown previously.<\/p>\r\n\r\n

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      Example<\/h3>\r\n

      Write the Lewis electron structure for each species.<\/p>\r\n\r\n

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      1. NCl3<\/sub><\/li>\r\n \t
      2. S2<\/sub>2\u2212<\/sup><\/li>\r\n \t
      3. NOCl<\/li>\r\n<\/ol>\r\n

        Given: <\/strong>chemical species<\/p>\r\n

        Asked for: <\/strong>Lewis electron structures<\/p>\r\n

        Strategy:<\/strong><\/p>\r\n

        Use the six-step procedure to write the Lewis electron structure for each species.\r\n[reveal-answer q=\"941055\"]Show Answer[\/reveal-answer]\r\n[hidden-answer a=\"941055\"]Nitrogen is less electronegative than chlorine, and halogen atoms are usually terminal, so nitrogen is the central atom. The nitrogen atom (group 15) has 5 valence electrons and each chlorine atom (group 17) has 7 valence electrons, for a total of 26 valence electrons. Using 2 electrons for each N\u2013Cl bond and adding three lone pairs to each Cl account for (3\u00a0\u00d7\u00a02)\u00a0+\u00a0(3\u00a0\u00d7\u00a02\u00a0\u00d7\u00a03)\u00a0=\u00a024 electrons. Rule 5 leads us to place the remaining 2 electrons on the central N:<\/p>\r\n

        \"579121db32b7b48c36ae54dcf17ee6d8.jpg\"<\/a>[\/hidden-answer]<\/p>\r\n\r\n<\/div>\r\n

        Nitrogen trichloride is an unstable oily liquid once used to bleach flour; this use is now prohibited in the United States.<\/p>\r\n\r\n

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        1. \r\n
          \"e9847c1f1728e285bb9dee914470d140.jpg\"<\/a><\/div><\/li>\r\n \t
        2. \r\n

          In a diatomic molecule or ion, we do not need to worry about a central atom. Each sulfur atom (group 16) contains 6 valence electrons, and we need to add 2 electrons for the \u22122 charge, giving a total of 14 valence electrons. Using 2 electrons for the S\u2013S bond, we arrange the remaining 12 electrons as three lone pairs on each sulfur, giving each S atom an octet of electrons:<\/p>\r\n\r\n

          \"6e69300f5796dbb06798c387bd7bcc89.jpg\"<\/a><\/div><\/li>\r\n \t
        3. \r\n

          Because nitrogen is less electronegative than oxygen or chlorine, it is the central atom. The N atom (group 15) has 5 valence electrons, the O atom (group 16) has 6 valence electrons, and the Cl atom (group 17) has 7 valence electrons, giving a total of 18 valence electrons. Placing one bonding pair of electrons between each pair of bonded atoms uses 4 electrons and gives the following:<\/p>\r\n\r\n

          \"dc6a5041bcdabb91cb727522339fe44b.jpg\"<\/a><\/div>\r\n

          Adding three lone pairs each to oxygen and to chlorine uses 12 more electrons, leaving 2 electrons to place as a lone pair on nitrogen:<\/p>\r\n\r\n

          \"312e28bcf7eb8ae248c0e7fa83ea0f22.jpg\"<\/a><\/div><\/li>\r\n \t
        4. \r\n

          Because this Lewis structure has only 6 electrons around the central nitrogen, a lone pair of electrons on a terminal atom must be used to form a bonding pair. We could use a lone pair on either O or Cl. Because we have seen many structures in which O forms a double bond but none with a double bond to Cl, it is reasonable to select a lone pair from O to give the following:<\/p>\r\n\r\n

          \"56192089dcd6a3ab3cfc307d5c85f9b0.jpg\"<\/a><\/div>\r\n

          All atoms now have octet configurations. This is the Lewis electron structure of nitrosyl chloride, a highly corrosive, reddish-orange gas.<\/p>\r\n\r\n

          \"5880b3fce5aac0e0cc466c9d707c9ee6.jpg\"<\/a><\/div><\/li>\r\n<\/ol>\r\n<\/div>\r\n<\/div>\r\n
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          Example<\/h3>\r\n

          Write Lewis electron structures for CO2<\/sub> and SCl2<\/sub>, a vile-smelling, unstable red liquid that is used in the manufacture of rubber.\r\n[reveal-answer q=\"594677\"]Show Answer[\/reveal-answer]\r\n[hidden-answer a=\"594677\"]<\/p>\r\n

          \"03bae19e91ea497c755e3fb932ed38d6.jpg\"<\/a>\"83b695097f238c87d66b81488d2c724a.jpg\"<\/a>\"5a3bdca4a8ddf4937091e34428109a33.jpg\"\"1f1706c3f1e84e24caca141d0bcf5843.jpg\"<\/p>\r\n[\/hidden-answer]\r\n\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n

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          <\/h3>\r\n

          It is sometimes possible to write more than one Lewis structure for a substance that does not violate the octet rule, as we saw for CH2<\/sub>O, but not every Lewis structure may be equally reasonable. In these situations, we can choose the most stable Lewis structure by considering the formal <\/a>charge<\/a> on the atoms, which is the difference between the number of valence electrons in the free atom and the number assigned to it in the Lewis electron structure. The formal charge is a way of computing the charge distribution within a Lewis structure; the sum of the formal charges on the atoms within a molecule or an ion must equal the overall charge on the molecule or ion. A formal charge does not<\/em> represent a true charge on an atom in a covalent bond but is simply used to predict the most likely structure when a compound has more than one valid Lewis structure.<\/p>\r\n

          To calculate formal charges, we assign electrons in the molecule to individual atoms according to these rules:<\/p>\r\n\r\n