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Chapter 2

Chapter 2. Polar Covalent Bonds; Acids and Bases. Introduction.

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Chapter 2

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  1. Chapter 2 Polar Covalent Bonds; Acids and Bases

  2. Introduction • The Lewis Model of Bonding: Atoms bond together in such a way that each atom participating in a chemical bond acquires a completed valence-shell electron configuration resembling that of the noble gas nearest it in the Periodic Table.

  3. Ionic Bond – is a chemical bond resulting from the electrostatic attraction between an anion and a cation. .. . .. .. .. . .. Na + F Na++ F- .. ..

  4. Covalent Bond – is a chemical bond formed by the sharing of electron pairs between atoms. 2 H· H2 • Based on the the degree of electron sharing, covalent bonds can be divided into: • Nonpolar covalent bond • Polar covalent bond

  5. I. Polar Covalent Bonds • Electronegativity • Dipole Moments

  6. Polar Covalent Bond – is a covalent bond with ionic character. • Bonding electrons are attracted more strongly by one atom than by the other • Electron distribution between atoms is not symmetrical

  7. A. Polar Covalent Bonds: Electronegativity • Electronegativity(EN): • is a measure of the force of an atom’s attraction for electrons it shares in a covalent bond with another atom. • is an intrinsic ability of an atom to attract the shared electrons in a covalent bond

  8. Bond Polarity and Electronegativity • EN of an atom is related to its ionization energy (IE) and electron affinity (EA). • Atom in ground state (g) + Energy  Atom+ (g) + e- DE = IE • Atom (g) + e- Atom-DE = EA • Differences in EN produce bond polarity. DEN SEN Bond Polarity ~DEN ~

  9. Bond Polarity and Electronegativity • EN is based on an arbitrary scale. • The most widely used scale of EN was devised by Linus Paulingin the 1930’s and is based on bond energies. • Important EN values: • F is the most electronegative (EN = 4.0) • Cs is the least electronegative (EN = 0.7) • C has an EN = 2.5

  10. The Periodic Table and Electronegativity

  11. The Periodic Table and Electronegativity • Metals on left side of periodic table attract electrons weakly, lower EN

  12. The Periodic Table and Electronegativity • Halogens and other reactive nonmetals on right side of periodic table attract electrons strongly, higher EN

  13. Classification of Covalent Bonds

  14. Classification of Covalent Bonds

  15. Classification of Covalent Bonds: Examples • C–H bonds are relatively nonpolar DEN = ENC – ENH = 2.5 - 2.1 = 0.4 • C-O and C-X bonds (more electronegative elements) are polar DEN = ENO – ENC = 3.5 - 2.5 = 1

  16. Bond Polarity and Inductive Effect • Bonding electrons are drawn toward electronegative atom • C acquires partial positive charge, + • Electronegative atom acquires partial negative charge, - The crossed arrow indicates the direction of the electron displacement

  17. Bond Polarity and Inductive Effect • Inductive Effect – is the shifting of electrons in a bond in response to EN of nearby atoms • Metals (Li and Mg) inductively donate e- • Reactive nonmetals (O and Cl) inductively withdraw e-

  18. Electrostatic Potential Maps • Electrostatic potential maps show calculated charge distributions • Colors indicate electron-rich (red) and electron-poor (blue) regions

  19. Practice Problem: Which element in each of the following pairs is more electronegative? • Li or H • B or Br • Cl or I • C or H

  20. Practice Problem: Use the d+/d- convention to indicate the direction of expected polarity for each of the bonds indicated • H3C — OH • H3C — MgBr • H3C — F • H3C — Br • H3C — NH2 • H3C — Li • H2N — H

  21. Practice Problem: Use the electronegativity values to rank the following bonds from least polar to most polar: • H3C — Li • H3C — K • H3C — F • H3C — MgBr • H3C — OH

  22. Practice Problem: Look at the following electrostatic potential map of methyl alcohol, and tell the direction of polarization of the C-O bond:

  23. B. Polar Covalent Bonds: Dipole Moments • Molecular Polarity • is the tendency of molecules as a whole to be polar • results from vector summation of individual bond polarities and lone-pair contributions • “Like dissolves like” • Strongly polar substances are soluble in polar solvents like water; nonpolar substances are insoluble in water.

  24. To predict whether a molecule is polar, determine: • if the molecule has polar bonds, and • the arrangements of these bonds in space

  25. Dipole moment (m)– is a measure of net molecular polarity, due to difference in summed charges = Q r magnitude of charge Q at either end of molecular dipole distance r between charges • It is expressed in Debyes (D) • 1 D = 3.336 x 10-30 Coulomb meter

  26. Calculating the dipole moment of an average bond: raverage covalent bond= 100pm Qelectron = 1.60 x 10-19 C The dipole moment (m) of an average covalent bond is 4.80 D

  27. Calculating Ionic Character: Chloromethane Given that r C-Cl = 178 pm and assuming C-H is negligible, then calculatedCH3Cl = 178 pm x 4.80 D = 8.5 D measured CH3Cl = 1.87D % ionic character = measured / calculated x 100= 22%

  28. Dipole Moments in Water and Ammonia • H2O and NH3 have large dipole moments: • ENO and ENN > ENH • Both O and N have lone-pair electrons oriented away from all nuclei

  29. Dipole Moments in Water and Ammonia • H2O and NH3 have large dipole moments: • ENO and ENN > ENH • Both O and N have lone-pair electrons oriented away from all nuclei

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

  31. Practice Problem: Carbon dioxide, CO2, has zero dipole moment even though carbon-oxygen bonds are strongly polarized. Explain.

  32. Practice Problem: Make three-dimensional drawings of the following molecules, and predict whether each has a dipole moment. If you expect a dipole moment, show its direction. • H2C = CH2 • CHCl3 • CH2Cl2 • H2C = CCl2

  33. II. Formal Charges and Resonance • Formal Charges • Resonance

  34. A. Formal Charges • Formal Charge - is the charge on an atom in a molecule or polyatomic ion

  35. Comparing the bonding of the atom in the molecule to the valence electron structure No formal charge No formal charge

  36. Calculating the formal charge

  37. If the atom has one more electron in the molecule, it is shown with a “-” charge • If the atom has one less electron, it is shown with a “+” charge • Neutral molecules with both a “+” and a “-” are dipolar

  38. Practice Problem: Dimethyl sulfoxide, a common solvent, has the structure indicated. Show why dimethyl sulfoxide must have formal charges on S and O.

  39. Practice Problem: Calculate formal charges for the nonhydrogen atoms in the following molecules: • Diazomethane, • Acetonitrile oxide, • Methyl isocyanide

  40. Practice Problem: Organic phosphates occur commonly among biological molecules. Calculate formal charges on the four O atoms in the methyl phosphate ion.

  41. B. Resonance • Some molecules have Lewis structures that cannot be shown with a single representation • In these cases we draw structures that contribute to the final structure but which differ in the position of the  bond(s) or lone pair(s)

  42. Resonance Forms • Resonance forms – are Lewis structures of the same molecule whose only difference is the placement of p and nonbonding valence electrons (= delocalized). • The atoms occupy the same place in the different forms • The connections between atoms are the same. • The resonance forms are connected by a double-headed arrow

  43. Resonance Hybrids • A structure with resonance forms does not alternate between the forms • Instead, it is a hybrid of the two resonance forms, so the structure is called a resonance hybrid

  44. Resonance Hybrids: Benzene • For example, benzene (C6H6) has two resonance forms with alternating double and single bonds • In the resonance hybrid, the actual structure, all its C-C bonds are equivalent, midway between double and single

  45. Rules for Resonance Forms1 • Individual resonance forms are imaginary- the real structure is a hybrid (only by knowing the contributors can you visualize the actual structure)

  46. Rules for Resonance Forms2 • Resonance forms differ only in the placement of their  or nonbonding electrons

  47. Rules for Resonance Forms3 • Different resonance forms of a substance don’t have to be equivalent

  48. Rules for Resonance Forms4 • Resonance forms must be valid Lewis structures: the octet rule applies

  49. Rules for Resonance Forms5 • The resonance hybrid is more stable than any individual resonance form • Resonance leads to stability. The larger the # of the resonance forms, the more stable the substance

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