{"id":1978,"date":"2015-04-22T20:17:48","date_gmt":"2015-04-22T20:17:48","guid":{"rendered":"https:\/\/courses.candelalearning.com\/oschemtemp\/?post_type=chapter&#038;p=1978"},"modified":"2015-08-31T16:20:30","modified_gmt":"2015-08-31T16:20:30","slug":"valence-bond-theory-3-formulas-missing","status":"publish","type":"chapter","link":"https:\/\/courses.lumenlearning.com\/suny-buffstate-chemistryformajorsxmaster\/chapter\/valence-bond-theory-3-formulas-missing\/","title":{"raw":"Valence Bond Theory","rendered":"Valence Bond Theory"},"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>Describe the formation of covalent bonds in terms of atomic orbital overlap<\/li>\r\n\t<li>Define and give examples of \u03c3 and \u03c0 bonds<\/li>\r\n<\/ul>\r\n<\/div>\r\nAs we know, a scientific theory is a strongly supported explanation for observed natural laws or large bodies of experimental data. For a theory to be accepted, it must explain experimental data and be able to predict behavior. For example, VSEPR theory has gained widespread acceptance because it predicts three-dimensional molecular shapes that are consistent with experimental data collected for thousands of different molecules. However, VSEPR theory does not provide an explanation of chemical bonding.\r\n<p id=\"fs-idp41602432\">There are successful theories that describe the electronic structure of atoms. We can use quantum mechanics to predict the specific regions around an atom where electrons are likely to be located: A spherical shape for an <em data-effect=\"italics\">s<\/em> orbital, a dumbbell shape for a <em data-effect=\"italics\">p<\/em> orbital, and so forth. However, these predictions only describe the orbitals around free atoms. When atoms bond to form molecules, atomic orbitals are not sufficient to describe the regions where electrons will be located in the molecule. A more complete understanding of electron distributions requires a model that can account for the electronic structure of molecules. One popular theory holds that a covalent bond forms when a pair of electrons is shared by two atoms and is simultaneously attracted by the nuclei of both atoms. In the following sections, we will discuss how such bonds are described by valence bond theory and hybridization.<\/p>\r\n<p id=\"fs-idp56452464\"><strong><span data-type=\"term\">Valence bond theory<\/span><\/strong> describes a covalent bond as the overlap of half-filled atomic orbitals (each containing a single electron) that yield a pair of electrons shared between the two bonded atoms. We say that orbitals on two different atoms <strong><span data-type=\"term\">overlap<\/span><\/strong> when a portion of one orbital and a portion of a second orbital occupy the same region of space. According to valence bond theory, a covalent bond results when two conditions are met: (1) an orbital on one atom overlaps an orbital on a second atom and (2) the single electrons in each orbital combine to form an electron pair. The mutual attraction between this negatively charged electron pair and the two atoms\u2019 positively charged nuclei serves to physically link the two atoms through a force we define as a covalent bond. The strength of a covalent bond depends on the extent of overlap of the orbitals involved. Orbitals that overlap extensively form bonds that are stronger than those that have less overlap.<\/p>\r\n<p id=\"fs-idp45162384\">The energy of the system depends on how much the orbitals overlap. Figure 1 illustrates how the sum of the energies of two hydrogen atoms (the colored curve) changes as they approach each other. When the atoms are far apart there is no overlap, and by convention we set the sum of the energies at zero. As the atoms move together, their orbitals begin to overlap. Each electron begins to feel the attraction of the nucleus in the other atom. In addition, the electrons begin to repel each other, as do the nuclei. While the atoms are still widely separated, the attractions are slightly stronger than the repulsions, and the energy of the system decreases. (A bond begins to form.) As the atoms move closer together, the overlap increases, so the attraction of the nuclei for the electrons continues to increase (as do the repulsions among electrons and between the nuclei). At some specific distance between the atoms, which varies depending on the atoms involved, the energy reaches its lowest (most stable) value. This optimum distance between the two bonded nuclei is the bond distance between the two atoms. The bond is stable because at this point, the attractive and repulsive forces combine to create the lowest possible energy configuration. If the distance between the nuclei were to decrease further, the repulsions between nuclei and the repulsions as electrons are confined in closer proximity to each other would become stronger than the attractive forces. The energy of the system would then rise (making the system destabilized), as shown at the far left of Figure 1.<\/p>\r\n\r\n\r\n[caption id=\"\" align=\"alignnone\" width=\"880\"]<img class=\"\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23211750\/CNX_Chem_08_01_Morse1.jpg\" alt=\"A pair of diagrams are shown and labeled \u201ca\u201d and \u201cb\u201d. Diagram a shows three consecutive images. The first image depicts two separated blurry circles, each labeled with a positive sign and the term \u201cH atom.\u201d The phrase written under them reads, \u201cSufficiently far apart to have no interaction.\u201d The second image shows the same two circles, but this time they are much closer together and are labeled, \u201cAtoms begin to interact as they move closer together.\u201d The third image shows the two circles overlapping, labeled, \u201cH subscript 2,\u201d and, \u201cOptimum distance to achieve lowest overall energy of system.\u201d Diagram b shows a graph on which the y-axis is labeled \u201cEnergy ( J ),\u201d and the x-axis is labeled, \u201cInternuclear distance ( p m ).\u201d The midpoint of the y-axis is labeled as zero. The curve on the graph begins at zero p m and high on the y-axis. The graph slopes downward steeply to a point far below the zero joule line on the y-axis and the lowest point reads \u201c0.74 p m\u201d and \u201cH bonded to H bond length.\u201d It is also labeled \u201c negative 7.24 times 10 superscript negative 19 J.\u201d The graph then rises again to zero J. The graph is accompanied by the same images from diagram a; the first image correlates to the point in the graph where it crosses the zero point on the y-axis, the third image where the graph is lowest.\" width=\"880\" height=\"747\" data-media-type=\"image\/jpeg\" \/> Figure 1. (a) The interaction of two hydrogen atoms changes as a function of distance. (b) The energy of the system changes as the atoms interact. The lowest (most stable) energy occurs at a distance of 74 pm, which is the bond length observed for the H<sub>2<\/sub> molecule.[\/caption]\r\n<p id=\"fs-idm5865744\">The bond energy is the difference between the energy minimum (which occurs at the bond distance) and the energy of the two separated atoms. This is the quantity of energy released when the bond is formed. Conversely, the same amount of energy is required to break the bond. For the H<sub>2<\/sub> molecule shown in Figure 1, at the bond distance of 74 pm the system is 7.24 \u00d7 10<sup>-19<\/sup> J lower in energy than the two separated hydrogen atoms. This may seem like a small number. However, we know from our earlier description of thermochemistry that bond energies are often discussed on a per-mole basis. For example, it requires 7.24 \u00d7 10<sup>-19<\/sup> J to break one H\u2013H bond, but it takes 4.36 \u00d7 10<sup>5<\/sup> J to break 1 mole of H\u2013H bonds. A comparison of some bond lengths and energies is shown in\u00a0Table 1. We can find many of these bonds in a variety of molecules, and this table provides average values. For example, breaking the first C\u2013H bond in CH<sub>4<\/sub> requires 439.3 kJ\/mol, while breaking the first C\u2013H bond in H\u2013CH<sub>2<\/sub>C<sub>6<\/sub>H<sub>5<\/sub> (a common paint thinner) requires 375.5 kJ\/mol.<\/p>\r\n\r\n<table id=\"fs-idp43355056\" summary=\"A table is shown that contains six columns and eleven rows. The header row reads \u201cBond,\u201d \u201cLength ( p m ),\u201d \u201cEnergy ( k J \/ mol ),\u201d \u201cBond,\u201d \u201cLength ( p m ),\u201d and \u201cEnergy ( k J \/ mol ).\u201d The first column contains the symbols: \u201cH bond H,\u201d \u201cH bond C,\u201d \u201cH bond N,\u201d \u201cH bond O,\u201d \u201cC bond C,\u201d \u201cC double bond C,\u201d \u201cC triple bond C,\u201d \u201cC bond N,\u201d \u201cC double bond N,\u201d and \u201cC triple bond N.\u201d The second column contains the numbers \u201c74,\u201d \u201c106.8,\u201d \u201c101.5,\u201d \u201c97.5,\u201d \u201c150.6,\u201d \u201c133.5,\u201d \u201c120.8,\u201d \u201c142.1,\u201d \u201c130.0,\u201d and \u201c116.1.\u201d The third column contains the numbers: \u201c436,\u201d \u201c413,\u201d \u201c391,\u201d \u201c467,\u201d \u201c347,\u201d \u201c614,\u201d \u201c839,\u201d \u201c305,\u201d \u201c615,\u201d and \u201c891.\u201d The fourth column contains the symbols: \u201cC bond O,\u201d \u201cC double bond O,\u201d \u201cC triple bond O,\u201d \u201cH bond C l,\u201d \u201cH bond B r,\u201d \u201cH bond I,\u201d \u201cO bond O,\u201d \u201cO double bond O,\u201d \u201cF bond F,\u201d and \u201cC l bond C l.\u201d The fifth column contains the numbers \u201c140.1,\u201d \u201c119.7,\u201d \u201c113.7,\u201d \u201c127.5,\u201d \u201c141.4,\u201d \u201c160.9,\u201d \u201c148,\u201d \u201c120.8,\u201d \u201c141.2,\u201d and \u201c198.8.\u201d The sixth column contains the numbers \u201c358,\u201d \u201c745,\u201d \u201c1072,\u201d \u201c431,\u201d \u201c366,\u201d \u201c298,\u201d \u201c146,\u201d \u201c498,\u201d \u201c159,\u201d and \u201c243.\u201d\"><caption><span data-type=\"title\">Table 1. Representative Bond Energies and Lengths<\/span><\/caption>\r\n<thead>\r\n<tr valign=\"top\">\r\n<th>Bond<\/th>\r\n<th>Length (pm)<\/th>\r\n<th>Energy (kJ\/mol)<\/th>\r\n<th>Bond<\/th>\r\n<th>Length (pm)<\/th>\r\n<th>Energy (kJ\/mol)<\/th>\r\n<\/tr>\r\n<\/thead>\r\n<tbody>\r\n<tr valign=\"top\">\r\n<td>H\u2013H<\/td>\r\n<td>74<\/td>\r\n<td>436<\/td>\r\n<td>C\u2013O<\/td>\r\n<td>140.1<\/td>\r\n<td>358<\/td>\r\n<\/tr>\r\n<tr valign=\"top\">\r\n<td>H\u2013C<\/td>\r\n<td>106.8<\/td>\r\n<td>413<\/td>\r\n<td>[latex]\\text{C}=\\text{O}[\/latex]<\/td>\r\n<td>119.7<\/td>\r\n<td>745<\/td>\r\n<\/tr>\r\n<tr valign=\"top\">\r\n<td>H\u2013N<\/td>\r\n<td>101.5<\/td>\r\n<td>391<\/td>\r\n<td>[latex]\\text{C}\\equiv \\text{O}[\/latex]<\/td>\r\n<td>113.7<\/td>\r\n<td>1072<\/td>\r\n<\/tr>\r\n<tr valign=\"top\">\r\n<td>H\u2013O<\/td>\r\n<td>97.5<\/td>\r\n<td>467<\/td>\r\n<td>H\u2013Cl<\/td>\r\n<td>127.5<\/td>\r\n<td>431<\/td>\r\n<\/tr>\r\n<tr valign=\"top\">\r\n<td>C\u2013C<\/td>\r\n<td>150.6<\/td>\r\n<td>347<\/td>\r\n<td>H\u2013Br<\/td>\r\n<td>141.4<\/td>\r\n<td>366<\/td>\r\n<\/tr>\r\n<tr valign=\"top\">\r\n<td>[latex]\\text{C}=\\text{C}[\/latex]<\/td>\r\n<td>133.5<\/td>\r\n<td>614<\/td>\r\n<td>H\u2013I<\/td>\r\n<td>160.9<\/td>\r\n<td>298<\/td>\r\n<\/tr>\r\n<tr valign=\"top\">\r\n<td>[latex]\\text{C}\\equiv \\text{C}[\/latex]<\/td>\r\n<td>120.8<\/td>\r\n<td>839<\/td>\r\n<td>O\u2013O<\/td>\r\n<td>148<\/td>\r\n<td>146<\/td>\r\n<\/tr>\r\n<tr valign=\"top\">\r\n<td>C\u2013N<\/td>\r\n<td>142.1<\/td>\r\n<td>305<\/td>\r\n<td>[latex]\\text{O}=\\text{O}[\/latex]<\/td>\r\n<td>120.8<\/td>\r\n<td>498<\/td>\r\n<\/tr>\r\n<tr valign=\"top\">\r\n<td>[latex]\\text{C}=\\text{N}[\/latex]<\/td>\r\n<td>130.0<\/td>\r\n<td>615<\/td>\r\n<td>F\u2013F<\/td>\r\n<td>141.2<\/td>\r\n<td>159<\/td>\r\n<\/tr>\r\n<tr valign=\"top\">\r\n<td>[latex]\\text{C}\\equiv \\text{N}[\/latex]<\/td>\r\n<td>116.1<\/td>\r\n<td>891<\/td>\r\n<td>Cl\u2013Cl<\/td>\r\n<td>198.8<\/td>\r\n<td>243<\/td>\r\n<\/tr>\r\n<\/tbody>\r\n<\/table>\r\n<p id=\"fs-idm10626416\">In addition to the distance between two orbitals, the orientation of orbitals also affects their overlap (other than for two <em data-effect=\"italics\">s<\/em> orbitals, which are spherically symmetric). Greater overlap is possible when orbitals are oriented such that they overlap on a direct line between the two nuclei. Figure 2 illustrates this for two <em data-effect=\"italics\">p<\/em> orbitals from different atoms; the overlap is greater when the orbitals overlap end to end rather than at an angle.<\/p>\r\n\r\n\r\n[caption id=\"\" align=\"alignnone\" width=\"650\"]<img src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23211752\/CNX_Chem_08_01_overlap1.jpg\" alt=\"Two diagrams are shown. Diagram a contains two molecules whose p orbitals, which are depicted as two balloon-shaped structures that meet together to form a peanut shape, are laid end over end, creating an area of overlap. In diagram b, the same two molecules are shown, but this time, they are laid out in a way so as to form a near-ninety degree angle. In this diagram, the ends of two of these peanut-shaped orbitals do not overlap nearly as much.\" width=\"650\" height=\"158\" data-media-type=\"image\/jpeg\" \/> Figure 2. (a) The overlap of two p orbitals is greatest when the orbitals are directed end to end. (b) Any other arrangement results in less overlap. The plus signs indicate the locations of the nuclei.[\/caption]\r\n<p id=\"fs-idp40817312\">The overlap of two <em data-effect=\"italics\">s<\/em> orbitals (as in H<sub>2<\/sub>), the overlap of an <em data-effect=\"italics\">s<\/em> orbital and a <em data-effect=\"italics\">p<\/em> orbital (as in HCl), and the end-to-end overlap of two <em data-effect=\"italics\">p<\/em> orbitals (as in Cl<sub>2<\/sub>) all produce <span data-type=\"term\">sigma bonds (\u03c3 bonds)<\/span>, as illustrated in Figure 3. A \u03c3 bond is a covalent bond in which the electron density is concentrated in the region along the internuclear axis; that is, a line between the nuclei would pass through the center of the overlap region. Single bonds in Lewis structures are described as \u03c3 bonds in valence bond theory.<\/p>\r\n\r\n\r\n[caption id=\"\" align=\"alignnone\" width=\"650\"]<img src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23211753\/CNX_Chem_08_01_sigma1.jpg\" alt=\"Three diagrams are shown and labeled \u201ca,\u201d \u201cb,\u201d and \u201cc.\u201d Diagram a shows two spherical orbitals lying side by side and overlapping. Diagram b shows one spherical and one peanut-shaped orbital lying near one another so that the spherical orbital overlaps with one end of the peanut-shaped orbital. Diagram c shows two peanut-shaped orbitals lying end to end so that one end of each orbital overlaps the other.\" width=\"650\" height=\"131\" data-media-type=\"image\/jpeg\" \/> Figure 3. Sigma (\u03c3) bonds form from the overlap of the following: (a) two s orbitals, (b) an s orbital and a p orbital, and (c) two p orbitals. The plus signs indicate the locations of the nuclei.[\/caption]\r\n\r\n[caption id=\"\" align=\"alignright\" width=\"325\"]<img src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23211754\/CNX_Chem_08_01_pi1.jpg\" alt=\"Two peanut-shaped orbitals are shown, lying vertically and parallel with one another. They overlap one another along the top and bottom of the orbital.\" width=\"325\" height=\"115\" data-media-type=\"image\/jpeg\" \/> Figure 4. Pi (\u03c0) bonds form from the side-by-side overlap of two p orbitals. The plus signs indicate the location of the nuclei.[\/caption]\r\n<p id=\"fs-idp92007424\">A <strong><span data-type=\"term\">pi bond (\u03c0 bond)<\/span><\/strong> is a type of covalent bond that results from the side-by-side overlap of two <em data-effect=\"italics\">p<\/em> orbitals, as illustrated in Figure 4. In a \u03c0 bond, the regions of orbital overlap lie on opposite sides of the internuclear axis. Along the axis itself, there is a <strong><span data-type=\"term\">node<\/span><\/strong>, that is, a plane with no probability of finding an electron.<\/p>\r\n<p id=\"fs-idp46373312\">While all single bonds are \u03c3 bonds, multiple bonds consist of both \u03c3 and \u03c0 bonds. As the Lewis structures in suggest, O<sub>2<\/sub> contains a double bond, and N<sub>2<\/sub> contains a triple bond. The double bond consists of one \u03c3 bond and one \u03c0 bond, and the triple bond consists of one \u03c3 bond and two \u03c0 bonds. Between any two atoms, the first bond formed will always be a \u03c3 bond, but there can only be one \u03c3 bond in any one location. In any multiple bond, there will be one \u03c3 bond, and the remaining one or two bonds will be \u03c0 bonds. These bonds are described in more detail later in this chapter.<\/p>\r\n<span id=\"fs-idp200689696\" data-type=\"media\" data-alt=\"A diagram contains three Lewis structures. The left most structure shows an H atom bonded to a C l atom by a single bond. The C l atom has three lone pairs of electrons. The phrase \u201cOne sigma bond No pi bonds\u201d is written below the drawing. The center structure shows two O atoms bonded by a double bond. The O atoms each have two lone pairs of electrons. The phrase \u201cOne sigma bond One pi bond\u201d is written below the drawing. The right most structure shows two N atoms bonded by a triple bond. Each N atom has a lone pairs of electrons. The phrase \u201cOne sigma bond Two pi bonds\u201d is written below the drawing.\">\r\n<img src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23211755\/CNX_Chem_08_01_bondtype_img1.jpg\" alt=\"A diagram contains three Lewis structures. The left most structure shows an H atom bonded to a C l atom by a single bond. The C l atom has three lone pairs of electrons. The phrase \u201cOne sigma bond No pi bonds\u201d is written below the drawing. The center structure shows two O atoms bonded by a double bond. The O atoms each have two lone pairs of electrons. The phrase \u201cOne sigma bond One pi bond\u201d is written below the drawing. The right most structure shows two N atoms bonded by a triple bond. Each N atom has a lone pairs of electrons. The phrase \u201cOne sigma bond Two pi bonds\u201d is written below the drawing.\" data-media-type=\"image\/jpeg\" \/><\/span>\r\n<p id=\"fs-idp228190304\">As seen in Table 1, an average carbon-carbon single bond is 347 kJ\/mol, while in a carbon-carbon double bond, the \u03c0 bond increases the bond strength by 267 kJ\/mol. Adding an additional \u03c0 bond causes a further increase of 225 kJ\/mol. We can see a similar pattern when we compare other \u03c3 and \u03c0 bonds. Thus, each individual \u03c0 bond is generally weaker than a corresponding \u03c3 bond between the same two atoms. In a \u03c3 bond, there is a greater degree of orbital overlap than in a \u03c0 bond.<\/p>\r\n\r\n<div id=\"fs-idp85049088\" data-type=\"example\">\r\n<div class=\"textbox shaded\">\r\n<h3>Example 1<\/h3>\r\n<h4 id=\"fs-idm26209776\"><strong><span data-type=\"title\">Counting \u03c3 and \u03c0 Bonds<\/span><\/strong>\r\n<span id=\"fs-idp98308768\" data-type=\"media\" data-alt=\"This figure shows a molecule composed of four carbon atoms. There is a double bond between carbons one and two and three and four, while a single bond holds carbon two and three together. Carbons one and four are also bonded to two hydrogens with a single bond while carbons two and three are each bonded to one hydrogen each by a single bond.\">\r\n<img src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23211757\/CNX_Chem_08_01_C4H6_img1.jpg\" alt=\"This figure shows a molecule composed of four carbon atoms. There is a double bond between carbons one and two and three and four, while a single bond holds carbon two and three together. Carbons one and four are also bonded to two hydrogens with a single bond while carbons two and three are each bonded to one hydrogen each by a single bond.\" data-media-type=\"image\/jpeg\" \/><\/span><\/h4>\r\n<p id=\"fs-idm57104384\">Butadiene, C<sub>6<\/sub>H<sub>6<\/sub>, is used to make synthetic rubber. Identify the number of \u03c3 and \u03c0 bonds contained in this molecule.<\/p>\r\n\r\n<h3 id=\"fs-idm9825456\"><span data-type=\"title\">Solution<\/span><\/h3>\r\nThere are six \u03c3 C\u2013H bonds and one \u03c3 C\u2013C bond, for a total of seven from the single bonds. There are two double bonds that each have a \u03c0 bond in addition to the \u03c3 bond. This gives a total nine \u03c3 and two \u03c0 bonds overall.\r\n<h4 id=\"fs-idp46378096\"><strong><span data-type=\"title\">Check Your Learning<\/span><\/strong><\/h4>\r\nIdentify each illustration as depicting a \u03c3 or \u03c0 bond:\r\n<p id=\"fs-idp138827456\">(a) side-by-side overlap of a 4<em data-effect=\"italics\">p<\/em> and a 2<em data-effect=\"italics\">p<\/em> orbital<\/p>\r\n<p id=\"fs-idp7133408\">(b) end-to-end overlap of a 4<em data-effect=\"italics\">p<\/em> and 4<em data-effect=\"italics\">p<\/em> orbital<\/p>\r\n<p id=\"fs-idp145728832\">(c) end-to-end overlap of a 4<em data-effect=\"italics\">p<\/em> and a 2<em data-effect=\"italics\">p<\/em> orbital<\/p>\r\n<span id=\"fs-idm5769888\" data-type=\"media\" data-alt=\"Three diagrams are shown and labeled \u201ca,\u201d \u201cb,\u201d and \u201cc.\u201d Diagram a depicts two peanut-shaped orbitals lying vertically side-by-side and overlapping. One orbital is smaller than the other. Diagram b shows two peanut-shaped orbitals lying end-to-end and overlapping. Diagram c shows two unequally sized peanut-shaped orbitals lying end-to-end and overlapping.\">\r\n<img class=\"\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23211758\/CNX_Chem_08_01_Exover_img1.jpg\" alt=\"Three diagrams are shown and labeled \u201ca,\u201d \u201cb,\u201d and \u201cc.\u201d Diagram a depicts two peanut-shaped orbitals lying vertically side-by-side and overlapping. One orbital is smaller than the other. Diagram b shows two peanut-shaped orbitals lying end-to-end and overlapping. Diagram c shows two unequally sized peanut-shaped orbitals lying end-to-end and overlapping.\" width=\"873\" height=\"137\" data-media-type=\"image\/jpeg\" \/><\/span>\r\n<div id=\"fs-idp42592272\" data-type=\"note\">\r\n<div data-type=\"title\"><\/div>\r\n<div style=\"text-align: right;\" data-type=\"title\"><strong>Answer<\/strong>:\u00a0(a) is a \u03c0 bond with a node along the axis connecting the nuclei while (b) and (c) are \u03c3 bonds that overlap along the axis.<\/div>\r\n<\/div>\r\n<\/div>\r\n<section id=\"fs-idm40478192\" class=\"summary\" data-depth=\"1\">\r\n<div id=\"post-355\" class=\"post-355 chapter type-chapter status-publish hentry type-1\">\r\n<div class=\"entry-content\">\r\n<div class=\"im_section\">\r\n<div class=\"im_section\">\r\n<div id=\"mclean-ch03_s01_s02_n01\" class=\"im_key_takeaways im_editable im_block\">\r\n<div class=\"bcc-box bcc-success\">\r\n<h2>Key Concepts and Summary<\/h2>\r\n<section>\r\n<div data-type=\"note\">\r\n<p id=\"fs-idp236281408\">Valence bond theory describes bonding as a consequence of the overlap of two separate atomic orbitals on different atoms that creates a region with one pair of electrons shared between the two atoms. When the orbitals overlap along an axis containing the nuclei, they form a \u03c3 bond. When they overlap in a fashion that creates a node along this axis, they form a \u03c0 bond.<\/p>\r\n\r\n<\/div>\r\n<\/section><\/div>\r\n<div class=\"bcc-box bcc-info\">\r\n<h3>Chemistry End of Chapter Exercises<\/h3>\r\n<section id=\"fs-idp52122896\" class=\"exercises\" data-depth=\"1\">\r\n<div id=\"fs-idp92773152\" data-type=\"exercise\">\r\n<div id=\"fs-idp39612288\" data-type=\"problem\">\r\n<ol>\r\n\t<li id=\"fs-idp78431584\">Explain how \u03c3 and \u03c0 bonds are similar and how they are different.<\/li>\r\n\t<li>Draw a curve that describes the energy of a system with H and Cl atoms at varying distances. Then, find the minimum energy of this curve two ways.\r\n<ol>\r\n\t<li>Use the bond energy found in Table 8.1\u00a0to calculate the energy for one single HCl bond (Hint: How many bonds are in a mole?)<\/li>\r\n\t<li>Use the enthalpy of reaction and the bond energies for H<sub>2<\/sub> and Cl<sub>2<\/sub> to solve for the energy of one mole of HCl bonds. [latex]{\\text{H}}_{2}\\left(g\\right)+{\\text{Cl}}_{2}\\left(g\\right)\\rightleftharpoons2\\text{HCl}\\left(g\\right)\\Delta{H}_{\\text{rxn}}^{\\textdegree }=\\text{-184.7 kJ\/mol}[\/latex]<\/li>\r\n<\/ol>\r\n<\/li>\r\n\t<li>Explain why bonds occur at specific average bond distances instead of the atoms approaching each other infinitely close.<\/li>\r\n\t<li>Use valence bond theory to explain the bonding in F<sub>2<\/sub>, HF, and ClBr. Sketch the overlap of the atomic orbitals involved in the bonds.<\/li>\r\n\t<li>Use valence bond theory to explain the bonding in O<sub>2<\/sub>. Sketch the overlap of the atomic orbitals involved in the bonds in O<sub>2<\/sub>.<\/li>\r\n\t<li>How many \u03c3 and \u03c0 bonds are present in the molecule HCN?<\/li>\r\n\t<li>A friend tells you N<sub>2<\/sub> has three \u03c0 bonds due to overlap of the three <em data-effect=\"italics\">p<\/em>-orbitals on each N atom. Do you agree?<\/li>\r\n\t<li>Draw the Lewis structures for CO<sub>2<\/sub> and CO, and predict the number of \u03c3 and \u03c0 bonds for each molecule. (a) CO<sub>2\u00a0<\/sub>(b) CO<\/li>\r\n<\/ol>\r\n<\/div>\r\n<\/div>\r\n<\/section><\/div>\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<div id=\"post-355\" class=\"post-355 chapter type-chapter status-publish hentry type-1\">\r\n<div class=\"entry-content\">\r\n<div class=\"im_section\">\r\n<div class=\"im_section\">\r\n<div id=\"mclean-ch03_s01_s02_n01\" class=\"im_key_takeaways im_editable im_block\">\r\n<div class=\"bcc-box bcc-info\">\r\n<h4>Selected Answers<\/h4>\r\n1.\u00a0Similarities: Both types of bonds result from overlap of atomic orbitals on adjacent atoms and contain a maximum of two electrons. Differences: \u03c3 bonds are stronger and result from end-to-end overlap and all single bonds are \u03c3 bonds; \u03c0 bonds between the same two atoms are weaker because they result from side-by-side overlap, and multiple bonds contain one or more \u03c0 bonds (in addition to a \u03c3 bond).\r\n<div data-type=\"newline\">3.\u00a0The specific average bond distance is the distance with the lowest energy. At distances less than the bond distance, the positive charges on the two nuclei repel each other, and the overall energy increases.<\/div>\r\n<div data-type=\"newline\"><\/div>\r\n<div data-type=\"newline\">5. Bonding: One \u03c3 bond and one \u03c0 bond. The <em data-effect=\"italics\">s<\/em> orbitals are filled and do not overlap. The <em data-effect=\"italics\">p<\/em> orbitals overlap along the axis to form a \u03c3 bond and side by side to form the \u03c0 bond\r\n.<img class=\"\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23211801\/CNX_Chem_08_01_O2bonds_img1.jpg\" alt=\"This figure shows the orbitals of two atoms being added together as they form bonds. The two atoms are shown separately on the right, each having two peanut-shaped orbitals lying perpendicularly to one another. A right-facing arrow shows that the two have moved closer together and now the upper and lower portions of the vertical peanut-shaped orbitals are shown as merging together above and below the plane of the molecule while the horizontal peanut-shaped orbitals are overlapping between the two nuclei.\" width=\"846\" height=\"222\" data-media-type=\"image\/jpeg\" \/>\r\n<div data-type=\"newline\">7. No, two of the <em data-effect=\"italics\">p<\/em> orbitals (one on each N) will be oriented end to end and will form a \u03c3 bond.<\/div>\r\n<div data-type=\"newline\">\r\n<div id=\"fs-idp108425904\" data-type=\"exercise\">\r\n<div id=\"fs-idp75120720\" data-type=\"solution\">\r\n<div data-type=\"newline\">\u00a0<img class=\"\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23211803\/CNX_Chem_08_01_N2LewStru_img1.jpg\" alt=\"Two nitrogen atoms are shown both in a Lewis structure and as a diagram. The Lewis structure depicts the two nitrogen atoms bonded by a triple bond. The diagram shows two nitrogen atoms and their three peanut-shaped p-orbitals. One of the orbitals lies horizontally and overlaps between the two nuclei. It is labeled, \u201csigma bond.\u201d The other two lie vertically and in the z-plane of the page. They overlap above and below and into and out of the page in relation to the nuclei. They are each labeled, \u201cpi bond.\u201d\" width=\"528\" height=\"208\" data-media-type=\"image\/jpeg\" \/><\/div>\r\n<div data-type=\"newline\"><\/div>\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<\/section>\r\n<div class=\"bcc-box bcc-success\"><section id=\"glossary\">\r\n<h3>Glossary<\/h3>\r\n<div data-type=\"definition\">\r\n<div id=\"fs-idm8143856\" data-type=\"definition\">\r\n<div data-type=\"glossary\">\r\n<p data-type=\"definition\"><strong><span data-type=\"term\">overlap\r\n<\/span><\/strong>coexistence of orbitals from two different atoms sharing the same region of space, leading to the formation of a covalent bond<\/p>\r\n<p data-type=\"definition\"><strong><span data-type=\"term\">node\r\n<\/span><\/strong>plane separating different lobes of orbitals, where the probability of finding an electron is zero<\/p>\r\n<p data-type=\"definition\"><strong><span data-type=\"term\">pi bond (\u03c0 bond)\r\n<\/span><\/strong>covalent bond formed by side-by-side overlap of atomic orbitals; the electron density is found on opposite sides of the internuclear axis<\/p>\r\n<p data-type=\"definition\"><strong><span data-type=\"term\">sigma bond (\u03c3 bond)\r\n<\/span><\/strong>covalent bond formed by overlap of atomic orbitals along the internuclear axis<\/p>\r\n<p data-type=\"definition\"><strong><span data-type=\"term\">valence bond theory\r\n<\/span><\/strong>description of bonding that involves atomic orbitals overlapping to form \u03c3 or \u03c0 bonds, within which pairs of electrons are shared<\/p>\r\n\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<\/section><\/div>\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>Describe the formation of covalent bonds in terms of atomic orbital overlap<\/li>\n<li>Define and give examples of \u03c3 and \u03c0 bonds<\/li>\n<\/ul>\n<\/div>\n<p>As we know, a scientific theory is a strongly supported explanation for observed natural laws or large bodies of experimental data. For a theory to be accepted, it must explain experimental data and be able to predict behavior. For example, VSEPR theory has gained widespread acceptance because it predicts three-dimensional molecular shapes that are consistent with experimental data collected for thousands of different molecules. However, VSEPR theory does not provide an explanation of chemical bonding.<\/p>\n<p id=\"fs-idp41602432\">There are successful theories that describe the electronic structure of atoms. We can use quantum mechanics to predict the specific regions around an atom where electrons are likely to be located: A spherical shape for an <em data-effect=\"italics\">s<\/em> orbital, a dumbbell shape for a <em data-effect=\"italics\">p<\/em> orbital, and so forth. However, these predictions only describe the orbitals around free atoms. When atoms bond to form molecules, atomic orbitals are not sufficient to describe the regions where electrons will be located in the molecule. A more complete understanding of electron distributions requires a model that can account for the electronic structure of molecules. One popular theory holds that a covalent bond forms when a pair of electrons is shared by two atoms and is simultaneously attracted by the nuclei of both atoms. In the following sections, we will discuss how such bonds are described by valence bond theory and hybridization.<\/p>\n<p id=\"fs-idp56452464\"><strong><span data-type=\"term\">Valence bond theory<\/span><\/strong> describes a covalent bond as the overlap of half-filled atomic orbitals (each containing a single electron) that yield a pair of electrons shared between the two bonded atoms. We say that orbitals on two different atoms <strong><span data-type=\"term\">overlap<\/span><\/strong> when a portion of one orbital and a portion of a second orbital occupy the same region of space. According to valence bond theory, a covalent bond results when two conditions are met: (1) an orbital on one atom overlaps an orbital on a second atom and (2) the single electrons in each orbital combine to form an electron pair. The mutual attraction between this negatively charged electron pair and the two atoms\u2019 positively charged nuclei serves to physically link the two atoms through a force we define as a covalent bond. The strength of a covalent bond depends on the extent of overlap of the orbitals involved. Orbitals that overlap extensively form bonds that are stronger than those that have less overlap.<\/p>\n<p id=\"fs-idp45162384\">The energy of the system depends on how much the orbitals overlap. Figure 1 illustrates how the sum of the energies of two hydrogen atoms (the colored curve) changes as they approach each other. When the atoms are far apart there is no overlap, and by convention we set the sum of the energies at zero. As the atoms move together, their orbitals begin to overlap. Each electron begins to feel the attraction of the nucleus in the other atom. In addition, the electrons begin to repel each other, as do the nuclei. While the atoms are still widely separated, the attractions are slightly stronger than the repulsions, and the energy of the system decreases. (A bond begins to form.) As the atoms move closer together, the overlap increases, so the attraction of the nuclei for the electrons continues to increase (as do the repulsions among electrons and between the nuclei). At some specific distance between the atoms, which varies depending on the atoms involved, the energy reaches its lowest (most stable) value. This optimum distance between the two bonded nuclei is the bond distance between the two atoms. The bond is stable because at this point, the attractive and repulsive forces combine to create the lowest possible energy configuration. If the distance between the nuclei were to decrease further, the repulsions between nuclei and the repulsions as electrons are confined in closer proximity to each other would become stronger than the attractive forces. The energy of the system would then rise (making the system destabilized), as shown at the far left of Figure 1.<\/p>\n<div style=\"width: 890px\" class=\"wp-caption alignnone\"><img loading=\"lazy\" decoding=\"async\" class=\"\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23211750\/CNX_Chem_08_01_Morse1.jpg\" alt=\"A pair of diagrams are shown and labeled \u201ca\u201d and \u201cb\u201d. Diagram a shows three consecutive images. The first image depicts two separated blurry circles, each labeled with a positive sign and the term \u201cH atom.\u201d The phrase written under them reads, \u201cSufficiently far apart to have no interaction.\u201d The second image shows the same two circles, but this time they are much closer together and are labeled, \u201cAtoms begin to interact as they move closer together.\u201d The third image shows the two circles overlapping, labeled, \u201cH subscript 2,\u201d and, \u201cOptimum distance to achieve lowest overall energy of system.\u201d Diagram b shows a graph on which the y-axis is labeled \u201cEnergy ( J ),\u201d and the x-axis is labeled, \u201cInternuclear distance ( p m ).\u201d The midpoint of the y-axis is labeled as zero. The curve on the graph begins at zero p m and high on the y-axis. The graph slopes downward steeply to a point far below the zero joule line on the y-axis and the lowest point reads \u201c0.74 p m\u201d and \u201cH bonded to H bond length.\u201d It is also labeled \u201c negative 7.24 times 10 superscript negative 19 J.\u201d The graph then rises again to zero J. The graph is accompanied by the same images from diagram a; the first image correlates to the point in the graph where it crosses the zero point on the y-axis, the third image where the graph is lowest.\" width=\"880\" height=\"747\" data-media-type=\"image\/jpeg\" \/><\/p>\n<p class=\"wp-caption-text\">Figure 1. (a) The interaction of two hydrogen atoms changes as a function of distance. (b) The energy of the system changes as the atoms interact. The lowest (most stable) energy occurs at a distance of 74 pm, which is the bond length observed for the H<sub>2<\/sub> molecule.<\/p>\n<\/div>\n<p id=\"fs-idm5865744\">The bond energy is the difference between the energy minimum (which occurs at the bond distance) and the energy of the two separated atoms. This is the quantity of energy released when the bond is formed. Conversely, the same amount of energy is required to break the bond. For the H<sub>2<\/sub> molecule shown in Figure 1, at the bond distance of 74 pm the system is 7.24 \u00d7 10<sup>-19<\/sup> J lower in energy than the two separated hydrogen atoms. This may seem like a small number. However, we know from our earlier description of thermochemistry that bond energies are often discussed on a per-mole basis. For example, it requires 7.24 \u00d7 10<sup>-19<\/sup> J to break one H\u2013H bond, but it takes 4.36 \u00d7 10<sup>5<\/sup> J to break 1 mole of H\u2013H bonds. A comparison of some bond lengths and energies is shown in\u00a0Table 1. We can find many of these bonds in a variety of molecules, and this table provides average values. For example, breaking the first C\u2013H bond in CH<sub>4<\/sub> requires 439.3 kJ\/mol, while breaking the first C\u2013H bond in H\u2013CH<sub>2<\/sub>C<sub>6<\/sub>H<sub>5<\/sub> (a common paint thinner) requires 375.5 kJ\/mol.<\/p>\n<table id=\"fs-idp43355056\" summary=\"A table is shown that contains six columns and eleven rows. The header row reads \u201cBond,\u201d \u201cLength ( p m ),\u201d \u201cEnergy ( k J \/ mol ),\u201d \u201cBond,\u201d \u201cLength ( p m ),\u201d and \u201cEnergy ( k J \/ mol ).\u201d The first column contains the symbols: \u201cH bond H,\u201d \u201cH bond C,\u201d \u201cH bond N,\u201d \u201cH bond O,\u201d \u201cC bond C,\u201d \u201cC double bond C,\u201d \u201cC triple bond C,\u201d \u201cC bond N,\u201d \u201cC double bond N,\u201d and \u201cC triple bond N.\u201d The second column contains the numbers \u201c74,\u201d \u201c106.8,\u201d \u201c101.5,\u201d \u201c97.5,\u201d \u201c150.6,\u201d \u201c133.5,\u201d \u201c120.8,\u201d \u201c142.1,\u201d \u201c130.0,\u201d and \u201c116.1.\u201d The third column contains the numbers: \u201c436,\u201d \u201c413,\u201d \u201c391,\u201d \u201c467,\u201d \u201c347,\u201d \u201c614,\u201d \u201c839,\u201d \u201c305,\u201d \u201c615,\u201d and \u201c891.\u201d The fourth column contains the symbols: \u201cC bond O,\u201d \u201cC double bond O,\u201d \u201cC triple bond O,\u201d \u201cH bond C l,\u201d \u201cH bond B r,\u201d \u201cH bond I,\u201d \u201cO bond O,\u201d \u201cO double bond O,\u201d \u201cF bond F,\u201d and \u201cC l bond C l.\u201d The fifth column contains the numbers \u201c140.1,\u201d \u201c119.7,\u201d \u201c113.7,\u201d \u201c127.5,\u201d \u201c141.4,\u201d \u201c160.9,\u201d \u201c148,\u201d \u201c120.8,\u201d \u201c141.2,\u201d and \u201c198.8.\u201d The sixth column contains the numbers \u201c358,\u201d \u201c745,\u201d \u201c1072,\u201d \u201c431,\u201d \u201c366,\u201d \u201c298,\u201d \u201c146,\u201d \u201c498,\u201d \u201c159,\u201d and \u201c243.\u201d\">\n<caption><span data-type=\"title\">Table 1. Representative Bond Energies and Lengths<\/span><\/caption>\n<thead>\n<tr valign=\"top\">\n<th>Bond<\/th>\n<th>Length (pm)<\/th>\n<th>Energy (kJ\/mol)<\/th>\n<th>Bond<\/th>\n<th>Length (pm)<\/th>\n<th>Energy (kJ\/mol)<\/th>\n<\/tr>\n<\/thead>\n<tbody>\n<tr valign=\"top\">\n<td>H\u2013H<\/td>\n<td>74<\/td>\n<td>436<\/td>\n<td>C\u2013O<\/td>\n<td>140.1<\/td>\n<td>358<\/td>\n<\/tr>\n<tr valign=\"top\">\n<td>H\u2013C<\/td>\n<td>106.8<\/td>\n<td>413<\/td>\n<td>[latex]\\text{C}=\\text{O}[\/latex]<\/td>\n<td>119.7<\/td>\n<td>745<\/td>\n<\/tr>\n<tr valign=\"top\">\n<td>H\u2013N<\/td>\n<td>101.5<\/td>\n<td>391<\/td>\n<td>[latex]\\text{C}\\equiv \\text{O}[\/latex]<\/td>\n<td>113.7<\/td>\n<td>1072<\/td>\n<\/tr>\n<tr valign=\"top\">\n<td>H\u2013O<\/td>\n<td>97.5<\/td>\n<td>467<\/td>\n<td>H\u2013Cl<\/td>\n<td>127.5<\/td>\n<td>431<\/td>\n<\/tr>\n<tr valign=\"top\">\n<td>C\u2013C<\/td>\n<td>150.6<\/td>\n<td>347<\/td>\n<td>H\u2013Br<\/td>\n<td>141.4<\/td>\n<td>366<\/td>\n<\/tr>\n<tr valign=\"top\">\n<td>[latex]\\text{C}=\\text{C}[\/latex]<\/td>\n<td>133.5<\/td>\n<td>614<\/td>\n<td>H\u2013I<\/td>\n<td>160.9<\/td>\n<td>298<\/td>\n<\/tr>\n<tr valign=\"top\">\n<td>[latex]\\text{C}\\equiv \\text{C}[\/latex]<\/td>\n<td>120.8<\/td>\n<td>839<\/td>\n<td>O\u2013O<\/td>\n<td>148<\/td>\n<td>146<\/td>\n<\/tr>\n<tr valign=\"top\">\n<td>C\u2013N<\/td>\n<td>142.1<\/td>\n<td>305<\/td>\n<td>[latex]\\text{O}=\\text{O}[\/latex]<\/td>\n<td>120.8<\/td>\n<td>498<\/td>\n<\/tr>\n<tr valign=\"top\">\n<td>[latex]\\text{C}=\\text{N}[\/latex]<\/td>\n<td>130.0<\/td>\n<td>615<\/td>\n<td>F\u2013F<\/td>\n<td>141.2<\/td>\n<td>159<\/td>\n<\/tr>\n<tr valign=\"top\">\n<td>[latex]\\text{C}\\equiv \\text{N}[\/latex]<\/td>\n<td>116.1<\/td>\n<td>891<\/td>\n<td>Cl\u2013Cl<\/td>\n<td>198.8<\/td>\n<td>243<\/td>\n<\/tr>\n<\/tbody>\n<\/table>\n<p id=\"fs-idm10626416\">In addition to the distance between two orbitals, the orientation of orbitals also affects their overlap (other than for two <em data-effect=\"italics\">s<\/em> orbitals, which are spherically symmetric). Greater overlap is possible when orbitals are oriented such that they overlap on a direct line between the two nuclei. Figure 2 illustrates this for two <em data-effect=\"italics\">p<\/em> orbitals from different atoms; the overlap is greater when the orbitals overlap end to end rather than at an angle.<\/p>\n<div style=\"width: 660px\" class=\"wp-caption alignnone\"><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23211752\/CNX_Chem_08_01_overlap1.jpg\" alt=\"Two diagrams are shown. Diagram a contains two molecules whose p orbitals, which are depicted as two balloon-shaped structures that meet together to form a peanut shape, are laid end over end, creating an area of overlap. In diagram b, the same two molecules are shown, but this time, they are laid out in a way so as to form a near-ninety degree angle. In this diagram, the ends of two of these peanut-shaped orbitals do not overlap nearly as much.\" width=\"650\" height=\"158\" data-media-type=\"image\/jpeg\" \/><\/p>\n<p class=\"wp-caption-text\">Figure 2. (a) The overlap of two p orbitals is greatest when the orbitals are directed end to end. (b) Any other arrangement results in less overlap. The plus signs indicate the locations of the nuclei.<\/p>\n<\/div>\n<p id=\"fs-idp40817312\">The overlap of two <em data-effect=\"italics\">s<\/em> orbitals (as in H<sub>2<\/sub>), the overlap of an <em data-effect=\"italics\">s<\/em> orbital and a <em data-effect=\"italics\">p<\/em> orbital (as in HCl), and the end-to-end overlap of two <em data-effect=\"italics\">p<\/em> orbitals (as in Cl<sub>2<\/sub>) all produce <span data-type=\"term\">sigma bonds (\u03c3 bonds)<\/span>, as illustrated in Figure 3. A \u03c3 bond is a covalent bond in which the electron density is concentrated in the region along the internuclear axis; that is, a line between the nuclei would pass through the center of the overlap region. Single bonds in Lewis structures are described as \u03c3 bonds in valence bond theory.<\/p>\n<div style=\"width: 660px\" class=\"wp-caption alignnone\"><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23211753\/CNX_Chem_08_01_sigma1.jpg\" alt=\"Three diagrams are shown and labeled \u201ca,\u201d \u201cb,\u201d and \u201cc.\u201d Diagram a shows two spherical orbitals lying side by side and overlapping. Diagram b shows one spherical and one peanut-shaped orbital lying near one another so that the spherical orbital overlaps with one end of the peanut-shaped orbital. Diagram c shows two peanut-shaped orbitals lying end to end so that one end of each orbital overlaps the other.\" width=\"650\" height=\"131\" data-media-type=\"image\/jpeg\" \/><\/p>\n<p class=\"wp-caption-text\">Figure 3. Sigma (\u03c3) bonds form from the overlap of the following: (a) two s orbitals, (b) an s orbital and a p orbital, and (c) two p orbitals. The plus signs indicate the locations of the nuclei.<\/p>\n<\/div>\n<div style=\"width: 335px\" class=\"wp-caption alignright\"><img loading=\"lazy\" decoding=\"async\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23211754\/CNX_Chem_08_01_pi1.jpg\" alt=\"Two peanut-shaped orbitals are shown, lying vertically and parallel with one another. They overlap one another along the top and bottom of the orbital.\" width=\"325\" height=\"115\" data-media-type=\"image\/jpeg\" \/><\/p>\n<p class=\"wp-caption-text\">Figure 4. Pi (\u03c0) bonds form from the side-by-side overlap of two p orbitals. The plus signs indicate the location of the nuclei.<\/p>\n<\/div>\n<p id=\"fs-idp92007424\">A <strong><span data-type=\"term\">pi bond (\u03c0 bond)<\/span><\/strong> is a type of covalent bond that results from the side-by-side overlap of two <em data-effect=\"italics\">p<\/em> orbitals, as illustrated in Figure 4. In a \u03c0 bond, the regions of orbital overlap lie on opposite sides of the internuclear axis. Along the axis itself, there is a <strong><span data-type=\"term\">node<\/span><\/strong>, that is, a plane with no probability of finding an electron.<\/p>\n<p id=\"fs-idp46373312\">While all single bonds are \u03c3 bonds, multiple bonds consist of both \u03c3 and \u03c0 bonds. As the Lewis structures in suggest, O<sub>2<\/sub> contains a double bond, and N<sub>2<\/sub> contains a triple bond. The double bond consists of one \u03c3 bond and one \u03c0 bond, and the triple bond consists of one \u03c3 bond and two \u03c0 bonds. Between any two atoms, the first bond formed will always be a \u03c3 bond, but there can only be one \u03c3 bond in any one location. In any multiple bond, there will be one \u03c3 bond, and the remaining one or two bonds will be \u03c0 bonds. These bonds are described in more detail later in this chapter.<\/p>\n<p><span id=\"fs-idp200689696\" data-type=\"media\" data-alt=\"A diagram contains three Lewis structures. The left most structure shows an H atom bonded to a C l atom by a single bond. The C l atom has three lone pairs of electrons. The phrase \u201cOne sigma bond No pi bonds\u201d is written below the drawing. The center structure shows two O atoms bonded by a double bond. The O atoms each have two lone pairs of electrons. The phrase \u201cOne sigma bond One pi bond\u201d is written below the drawing. The right most structure shows two N atoms bonded by a triple bond. Each N atom has a lone pairs of electrons. The phrase \u201cOne sigma bond Two pi bonds\u201d is written below the drawing.\"><br \/>\n<img decoding=\"async\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23211755\/CNX_Chem_08_01_bondtype_img1.jpg\" alt=\"A diagram contains three Lewis structures. The left most structure shows an H atom bonded to a C l atom by a single bond. The C l atom has three lone pairs of electrons. The phrase \u201cOne sigma bond No pi bonds\u201d is written below the drawing. The center structure shows two O atoms bonded by a double bond. The O atoms each have two lone pairs of electrons. The phrase \u201cOne sigma bond One pi bond\u201d is written below the drawing. The right most structure shows two N atoms bonded by a triple bond. Each N atom has a lone pairs of electrons. The phrase \u201cOne sigma bond Two pi bonds\u201d is written below the drawing.\" data-media-type=\"image\/jpeg\" \/><\/span><\/p>\n<p id=\"fs-idp228190304\">As seen in Table 1, an average carbon-carbon single bond is 347 kJ\/mol, while in a carbon-carbon double bond, the \u03c0 bond increases the bond strength by 267 kJ\/mol. Adding an additional \u03c0 bond causes a further increase of 225 kJ\/mol. We can see a similar pattern when we compare other \u03c3 and \u03c0 bonds. Thus, each individual \u03c0 bond is generally weaker than a corresponding \u03c3 bond between the same two atoms. In a \u03c3 bond, there is a greater degree of orbital overlap than in a \u03c0 bond.<\/p>\n<div id=\"fs-idp85049088\" data-type=\"example\">\n<div class=\"textbox shaded\">\n<h3>Example 1<\/h3>\n<h4 id=\"fs-idm26209776\"><strong><span data-type=\"title\">Counting \u03c3 and \u03c0 Bonds<\/span><\/strong><br \/>\n<span id=\"fs-idp98308768\" data-type=\"media\" data-alt=\"This figure shows a molecule composed of four carbon atoms. There is a double bond between carbons one and two and three and four, while a single bond holds carbon two and three together. Carbons one and four are also bonded to two hydrogens with a single bond while carbons two and three are each bonded to one hydrogen each by a single bond.\"><br \/>\n<img decoding=\"async\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23211757\/CNX_Chem_08_01_C4H6_img1.jpg\" alt=\"This figure shows a molecule composed of four carbon atoms. There is a double bond between carbons one and two and three and four, while a single bond holds carbon two and three together. Carbons one and four are also bonded to two hydrogens with a single bond while carbons two and three are each bonded to one hydrogen each by a single bond.\" data-media-type=\"image\/jpeg\" \/><\/span><\/h4>\n<p id=\"fs-idm57104384\">Butadiene, C<sub>6<\/sub>H<sub>6<\/sub>, is used to make synthetic rubber. Identify the number of \u03c3 and \u03c0 bonds contained in this molecule.<\/p>\n<h3 id=\"fs-idm9825456\"><span data-type=\"title\">Solution<\/span><\/h3>\n<p>There are six \u03c3 C\u2013H bonds and one \u03c3 C\u2013C bond, for a total of seven from the single bonds. There are two double bonds that each have a \u03c0 bond in addition to the \u03c3 bond. This gives a total nine \u03c3 and two \u03c0 bonds overall.<\/p>\n<h4 id=\"fs-idp46378096\"><strong><span data-type=\"title\">Check Your Learning<\/span><\/strong><\/h4>\n<p>Identify each illustration as depicting a \u03c3 or \u03c0 bond:<\/p>\n<p id=\"fs-idp138827456\">(a) side-by-side overlap of a 4<em data-effect=\"italics\">p<\/em> and a 2<em data-effect=\"italics\">p<\/em> orbital<\/p>\n<p id=\"fs-idp7133408\">(b) end-to-end overlap of a 4<em data-effect=\"italics\">p<\/em> and 4<em data-effect=\"italics\">p<\/em> orbital<\/p>\n<p id=\"fs-idp145728832\">(c) end-to-end overlap of a 4<em data-effect=\"italics\">p<\/em> and a 2<em data-effect=\"italics\">p<\/em> orbital<\/p>\n<p><span id=\"fs-idm5769888\" data-type=\"media\" data-alt=\"Three diagrams are shown and labeled \u201ca,\u201d \u201cb,\u201d and \u201cc.\u201d Diagram a depicts two peanut-shaped orbitals lying vertically side-by-side and overlapping. One orbital is smaller than the other. Diagram b shows two peanut-shaped orbitals lying end-to-end and overlapping. Diagram c shows two unequally sized peanut-shaped orbitals lying end-to-end and overlapping.\"><br \/>\n<img loading=\"lazy\" decoding=\"async\" class=\"\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23211758\/CNX_Chem_08_01_Exover_img1.jpg\" alt=\"Three diagrams are shown and labeled \u201ca,\u201d \u201cb,\u201d and \u201cc.\u201d Diagram a depicts two peanut-shaped orbitals lying vertically side-by-side and overlapping. One orbital is smaller than the other. Diagram b shows two peanut-shaped orbitals lying end-to-end and overlapping. Diagram c shows two unequally sized peanut-shaped orbitals lying end-to-end and overlapping.\" width=\"873\" height=\"137\" data-media-type=\"image\/jpeg\" \/><\/span><\/p>\n<div id=\"fs-idp42592272\" data-type=\"note\">\n<div data-type=\"title\"><\/div>\n<div style=\"text-align: right;\" data-type=\"title\"><strong>Answer<\/strong>:\u00a0(a) is a \u03c0 bond with a node along the axis connecting the nuclei while (b) and (c) are \u03c3 bonds that overlap along the axis.<\/div>\n<\/div>\n<\/div>\n<section id=\"fs-idm40478192\" class=\"summary\" data-depth=\"1\">\n<div id=\"post-355\" class=\"post-355 chapter type-chapter status-publish hentry type-1\">\n<div class=\"entry-content\">\n<div class=\"im_section\">\n<div class=\"im_section\">\n<div id=\"mclean-ch03_s01_s02_n01\" class=\"im_key_takeaways im_editable im_block\">\n<div class=\"bcc-box bcc-success\">\n<h2>Key Concepts and Summary<\/h2>\n<section>\n<div data-type=\"note\">\n<p id=\"fs-idp236281408\">Valence bond theory describes bonding as a consequence of the overlap of two separate atomic orbitals on different atoms that creates a region with one pair of electrons shared between the two atoms. When the orbitals overlap along an axis containing the nuclei, they form a \u03c3 bond. When they overlap in a fashion that creates a node along this axis, they form a \u03c0 bond.<\/p>\n<\/div>\n<\/section>\n<\/div>\n<div class=\"bcc-box bcc-info\">\n<h3>Chemistry End of Chapter Exercises<\/h3>\n<section id=\"fs-idp52122896\" class=\"exercises\" data-depth=\"1\">\n<div id=\"fs-idp92773152\" data-type=\"exercise\">\n<div id=\"fs-idp39612288\" data-type=\"problem\">\n<ol>\n<li id=\"fs-idp78431584\">Explain how \u03c3 and \u03c0 bonds are similar and how they are different.<\/li>\n<li>Draw a curve that describes the energy of a system with H and Cl atoms at varying distances. Then, find the minimum energy of this curve two ways.\n<ol>\n<li>Use the bond energy found in Table 8.1\u00a0to calculate the energy for one single HCl bond (Hint: How many bonds are in a mole?)<\/li>\n<li>Use the enthalpy of reaction and the bond energies for H<sub>2<\/sub> and Cl<sub>2<\/sub> to solve for the energy of one mole of HCl bonds. [latex]{\\text{H}}_{2}\\left(g\\right)+{\\text{Cl}}_{2}\\left(g\\right)\\rightleftharpoons2\\text{HCl}\\left(g\\right)\\Delta{H}_{\\text{rxn}}^{\\textdegree }=\\text{-184.7 kJ\/mol}[\/latex]<\/li>\n<\/ol>\n<\/li>\n<li>Explain why bonds occur at specific average bond distances instead of the atoms approaching each other infinitely close.<\/li>\n<li>Use valence bond theory to explain the bonding in F<sub>2<\/sub>, HF, and ClBr. Sketch the overlap of the atomic orbitals involved in the bonds.<\/li>\n<li>Use valence bond theory to explain the bonding in O<sub>2<\/sub>. Sketch the overlap of the atomic orbitals involved in the bonds in O<sub>2<\/sub>.<\/li>\n<li>How many \u03c3 and \u03c0 bonds are present in the molecule HCN?<\/li>\n<li>A friend tells you N<sub>2<\/sub> has three \u03c0 bonds due to overlap of the three <em data-effect=\"italics\">p<\/em>-orbitals on each N atom. Do you agree?<\/li>\n<li>Draw the Lewis structures for CO<sub>2<\/sub> and CO, and predict the number of \u03c3 and \u03c0 bonds for each molecule. (a) CO<sub>2\u00a0<\/sub>(b) CO<\/li>\n<\/ol>\n<\/div>\n<\/div>\n<\/section>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<div id=\"post-355\" class=\"post-355 chapter type-chapter status-publish hentry type-1\">\n<div class=\"entry-content\">\n<div class=\"im_section\">\n<div class=\"im_section\">\n<div id=\"mclean-ch03_s01_s02_n01\" class=\"im_key_takeaways im_editable im_block\">\n<div class=\"bcc-box bcc-info\">\n<h4>Selected Answers<\/h4>\n<p>1.\u00a0Similarities: Both types of bonds result from overlap of atomic orbitals on adjacent atoms and contain a maximum of two electrons. Differences: \u03c3 bonds are stronger and result from end-to-end overlap and all single bonds are \u03c3 bonds; \u03c0 bonds between the same two atoms are weaker because they result from side-by-side overlap, and multiple bonds contain one or more \u03c0 bonds (in addition to a \u03c3 bond).<\/p>\n<div data-type=\"newline\">3.\u00a0The specific average bond distance is the distance with the lowest energy. At distances less than the bond distance, the positive charges on the two nuclei repel each other, and the overall energy increases.<\/div>\n<div data-type=\"newline\"><\/div>\n<div data-type=\"newline\">5. Bonding: One \u03c3 bond and one \u03c0 bond. The <em data-effect=\"italics\">s<\/em> orbitals are filled and do not overlap. The <em data-effect=\"italics\">p<\/em> orbitals overlap along the axis to form a \u03c3 bond and side by side to form the \u03c0 bond<br \/>\n.<img loading=\"lazy\" decoding=\"async\" class=\"\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23211801\/CNX_Chem_08_01_O2bonds_img1.jpg\" alt=\"This figure shows the orbitals of two atoms being added together as they form bonds. The two atoms are shown separately on the right, each having two peanut-shaped orbitals lying perpendicularly to one another. A right-facing arrow shows that the two have moved closer together and now the upper and lower portions of the vertical peanut-shaped orbitals are shown as merging together above and below the plane of the molecule while the horizontal peanut-shaped orbitals are overlapping between the two nuclei.\" width=\"846\" height=\"222\" data-media-type=\"image\/jpeg\" \/><\/p>\n<div data-type=\"newline\">7. No, two of the <em data-effect=\"italics\">p<\/em> orbitals (one on each N) will be oriented end to end and will form a \u03c3 bond.<\/div>\n<div data-type=\"newline\">\n<div id=\"fs-idp108425904\" data-type=\"exercise\">\n<div id=\"fs-idp75120720\" data-type=\"solution\">\n<div data-type=\"newline\">\u00a0<img loading=\"lazy\" decoding=\"async\" class=\"\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23211803\/CNX_Chem_08_01_N2LewStru_img1.jpg\" alt=\"Two nitrogen atoms are shown both in a Lewis structure and as a diagram. The Lewis structure depicts the two nitrogen atoms bonded by a triple bond. The diagram shows two nitrogen atoms and their three peanut-shaped p-orbitals. One of the orbitals lies horizontally and overlaps between the two nuclei. It is labeled, \u201csigma bond.\u201d The other two lie vertically and in the z-plane of the page. They overlap above and below and into and out of the page in relation to the nuclei. They are each labeled, \u201cpi bond.\u201d\" width=\"528\" height=\"208\" data-media-type=\"image\/jpeg\" \/><\/div>\n<div data-type=\"newline\"><\/div>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n<\/section>\n<div class=\"bcc-box bcc-success\">\n<section id=\"glossary\">\n<h3>Glossary<\/h3>\n<div data-type=\"definition\">\n<div id=\"fs-idm8143856\" data-type=\"definition\">\n<div data-type=\"glossary\">\n<p data-type=\"definition\"><strong><span data-type=\"term\">overlap<br \/>\n<\/span><\/strong>coexistence of orbitals from two different atoms sharing the same region of space, leading to the formation of a covalent bond<\/p>\n<p data-type=\"definition\"><strong><span data-type=\"term\">node<br \/>\n<\/span><\/strong>plane separating different lobes of orbitals, where the probability of finding an electron is zero<\/p>\n<p data-type=\"definition\"><strong><span data-type=\"term\">pi bond (\u03c0 bond)<br \/>\n<\/span><\/strong>covalent bond formed by side-by-side overlap of atomic orbitals; the electron density is found on opposite sides of the internuclear axis<\/p>\n<p data-type=\"definition\"><strong><span data-type=\"term\">sigma bond (\u03c3 bond)<br \/>\n<\/span><\/strong>covalent bond formed by overlap of atomic orbitals along the internuclear axis<\/p>\n<p data-type=\"definition\"><strong><span data-type=\"term\">valence bond theory<br \/>\n<\/span><\/strong>description of bonding that involves atomic orbitals overlapping to form \u03c3 or \u03c0 bonds, within which pairs of electrons are shared<\/p>\n<\/div>\n<\/div>\n<\/div>\n<\/section>\n<\/div>\n<\/div>\n\n\t\t\t <section class=\"citations-section\" role=\"contentinfo\">\n\t\t\t <h3>Candela Citations<\/h3>\n\t\t\t\t\t <div>\n\t\t\t\t\t\t <div id=\"citation-list-1978\">\n\t\t\t\t\t\t\t <div class=\"licensing\"><div class=\"license-attribution-dropdown-subheading\">CC licensed content, Shared previously<\/div><ul class=\"citation-list\"><li>Chemistry. <strong>Provided by<\/strong>: OpenStax College. <strong>Located at<\/strong>: <a target=\"_blank\" href=\"http:\/\/openstaxcollege.org\">http:\/\/openstaxcollege.org<\/a>. <strong>License<\/strong>: <em><a target=\"_blank\" rel=\"license\" href=\"https:\/\/creativecommons.org\/licenses\/by\/4.0\/\">CC BY: Attribution<\/a><\/em>. <strong>License Terms<\/strong>: Download for free at https:\/\/openstaxcollege.org\/textbooks\/chemistry\/get<\/li><\/ul><\/div>\n\t\t\t\t\t\t <\/div>\n\t\t\t\t\t <\/div>\n\t\t\t <\/section>","protected":false},"author":5,"menu_order":2,"template":"","meta":{"_candela_citation":"[{\"type\":\"cc\",\"description\":\"Chemistry\",\"author\":\"\",\"organization\":\"OpenStax College\",\"url\":\"http:\/\/openstaxcollege.org\",\"project\":\"\",\"license\":\"cc-by\",\"license_terms\":\"Download for free at https:\/\/openstaxcollege.org\/textbooks\/chemistry\/get\"}]","CANDELA_OUTCOMES_GUID":"","pb_show_title":"on","pb_short_title":"","pb_subtitle":"","pb_authors":[],"pb_section_license":""},"chapter-type":[],"contributor":[],"license":[],"class_list":["post-1978","chapter","type-chapter","status-publish","hentry"],"part":3005,"_links":{"self":[{"href":"https:\/\/courses.lumenlearning.com\/suny-buffstate-chemistryformajorsxmaster\/wp-json\/pressbooks\/v2\/chapters\/1978","targetHints":{"allow":["GET"]}}],"collection":[{"href":"https:\/\/courses.lumenlearning.com\/suny-buffstate-chemistryformajorsxmaster\/wp-json\/pressbooks\/v2\/chapters"}],"about":[{"href":"https:\/\/courses.lumenlearning.com\/suny-buffstate-chemistryformajorsxmaster\/wp-json\/wp\/v2\/types\/chapter"}],"author":[{"embeddable":true,"href":"https:\/\/courses.lumenlearning.com\/suny-buffstate-chemistryformajorsxmaster\/wp-json\/wp\/v2\/users\/5"}],"version-history":[{"count":11,"href":"https:\/\/courses.lumenlearning.com\/suny-buffstate-chemistryformajorsxmaster\/wp-json\/pressbooks\/v2\/chapters\/1978\/revisions"}],"predecessor-version":[{"id":5180,"href":"https:\/\/courses.lumenlearning.com\/suny-buffstate-chemistryformajorsxmaster\/wp-json\/pressbooks\/v2\/chapters\/1978\/revisions\/5180"}],"part":[{"href":"https:\/\/courses.lumenlearning.com\/suny-buffstate-chemistryformajorsxmaster\/wp-json\/pressbooks\/v2\/parts\/3005"}],"metadata":[{"href":"https:\/\/courses.lumenlearning.com\/suny-buffstate-chemistryformajorsxmaster\/wp-json\/pressbooks\/v2\/chapters\/1978\/metadata\/"}],"wp:attachment":[{"href":"https:\/\/courses.lumenlearning.com\/suny-buffstate-chemistryformajorsxmaster\/wp-json\/wp\/v2\/media?parent=1978"}],"wp:term":[{"taxonomy":"chapter-type","embeddable":true,"href":"https:\/\/courses.lumenlearning.com\/suny-buffstate-chemistryformajorsxmaster\/wp-json\/pressbooks\/v2\/chapter-type?post=1978"},{"taxonomy":"contributor","embeddable":true,"href":"https:\/\/courses.lumenlearning.com\/suny-buffstate-chemistryformajorsxmaster\/wp-json\/wp\/v2\/contributor?post=1978"},{"taxonomy":"license","embeddable":true,"href":"https:\/\/courses.lumenlearning.com\/suny-buffstate-chemistryformajorsxmaster\/wp-json\/wp\/v2\/license?post=1978"}],"curies":[{"name":"wp","href":"https:\/\/api.w.org\/{rel}","templated":true}]}}