1. Organic Chemistry 4 th Edition Paula Yurkanis Bruice Chapter 1 Electronic Structure and Bonding Acids and Bases Irene Lee Case Western Reserve University Cleveland, OH ©2004, Prentice Hall

2. Organic compounds are compounds containing carbon Carbon neither readily gives up nor readily accepts electrons Carbon shares electrons with other carbon atoms as well as with several different kinds of atoms Organic Chemistry

3. The Structure of an Atom An atom consists of electrons, positively charged protons, and neutral neutrons Electrons form chemical bonds Atomic number: numbers of protons in its nucleus Mass number: the sum of the protons and neutrons of an atom Isotopes have the same atomic number but different mass numbers The atomic weight: the average weighted mass of its atoms Molecular weight: the sum of the atomic weights of all the atoms in the molecule

4. The Distribution of Electrons in an Atom Quantum mechanics uses the mathematical equation of wave motions to characterize the motion of an electron around a nucleus Wave functions or orbitals tell us the energy of the electron and the volume of space around the nucleus where an electron is most likely to be found The atomic orbital closer to the nucleus has the lowest energy Degenerate orbitals have the same energy

5. Table 1.1

7. The Aufbau principle: electrons occupy the orbitals with the lowest energy first The Pauli exclusion principle: only two electrons can occupy one atomic orbital and the two electrons have opposite spin Hund’s rule: electrons will occupy empty degenerated orbitals before pairing up in the same orbital

8. Ionic compounds are formed when an electropositive element transfers electron(s) to an electronegative element

9. Covalent Compounds Equal sharing of electrons: nonpolar covalent bond (e.g., H 2 ) Sharing of electrons between atoms of different electronegativities: polar covalent bond (e.g., HF)

10. Electrostatic Potential Maps

11. A polar bond has a negative end and a positive end dipole moment (D) =  = e x d (e) : magnitude of the charge on the atom (d) : distance between the two charges A Dipole

12. Lewis Structure Formal charge = number of valence electrons – (number of lone pair electrons +1/2 number of bonding electrons)

14. Important Bond Numbers HFIClBr one bond O two bonds N three bonds C four bonds

15. The s Orbitals

16. The p Orbitals

17. Molecular Orbitals Molecular orbitals belong to the whole molecule  bond: formed by overlapping of two s orbitals Bond strength/bond dissociation: energy required to break a bond or energy released to form a bond

19. In-phase overlap forms a bonding MO; out-of-phase overlap forms an antibonding MO

20. Sigma bond (  ) is formed by end-on overlap of two p orbitals A  bond is stronger than a  bond

21. Pi bond (  ) is formed by sideways overlap of two parallel p orbitals

22. Bonding in Methane and Ethane: Single Bonds Hybridization of orbitals:

23. The orbitals used in bond formation determine the bond angles Tetrahedral bond angle: 109.5 ° Electron pairs spread themselves into space as far from each other as possible

24. Hybrid Orbitals of Ethane

25. Bonding in Ethene: A Double Bond

26. The bond angle in the sp 2 carbon is 120 ° The sp 2 carbon is the trigonal planar carbon An sp-Hybridized Carbon

27. Bonding in Ethyne: A Triple Bond Bond angle of the sp carbon: 180 ° A triple bond consists of one  bond and two  bonds

28. Bonding in the Methyl Cation

29. Bonding in the Methyl Radical

30. Bonding in the Methyl Anion

31. Bonding in Water

32. Bonding in Ammonia and in the Ammonium Ion

33. Bonding in Hydrogen Halides

34. Summary A  bond is weaker than a  bond The greater the electron density in the region of orbital overlap, the stronger is the bond The more s character, the shorter and stronger is the bond The more s character, the larger is the bond angle

35. The vector sum of the magnitude and the direction of the individual bond dipole determines the overall dipole moment of a molecule Molecular Dipole Moment

36. Brønsted–Lowry Acids and Bases Acid donates a proton Base accepts a proton Strong reacts to give weak The weaker the base, the stronger is its conjugate acid Stable bases are weak bases

37. An Acid/Base Equilibrium K a : The acid dissociation constant

38. The Henderson–Hasselbalch Equation A compound will exist primarily in its acidic form at a pH < its pK a A compound will exist primarily in its basic form at a pH > its pK a A buffer solution maintains a nearly constant pH upon addition of small amount of acid or base The pH indicates the concentration of hydrogen ions (H + )

39. When atoms are very different in size, the stronger acid will have its proton attached to the largest atom

40. When atoms are similar in size, the stronger acid will have its proton attached to the more electronegative atom Inductive electron withdrawal increases the acidity of a conjugate acid

41. Acetic acid is more acidic than ethanol The delocalized electrons in acetic acid are shared by more than two atoms, thereby stabilizing the conjugated base

42. Lewis acid: non-proton-donating acid; will accept two electrons Lewis base: electron pair donors Lewis Acids and Bases