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2.1 Polar Covalent Bonds: Electronegativity

Covalent bonds can have ionic character These are polar covalent bonds Bonding electrons attracted more strongly by one atom than by the other Electron distribution between atoms is not symmetrical. 2.1 Polar Covalent Bonds: Electronegativity. Bond Polarity and Electronegativity.

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2.1 Polar Covalent Bonds: Electronegativity

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  1. Covalent bonds can have ionic character These are polar covalent bonds Bonding electrons attracted more strongly by one atom than by the other Electron distribution between atoms is not symmetrical 2.1 Polar Covalent Bonds: Electronegativity

  2. Bond Polarity and Electronegativity • Electronegativity (EN): intrinsic ability of an atom to attract the shared electrons in a covalent bond • Differences in EN produce bond polarity • Arbitrary scale. As shown in Figure 2.2, electronegativities are based on an arbitrary scale • F is most electronegative (EN = 4.0), Cs is least (EN = 0.7) • Metals on left side of periodic table attract electrons weakly, lower EN • Halogens and other reactive nonmetals on right side of periodic table attract electrons strongly, higher electronegativities • EN of C = 2.5

  3. The Periodic Table and Electronegativity

  4. Bond Polarity and Inductive Effect • Nonpolar Covalent Bonds: atoms with similar EN • Polar Covalent Bonds: Difference in EN of atoms < 2 • Ionic Bonds: Difference in EN > 2 • C–H bonds, relatively nonpolar C-O, C-X bonds (more electronegative elements) are polar • Bonding electrons toward electronegative atom • C acquires partial positive charge, + • Electronegative atom acquires partial negative charge, - • Inductive effect: shifting of electrons in a bond in response to EN of nearby atoms

  5. Electrostatic Potential Maps • Electrostatic potential maps show calculated charge distributions • Colors indicate electron-rich (red) and electron-poor (blue) regions • Arrows indicate direction of bond polarity

  6. 2.2 Polar Covalent Bonds: Dipole Moments • Molecules as a whole are often polar from vector summation of individual bond polarities and lone-pair contributions • Strongly polar substances soluble in polar solvents like water; nonpolar substances are insoluble in water. • Dipole moment () - Net molecular polarity, due to difference in summed charges •  - magnitude of charge Q at end of molecular dipole times distance r between charges •  = Qr, in debyes (D), 1 D = 3.336  1030 coulomb meter • length of an average covalent bond, the dipole moment would be 1.60  1029 Cm, or 4.80 D.

  7. Dipole Moments in Water and Ammonia • Large dipole moments • EN of O and N > H • Both O and N have lone-pair electrons oriented away from all nuclei

  8. Absence of Dipole Moments • In symmetrical molecules, the dipole moments of each bond has one in the opposite direction • The effects of the local dipoles cancel each other

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