1. Introduction to Organic Chemistry

2. What is Organic Chemistry?
Living things are made of organic chemicals (carbon-based compounds)
Proteins that make up hair
DNA, controls genetic make-up
Foods, medicines

3. Origins of Organic Chemistry
Foundations of organic chemistry from mid-1700’s.
Compounds obtained from plants, animals hard to isolate, and purify.
Compounds also decomposed more easily.
Torben Bergman (1770) first to make distinction between organic and inorganic chemistry.
It was thought that organic compounds must contain some “vital force” because they were from living sources.

4. Origins of Organic Chemistry
Because of “vital force”, it was thought that organic compounds could not be synthesized in laboratory like inorganic compounds.
1816, Chevreul showed that not to be the case, he could prepare soap from animal fat and an alkali and glycerol is a product
1828, Woehler showed that it was possible to convert inorganic salt ammonium cyanate into organic substance “urea”

5. Origins of Organic Chemistry
Organic chemistry is study of carbon compounds.
Why is it so special?
90% of more than 30 million chemical compounds contain carbon.
Examination of carbon in periodic chart answers some of these questions.
Carbon is group 4A element, it can share 4 valence electrons and form 4 covalent bonds.

6. Isomers

7. Abundance of Organic Compounds
Why are there so many more organic compounds than inorganic?
Carbon has unique bonding characteristics
Strong, covalent bonds with C and H
Isomerism
Groups of carbon atoms can form more than one unique compound

8. Electrons, Bonds, Lewis Structures

9. Development of Chemical Bonding Theory
Kekulé and Couper independently observed that carbon always has four bonds
van't Hoff and Le Bel proposed that the four bonds of carbon have specific spatial directions
Atoms surround carbon as corners of a tetrahedron

10. Development of Chemical Bonding Theory
Atoms form bonds because the compound that results is more stable than the separate atoms
Ionic bonds in salts form as a result of electron transfers
Organic compounds have covalent bonds from sharing electrons (G. N. Lewis, 1916)

11. Development of Chemical Bonding Theory
Atoms with one, two, or three valence electrons form one, two, or three bonds.
Atoms with four or more valence electrons form as many bonds as they need electrons to fill the s and p levels of their valence shells to reach a stable octet.
Carbon has four valence electrons (2s2 2p2), forming four bonds (CH4).

12. Development of Chemical Bonding Theory
Nitrogen has five valence electrons (2s2 2p3) but forms only three bonds (NH3).
Oxygen has six valence electrons (2s2 2p4) but forms two bonds (H2O)

13. Development of Chemical Bonding Theory

14. Non-Bonding Electrons
Valence electrons not used in bonding are called nonbonding electrons, or lone-pair electrons
Nitrogen atom in ammonia (NH3)
Shares six valence electrons in three covalent bonds and remaining two valence electrons are nonbonding lone pair

15. Formal Charge

16. Formal Charges Neutral molecules with both a “+” and a “-” are dipolar
Partial Charge vs. Formal Charge
Partial charge is a real value
Formal charge may or may not correspond to a real charge
Atoms with FC usually bear at least partial charge ( positive or negative)
FC helps us determine overall charge distribution and is useful for understanding reaction mechanisms
Neutral molecules with both a “+” and a “-” are dipolar

17. How to Determine FC
FC = [# of valence e-] – [non-bonding e-] – [shared e-/2]

18. How to Determine FC
FC = [# of valence e-] – [non-bonding e-] – [shared e-/2]

19. How to Determine FC
FC = [# of valence e-] – [non-bonding e-] – [shared e-/2]

20. Formal Charge for Dimethyl Sulfoxide
Atomic sulfur has 6 valence electrons.
Dimethyl sulfoxide sulfur has only 5.
It has lost an electron and has positive charge.
Oxygen atom in DMSO has gained electron and has negative charge.

21. Formal Charges (Continued)

22. Induction, Polar Covalent Bonds, Electronegativity

23. 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

24. 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
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

25. The Periodic Table and Electronegativity

26. Bond Polarity and Inductive Effect
Nonpolar Covalent Bonds: atoms with similar EN
Polar Covalent Bonds: Difference in EN of atoms < 1.7
Ionic Bonds: Difference in EN > 1.7
C–H bonds, relatively nonpolar C-O, C-X bonds (more electronegative elements) are polar
When C bonds with more EN 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

27. 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

28. Atomic Orbitals

29. Atomic Structure Structure of an atom
Positively charged nucleus (very dense, protons and neutrons) and small (10-15 m)
Negatively charged electrons are in a cloud (10-10 m) around nucleus
Diameter is about 2  m (200 picometers (pm)) [the unit ångström (Å) is m = 100 pm]

30. Shells, Subshells, Orbitals
The number of subshells in a shell = shell number
The first subshell s has 1 orbital. Each successive subshell adds 2 more orbitals (1, 3, 5, 7, etc).
Each orbital can hold only 2 electrons of opposite spin.
An atom with n = 3 also includes all subshells and orbitals for n < 3:
1s, 2s, 2p, 3s, 3p, 3d

31. Atomic Structure: Orbitals
Quantum mechanics: describes electron energies and locations by a wave equation
Wave function solution of wave equation
Each wave function is an orbital, ψ
A plot of ψ describes where electron most likely to be
Electron cloud has no specific boundary so we show most probable area, i.e., this is a probability function.

32. Shapes of Atomic Orbitals for Electrons
Four different kinds of orbitals for electrons based on those derived for a hydrogen atom
Denoted s, p, d, and f
s and p orbitals most important in organic and biological chemistry
s orbitals: spherical, nucleus at center
p orbitals: dumbbell-shaped, nucleus at middle
d orbitals: elongated dumbbell-shaped, nucleus at center

33. Orbitals and Shells (Continued)
Orbitals are grouped in shells of increasing size and energy
Different shells contain different numbers and kinds of orbitals
Each orbital can be occupied by two electrons

34. Orbitals and Shells (Continued)
First shell contains one s orbital, denoted 1s, holds only two electrons
Second shell contains one s orbital (2s) and three p orbitals (2p), eight electrons
Third shell contains an s orbital (3s), three p orbitals (3p), and five d orbitals (3d), 18 electrons

35. P-Orbitals
In each shell there are three perpendicular p orbitals, px, py, and pz, of equal energy
Lobes of a p orbital are separated by region of zero electron density, a node

36. Bonding Characteristics of Carbon
Q: If 2s electrons are already paired, with only 2 2p electrons unpaired, how does carbon form 4 covalent bonds? C 2p
Valence shell electrons 2s 1s

37. How are electrons arranged?
Aufbau Principle
Electrons fill orbitals starting at the lowest available (possible) energy states before filling higher states (e.g. 1s before 2s).
Sometimes a low energy subshell has lower energy than upper subshell of preceding shell (e.g., 4s fills before 3d).
Pauli exclusion principle
QM principle: no two identical fermions (particles with half-integer spin) may occupy the same quantum state simultaneously (why paired electrons have different spin).
Hund's rule
Every orbital in a subshell is singly occupied with one electron before any one orbital is doubly occupied, and all electrons in singly occupied orbitals have the same spin.

38. Valence Bond Theory

39. Describing Chemical Bonds: Valence Bond Theory
Covalent bond forms when two atoms approach each other closely so that a singly occupied orbital on one atom overlaps a singly occupied orbital on the other atom
Two models to describe covalent bonding.
Valence Bond Theory:
Electrons are paired in the overlapping orbitals and are attracted to nuclei of both atoms
H–H bond results from the overlap of two singly occupied hydrogen 1s orbitals
H-H bond is cylindrically symmetrical, sigma (s) bond

40. Bond Energy Reaction 2 H·  H2 releases 436 kJ/mol
Product has 436 kJ/mol less energy than two atoms: H–H has bond strength of 436 kJ/mol. (1 kJ = kcal; 1 kcal = kJ)

41. Bond Energy Distance between nuclei that leads to maximum stability
If too close, they repel because both are positively charged
If too far apart, bonding is weak

42. Dipole Moments, Molecular Polarity

43. 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 are 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 =  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.

44. 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

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

46. Acids and Bases, pKa

47. Acid/Base Definitions
Brønsted-Lowry Acids and Bases
Acid donates proton
Base accept proton
Conjugate acid-base pair
acid
base
acid
base
Lewis Acids and Bases
Acid accepts electron pair
Base donates electron pair

48. Curve Arrow Notation, Reaction Mechanisms
Bond-Breaking
Bond-Making

49. The Reaction of Acid with Base
Hydronium ion, product when base H2O gains a proton
HCl donates a proton to water molecule, yielding hydronium ion (H3O+) [conjugate acid] and Cl [conjugate base]
The reverse is also a Brønsted acid–base reaction of the conjugate acid and conjugate base

50. Acidity/Basicity H+ OH- pH = -log[H+]
acidic
neutral
basic 7 14 pH
pH = -log[H+]
The pH of solution determines form of carboxylic acid
Ex. Carboxylate ion predominates at pH 7.4 (physiological pH)
Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7th Edition, 2011

51. Acid and Base Strength
Strong acids/bases: Dissociate completely in water
Weak acids/bases: Dissociate incompletely in water
Strength of acid can be related to acid dissociation constant (Ka)
Stronger acids have larger Ka, lower p Ka values.
Ka ranges from 1015 for the strongest acids to very small values (10-60) for the weakest
[H3O+], [B-] dissociated acid components
[HB] undissociated acid

52. pKa’s of Some Common Acids

53. pKa – the Acid Strength Scale
pKa = –log Ka
The free energy in an equilibrium is related to –log of Keq (DG = –RT log Keq)
A smaller value of pKa indicates a stronger acid and is proportional to the energy difference between products and reactants
The pKa of water is 15.74

54. Predicting Acid–Base Reactions from pKa Values
pKa values are related as logarithms to equilibrium constants
Useful for predicting whether a given acid-base reaction will take place
The difference in two pKa values is the log of the ratio of equilibrium constants, and can be used to calculate the extent of transfer
The stronger base holds the proton more tightly

55. Molecular Models
In Class Assignment 01

56. 2.12 Molecular Models Organic chemistry is 3-D space
Molecular shape is critical in determining the chemistry a compound undergoes in the lab, and in living organisms

57. 2.12 Molecular Models
Build the following compounds with your molecular modeling kit and look at the geometry:
Hexane
2-methylhexane
Benzene
ethyne