Organic Chemistry 4th Edition Paula Yurkanis Bruice Chapter 1: Electronic Structure and Bonding Acids and Bases

1. Organic Chemistry
4th
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., H2)
• 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) = µ = 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
H F ICl Brone bond
Otwo bonds
Nthree bonds
Cfour bonds

15. The s Orbitals

16. The p Orbitals

17. Molecular Orbitals
• Molecular orbitals belong to the whole molecule
• σ 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 (σ) is formed by end-on overlap of two
p orbitals
A σ bond is stronger than a π bond

21. Pi bond (π) 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 sp2
carbon is 120°
• The sp2
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 σ bond and two π 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 π bond is weaker than a σ 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
Ka: The acid dissociation constant
H2O + HA H3O+
+ A-
[H3O+
][A-
]
Ka =
[H2O][HA]
pKa = -log Ka

38. The Henderson–Hasselbalch Equation
• A compound will exist primarily in its acidic form at a pH
< its pKa
• A compound will exist primarily in its basic form at a pH
> its pKa
• A buffer solution maintains a nearly constant pH upon
addition of small amount of acid or base
[ ]
[ ]−
+=
A
HA
logpHp aK
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
CH3COH
O
CH3CH2OH
pKa = 4.76 pKa = 15.9
acetic acid ethanol
The delocalized electrons in acetic acid are shared by
more than two atoms, thereby stabilizing the conjugated
base
CH3CO-
O
CH3CO-
O

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