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2. Polar Covalent Bonds: Acids and Bases

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  1. 2. Polar Covalent Bonds: Acids and Bases Why this chapter? Description of basic ways chemists account for chemical reactivity. Establish foundation for understanding specific reactions discussed in subsequent chapters.

  2. 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

  3. 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

  4. The Periodic Table and Electronegativity

  5. 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

  6. Polarity, Shape CO2 Electronegativity Difference < 0.5 Nonpolar 0.5-1.7 Polar >1.8 Ionic O C O 3.5 2.5 3.5 O=C=O d- d+ d- Polar Covalent Bonds Linear Shape (180o) Nonpolar Compound

  7. Polarity, Shape e-’s in 2 directions = 180o O=C=O Linear d- d+ d- Nonpolar Compound e-’s in 3 directions = 120o d- Trigonal planar Polar Compound d+

  8. F B F F Polarity, Shape BF3 d- 4.0 d+ F 2.0 d- d- F F B 4.0 4.0 (120o) Trigonal Planar Polar Covalent Bonds Nonpolar Compound

  9. H H C H H C H H Cl Cl O H H H-O-H Polarity, Shape e-’s in 4 directions = 109.5o d+ d- Tetrahedral d+ 4 directions = 109.5o d- d+ Bent

  10. N N H H H H H H Polarity, Shape e-’s in 4 directions = 109.5o d- d+ d+ d+ Pyramidal (109.5o) Tetrahedral Configuration of Electrons Trigonal Pyramid Configuration of Atoms

  11. Tetrahedral electron-pair Geometries Tetrahedral Pyramidal Bent

  12. 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

  13. 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.

  14. 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

  15. 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

  16. Which of the above molecules is the most polar? Learning Check: • A • B • C • D • It cannot be determined

  17. Which of the above molecules is the most polar? Solution: • A • B • C • D • It cannot be determined

  18. Learning Check: Which of the following molecules is nonpolar? • H2O • CH3CN • CHCl3 • CH2Cl2 • CO2

  19. Solution: Which of the following molecules is nonpolar? • H2O • CH3CN • CHCl3 • CH2Cl2 • CO2

  20. Which of the above molecules has (have) a dipole moment? Learning Check: • A only • B only • C only • A and B • B and C

  21. Which of the above molecules has (have) a dipole moment? Solution: • A only • B only • C only • A and B • B and C

  22. 2.3 Formal Charges • Sometimes it is necessary to have structures with formal charges on individual atoms • We compare the bonding of the atom in the molecule to the valence electron structure • 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

  23. Formal Charge for Dimethyl Sulfoxide • Atomic sulfur has 6 valence electrons. Dimethyl suloxide sulfur has only 5. • It has lost an electron and has positive charge. • Oxygen atom in DMSO has gained electron and has (-) charge.

  24. What are the hybridization and a formal charge, respectively, on boron in BH4–? Learning Check: • sp3, 0 • sp3, –1 • sp3, +1 • sp2, –1 • sp2, +1

  25. What are the hybridization and a formal charge, respectively, on boron in BH4–? Solution: • sp3, 0 • sp3, –1 • sp3, +1 • sp2, –1 • sp2, +1

  26. 2.4 Resonance • Some molecules are have 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) • Such a structure is delocalized and is represented by resonance forms • The resonance forms are connected by a double-headed arrow

  27. 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 • 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

  28. 2.5 Rules for Resonance Forms • Individual resonance forms are imaginary - the real structure is a hybrid (only by knowing the contributors can you visualize the actual structure) • Resonance forms differ only in the placement of their  or nonbonding electrons • Different resonance forms of a substance don’t have to be equivalent • Resonance forms must be valid Lewis structures: the octet rule applies • The resonance hybrid is more stable than any individual resonance form would be

  29. Curved Arrows and Resonance Forms • We can imagine that electrons move in pairs to convert from one resonance form to another • A curved arrow shows that a pair of electrons moves from the atom or bond at the tail of the arrow to the atom or bond at the head of the arrow

  30. 2.6 Drawing Resonance Forms • Any three-atom grouping with a multiple bond has two resonance forms

  31. Different Atoms in Resonance Forms • Sometimes resonance forms involve different atom types as well as locations • The resulting resonance hybrid has properties associated with both types of contributors • The types may contribute unequally • The “enolate” derived from acetone is a good illustration, with delocalization between carbon and oxygen

  32. 2,4-Pentanedione • The anion derived from 2,4-pentanedione • Lone pair of electrons and a formal negative charge on the central carbon atom, next to a C=O bond on the left and on the right • Three resonance structures result

  33. Which of the following statements is false concerning resonance structures? Learning Check: • The arrangement of nuclei in all contributing structures must be the same. • The arrangement of electrons in each contributing structure is different. • Each atom in a contributing structure must have a completed valence shell. • The contributing structures may have different energies. • All contributing structures must have the correct number of valence electrons.

  34. Which of the following statements is false concerning resonance structures? Solution: • The arrangement of nuclei in all contributing structures must be the same. • The arrangement of electrons in each contributing structure is different. • Each atom in a contributing structure must have a completed valence shell. • The contributing structures may have different energies. • All contributing structures must have the correct number of valence electrons.

  35. Consider all equivalent resonance structures for the carbonate anion shown below. If the bond order is defined as the number of electron pairs shared between two atoms, what is the average bond order for the C-O1 bond? Learning Check: • 2 • 3/2 • 4/3 • 1 • 1/4

  36. Consider all equivalent resonance structures for the carbonate anion shown below. If the bond order is defined as the number of electron pairs shared between two atoms, what is the average bond order for the C-O1 bond? Solution: • 2 • 3/2 • 4/3 • 1 • 1/4

  37. Consider all equivalent resonance structures for carbonate anion shown below. What is the formal charge on O1? Learning Check: • –1 • –2/3 • –1/2 • –1/3 • 0

  38. Consider all equivalent resonance structures for carbonate anion shown below. What is the formal charge on O1? Solution: • –1 • –2/3 • –1/2 • –1/3 • 0

  39. Anions with the negative charge residing on carbons are called carbanions. In the methyl anion the carbon is sp3 hybridized. What is the hybridization of carbon 1 in the allyl anion? Learning Check: • sp3 • sp2 • sp • sp3d • carbon 1 is not hybridized

  40. Anions with the negative charge residing on carbons are called carbanions. In the methyl anion the carbon is sp3 hybridized. What is the hybridization of carbon 1 in the allyl anion? Solution: • sp3 • sp2 • sp • sp3d • carbon 1 is not hybridized

  41. How many uncharged resonance structures exist for the following molecule? Learning Check: • Only the one shown above • 2 • 3 • 4 • 5

  42. How many uncharged resonance structures exist for the following molecule? Solution: • Only the one shown above • 2 • 3 • 4 • 5

  43. 2.7 Acids and Bases: The Brønsted–Lowry Definition • The terms “acid” and “base” can have different meanings in different contexts • For that reason, we specify the usage with more complete terminology • The idea that acids are solutions containing a lot of “H+” and bases are solutions containing a lot of “OH-” is not very useful in organic chemistry • Instead, Brønsted–Lowry theory defines acids and bases by their role in reactions that transfer protons (H+) between donors and acceptors

  44. Bronsted-Lowry theory Acid = a Proton donor donates a proton, H+, to solvent Example: HCl HCl + H2O H3O++ Cl- hydronium ion

  45. Bronsted-Lowry theory Example: HNO3 HNO3 + H2O H3O+ + NO3- HNO3 is the acid: it donates the proton.

  46. Bronsted-Lowry theory Base =a proton acceptor accepts a Proton, H+, from solvent NH3 + H2O NH4+ + OH- Accepted H+ NH3 is the Base: it takes the proton. Hydroxide ions (OH-) form when water donates the proton.

  47. CommonAcids • H2SO4 • HCl • H3PO4 • H2CO3 • HC2H3O2 • H3C6O7H8 • HC6O6H7 • H2C4O6H4 • H2C9O4H8 • Sulfuric Acid • Hydrochloric Acid • Phosphoric Acid • Carbonic Acid • Acetic Acid • Citric Acid • Ascorbic Acid • Tartaric Acid • Acetyl Salicylic Acid • Battery Acid • Stomach Acid • Coca Cola • Carbonated Water • Vinegar • Citrus fruits • Vitamin C • Grapes • Aspirin

  48. Naming Acids • Binary Acids = (-ideending) Hydro- ______-ic Acid HCl = Hydrogen Chloride = Hydrochloric Acid • Oxoacids = H, Nonmetal & O -ate ending (The higher number of O’s) ______________-icAcid HNO3 = Hydrogen Nitrate = NitricAcid -ite ending (The lower number of O’s) ______________-ousAcid HNO2 = Hydrogen Nitrite = NitrousAcid

  49. Naming Acids Dihydrogen Sulfate SulfuricAcid Dihydrogen Sulfite Sulfurous Acid Dihydrogen Sulfide Hydrosulfuric Acid Trihydrogen Phosphate PhosphoricAcid Hydrogen Bromide Hydrobromic Acid Dihydrogen carbonate CarbonicAcid Hydrogen Nitrite Nitrous Acid Hydrogen Nitrate NitricAcid HC2H3O2 Hydrogen Acetate AceticAcid