CHAPTER ONE Chemical Bonding Because the charge on an electron is 4.80X 10 electrostatic units(esu) and tances within a molecule typically fall in the 10 cm range, molecular dipole are on the order of 10 8 esu- cm. In order to simplify the reporting of dipole this value of 10 esu' cm is defined as a debye, D. Thus the experimentally determined honor of Peter Debye, a dipole moment of hydrogen fluoride, 1.7 X 10esucm is stated as 1.7 Dutch scientist who did Table 1.3 lists the dipole moments of various bond types. For H-F, H-Cl, ortant work in many H-Br, and H-I these"bond dipoles"are really molecular dipole moments. A polar was awarded the Nobel Prize molecule has a dipole moment, a nonpolar one does not. Thus, all of the hydrogen halides are polar molecules. In order to be polar, a molecule must have polar bonds, bu can't have a shape that causes all the individual bond dipoles to cancel. We will have more to say about this in Section 1. ll after we have developed a feeling for the three- dimensional shapes of molecules The bond dipoles in Table 1.3 depend on the difference in electronegativity of the bonded atoms and on the bond distance. The polarity of a C-H bond is relatively low; substantially less than a C-O bond, for example. Don't lose sight of an even more important difference between a C-H bond and a C-O bond, and that is the direction of the dipole moment. In a C-H bond the electrons are drawn away from H, toward C In a C-o bond, electrons are drawn from C toward O. As we'll see in later chap ters, the kinds of reactions that a substance undergoes can often be related to the size and direction of key bond dipoles 1.6 FORMAL CHARGE Lewis structures frequently contain atoms that bear a positive or negative charge. molecule as a whole is neutral, the sum of its positive charges must equal the sum of its negative charges. An example is nitric acid, HNO As written, the structural formula for nitric acid depicts different bonding patterns for its three oxygens. One oxygen is doubly bonded to nitrogen, another is singly bonded TABLE 1.3 Selected Bond Dipole Moments Bond* Dipole moment, D Bond* Dipole moment, D H—F C-F H—C H—C C≡N 3.6 irection of the dipole moment is toward the more electronegative atom In the listed examples en and carbon are the positive ends of the dipoles. Carbon is the negative end of the dipole ted with the C-H bond Back Forward Main MenuToc Study Guide ToC Student o MHHE WebsiteBecause the charge on an electron is 4.80  1010 electrostatic units (esu) and the dis￾tances within a molecule typically fall in the 108 cm range, molecular dipole moments are on the order of 1018 esu·cm. In order to simplify the reporting of dipole moments this value of 1018 esu cm is defined as a debye, D. Thus the experimentally determined dipole moment of hydrogen fluoride, 1.7  1018 esu cm is stated as 1.7 D. Table 1.3 lists the dipole moments of various bond types. For H±F, H±Cl, H±Br, and H±I these “bond dipoles” are really molecular dipole moments. A polar molecule has a dipole moment, a nonpolar one does not. Thus, all of the hydrogen halides are polar molecules. In order to be polar, a molecule must have polar bonds, but can’t have a shape that causes all the individual bond dipoles to cancel. We will have more to say about this in Section 1.11 after we have developed a feeling for the three￾dimensional shapes of molecules. The bond dipoles in Table 1.3 depend on the difference in electronegativity of the bonded atoms and on the bond distance. The polarity of a C±H bond is relatively low; substantially less than a C±O bond, for example. Don’t lose sight of an even more important difference between a C±H bond and a C±O bond, and that is the direction of the dipole moment. In a C±H bond the electrons are drawn away from H, toward C. In a C±O bond, electrons are drawn from C toward O. As we’ll see in later chap￾ters, the kinds of reactions that a substance undergoes can often be related to the size and direction of key bond dipoles. 1.6 FORMAL CHARGE Lewis structures frequently contain atoms that bear a positive or negative charge. If the molecule as a whole is neutral, the sum of its positive charges must equal the sum of its negative charges. An example is nitric acid, HNO3: As written, the structural formula for nitric acid depicts different bonding patterns for its three oxygens. One oxygen is doubly bonded to nitrogen, another is singly bonded H±O±N O O   œ ± 16 CHAPTER ONE Chemical Bonding TABLE 1.3 Selected Bond Dipole Moments Bond* H±F H±Cl H±Br H±I H±C H±N H±O Dipole moment, D 1.7 1.1 0.8 0.4 0.3 1.3 1.5 Bond* C±F C±O C±N CœO CœN CPN Dipole moment, D 1.4 0.7 0.4 2.4 1.4 3.6 *The direction of the dipole moment is toward the more electronegative atom. In the listed examples hydrogen and carbon are the positive ends of the dipoles. Carbon is the negative end of the dipole associated with the C±H bond. The debye unit is named in honor of Peter Debye, a Dutch scientist who did im￾portant work in many areas of chemistry and physics and was awarded the Nobel Prize in chemistry in 1936. Back Forward Main Menu TOC Study Guide TOC Student OLC MHHE Website