{"id":1682,"date":"2019-01-02T22:09:12","date_gmt":"2019-01-02T22:09:12","guid":{"rendered":"https:\/\/courses.lumenlearning.com\/suny-mcc-introductorychemistry\/?post_type=chapter&#038;p=1682"},"modified":"2019-01-04T03:35:24","modified_gmt":"2019-01-04T03:35:24","slug":"polarity-of-molecules","status":"publish","type":"chapter","link":"https:\/\/courses.lumenlearning.com\/suny-mcc-introductorychemistry\/chapter\/polarity-of-molecules\/","title":{"raw":"6.4 Polarity of Molecules","rendered":"6.4 Polarity of Molecules"},"content":{"raw":"<div class=\"textbox learning-objectives\">\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<\/li>\r\n \t<li>Define electronegativity and assess the polarity of covalent bonds<\/li>\r\n<\/ul>\r\n<\/div>\r\n<h2>Lewis Structures<\/h2>\r\nWe also use Lewis symbols to indicate the formation of covalent bonds, which are shown in<strong> Lewis structures<\/strong>, drawings that describe the bonding in molecules and polyatomic ions. For example, when two chlorine atoms form a chlorine molecule, they share one pair of electrons:\r\n\r\n<img class=\"aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23211358\/CNX_Chem_07_03_C12dot_img1.jpg\" alt=\"A Lewis dot diagram shows a reaction. Two chlorine symbols, each surrounded by seven dots are separated by a plus sign. The dots on the first atom are all black and the dots on the second atom are all read. The phrase, \u201cChlorine atoms\u201d is written below. A right-facing arrow points to two chlorine symbols, each with six dots surrounding their outer edges and a shared pair of dots in between. One of the shared dots is black and one is red. The phrase, \u201cChlorine molecule\u201d is written below.\" \/>\r\n\r\nThe Lewis structure indicates that each Cl atom has three pairs of electrons that are not used in bonding (called <strong>lone pairs<\/strong>) and one shared pair of electrons (written between the atoms). A dash (or line) is sometimes used to indicate a shared pair of electrons:\r\n\r\nA single shared pair of electrons is called a <strong>single bond<\/strong>. Each Cl atom interacts with eight valence electrons: the six in the lone pairs and the two in the single bond.\r\n<h2>Pure vs. Polar Covalent Bonds<\/h2>\r\nIf the atoms that form a covalent bond are identical, as in H<sub>2<\/sub>, Cl<sub>2<\/sub>, and other diatomic molecules, then the electrons in the bond must be shared equally. We refer to this as a <strong>pure covalent bond<\/strong>. Electrons shared in pure covalent bonds have an equal probability of being near each nucleus.\r\n\r\n<img class=\"aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23211400\/CNX_Chem_07_03_Cl2dash_img1.jpg\" alt=\"Two Lewis structures are shown. The left-hand structure shows two H atoms connected by a single bond. The right-hand structure shows two C l atoms connected by a single bond and each surrounded by six dots.\" \/>\r\n\r\nIn the case of Cl<sub>2<\/sub>, each atom starts off with seven valence electrons, and each Cl shares one electron with the other, forming one covalent bond:\r\n\r\n[latex]\\text{Cl}+\\text{Cl}\\rightarrow{\\text{Cl}}_{2}[\/latex]\r\n\r\nThe total number of electrons around each individual atom consists of six nonbonding electrons and two shared (i.e., bonding) electrons for eight total electrons, matching the number of valence electrons in the noble gas argon. Since the bonding atoms are identical, Cl<sub>2<\/sub> also features a pure covalent bond.\r\n\r\n[caption id=\"\" align=\"alignright\" width=\"300\"]<img src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23211346\/CNX_Chem_07_02_HClBond1.jpg\" alt=\"Two diagrams are shown and labeled \u201ca\u201d and \u201cb.\u201d Diagram a shows a small sphere labeled, \u201cH\u201d and a larger sphere labeled, \u201cC l\u201d that overlap slightly. Both spheres have a small dot in the center. Diagram b shows an H bonded to a C l with a single bond. A dipole and a positive sign are written above the H and a dipole and negative sign are written above the C l. An arrow points toward the C l with a plus sign on the end furthest from the arrow\u2019s head near the H.\" width=\"300\" height=\"206\" \/> Figure 2. (a) The distribution of electron density in the HCl molecule is uneven. The electron density is greater around the chlorine nucleus. The small, black dots indicate the location of the hydrogen and chlorine nuclei in the molecule. (b) Symbols \u03b4+ and \u03b4\u2013 indicate the polarity of the H\u2013Cl bond.[\/caption]\r\n\r\nWhen the atoms linked by a covalent bond are different, the bonding electrons are shared, but no longer equally. Instead, the bonding electrons are more attracted to one atom than the other, giving rise to a shift of electron density toward that atom. This unequal distribution of electrons is known as a <strong>polar covalent bond<\/strong>, characterized by a partial positive charge on one atom and a partial negative charge on the other. The atom that attracts the electrons more strongly acquires the partial negative charge and vice versa. For example, the electrons in the H\u2013Cl bond of a hydrogen chloride molecule spend more time near the chlorine atom than near the hydrogen atom. Thus, in an HCl molecule, the chlorine atom carries a partial negative charge and the hydrogen atom has a partial positive charge.\u00a0Figure 2 shows the distribution of electrons in the H\u2013Cl bond. Note that the shaded area around Cl is much larger than it is around H. Compare this to Figure 1, which shows the even distribution of electrons in the H<sub>2<\/sub> nonpolar bond.\r\n\r\nWe sometimes designate the positive and negative atoms in a polar covalent bond using a lowercase Greek letter \u201cdelta,\u201d \u03b4, with a plus sign or minus sign to indicate whether the atom has a partial positive charge (\u03b4+) or a partial negative charge (\u03b4\u2013). This symbolism is shown for the H\u2013Cl molecule in Figure 2.\r\n<h2>Electronegativity<\/h2>\r\nWhether a bond is nonpolar or polar covalent is determined by a property of the bonding atoms called <strong>electronegativity<\/strong>. Electronegativity is a measure of the tendency of an atom to attract electrons (or electron density) towards itself. It determines how the shared electrons are distributed between the two atoms in a bond. The more strongly an atom attracts the electrons in its bonds, the larger its electronegativity. Electrons in a polar covalent bond are shifted toward the more electronegative atom; thus, the more electronegative atom is the one with the partial negative charge. The greater the difference in electronegativity, the more polarized the electron distribution and the larger the partial charges of the atoms.\r\n\r\nFigure 3 shows the electronegativity values of the elements as proposed by one of the most famous chemists of the twentieth century: Linus Pauling (Figure 4). In general, electronegativity increases from left to right across a period in the periodic table and decreases down a group. Thus, the nonmetals, which lie in the upper right, tend to have the highest electronegativities, with fluorine the most electronegative element of all (EN = 4.0). Metals tend to be less electronegative elements, and the group 1 metals have the lowest electronegativities. Note that noble gases are excluded from this figure because these atoms usually do not share electrons with others atoms since they have a full valence shell. (While noble gas compounds such as XeO<sub>2<\/sub> do exist, they can only be formed under extreme conditions, and thus they do not fit neatly into the general model of electronegativity.)\r\n\r\n[caption id=\"\" align=\"alignnone\" width=\"880\"]<img src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23211347\/CNX_Chem_07_02_ENTable1.jpg\" alt=\"Part of the periodic table is shown. A downward-facing arrow is drawn to the left of the table and labeled, \u201cDecreasing electronegativity,\u201d while a right-facing arrow is drawn above the table and labeled \u201cIncreasing electronegativity.\u201d The electronegativity for almost all the elements is given.\" width=\"880\" height=\"377\" \/> Figure 3. The electronegativity values derived by Pauling follow predictable periodic trends with the higher electronegativities toward the upper right of the periodic table.[\/caption]\r\n\r\n<div class=\"textbox shaded\">\r\n<h3>Portrait of a Chemist: Linus Pauling<\/h3>\r\nLinus Pauling, shown in Figure 4, is the only person to have received two unshared (individual) Nobel Prizes: one for chemistry in 1954 for his work on the nature of chemical bonds and one for peace in 1962 for his opposition to weapons of mass destruction. He developed many of the theories and concepts that are foundational to our current understanding of chemistry, including electronegativity and resonance structures.\r\n\r\n[caption id=\"\" align=\"aligncenter\" width=\"325\"]<img src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23211349\/CNX_Chem_07_02_Pauling1.jpg\" alt=\"A photograph of Linus Pauling is shown.\" width=\"325\" height=\"394\" \/> Figure 4. Linus Pauling (1901\u20131994) made many important contributions to the field of chemistry. He was also a prominent activist, publicizing issues related to health and nuclear weapons.[\/caption]\r\n\r\nPauling also contributed to many other fields besides chemistry. His research on sickle cell anemia revealed the cause of the disease\u2014the presence of a genetically inherited abnormal protein in the blood\u2014and paved the way for the field of molecular genetics. His work was also pivotal in curbing the testing of nuclear weapons; he proved that radioactive fallout from nuclear testing posed a public health risk.\r\n\r\n<\/div>\r\n<h3>Electronegativity and Bond Type<\/h3>\r\nThe absolute value of the difference in electronegativity (\u0394EN) of two bonded atoms provides a rough measure of the polarity to be expected in the bond and, thus, the bond type. When the difference is very small or zero, the bond is covalent and nonpolar. When it is large, the bond is polar covalent or ionic. The absolute values of the electronegativity differences between the atoms in the bonds H\u2013H, H\u2013Cl, and Na\u2013Cl are 0 (nonpolar), 0.9 (polar covalent), and 2.1 (ionic), respectively. The degree to which electrons are shared between atoms varies from completely equal (pure covalent bonding) to not at all (ionic bonding). Figure 5 shows the relationship between electronegativity difference and bond type.\r\n\r\n[caption id=\"attachment_1667\" align=\"aligncenter\" width=\"627\"]<img class=\"wp-image-1667 size-full\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/2835\/2018\/08\/29173516\/Polar-bonds.jpg\" alt=\"Two flow charts and table are shown. The first flow chart is labeled, \u201cElectronegativity difference between bonding atoms.\u201d Below this label are three rounded text bubbles, connected by a downward-facing arrow, labeled, \u201cZero,\u201d \u201cIntermediate,\u201d and \u201cLarge,\u201d respectively. The second flow chart is labeled, \u201cBond type.\u201d Below this label are three rounded text bubbles, connected by a downward-facing arrow, labeled, \u201cPure covalent,\u201d \u201cPolar covalent,\u201d and \u201cIonic,\u201d respectively. A double ended arrow is written vertically to the right of the flow charts and labeled, \u201cCovalent character decreases; ionic character increases.\u201d The table is made up of two columns and four rows. The header line is labeled \u201cBond type\u201d and \u201cElectronegativity difference.\u201d The left column contains the phrases \u201cPure covalent,\u201d \u201cPolar covalent,\u201d and \u201cIonic,\u201d while the right column contains the values \u201cless than 0.4,\u201d \u201cbetween 0.4 and 1.8,\u201d and \u201cgreater than 1.8.\u201d\" width=\"627\" height=\"312\" \/> Figure 5. As the electronegativity difference increases between two atoms, the bond becomes more ionic.[\/caption]\r\n\r\nA rough approximation of the electronegativity differences associated with covalent, polar covalent, and ionic bonds is shown in Figure 5. This table is just a general guide, however, with many exceptions. For example, the H and F atoms in HF have an electronegativity difference of 1.9, and the N and H atoms in NH<sub>3<\/sub> a difference of 0.9, yet both of these compounds form bonds that are considered polar covalent. Likewise, the Na and Cl atoms in NaCl have an electronegativity difference of 2.1, and the Mn and I atoms in MnI<sub>2<\/sub> have a difference of 1.0, yet both of these substances form ionic compounds.\r\n\r\nThe best guide to the covalent or ionic character of a bond is to consider the types of atoms involved and their relative positions in the periodic table. Bonds between two nonmetals are generally covalent; bonding between a metal and a nonmetal is often ionic.\r\n\r\nSome compounds contain both covalent and ionic bonds. The atoms in polyatomic ions, such as OH<sup>\u2013<\/sup>, [latex]{\\text{NO}}_{3}^{-},[\/latex] and [latex]{\\text{NH}}_{4}^{\\text{+}},[\/latex] are held together by polar covalent bonds. However, these polyatomic ions form ionic compounds by combining with ions of opposite charge. For example, potassium nitrate, KNO<sub>3<\/sub>, contains the K<sup>+<\/sup> cation and the polyatomic [latex]{\\text{NO}}_{3}^{-}[\/latex] anion. Thus, bonding in potassium nitrate is ionic, resulting from the electrostatic attraction between the ions K<sup>+<\/sup> and [latex]{\\text{NO}}_{3}^{-},[\/latex] as well as covalent between the nitrogen and oxygen atoms in [latex]{\\text{NO}}_{3}^{-}[\/latex].\r\n<div class=\"textbox examples\">\r\n<h3>Example 1:\u00a0Electronegativity and Bond Polarity<\/h3>\r\nBond polarities play an important role in determining the structure of proteins. Using the electronegativity values in Figure 7.6, arrange the following covalent bonds\u2014all commonly found in amino acids\u2014in order of increasing polarity. Then designate the positive and negative atoms using the symbols \u03b4+ and \u03b4\u2013:\r\n\r\nC\u2013H, C\u2013N, C\u2013O, N\u2013H, O\u2013H, S\u2013H\r\n\r\n[reveal-answer q=\"862588\"]Show Answer[\/reveal-answer]\r\n[hidden-answer a=\"862588\"]\r\n\r\nThe polarity of these bonds increases as the absolute value of the electronegativity difference increases. The atom with the \u03b4\u2013 designation is the more electronegative of the two. Table 1\u00a0shows these bonds in order of increasing polarity.\r\n<table summary=\"This table has three columns and seven rows. The first row is a header row and it labels each column. The first column header is, \u201cBond,\u201d the second is, \u201ccapital delta E N,\u201d and the third is, \u201cPolarity.\u201d Under the \u201cBond\u201d column are the following: C bonds to H with a single bond; S bonds to H with a single bond; C bonds to N with a single bond; N bonds to H with a single bond; C bonds to O with a single bond; and O bonds to H with a single bond. Under the \u201ccapital delta E N\u201d columna are the values: 0.4; 0.4; 0.5; 0.9; 1.0; and 1.4. Under the \u201cPolarity\u201d column are the follwoing: C bonds to H with a single bond, there is a lowercase delta negative sign above the C and a lowercase delta positive sign over H; S bonds to H with a single bond, there is a lowercase delta negative sign over the S and a lowercase delta positive sign over the H; C bonds to N with a single bond, there is a lowercase delta positive sign over the C and a lowercase delta negative sign over the N; N bonds to H with a single bond, there is a lowercase delta negative sign over the N and a lowercase delta positive sign over the H; C bonds to O with a single bond, there is a lowercase delta positive sign over C and a lowercase delta negative sign over the O; and O bonds to H with a single bond, there is a lowercase delta negative sign over the O and a lowercase delta positive sign over the H.\">\r\n<thead>\r\n<tr>\r\n<th colspan=\"3\">Table 1. Bond Polarity and Electronegativity Difference<\/th>\r\n<\/tr>\r\n<\/thead>\r\n<tbody>\r\n<tr>\r\n<th>Bond<\/th>\r\n<th>\u0394EN<\/th>\r\n<th>Polarity<\/th>\r\n<\/tr>\r\n<tr>\r\n<td>C\u2013H<\/td>\r\n<td>0.4<\/td>\r\n<td>[latex]\\stackrel{\\delta -}{\\text{C}}-\\stackrel{\\delta \\text{+}}{\\text{H}}[\/latex]<\/td>\r\n<\/tr>\r\n<tr>\r\n<td>S\u2013H<\/td>\r\n<td>0.4<\/td>\r\n<td>[latex]\\stackrel{\\delta -}{\\text{S}}-\\stackrel{\\delta \\text{+}}{\\text{H}}[\/latex]<\/td>\r\n<\/tr>\r\n<tr>\r\n<td>C\u2013N<\/td>\r\n<td>0.5<\/td>\r\n<td>[latex]\\stackrel{\\delta \\text{+}}{\\text{C}}-\\stackrel{\\delta -}{\\text{N}}[\/latex]<\/td>\r\n<\/tr>\r\n<tr>\r\n<td>N\u2013H<\/td>\r\n<td>0.9<\/td>\r\n<td>[latex]\\stackrel{\\delta -}{\\text{N}}-\\stackrel{\\delta \\text{+}}{\\text{H}}[\/latex]<\/td>\r\n<\/tr>\r\n<tr>\r\n<td>C\u2013O<\/td>\r\n<td>1.0<\/td>\r\n<td>[latex]\\stackrel{\\delta \\text{+}}{\\text{C}}-\\stackrel{\\delta -}{\\text{O}}[\/latex]<\/td>\r\n<\/tr>\r\n<tr>\r\n<td>O\u2013H<\/td>\r\n<td>1.4<\/td>\r\n<td>[latex]\\stackrel{\\delta -}{\\text{O}}-\\stackrel{\\delta \\text{+}}{\\text{H}}[\/latex]<\/td>\r\n<\/tr>\r\n<\/tbody>\r\n<\/table>\r\n[\/hidden-answer]\r\n\r\n<strong>Check Your Learning<\/strong>\r\n\r\nSilicones are polymeric compounds containing, among others, the following types of covalent bonds: Si\u2013O, Si\u2013C, C\u2013H, and C\u2013C. Using the electronegativity values in Figure 3, arrange the bonds in order of increasing polarity and designate the positive and negative atoms using the symbols \u03b4+ and \u03b4\u2013.\r\n[reveal-answer q=\"113211\"]Show Answer[\/reveal-answer]\r\n[hidden-answer a=\"113211\"]\r\n<table summary=\"This table has three columns and five rows. The first row is a header row and it labels each column. The first column header is, \u201cBond,\u201d the second column header is, \u201cElectronegativity Difference,\u201d and the third column header is, \u201cPolarity.\u201d Under the column \u201cBond\u201d are the following: C bonds to C with a single bond; C bonds to H with a single bond; S i bonds to C with a single bond; and S i bonds to O with a single bond. Under the column \u201cElectronegativity Difference\u201d are the values: 0.0; 0.4; 0.7; and 1.7. Under the column \u201cPolarity\u201d are the following: nonpolar, C bonds to H with a single bond, there is a lowercase delta negative sign over C and a lowercase delta positive sign over H; S i bonds to C with a single bond, there is a lowercase delta positive sign over S i and a lowercase delta negative sign over C; and S i bonds to O with a single bond, there is a lowercase delta positive sign over S i and a lowercase delta negative sign over O.\">\r\n<thead>\r\n<tr>\r\n<th>Bond<\/th>\r\n<th>Electronegativity Difference<\/th>\r\n<th>Polarity<\/th>\r\n<\/tr>\r\n<\/thead>\r\n<tbody>\r\n<tr>\r\n<td>C\u2013C<\/td>\r\n<td>0.0<\/td>\r\n<td>nonpolar<\/td>\r\n<\/tr>\r\n<tr>\r\n<td>C\u2013H<\/td>\r\n<td>0.4<\/td>\r\n<td>[latex]\\stackrel{\\delta -}{\\text{C}}-\\stackrel{\\delta \\text{+}}{\\text{H}}[\/latex]<\/td>\r\n<\/tr>\r\n<tr>\r\n<td>Si\u2013C<\/td>\r\n<td>0.7<\/td>\r\n<td>[latex]\\stackrel{\\delta \\text{+}}{\\text{Si}}-\\stackrel{\\delta -}{\\text{C}}[\/latex]<\/td>\r\n<\/tr>\r\n<tr>\r\n<td>Si\u2013O<\/td>\r\n<td>1.7<\/td>\r\n<td>[latex]\\stackrel{\\delta \\text{+}}{\\text{Si}}-\\stackrel{\\delta -}{\\text{O}}[\/latex]<\/td>\r\n<\/tr>\r\n<\/tbody>\r\n<\/table>\r\n[\/hidden-answer]\r\n\r\n<\/div>\r\n<h2>Molecular Polarity and Dipole Moment<\/h2>\r\nAs discussed previously, polar covalent bonds connect two atoms with differing electronegativities, leaving one atom with a partial positive charge (\u0394+) and the other atom with a partial negative charge (\u0394\u2013), as the electrons are pulled toward the more electronegative atom. This separation of charge gives rise to a bond dipole moment. The magnitude of a <strong>bond dipole moment<\/strong> is represented by the Greek letter mu (<em>\u00b5<\/em>) and is given by the formula shown below, where Q is the magnitude of the partial charges (determined by the electronegativity difference) and r is the distance between the charges:\r\n\r\n[latex]\\mu =\\text{Qr}[\/latex]\r\n\r\nThis bond moment can be represented as a <strong>vector<\/strong>, a quantity having both direction and magnitude (Figure 13). Dipole vectors are shown as arrows pointing along the bond from the less electronegative atom toward the more electronegative atom. A small plus sign is drawn on the less electronegative end to indicate the partially positive end of the bond. The length of the arrow is proportional to the magnitude of the electronegativity difference between the two atoms.\r\n\r\n[caption id=\"attachment_455\" align=\"aligncenter\" width=\"650\"]<img src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23205156\/CNX_Chem_07_06_BondVector.jpg\" alt=\"Two images are shown and labeled, \u201ca\u201d and \u201cb.\u201d Image a shows a large sphere labeled, \u201cC,\u201d a left-facing arrow with a crossed end, and a smaller sphere labeled \u201cH.\u201d Image b shows a large sphere labeled, \u201cB,\u201d a right-facing arrow with a crossed end, and a smaller sphere labeled \u201cF.\u201d\" width=\"650\" height=\"172\" \/> Figure 13. (a) There is a small difference in electronegativity between C and H, represented as a short vector. (b) The electronegativity difference between B and F is much larger, so the vector representing the bond moment is much longer.[\/caption]\r\n\r\nA whole molecule may also have a separation of charge, depending on its molecular structure and the polarity of each of its bonds. If such a charge separation exists, the molecule is said to be a <strong>polar molecule<\/strong> (or dipole); otherwise the molecule is said to be nonpolar. The <strong>dipole moment <\/strong>measures the extent of net charge separation in the molecule as a whole. We determine the dipole moment by adding the bond moments in three-dimensional space, taking into account the molecular structure.\r\n\r\nFor diatomic molecules, there is only one bond, so its bond dipole moment determines the molecular polarity. Homonuclear diatomic molecules such as Br<sub>2<\/sub> and N<sub>2<\/sub> have no difference in electronegativity, so their dipole moment is zero. For heteronuclear molecules such as CO, there is a small dipole moment. For HF, there is a larger dipole moment because there is a larger difference in electronegativity.\r\n\r\nWhen a molecule contains more than one bond, the geometry must be taken into account. If the bonds in a molecule are arranged such that their bond moments cancel (vector sum equals zero), then the molecule is nonpolar. This is the situation in CO<sub>2<\/sub> (Figure 14). Each of the bonds is polar, but the molecule as a whole is nonpolar. From the Lewis structure, and using VSEPR theory, we determine that the CO<sub>2<\/sub> molecule is linear with polar C=O bonds on opposite sides of the carbon atom. The bond moments cancel because they are pointed in opposite directions. In the case of the water molecule (Figure 14), the Lewis structure again shows that there are two bonds to a central atom, and the electronegativity difference again shows that each of these bonds has a nonzero bond moment. In this case, however, the molecular structure is bent because of the lone pairs on O, and the two bond moments do not cancel. Therefore, water does have a net dipole moment and is a polar molecule (dipole).\r\n\r\n[caption id=\"attachment_456\" align=\"aligncenter\" width=\"650\"]<img src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23205157\/CNX_Chem_07_06_CO2H2Odip.jpg\" alt=\"Two images are shown and labeled, \u201ca\u201d and \u201cb.\u201d Image a shows a carbon atom bonded to two oxygen atoms in a ball-and-stick representation. Two arrows face away from the center of the molecule in opposite directions and are drawn horizontally like the molecule. These arrows are labeled, \u201cBond moments,\u201d and the image is labeled, \u201cOverall dipole moment equals 0.\u201d Image b shows an oxygen atom bonded to two hydrogen atoms in a downward-facing v-shaped arrangement. An upward-facing, vertical arrow is drawn below the molecule while two upward and inward facing arrows are drawn above the molecule. The upper arrows are labeled, \u201cBond moments,\u201d while the image is labeled, \u201cOverall dipole moment.\u201d\" width=\"650\" height=\"396\" \/> Figure 14. The overall dipole moment of a molecule depends on the individual bond dipole moments and how they are arranged. (a) Each CO bond has a bond dipole moment, but they point in opposite directions so that the net CO<sub>2<\/sub> molecule is nonpolar. (b) In contrast, water is polar because the OH bond moments do not cancel out.[\/caption]\r\n\r\nThe OCS molecule has a structure similar to CO<sub>2<\/sub>, but a sulfur atom has replaced one of the oxygen atoms. To determine if this molecule is polar, we draw the molecular structure. VSEPR theory predicts a linear molecule:\r\n\r\n<img class=\"aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23205159\/CNX_Chem_07_06_OSC_img.jpg\" alt=\"An image shows a carbon atom double bonded to a sulfur atom and an oxygen atom which are arranged in a horizontal plane. Two arrows face away from the center of the molecule in opposite directions and are drawn horizontally like the molecule. The left-facing arrow is larger than the right-facing arrow. These arrows are labeled, \u201cBond moments,\u201d and a left-facing arrow below the molecule is labeled, \u201cOverall dipole moment.\u201d\" width=\"325\" height=\"223\" \/>\r\n\r\nAlthough the C\u2013O bond is polar, C and S have the same electronegativity values as shown in Figure 15, so there is no C\u2013S dipole. Thus, the two bonds do not have of the same bond dipole moment, and the bond moments do not cancel. Because oxygen is more electronegative than sulfur, the oxygen end of the molecule is the negative end.\r\n\r\n[caption id=\"attachment_263\" align=\"aligncenter\" width=\"1024\"]<img src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23204815\/CNX_Chem_07_02_ENTable-1024x438.jpg\" alt=\"Part of the periodic table is shown. A downward-facing arrow is drawn to the left of the table and labeled, \u201cDecreasing electronegativity,\u201d while a right-facing arrow is drawn above the table and labeled \u201cIncreasing electronegativity.\u201d The electronegativity for almost all the elements is given.\" width=\"1024\" height=\"438\" \/> Figure 15. The electronegativity values derived by Pauling follow predictable periodic trends with the higher electronegativities toward the upper right of the periodic table.[\/caption]\r\n\r\nChloromethane, CH<sub>3<\/sub>Cl, is another example of a polar molecule. Although the polar C\u2013Cl and C\u2013H bonds are arranged in a tetrahedral geometry, the C\u2013Cl bonds have a larger bond moment than the C\u2013H bond, and the bond moments do not completely cancel each other. All of the dipoles have a upward component in the orientation shown, since carbon is more electronegative than hydrogen and less electronegative than chlorine:\r\n\r\n<a href=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/915\/2017\/02\/02152658\/CNX_Chem_07_06_CHCl3_img-1.jpg\"><img class=\"size-full wp-image-6141 aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/915\/2017\/02\/02152658\/CNX_Chem_07_06_CHCl3_img-1.jpg\" alt=\"\" width=\"114\" height=\"132\" \/><\/a>\r\n\r\nWhen we examine the highly symmetrical molecules BF<sub>3<\/sub> (trigonal planar), CH<sub>4<\/sub> (tetrahedral), PF<sub>5<\/sub> (trigonal bipyramidal), and SF<sub>6<\/sub> (octahedral), in which all the polar bonds are identical, the molecules are nonpolar. The bonds in these molecules are arranged such that their dipoles cancel. However, just because a molecule contains identical bonds does not mean that the dipoles will always cancel. Many molecules that have identical bonds and lone pairs on the central atoms have bond dipoles that do not cancel. Examples include H<sub>2<\/sub>S and NH<sub>3<\/sub>. A hydrogen atom is at the positive end and a nitrogen or sulfur atom is at the negative end of the polar bonds in these molecules:\r\n\r\n<img class=\"aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23205201\/CNX_Chem_07_06_SH2NH3_img.jpg\" alt=\"Two Lewis structures are shown. The left structure shows a sulfur atom with two lone pairs of electrons single bonded to two hydrogen atoms. Near the sulfur is a dipole symbol with a superscripted negative sign. Near each hydrogen is a dipole symbol with a superscripted positive sign. The right structure shows a nitrogen atom with one lone pair of electrons single bonded to three hydrogen atoms. Near the nitrogen is a dipole symbol with a superscripted negative sign. Near each hydrogen is a dipole symbol with a superscripted positive sign.\" width=\"650\" height=\"126\" \/>\r\n\r\nTo summarize, to be polar, a molecule must:\r\n<ol>\r\n \t<li>Contain at least one polar covalent bond.<\/li>\r\n \t<li>Have a molecular structure such that the sum of the vectors of each bond dipole moment does not cancel.<\/li>\r\n<\/ol>\r\n<h2>Properties of Polar Molecules<\/h2>\r\nPolar molecules tend to align when placed in an electric field with the positive end of the molecule oriented toward the negative plate and the negative end toward the positive plate (Figure 16). We can use an electrically charged object to attract polar molecules, but nonpolar molecules are not attracted. Also, polar solvents are better at dissolving polar substances, and nonpolar solvents are better at dissolving nonpolar substances.\r\n\r\n[caption id=\"attachment_460\" align=\"aligncenter\" width=\"801\"]<img src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23205202\/CNX_Chem_07_06_Dipolfield-1024x628.jpg\" alt=\"Two diagrams are shown and labeled, \u201ca\u201d and \u201cb.\u201d Diagram a shows two vertical, black lines. The left line is labeled with a negative sign and the right with a positive sign. There are five molecules in between. The molecules are separate from one another and are composed of a hydrogen atom bonded to a fluorine atom. The fluorine atom is labeled with a dipole symbol and a superscripted negative sign while the hydrogen atom is labeled with a dipole symbol and a superscripted positive sign. The molecules are randomly oriented in the space. The right diagram is also bracketed by two vertical, lines, but this time the line labeled as negative is red and the line labeled as positive is blue. The same molecules are present, but this time they are all facing horizontally, with the hydrogen-end of each molecule facing toward the red line.\" width=\"801\" height=\"491\" \/> Figure 16. (a) Molecules are always randomly distributed in the liquid state in the absence of an electric field. (b) When an electric field is applied, polar molecules like HF will align to the dipoles with the field direction.[\/caption]\r\n\r\n<div class=\"textbox\">The <a href=\"http:\/\/phet.colorado.edu\/en\/simulation\/molecule-polarity\" target=\"_blank\" rel=\"noopener noreferrer\">PhET molecule polarity simulation<\/a> provides many ways to explore dipole moments of bonds and molecules.<\/div>\r\n<div class=\"textbox examples\">\r\n<h3>Example 2: Polarity Simulations<\/h3>\r\nOpen the <a href=\"http:\/\/phet.colorado.edu\/en\/simulation\/molecule-polarity\" target=\"_blank\" rel=\"noopener noreferrer\">PhET molecule polarity simulation <\/a>and select the \u201cThree Atoms\u201d tab at the top. This should display a molecule ABC with three electronegativity adjustors. You can display or hide the bond moments, molecular dipoles, and partial charges at the right. Turning on the Electric Field will show whether the molecule moves when exposed to a field, similar to Figure 15.\r\n\r\nUse the electronegativity controls to determine how the molecular dipole will look for the starting bent molecule if:\r\n<ol>\r\n \t<li>A and C are very electronegative and B is in the middle of the range.<\/li>\r\n \t<li>A is very electronegative, and B and C are not.<\/li>\r\n<\/ol>\r\n[reveal-answer q=\"264495\"]Show Answer[\/reveal-answer]\r\n[hidden-answer a=\"264495\"]\r\n<ol>\r\n \t<li>Molecular dipole moment points immediately between A and C.<\/li>\r\n \t<li>Molecular dipole moment points along the A\u2013B bond, toward A.<\/li>\r\n<\/ol>\r\n[\/hidden-answer]\r\n<h4><strong>Check Your Learning<\/strong><\/h4>\r\nDetermine the partial charges that will give the largest possible bond dipoles.\r\n\r\n[reveal-answer q=\"192987\"]Show Answer[\/reveal-answer]\r\n[hidden-answer a=\"192987\"]The largest bond moments will occur with the largest partial charges. The two solutions above represent how unevenly the electrons are shared in the bond. The bond moments will be maximized when the electronegativity difference is greatest. The controls for A and C should be set to one extreme, and B should be set to the opposite extreme. Although the magnitude of the bond moment will not change based on whether B is the most electronegative or the least, the direction of the bond moment will.[\/hidden-answer]\r\n\r\n<\/div>\r\n<div class=\"textbox key-takeaways\">\r\n<h3>Key Concepts and Summary<\/h3>\r\nA dipole moment measures a separation of charge. For one bond, the bond dipole moment is determined by the difference in electronegativity between the two atoms. For a molecule, the overall dipole moment is determined by both the individual bond moments and how these dipoles are arranged in the molecular structure. Polar molecules (those with an appreciable dipole moment) interact with electric fields, whereas nonpolar molecules do not.\r\n\r\n<\/div>\r\n<div class=\"textbox exercises\">\r\n<h3>Exercises<\/h3>\r\n<ol>\r\n \t<li>Explain how a molecule that contains polar bonds can be nonpolar.\r\n<ol>\r\n \t<li><\/li>\r\n<\/ol>\r\n<\/li>\r\n \t<li>Which of the following molecules contain polar bonds? Which of these molecules and ions have dipole moments?\r\n<ol>\r\n \t<li>PCl<sub>3 <\/sub><\/li>\r\n \t<li>CS<sub>2 <\/sub><\/li>\r\n \t<li>CCl<sub>2<\/sub>F<sub>2 <\/sub><\/li>\r\n \t<li>ClNO (N is the central atom)<\/li>\r\n \t<li>CF<sub>4 <\/sub><\/li>\r\n \t<li>CH<sub>3<\/sub>Cl<\/li>\r\n \t<li>H<sub>2<\/sub>CO<\/li>\r\n<\/ol>\r\n<\/li>\r\n \t<li>Use the <a href=\"http:\/\/phet.colorado.edu\/en\/simulation\/molecule-polarity\" target=\"_blank\" rel=\"noopener noreferrer\">PhET simulation<\/a> to perform the following exercises for a two-atom molecule:\r\n<ol>\r\n \t<li>Adjust the electronegativity value so the bond dipole is pointing toward B. Then determine what the electronegativity values must be to switch the dipole so that it points toward A.<\/li>\r\n \t<li>With a partial positive charge on A, turn on the electric field and describe what happens.<\/li>\r\n \t<li>With a small partial negative charge on A, turn on the electric field and describe what happens.<\/li>\r\n \t<li>Reset all, and then with a large partial negative charge on A, turn on the electric field and describe what happens.<\/li>\r\n<\/ol>\r\n<\/li>\r\n \t<li>Use the <a href=\"http:\/\/phet.colorado.edu\/en\/simulation\/molecule-polarity\" target=\"_blank\" rel=\"noopener noreferrer\">PhET simulation<\/a> to perform the following exercises for a real molecule. You may need to rotate the molecules in three dimensions to see certain dipoles.\r\n<ol>\r\n \t<li>Look at the bond dipoles for NH<sub>3<\/sub>. Use these dipoles to predict whether N or H is more electronegative.<\/li>\r\n \t<li>Predict whether there should be a molecular dipole for NH<sub>3<\/sub> and, if so, in which direction it will point. Check the molecular dipole box to test your hypothesis.<\/li>\r\n<\/ol>\r\n<\/li>\r\n<\/ol>\r\n<\/div>\r\n<div class=\"textbox key-takeaways\">\r\n<h3>Key Concepts and Summary<\/h3>\r\nCovalent bonds form when electrons are shared between atoms and are attracted by the nuclei of both atoms. In pure covalent bonds, the electrons are shared equally. In polar covalent bonds, the electrons are shared unequally, as one atom exerts a stronger force of attraction on the electrons than the other. The ability of an atom to attract a pair of electrons in a chemical bond is called its electronegativity. The difference in electronegativity between two atoms determines how polar a bond will be. In a diatomic molecule with two identical atoms, there is no difference in electronegativity, so the bond is nonpolar or pure covalent. When the electronegativity difference is very large, as is the case between metals and nonmetals, the bonding is characterized as ionic.\r\n\r\n<\/div>\r\n<h2>Glossary<\/h2>\r\n<strong>bond dipole moment: <\/strong>separation of charge in a bond that depends on the difference in electronegativity and the bond distance represented by partial charges or a vector\r\n\r\n<strong>bond length:\u00a0<\/strong>distance between the nuclei of two bonded atoms at which the lowest potential energy is achieved\r\n\r\n<strong>covalent bond:\u00a0<\/strong>bond formed when electrons are shared between atoms\r\n\r\n<strong>dipole moment: <\/strong>property of a molecule that describes the separation of charge determined by the sum of the individual bond moments based on the molecular structure\r\n\r\n<strong>electronegativity:\u00a0<\/strong>tendency of an atom to attract electrons in a bond to itself\r\n\r\n<strong>polar covalent bond:\u00a0<\/strong>covalent bond between atoms of different electronegativities; a covalent bond with a positive end and a negative end\r\n\r\n<strong>polar molecule: <\/strong>(also, dipole) molecule with an overall dipole moment\r\n\r\n<strong>pure covalent bond:\u00a0<\/strong>(also, nonpolar covalent bond) covalent bond between atoms of identical electronegativities","rendered":"<div class=\"textbox learning-objectives\">\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<\/li>\n<li>Define electronegativity and assess the polarity of covalent bonds<\/li>\n<\/ul>\n<\/div>\n<h2>Lewis Structures<\/h2>\n<p>We also use Lewis symbols to indicate the formation of covalent bonds, which are shown in<strong> Lewis structures<\/strong>, drawings that describe the bonding in molecules and polyatomic ions. For example, when two chlorine atoms form a chlorine molecule, they share one pair of electrons:<\/p>\n<p><img decoding=\"async\" class=\"aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23211358\/CNX_Chem_07_03_C12dot_img1.jpg\" alt=\"A Lewis dot diagram shows a reaction. Two chlorine symbols, each surrounded by seven dots are separated by a plus sign. The dots on the first atom are all black and the dots on the second atom are all read. The phrase, \u201cChlorine atoms\u201d is written below. A right-facing arrow points to two chlorine symbols, each with six dots surrounding their outer edges and a shared pair of dots in between. One of the shared dots is black and one is red. The phrase, \u201cChlorine molecule\u201d is written below.\" \/><\/p>\n<p>The Lewis structure indicates that each Cl atom has three pairs of electrons that are not used in bonding (called <strong>lone pairs<\/strong>) and one shared pair of electrons (written between the atoms). A dash (or line) is sometimes used to indicate a shared pair of electrons:<\/p>\n<p>A single shared pair of electrons is called a <strong>single bond<\/strong>. Each Cl atom interacts with eight valence electrons: the six in the lone pairs and the two in the single bond.<\/p>\n<h2>Pure vs. Polar Covalent Bonds<\/h2>\n<p>If the atoms that form a covalent bond are identical, as in H<sub>2<\/sub>, Cl<sub>2<\/sub>, and other diatomic molecules, then the electrons in the bond must be shared equally. We refer to this as a <strong>pure covalent bond<\/strong>. Electrons shared in pure covalent bonds have an equal probability of being near each nucleus.<\/p>\n<p><img decoding=\"async\" class=\"aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23211400\/CNX_Chem_07_03_Cl2dash_img1.jpg\" alt=\"Two Lewis structures are shown. The left-hand structure shows two H atoms connected by a single bond. The right-hand structure shows two C l atoms connected by a single bond and each surrounded by six dots.\" \/><\/p>\n<p>In the case of Cl<sub>2<\/sub>, each atom starts off with seven valence electrons, and each Cl shares one electron with the other, forming one covalent bond:<\/p>\n<p>[latex]\\text{Cl}+\\text{Cl}\\rightarrow{\\text{Cl}}_{2}[\/latex]<\/p>\n<p>The total number of electrons around each individual atom consists of six nonbonding electrons and two shared (i.e., bonding) electrons for eight total electrons, matching the number of valence electrons in the noble gas argon. Since the bonding atoms are identical, Cl<sub>2<\/sub> also features a pure covalent bond.<\/p>\n<div style=\"width: 310px\" 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\/23211346\/CNX_Chem_07_02_HClBond1.jpg\" alt=\"Two diagrams are shown and labeled \u201ca\u201d and \u201cb.\u201d Diagram a shows a small sphere labeled, \u201cH\u201d and a larger sphere labeled, \u201cC l\u201d that overlap slightly. Both spheres have a small dot in the center. Diagram b shows an H bonded to a C l with a single bond. A dipole and a positive sign are written above the H and a dipole and negative sign are written above the C l. An arrow points toward the C l with a plus sign on the end furthest from the arrow\u2019s head near the H.\" width=\"300\" height=\"206\" \/><\/p>\n<p class=\"wp-caption-text\">Figure 2. (a) The distribution of electron density in the HCl molecule is uneven. The electron density is greater around the chlorine nucleus. The small, black dots indicate the location of the hydrogen and chlorine nuclei in the molecule. (b) Symbols \u03b4+ and \u03b4\u2013 indicate the polarity of the H\u2013Cl bond.<\/p>\n<\/div>\n<p>When the atoms linked by a covalent bond are different, the bonding electrons are shared, but no longer equally. Instead, the bonding electrons are more attracted to one atom than the other, giving rise to a shift of electron density toward that atom. This unequal distribution of electrons is known as a <strong>polar covalent bond<\/strong>, characterized by a partial positive charge on one atom and a partial negative charge on the other. The atom that attracts the electrons more strongly acquires the partial negative charge and vice versa. For example, the electrons in the H\u2013Cl bond of a hydrogen chloride molecule spend more time near the chlorine atom than near the hydrogen atom. Thus, in an HCl molecule, the chlorine atom carries a partial negative charge and the hydrogen atom has a partial positive charge.\u00a0Figure 2 shows the distribution of electrons in the H\u2013Cl bond. Note that the shaded area around Cl is much larger than it is around H. Compare this to Figure 1, which shows the even distribution of electrons in the H<sub>2<\/sub> nonpolar bond.<\/p>\n<p>We sometimes designate the positive and negative atoms in a polar covalent bond using a lowercase Greek letter \u201cdelta,\u201d \u03b4, with a plus sign or minus sign to indicate whether the atom has a partial positive charge (\u03b4+) or a partial negative charge (\u03b4\u2013). This symbolism is shown for the H\u2013Cl molecule in Figure 2.<\/p>\n<h2>Electronegativity<\/h2>\n<p>Whether a bond is nonpolar or polar covalent is determined by a property of the bonding atoms called <strong>electronegativity<\/strong>. Electronegativity is a measure of the tendency of an atom to attract electrons (or electron density) towards itself. It determines how the shared electrons are distributed between the two atoms in a bond. The more strongly an atom attracts the electrons in its bonds, the larger its electronegativity. Electrons in a polar covalent bond are shifted toward the more electronegative atom; thus, the more electronegative atom is the one with the partial negative charge. The greater the difference in electronegativity, the more polarized the electron distribution and the larger the partial charges of the atoms.<\/p>\n<p>Figure 3 shows the electronegativity values of the elements as proposed by one of the most famous chemists of the twentieth century: Linus Pauling (Figure 4). In general, electronegativity increases from left to right across a period in the periodic table and decreases down a group. Thus, the nonmetals, which lie in the upper right, tend to have the highest electronegativities, with fluorine the most electronegative element of all (EN = 4.0). Metals tend to be less electronegative elements, and the group 1 metals have the lowest electronegativities. Note that noble gases are excluded from this figure because these atoms usually do not share electrons with others atoms since they have a full valence shell. (While noble gas compounds such as XeO<sub>2<\/sub> do exist, they can only be formed under extreme conditions, and thus they do not fit neatly into the general model of electronegativity.)<\/p>\n<div style=\"width: 890px\" 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\/23211347\/CNX_Chem_07_02_ENTable1.jpg\" alt=\"Part of the periodic table is shown. A downward-facing arrow is drawn to the left of the table and labeled, \u201cDecreasing electronegativity,\u201d while a right-facing arrow is drawn above the table and labeled \u201cIncreasing electronegativity.\u201d The electronegativity for almost all the elements is given.\" width=\"880\" height=\"377\" \/><\/p>\n<p class=\"wp-caption-text\">Figure 3. The electronegativity values derived by Pauling follow predictable periodic trends with the higher electronegativities toward the upper right of the periodic table.<\/p>\n<\/div>\n<div class=\"textbox shaded\">\n<h3>Portrait of a Chemist: Linus Pauling<\/h3>\n<p>Linus Pauling, shown in Figure 4, is the only person to have received two unshared (individual) Nobel Prizes: one for chemistry in 1954 for his work on the nature of chemical bonds and one for peace in 1962 for his opposition to weapons of mass destruction. He developed many of the theories and concepts that are foundational to our current understanding of chemistry, including electronegativity and resonance structures.<\/p>\n<div style=\"width: 335px\" class=\"wp-caption aligncenter\"><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\/23211349\/CNX_Chem_07_02_Pauling1.jpg\" alt=\"A photograph of Linus Pauling is shown.\" width=\"325\" height=\"394\" \/><\/p>\n<p class=\"wp-caption-text\">Figure 4. Linus Pauling (1901\u20131994) made many important contributions to the field of chemistry. He was also a prominent activist, publicizing issues related to health and nuclear weapons.<\/p>\n<\/div>\n<p>Pauling also contributed to many other fields besides chemistry. His research on sickle cell anemia revealed the cause of the disease\u2014the presence of a genetically inherited abnormal protein in the blood\u2014and paved the way for the field of molecular genetics. His work was also pivotal in curbing the testing of nuclear weapons; he proved that radioactive fallout from nuclear testing posed a public health risk.<\/p>\n<\/div>\n<h3>Electronegativity and Bond Type<\/h3>\n<p>The absolute value of the difference in electronegativity (\u0394EN) of two bonded atoms provides a rough measure of the polarity to be expected in the bond and, thus, the bond type. When the difference is very small or zero, the bond is covalent and nonpolar. When it is large, the bond is polar covalent or ionic. The absolute values of the electronegativity differences between the atoms in the bonds H\u2013H, H\u2013Cl, and Na\u2013Cl are 0 (nonpolar), 0.9 (polar covalent), and 2.1 (ionic), respectively. The degree to which electrons are shared between atoms varies from completely equal (pure covalent bonding) to not at all (ionic bonding). Figure 5 shows the relationship between electronegativity difference and bond type.<\/p>\n<div id=\"attachment_1667\" style=\"width: 637px\" class=\"wp-caption aligncenter\"><img loading=\"lazy\" decoding=\"async\" aria-describedby=\"caption-attachment-1667\" class=\"wp-image-1667 size-full\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/2835\/2018\/08\/29173516\/Polar-bonds.jpg\" alt=\"Two flow charts and table are shown. The first flow chart is labeled, \u201cElectronegativity difference between bonding atoms.\u201d Below this label are three rounded text bubbles, connected by a downward-facing arrow, labeled, \u201cZero,\u201d \u201cIntermediate,\u201d and \u201cLarge,\u201d respectively. The second flow chart is labeled, \u201cBond type.\u201d Below this label are three rounded text bubbles, connected by a downward-facing arrow, labeled, \u201cPure covalent,\u201d \u201cPolar covalent,\u201d and \u201cIonic,\u201d respectively. A double ended arrow is written vertically to the right of the flow charts and labeled, \u201cCovalent character decreases; ionic character increases.\u201d The table is made up of two columns and four rows. The header line is labeled \u201cBond type\u201d and \u201cElectronegativity difference.\u201d The left column contains the phrases \u201cPure covalent,\u201d \u201cPolar covalent,\u201d and \u201cIonic,\u201d while the right column contains the values \u201cless than 0.4,\u201d \u201cbetween 0.4 and 1.8,\u201d and \u201cgreater than 1.8.\u201d\" width=\"627\" height=\"312\" \/><\/p>\n<p id=\"caption-attachment-1667\" class=\"wp-caption-text\">Figure 5. As the electronegativity difference increases between two atoms, the bond becomes more ionic.<\/p>\n<\/div>\n<p>A rough approximation of the electronegativity differences associated with covalent, polar covalent, and ionic bonds is shown in Figure 5. This table is just a general guide, however, with many exceptions. For example, the H and F atoms in HF have an electronegativity difference of 1.9, and the N and H atoms in NH<sub>3<\/sub> a difference of 0.9, yet both of these compounds form bonds that are considered polar covalent. Likewise, the Na and Cl atoms in NaCl have an electronegativity difference of 2.1, and the Mn and I atoms in MnI<sub>2<\/sub> have a difference of 1.0, yet both of these substances form ionic compounds.<\/p>\n<p>The best guide to the covalent or ionic character of a bond is to consider the types of atoms involved and their relative positions in the periodic table. Bonds between two nonmetals are generally covalent; bonding between a metal and a nonmetal is often ionic.<\/p>\n<p>Some compounds contain both covalent and ionic bonds. The atoms in polyatomic ions, such as OH<sup>\u2013<\/sup>, [latex]{\\text{NO}}_{3}^{-},[\/latex] and [latex]{\\text{NH}}_{4}^{\\text{+}},[\/latex] are held together by polar covalent bonds. However, these polyatomic ions form ionic compounds by combining with ions of opposite charge. For example, potassium nitrate, KNO<sub>3<\/sub>, contains the K<sup>+<\/sup> cation and the polyatomic [latex]{\\text{NO}}_{3}^{-}[\/latex] anion. Thus, bonding in potassium nitrate is ionic, resulting from the electrostatic attraction between the ions K<sup>+<\/sup> and [latex]{\\text{NO}}_{3}^{-},[\/latex] as well as covalent between the nitrogen and oxygen atoms in [latex]{\\text{NO}}_{3}^{-}[\/latex].<\/p>\n<div class=\"textbox examples\">\n<h3>Example 1:\u00a0Electronegativity and Bond Polarity<\/h3>\n<p>Bond polarities play an important role in determining the structure of proteins. Using the electronegativity values in Figure 7.6, arrange the following covalent bonds\u2014all commonly found in amino acids\u2014in order of increasing polarity. Then designate the positive and negative atoms using the symbols \u03b4+ and \u03b4\u2013:<\/p>\n<p>C\u2013H, C\u2013N, C\u2013O, N\u2013H, O\u2013H, S\u2013H<\/p>\n<div class=\"qa-wrapper\" style=\"display: block\"><span class=\"show-answer collapsed\" style=\"cursor: pointer\" data-target=\"q862588\">Show Answer<\/span><\/p>\n<div id=\"q862588\" class=\"hidden-answer\" style=\"display: none\">\n<p>The polarity of these bonds increases as the absolute value of the electronegativity difference increases. The atom with the \u03b4\u2013 designation is the more electronegative of the two. Table 1\u00a0shows these bonds in order of increasing polarity.<\/p>\n<table summary=\"This table has three columns and seven rows. The first row is a header row and it labels each column. The first column header is, \u201cBond,\u201d the second is, \u201ccapital delta E N,\u201d and the third is, \u201cPolarity.\u201d Under the \u201cBond\u201d column are the following: C bonds to H with a single bond; S bonds to H with a single bond; C bonds to N with a single bond; N bonds to H with a single bond; C bonds to O with a single bond; and O bonds to H with a single bond. Under the \u201ccapital delta E N\u201d columna are the values: 0.4; 0.4; 0.5; 0.9; 1.0; and 1.4. Under the \u201cPolarity\u201d column are the follwoing: C bonds to H with a single bond, there is a lowercase delta negative sign above the C and a lowercase delta positive sign over H; S bonds to H with a single bond, there is a lowercase delta negative sign over the S and a lowercase delta positive sign over the H; C bonds to N with a single bond, there is a lowercase delta positive sign over the C and a lowercase delta negative sign over the N; N bonds to H with a single bond, there is a lowercase delta negative sign over the N and a lowercase delta positive sign over the H; C bonds to O with a single bond, there is a lowercase delta positive sign over C and a lowercase delta negative sign over the O; and O bonds to H with a single bond, there is a lowercase delta negative sign over the O and a lowercase delta positive sign over the H.\">\n<thead>\n<tr>\n<th colspan=\"3\">Table 1. Bond Polarity and Electronegativity Difference<\/th>\n<\/tr>\n<\/thead>\n<tbody>\n<tr>\n<th>Bond<\/th>\n<th>\u0394EN<\/th>\n<th>Polarity<\/th>\n<\/tr>\n<tr>\n<td>C\u2013H<\/td>\n<td>0.4<\/td>\n<td>[latex]\\stackrel{\\delta -}{\\text{C}}-\\stackrel{\\delta \\text{+}}{\\text{H}}[\/latex]<\/td>\n<\/tr>\n<tr>\n<td>S\u2013H<\/td>\n<td>0.4<\/td>\n<td>[latex]\\stackrel{\\delta -}{\\text{S}}-\\stackrel{\\delta \\text{+}}{\\text{H}}[\/latex]<\/td>\n<\/tr>\n<tr>\n<td>C\u2013N<\/td>\n<td>0.5<\/td>\n<td>[latex]\\stackrel{\\delta \\text{+}}{\\text{C}}-\\stackrel{\\delta -}{\\text{N}}[\/latex]<\/td>\n<\/tr>\n<tr>\n<td>N\u2013H<\/td>\n<td>0.9<\/td>\n<td>[latex]\\stackrel{\\delta -}{\\text{N}}-\\stackrel{\\delta \\text{+}}{\\text{H}}[\/latex]<\/td>\n<\/tr>\n<tr>\n<td>C\u2013O<\/td>\n<td>1.0<\/td>\n<td>[latex]\\stackrel{\\delta \\text{+}}{\\text{C}}-\\stackrel{\\delta -}{\\text{O}}[\/latex]<\/td>\n<\/tr>\n<tr>\n<td>O\u2013H<\/td>\n<td>1.4<\/td>\n<td>[latex]\\stackrel{\\delta -}{\\text{O}}-\\stackrel{\\delta \\text{+}}{\\text{H}}[\/latex]<\/td>\n<\/tr>\n<\/tbody>\n<\/table>\n<\/div>\n<\/div>\n<p><strong>Check Your Learning<\/strong><\/p>\n<p>Silicones are polymeric compounds containing, among others, the following types of covalent bonds: Si\u2013O, Si\u2013C, C\u2013H, and C\u2013C. Using the electronegativity values in Figure 3, arrange the bonds in order of increasing polarity and designate the positive and negative atoms using the symbols \u03b4+ and \u03b4\u2013.<\/p>\n<div class=\"qa-wrapper\" style=\"display: block\"><span class=\"show-answer collapsed\" style=\"cursor: pointer\" data-target=\"q113211\">Show Answer<\/span><\/p>\n<div id=\"q113211\" class=\"hidden-answer\" style=\"display: none\">\n<table summary=\"This table has three columns and five rows. The first row is a header row and it labels each column. The first column header is, \u201cBond,\u201d the second column header is, \u201cElectronegativity Difference,\u201d and the third column header is, \u201cPolarity.\u201d Under the column \u201cBond\u201d are the following: C bonds to C with a single bond; C bonds to H with a single bond; S i bonds to C with a single bond; and S i bonds to O with a single bond. Under the column \u201cElectronegativity Difference\u201d are the values: 0.0; 0.4; 0.7; and 1.7. Under the column \u201cPolarity\u201d are the following: nonpolar, C bonds to H with a single bond, there is a lowercase delta negative sign over C and a lowercase delta positive sign over H; S i bonds to C with a single bond, there is a lowercase delta positive sign over S i and a lowercase delta negative sign over C; and S i bonds to O with a single bond, there is a lowercase delta positive sign over S i and a lowercase delta negative sign over O.\">\n<thead>\n<tr>\n<th>Bond<\/th>\n<th>Electronegativity Difference<\/th>\n<th>Polarity<\/th>\n<\/tr>\n<\/thead>\n<tbody>\n<tr>\n<td>C\u2013C<\/td>\n<td>0.0<\/td>\n<td>nonpolar<\/td>\n<\/tr>\n<tr>\n<td>C\u2013H<\/td>\n<td>0.4<\/td>\n<td>[latex]\\stackrel{\\delta -}{\\text{C}}-\\stackrel{\\delta \\text{+}}{\\text{H}}[\/latex]<\/td>\n<\/tr>\n<tr>\n<td>Si\u2013C<\/td>\n<td>0.7<\/td>\n<td>[latex]\\stackrel{\\delta \\text{+}}{\\text{Si}}-\\stackrel{\\delta -}{\\text{C}}[\/latex]<\/td>\n<\/tr>\n<tr>\n<td>Si\u2013O<\/td>\n<td>1.7<\/td>\n<td>[latex]\\stackrel{\\delta \\text{+}}{\\text{Si}}-\\stackrel{\\delta -}{\\text{O}}[\/latex]<\/td>\n<\/tr>\n<\/tbody>\n<\/table>\n<\/div>\n<\/div>\n<\/div>\n<h2>Molecular Polarity and Dipole Moment<\/h2>\n<p>As discussed previously, polar covalent bonds connect two atoms with differing electronegativities, leaving one atom with a partial positive charge (\u0394+) and the other atom with a partial negative charge (\u0394\u2013), as the electrons are pulled toward the more electronegative atom. This separation of charge gives rise to a bond dipole moment. The magnitude of a <strong>bond dipole moment<\/strong> is represented by the Greek letter mu (<em>\u00b5<\/em>) and is given by the formula shown below, where Q is the magnitude of the partial charges (determined by the electronegativity difference) and r is the distance between the charges:<\/p>\n<p>[latex]\\mu =\\text{Qr}[\/latex]<\/p>\n<p>This bond moment can be represented as a <strong>vector<\/strong>, a quantity having both direction and magnitude (Figure 13). Dipole vectors are shown as arrows pointing along the bond from the less electronegative atom toward the more electronegative atom. A small plus sign is drawn on the less electronegative end to indicate the partially positive end of the bond. The length of the arrow is proportional to the magnitude of the electronegativity difference between the two atoms.<\/p>\n<div id=\"attachment_455\" style=\"width: 660px\" class=\"wp-caption aligncenter\"><img loading=\"lazy\" decoding=\"async\" aria-describedby=\"caption-attachment-455\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23205156\/CNX_Chem_07_06_BondVector.jpg\" alt=\"Two images are shown and labeled, \u201ca\u201d and \u201cb.\u201d Image a shows a large sphere labeled, \u201cC,\u201d a left-facing arrow with a crossed end, and a smaller sphere labeled \u201cH.\u201d Image b shows a large sphere labeled, \u201cB,\u201d a right-facing arrow with a crossed end, and a smaller sphere labeled \u201cF.\u201d\" width=\"650\" height=\"172\" \/><\/p>\n<p id=\"caption-attachment-455\" class=\"wp-caption-text\">Figure 13. (a) There is a small difference in electronegativity between C and H, represented as a short vector. (b) The electronegativity difference between B and F is much larger, so the vector representing the bond moment is much longer.<\/p>\n<\/div>\n<p>A whole molecule may also have a separation of charge, depending on its molecular structure and the polarity of each of its bonds. If such a charge separation exists, the molecule is said to be a <strong>polar molecule<\/strong> (or dipole); otherwise the molecule is said to be nonpolar. The <strong>dipole moment <\/strong>measures the extent of net charge separation in the molecule as a whole. We determine the dipole moment by adding the bond moments in three-dimensional space, taking into account the molecular structure.<\/p>\n<p>For diatomic molecules, there is only one bond, so its bond dipole moment determines the molecular polarity. Homonuclear diatomic molecules such as Br<sub>2<\/sub> and N<sub>2<\/sub> have no difference in electronegativity, so their dipole moment is zero. For heteronuclear molecules such as CO, there is a small dipole moment. For HF, there is a larger dipole moment because there is a larger difference in electronegativity.<\/p>\n<p>When a molecule contains more than one bond, the geometry must be taken into account. If the bonds in a molecule are arranged such that their bond moments cancel (vector sum equals zero), then the molecule is nonpolar. This is the situation in CO<sub>2<\/sub> (Figure 14). Each of the bonds is polar, but the molecule as a whole is nonpolar. From the Lewis structure, and using VSEPR theory, we determine that the CO<sub>2<\/sub> molecule is linear with polar C=O bonds on opposite sides of the carbon atom. The bond moments cancel because they are pointed in opposite directions. In the case of the water molecule (Figure 14), the Lewis structure again shows that there are two bonds to a central atom, and the electronegativity difference again shows that each of these bonds has a nonzero bond moment. In this case, however, the molecular structure is bent because of the lone pairs on O, and the two bond moments do not cancel. Therefore, water does have a net dipole moment and is a polar molecule (dipole).<\/p>\n<div id=\"attachment_456\" style=\"width: 660px\" class=\"wp-caption aligncenter\"><img loading=\"lazy\" decoding=\"async\" aria-describedby=\"caption-attachment-456\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23205157\/CNX_Chem_07_06_CO2H2Odip.jpg\" alt=\"Two images are shown and labeled, \u201ca\u201d and \u201cb.\u201d Image a shows a carbon atom bonded to two oxygen atoms in a ball-and-stick representation. Two arrows face away from the center of the molecule in opposite directions and are drawn horizontally like the molecule. These arrows are labeled, \u201cBond moments,\u201d and the image is labeled, \u201cOverall dipole moment equals 0.\u201d Image b shows an oxygen atom bonded to two hydrogen atoms in a downward-facing v-shaped arrangement. An upward-facing, vertical arrow is drawn below the molecule while two upward and inward facing arrows are drawn above the molecule. The upper arrows are labeled, \u201cBond moments,\u201d while the image is labeled, \u201cOverall dipole moment.\u201d\" width=\"650\" height=\"396\" \/><\/p>\n<p id=\"caption-attachment-456\" class=\"wp-caption-text\">Figure 14. The overall dipole moment of a molecule depends on the individual bond dipole moments and how they are arranged. (a) Each CO bond has a bond dipole moment, but they point in opposite directions so that the net CO<sub>2<\/sub> molecule is nonpolar. (b) In contrast, water is polar because the OH bond moments do not cancel out.<\/p>\n<\/div>\n<p>The OCS molecule has a structure similar to CO<sub>2<\/sub>, but a sulfur atom has replaced one of the oxygen atoms. To determine if this molecule is polar, we draw the molecular structure. VSEPR theory predicts a linear molecule:<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" class=\"aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23205159\/CNX_Chem_07_06_OSC_img.jpg\" alt=\"An image shows a carbon atom double bonded to a sulfur atom and an oxygen atom which are arranged in a horizontal plane. Two arrows face away from the center of the molecule in opposite directions and are drawn horizontally like the molecule. The left-facing arrow is larger than the right-facing arrow. These arrows are labeled, \u201cBond moments,\u201d and a left-facing arrow below the molecule is labeled, \u201cOverall dipole moment.\u201d\" width=\"325\" height=\"223\" \/><\/p>\n<p>Although the C\u2013O bond is polar, C and S have the same electronegativity values as shown in Figure 15, so there is no C\u2013S dipole. Thus, the two bonds do not have of the same bond dipole moment, and the bond moments do not cancel. Because oxygen is more electronegative than sulfur, the oxygen end of the molecule is the negative end.<\/p>\n<div id=\"attachment_263\" style=\"width: 1034px\" class=\"wp-caption aligncenter\"><img loading=\"lazy\" decoding=\"async\" aria-describedby=\"caption-attachment-263\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23204815\/CNX_Chem_07_02_ENTable-1024x438.jpg\" alt=\"Part of the periodic table is shown. A downward-facing arrow is drawn to the left of the table and labeled, \u201cDecreasing electronegativity,\u201d while a right-facing arrow is drawn above the table and labeled \u201cIncreasing electronegativity.\u201d The electronegativity for almost all the elements is given.\" width=\"1024\" height=\"438\" \/><\/p>\n<p id=\"caption-attachment-263\" class=\"wp-caption-text\">Figure 15. The electronegativity values derived by Pauling follow predictable periodic trends with the higher electronegativities toward the upper right of the periodic table.<\/p>\n<\/div>\n<p>Chloromethane, CH<sub>3<\/sub>Cl, is another example of a polar molecule. Although the polar C\u2013Cl and C\u2013H bonds are arranged in a tetrahedral geometry, the C\u2013Cl bonds have a larger bond moment than the C\u2013H bond, and the bond moments do not completely cancel each other. All of the dipoles have a upward component in the orientation shown, since carbon is more electronegative than hydrogen and less electronegative than chlorine:<\/p>\n<p><a href=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/915\/2017\/02\/02152658\/CNX_Chem_07_06_CHCl3_img-1.jpg\"><img loading=\"lazy\" decoding=\"async\" class=\"size-full wp-image-6141 aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images\/wp-content\/uploads\/sites\/915\/2017\/02\/02152658\/CNX_Chem_07_06_CHCl3_img-1.jpg\" alt=\"\" width=\"114\" height=\"132\" \/><\/a><\/p>\n<p>When we examine the highly symmetrical molecules BF<sub>3<\/sub> (trigonal planar), CH<sub>4<\/sub> (tetrahedral), PF<sub>5<\/sub> (trigonal bipyramidal), and SF<sub>6<\/sub> (octahedral), in which all the polar bonds are identical, the molecules are nonpolar. The bonds in these molecules are arranged such that their dipoles cancel. However, just because a molecule contains identical bonds does not mean that the dipoles will always cancel. Many molecules that have identical bonds and lone pairs on the central atoms have bond dipoles that do not cancel. Examples include H<sub>2<\/sub>S and NH<sub>3<\/sub>. A hydrogen atom is at the positive end and a nitrogen or sulfur atom is at the negative end of the polar bonds in these molecules:<\/p>\n<p><img loading=\"lazy\" decoding=\"async\" class=\"aligncenter\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23205201\/CNX_Chem_07_06_SH2NH3_img.jpg\" alt=\"Two Lewis structures are shown. The left structure shows a sulfur atom with two lone pairs of electrons single bonded to two hydrogen atoms. Near the sulfur is a dipole symbol with a superscripted negative sign. Near each hydrogen is a dipole symbol with a superscripted positive sign. The right structure shows a nitrogen atom with one lone pair of electrons single bonded to three hydrogen atoms. Near the nitrogen is a dipole symbol with a superscripted negative sign. Near each hydrogen is a dipole symbol with a superscripted positive sign.\" width=\"650\" height=\"126\" \/><\/p>\n<p>To summarize, to be polar, a molecule must:<\/p>\n<ol>\n<li>Contain at least one polar covalent bond.<\/li>\n<li>Have a molecular structure such that the sum of the vectors of each bond dipole moment does not cancel.<\/li>\n<\/ol>\n<h2>Properties of Polar Molecules<\/h2>\n<p>Polar molecules tend to align when placed in an electric field with the positive end of the molecule oriented toward the negative plate and the negative end toward the positive plate (Figure 16). We can use an electrically charged object to attract polar molecules, but nonpolar molecules are not attracted. Also, polar solvents are better at dissolving polar substances, and nonpolar solvents are better at dissolving nonpolar substances.<\/p>\n<div id=\"attachment_460\" style=\"width: 811px\" class=\"wp-caption aligncenter\"><img loading=\"lazy\" decoding=\"async\" aria-describedby=\"caption-attachment-460\" src=\"https:\/\/s3-us-west-2.amazonaws.com\/courses-images-archive-read-only\/wp-content\/uploads\/sites\/887\/2015\/04\/23205202\/CNX_Chem_07_06_Dipolfield-1024x628.jpg\" alt=\"Two diagrams are shown and labeled, \u201ca\u201d and \u201cb.\u201d Diagram a shows two vertical, black lines. The left line is labeled with a negative sign and the right with a positive sign. There are five molecules in between. The molecules are separate from one another and are composed of a hydrogen atom bonded to a fluorine atom. The fluorine atom is labeled with a dipole symbol and a superscripted negative sign while the hydrogen atom is labeled with a dipole symbol and a superscripted positive sign. The molecules are randomly oriented in the space. The right diagram is also bracketed by two vertical, lines, but this time the line labeled as negative is red and the line labeled as positive is blue. The same molecules are present, but this time they are all facing horizontally, with the hydrogen-end of each molecule facing toward the red line.\" width=\"801\" height=\"491\" \/><\/p>\n<p id=\"caption-attachment-460\" class=\"wp-caption-text\">Figure 16. (a) Molecules are always randomly distributed in the liquid state in the absence of an electric field. (b) When an electric field is applied, polar molecules like HF will align to the dipoles with the field direction.<\/p>\n<\/div>\n<div class=\"textbox\">The <a href=\"http:\/\/phet.colorado.edu\/en\/simulation\/molecule-polarity\" target=\"_blank\" rel=\"noopener noreferrer\">PhET molecule polarity simulation<\/a> provides many ways to explore dipole moments of bonds and molecules.<\/div>\n<div class=\"textbox examples\">\n<h3>Example 2: Polarity Simulations<\/h3>\n<p>Open the <a href=\"http:\/\/phet.colorado.edu\/en\/simulation\/molecule-polarity\" target=\"_blank\" rel=\"noopener noreferrer\">PhET molecule polarity simulation <\/a>and select the \u201cThree Atoms\u201d tab at the top. This should display a molecule ABC with three electronegativity adjustors. You can display or hide the bond moments, molecular dipoles, and partial charges at the right. Turning on the Electric Field will show whether the molecule moves when exposed to a field, similar to Figure 15.<\/p>\n<p>Use the electronegativity controls to determine how the molecular dipole will look for the starting bent molecule if:<\/p>\n<ol>\n<li>A and C are very electronegative and B is in the middle of the range.<\/li>\n<li>A is very electronegative, and B and C are not.<\/li>\n<\/ol>\n<div class=\"qa-wrapper\" style=\"display: block\"><span class=\"show-answer collapsed\" style=\"cursor: pointer\" data-target=\"q264495\">Show Answer<\/span><\/p>\n<div id=\"q264495\" class=\"hidden-answer\" style=\"display: none\">\n<ol>\n<li>Molecular dipole moment points immediately between A and C.<\/li>\n<li>Molecular dipole moment points along the A\u2013B bond, toward A.<\/li>\n<\/ol>\n<\/div>\n<\/div>\n<h4><strong>Check Your Learning<\/strong><\/h4>\n<p>Determine the partial charges that will give the largest possible bond dipoles.<\/p>\n<div class=\"qa-wrapper\" style=\"display: block\"><span class=\"show-answer collapsed\" style=\"cursor: pointer\" data-target=\"q192987\">Show Answer<\/span><\/p>\n<div id=\"q192987\" class=\"hidden-answer\" style=\"display: none\">The largest bond moments will occur with the largest partial charges. The two solutions above represent how unevenly the electrons are shared in the bond. The bond moments will be maximized when the electronegativity difference is greatest. The controls for A and C should be set to one extreme, and B should be set to the opposite extreme. Although the magnitude of the bond moment will not change based on whether B is the most electronegative or the least, the direction of the bond moment will.<\/div>\n<\/div>\n<\/div>\n<div class=\"textbox key-takeaways\">\n<h3>Key Concepts and Summary<\/h3>\n<p>A dipole moment measures a separation of charge. For one bond, the bond dipole moment is determined by the difference in electronegativity between the two atoms. For a molecule, the overall dipole moment is determined by both the individual bond moments and how these dipoles are arranged in the molecular structure. Polar molecules (those with an appreciable dipole moment) interact with electric fields, whereas nonpolar molecules do not.<\/p>\n<\/div>\n<div class=\"textbox exercises\">\n<h3>Exercises<\/h3>\n<ol>\n<li>Explain how a molecule that contains polar bonds can be nonpolar.\n<ol>\n<li><\/li>\n<\/ol>\n<\/li>\n<li>Which of the following molecules contain polar bonds? Which of these molecules and ions have dipole moments?\n<ol>\n<li>PCl<sub>3 <\/sub><\/li>\n<li>CS<sub>2 <\/sub><\/li>\n<li>CCl<sub>2<\/sub>F<sub>2 <\/sub><\/li>\n<li>ClNO (N is the central atom)<\/li>\n<li>CF<sub>4 <\/sub><\/li>\n<li>CH<sub>3<\/sub>Cl<\/li>\n<li>H<sub>2<\/sub>CO<\/li>\n<\/ol>\n<\/li>\n<li>Use the <a href=\"http:\/\/phet.colorado.edu\/en\/simulation\/molecule-polarity\" target=\"_blank\" rel=\"noopener noreferrer\">PhET simulation<\/a> to perform the following exercises for a two-atom molecule:\n<ol>\n<li>Adjust the electronegativity value so the bond dipole is pointing toward B. Then determine what the electronegativity values must be to switch the dipole so that it points toward A.<\/li>\n<li>With a partial positive charge on A, turn on the electric field and describe what happens.<\/li>\n<li>With a small partial negative charge on A, turn on the electric field and describe what happens.<\/li>\n<li>Reset all, and then with a large partial negative charge on A, turn on the electric field and describe what happens.<\/li>\n<\/ol>\n<\/li>\n<li>Use the <a href=\"http:\/\/phet.colorado.edu\/en\/simulation\/molecule-polarity\" target=\"_blank\" rel=\"noopener noreferrer\">PhET simulation<\/a> to perform the following exercises for a real molecule. You may need to rotate the molecules in three dimensions to see certain dipoles.\n<ol>\n<li>Look at the bond dipoles for NH<sub>3<\/sub>. Use these dipoles to predict whether N or H is more electronegative.<\/li>\n<li>Predict whether there should be a molecular dipole for NH<sub>3<\/sub> and, if so, in which direction it will point. Check the molecular dipole box to test your hypothesis.<\/li>\n<\/ol>\n<\/li>\n<\/ol>\n<\/div>\n<div class=\"textbox key-takeaways\">\n<h3>Key Concepts and Summary<\/h3>\n<p>Covalent bonds form when electrons are shared between atoms and are attracted by the nuclei of both atoms. In pure covalent bonds, the electrons are shared equally. In polar covalent bonds, the electrons are shared unequally, as one atom exerts a stronger force of attraction on the electrons than the other. The ability of an atom to attract a pair of electrons in a chemical bond is called its electronegativity. The difference in electronegativity between two atoms determines how polar a bond will be. In a diatomic molecule with two identical atoms, there is no difference in electronegativity, so the bond is nonpolar or pure covalent. When the electronegativity difference is very large, as is the case between metals and nonmetals, the bonding is characterized as ionic.<\/p>\n<\/div>\n<h2>Glossary<\/h2>\n<p><strong>bond dipole moment: <\/strong>separation of charge in a bond that depends on the difference in electronegativity and the bond distance represented by partial charges or a vector<\/p>\n<p><strong>bond length:\u00a0<\/strong>distance between the nuclei of two bonded atoms at which the lowest potential energy is achieved<\/p>\n<p><strong>covalent bond:\u00a0<\/strong>bond formed when electrons are shared between atoms<\/p>\n<p><strong>dipole moment: <\/strong>property of a molecule that describes the separation of charge determined by the sum of the individual bond moments based on the molecular structure<\/p>\n<p><strong>electronegativity:\u00a0<\/strong>tendency of an atom to attract electrons in a bond to itself<\/p>\n<p><strong>polar covalent bond:\u00a0<\/strong>covalent bond between atoms of different electronegativities; a covalent bond with a positive end and a negative end<\/p>\n<p><strong>polar molecule: <\/strong>(also, dipole) molecule with an overall dipole moment<\/p>\n<p><strong>pure covalent bond:\u00a0<\/strong>(also, nonpolar covalent bond) covalent bond between atoms of identical electronegativities<\/p>\n","protected":false},"author":6181,"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-1682","chapter","type-chapter","status-publish","hentry"],"part":352,"_links":{"self":[{"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-introductorychemistry\/wp-json\/pressbooks\/v2\/chapters\/1682","targetHints":{"allow":["GET"]}}],"collection":[{"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-introductorychemistry\/wp-json\/pressbooks\/v2\/chapters"}],"about":[{"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-introductorychemistry\/wp-json\/wp\/v2\/types\/chapter"}],"author":[{"embeddable":true,"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-introductorychemistry\/wp-json\/wp\/v2\/users\/6181"}],"version-history":[{"count":6,"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-introductorychemistry\/wp-json\/pressbooks\/v2\/chapters\/1682\/revisions"}],"predecessor-version":[{"id":1706,"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-introductorychemistry\/wp-json\/pressbooks\/v2\/chapters\/1682\/revisions\/1706"}],"part":[{"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-introductorychemistry\/wp-json\/pressbooks\/v2\/parts\/352"}],"metadata":[{"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-introductorychemistry\/wp-json\/pressbooks\/v2\/chapters\/1682\/metadata\/"}],"wp:attachment":[{"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-introductorychemistry\/wp-json\/wp\/v2\/media?parent=1682"}],"wp:term":[{"taxonomy":"chapter-type","embeddable":true,"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-introductorychemistry\/wp-json\/pressbooks\/v2\/chapter-type?post=1682"},{"taxonomy":"contributor","embeddable":true,"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-introductorychemistry\/wp-json\/wp\/v2\/contributor?post=1682"},{"taxonomy":"license","embeddable":true,"href":"https:\/\/courses.lumenlearning.com\/suny-mcc-introductorychemistry\/wp-json\/wp\/v2\/license?post=1682"}],"curies":[{"name":"wp","href":"https:\/\/api.w.org\/{rel}","templated":true}]}}