1. Chapter 1 Structure and Bonding; Acids and Bases
Suggested Problems: 26a,b,28,32,33,39,40,42,44,47-8,51-2,53,56-62,71

2. What is Organic Chemistry?
Chemistry of carbon
Living things are made of organic chemicals
Proteins that make up hair, enzymes
DNA, controls genetic make-up
Foods – sugars, fats
Medicines, plastics

3. Origins of Organic Chemistry
Foundations of organic chemistry from mid-1700’s.
Compounds obtained from plants, animals were low melting - hard to isolate and purify, decomposed easily.
It was thought that organic compounds must contain some “vital force” because they were from living sources.
By the mid-1800’s it was apparent that there was no fundamental difference between organic and inorganic substances.
Friedrich Wohler – ammonium cyanate to urea

4. 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.

5. 1.1 Atomic Structure Structure of an atom
Positively charged nucleus (protons and neutrons) and small (10-15 m) – contains most of mass of atom
Negatively charged electrons are in a cloud (10-10 m) around nucleus – take up most of volume
Diameter is about 2  m (200 picometers (pm)) [the unit ångström (Å) is m = 100 pm]

6. Atomic Number and Atomic Mass
The atomic number (Z) is the number of protons in the atom's nucleus
The mass number (A) is the number of protons plus neutrons
All the atoms of a given element have the same atomic number
Isotopes are atoms of the same element that have different numbers of neutrons and therefore different mass numbers
The atomic mass (atomic weight) of an element is the weighted average mass in atomic mass units (amu) of an element’s naturally occurring isotopes

7. Orbitals Electrons are distributed in 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.
An orbital can be thought of as defining a region of space around the nucleus where the electron can most likely be found.

8. Shapes of Atomic Orbitals for Electrons
Four different kinds of orbitals each with a different shape
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

9. 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

10. 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

11. 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

12. 1.2 Atomic Structure: Electron Configurations
Ground-state electron configuration (lowest energy arrangement) of an atom lists orbitals occupied by its electrons. Rules:
1. Lowest-energy orbitals fill first: 1s  2s  2p  3s  3p  4s  3d
2. Electrons act as if they were spinning around an axis. Electron spin can have only two orientations, up  and down . Only two electrons can occupy an orbital, and they must be of opposite spin (Pauli exclusion principle).
3. If two or more empty orbitals of equal energy are available, one electron occupies each with spins parallel until all orbitals are half full (Hund's rule).

13. 1.2 Atomic Structure: Electron Configurations
Problem How many electrons does each of the following elements have in its outermost electron shell?
a) Potassium b) Calcium c) Aluminum
c) Group 3 – therefore 3
Problem 1.2 Give the ground-state electron configuration of the following elements:
a) Boron b) Phosphorus c) Oxygen d) Argon
b) 1s22s22p63s23p3

14. 1.3 Development of Chemical Bonding Theory
Kekulé and Couper independently observed that carbon always has four bonds (tetravalent)
van't Hoff and Le Bel proposed that the four bonds of carbon have specific spatial directions (3-dimensional)
Atoms surround carbon as corners of a tetrahedron

15. 1.3 Development of Chemical Bonding Theory
Problem 1.3 Draw a molecule of chloromethane, CH3Cl, using solid, wedged, and dashed lines to show its tetrahedral geometry.
Problem 1.4 Convert the molecule of ethane, C2H6, into a structure that uses wedged, normal, and dashed lines to represent three-dimensionality.

16. 1.4 The Nature of Chemical Bonds
Atoms form bonds because the compound that results is more stable and lower in energy 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)

17. The Nature of Chemical Bonds
Lewis structures (electron dot) show valence electrons of an atom as dots
Hydrogen has one dot, representing its 1s electron
Carbon has four dots (2s2 2p2)
Kekulé structures (line-bond structures) have a line drawn between two atoms indicating a 2 electron covalent bond.
An octet of electrons in valence shell is especially stable – octet rule (two for hydrogen)
Atoms strive for noble gas configuration

18. The Nature of Chemical Bonds
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 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).

19. The Nature of Chemical Bonds
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)

20. The Nature of Chemical Bonds

21. 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

22. Non-Bonding Electrons
Problem 1.5 What are likely formulas for the following molecules?
d) SiF?
Problem 1.6 Write both electron-dot and line-bond structures for the following molecules, showing all nonbonded electrons.
c) CH3NH2, methylamine

23. 1.5 Forming Covalent 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
Valence bond theory: model 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

24. 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)
It takes energy to break a bond; energy is released in forming it.

25. Bond Energy
If too close, nuclei repel each other because both are positively charged
If too far apart, bonding is weak
Bond length –distance between nuclei that leads to maximum stability

26. 1.6 sp3 Orbitals and the Structure of Methane
Carbon has 4 valence electrons (2s2 2p2)
In CH4, all C–H bonds are identical (tetrahedral)
sp3 hybrid orbitals: s orbital and three p orbitals combine to form four equivalent, unsymmetrical, tetrahedral orbitals (sppp = sp3), Pauling (1931)

27. The Structure of Methane
sp3 orbitals on C overlap with 1s orbitals on 4 H atoms to form four identical C-H bonds
Each C–H bond has a strength of 439 kJ/mol and length of 109 pm
Bond angle: each H–C–H is 109.5º, the tetrahedral angle.

28. The Carbon in Methane is sp3
Carbon is tetrahedral.
The tetrahedral bond angle is 109.5°.

29. The Structure of Methane
Problem 1.8 Draw a tetrahedral representation of tetrachloromethane, CCl4, using the standard convention of solid, dashed, and wedged lines.
Problem 1.9 Why do you think a C-H bond (109 pm) is longer than an H-H bond (74 pm)?
The C-H bond involves an sp3 orbital on carbon. A p orbital extends further from the nucleus; therefore an sp3 hybrid orbital will extend further from the carbon nucleus.

30. 1.7 sp3 Orbitals and the Structure of Ethane
Two C’s bond to each other by overlap of an sp3 orbital from each carbon
Three sp3 orbitals on each C overlap with H 1s orbitals to form six C–H bonds
All bond angles of ethane are tetrahedral

31. 1.7 sp3 Orbitals and the Structure of Ethane
Problem Draw a line-bond structure for propane, CH3CH2CH3. Predict the value of each bond angle, and indicate the overall shape of the molecule.
Angles ~ 109.5o

32. 1.8 Other kinds of hybridization: sp2 and sp
sp2 hybrid orbitals: one 2s orbital combines with two 2p orbitals, giving 3 orbitals (spp = sp2). This results in a double bond.
sp2 orbitals are in a plane with120º angles
Remaining p orbital is perpendicular to the plane – will be used to form a p (pi) bond

33. Bonding in Ethene

34. Bonds From sp2 Hybrid Orbitals
Two sp2-hybridized orbitals with head-on overlap to form a sigma (s) bond
p orbitals overlap side-to-side to formation a pi () bond
sp2–sp2 s bond and 2p–2p  bond result in sharing four electrons and formation of C-C double bond
Electrons in the s bond are centered between nuclei
Electrons in the  bond occupy regions above and below the internuclear axis

35. Structure of Ethylene H atoms form s bonds with four sp2 orbitals
H–C–H and H–C–C bond angles of about 120°
C–C double bond in ethylene shorter and stronger than single bond in ethane
Ethylene C=C bond length 134 pm (C–C 154 pm)

36. The Carbons in Ethene are sp2

37. Bonds From sp Hybrid Orbitals
C-C a triple bond sharing six electrons
Carbon 2s orbital hybridizes with a single p orbital giving two sp hybrids
two p orbitals remain unchanged
sp orbitals are linear, 180º apart on x-axis
Two p orbitals are perpendicular on the y-axis and the z-axis

38. Bonding in Ethyne

39. Orbitals of Acetylene
Two sp hybrid orbitals from each C form sp–sp s bond
pz orbitals from each C form a pz–pz  bond by sideways overlap and py orbitals overlap similarly

40. The Carbons in Ethyne are sp

41. Bonding in Acetylene Sharing of six electrons forms C º C
Two sp orbitals form s bonds with hydrogens

42. Comparison of C–C and C–H Bonds in Methane, Ethane, Ethylene, and Acetylene
Triple bond shorter, stronger than double bond which is shorter, stronger than single bond.

43. Comparison of C–C and C–H Bonds in Methane, Ethane, Ethylene, and Acetylene
Problem 1.11 Draw both an electron-dot and line bond structure for acetaldehyde, CH3CHO.
Problem 1.12,13 Draw a line-bond structure for propene, CH3CH=CH2 and for propyne, CH3C≣CH. Indicate the hybridization of each carbon , and predict the value of each bond angle. sp
sp2
sp2
sp3
sp3
sp3 – 109.5o; sp2 – 120o; sp – 180o

44. 1.9 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

45. 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

46. The Periodic Table and Electronegativity
increases
decreases

47. 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-N, C-X bonds (more electronegative elements) are polar
Bonding electrons closer to 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

48. 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

49. Electrostatic Potential Maps
Problem 1.17 Use the d+/d- convention to indicate the direction of expected polarity for each of the bonds shown:
H3C-Br b) H3C-NH2 c) H2N-H
d) H3C-SH e) H3C-MgBr f) H3C-F d+ d- d+ d- d- d+ d+ d-
same d- d+
Problem 1.19 Indicate the direction of polarization of the C-Cl bond in chloromethane.

50. 1.10 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

51. 1.10 Brønsted Acids and Bases
“Brønsted-Lowry” is usually shortened to “Brønsted”
A Brønsted acid is a substance that donates a hydrogen cation (H+)
A Brønsted base is a substance that accepts the H+
“proton” is a synonym for H+ - loss of an electron from H leaving the bare nucleus—a proton

52. Acid and Base Strength
The acidity constant (Ka) for the reaction of an acid (HA) with water to form hydronium ion and the conjugate base (A-) is a measure related to the strength of the acid
Stronger acids have larger Ka
Note that brackets [ ] indicate concentration, moles per liter, M.

53. The Acidity Constant for Water
The concentration of water as a solvent does not change significantly when it is protonated
The molecular weight of H2O is 18 and one liter weighs 1000 grams, so the concentration is ~ 55.6 M at 25°
Allows for the comparison of other weak acids with water.
Ka ranges from 1015 for the strongest acids to very small values (10-60) for the weakest
Depending on what it is in solution with, water can serve as an acid or base

54. pKa – the Acid Strength Scale
pKa = –log Ka
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

55. pKa’s of Some Common Acids

56. 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 stronger base holds the proton more tightly
Look for the lower pKa acid – it wins.

57. Predicting Acid–Base Reactions from pKa Values
Problem 1.20 Formic acid, HCO2H, has pKa = 3.75, and picric acid, C6H3N3O7, has pKa = 0.38.
What is the Ka of each? Which is stronger?
Ka = 10-pKa
Formic = 1.78 x 10-4; picric = 0.42 (stronger)
Problem 1.21 Amide ion, H2N-, is a stronger base than hydroxide on, HO-. Which is the stronger acid, H2N-H (ammonia) or HO-H (water)?
If amide ion is stronger base, it’s conjugate acid must be weaker. So water is the stronger acid.
Problem 1.22 Will the following reaction take place according to the pKa data in Table 1.2?
HCN + CH3CO2-Na Na+CN- + CH3CO2H
pKa

58. 1.11 Organic Acids and Organic Bases
characterized by the presence of positively polarized hydrogen atom

59. Organic Acids
Those that lose a proton from O–H, such as methanol and acetic acid
Those that lose a proton from C–H, usually from a carbon atom next to a C=O double bond (O=C–C–H)

60. Organic Bases
Have an atom with a lone pair of electrons that can bond to H+
Nitrogen-containing compounds derived from ammonia are the most common organic bases
Oxygen-containing compounds can react as bases when with a strong acid or as acids with strong bases

61. 1.12 Acids and Bases: The Lewis Definition
Lewis acids are electron pair acceptors and Lewis bases are electron pair donors
Brønsted acids are not Lewis acids because they cannot accept an electron pair directly (only a proton would be a Lewis acid)
The Lewis definition leads to a general description of many reaction patterns but there is no scale of strengths as in the Brønsted definition of pKa

62. Illustration of Curved Arrows in Following Lewis Acid-Base Reactions
Arrow reflects
Movement of
electrons

63. Lewis Bases
Lewis bases can accept protons as well as Lewis acids, therefore the definition encompasses that for Brønsted bases
Most oxygen- and nitrogen-containing organic compounds are Lewis bases because they have lone pairs of electrons
Some compounds can act as both acids and bases, depending on the reaction

64. Summary s orbitals are spherical, p orbitals are dumbbell-shaped
Organic chemistry – chemistry of carbon compounds
Atom: charged nucleus containing positively charged protons and netrually charged neutrons surrounded by negatively charged electrons
Electronic structure of an atom described by wave equation
Electrons occupy orbitals around the nucleus.
Different orbitals have different energy levels and different shapes
s orbitals are spherical, p orbitals are dumbbell-shaped
Covalent bonds - electron pair is shared between atoms
Valence bond theory - electron sharing occurs by overlap of two atomic orbitals

65. Summary (Continued)
Sigma (s) bonds - Circular cross-section and are formed by head-on interaction
Pi () bonds - “dumbbell” shape from sideways interaction of p orbitals
Carbon uses hybrid orbitals to form bonds in organic molecules.
In single bonds with tetrahedral geometry, carbon has four sp3 hybrid orbitals
In double bonds with planar geometry, carbon uses three equivalent sp2 hybrid orbitals and one unhybridized p orbital
Carbon uses two equivalent sp hybrid orbitals to form a triple bond with linear geometry, with two unhybridized p orbitals
Organic molecules often have polar covalent bonds as a result of unsymmetrical electron sharing caused by differences in the electronegativity of atoms

66. Summary A Brønsted(–Lowry) acid donates a proton
A Brønsted(–Lowry) base accepts a proton
The strength of a Brønsted acid is related to the negative logarithm of the acidity constant, pKa. Weaker acids have higher pKa’s
A Lewis acid has an empty orbital that can accept an electron pair
A Lewis base can donate an unshared electron pair

67. Let’s Work a Problem
Rank the following substances in order of increasing acidity.

68. Answer
The lower the value of the pKa, the more acidic the molecule is:

69. Let’s Work a Problem
Draw an electron-dot structure for acetonitrile, C2H3N, which contains a carbon-nitrogen triple bond. How many electrons does the nitrogen atom have in its outer shell ? How many are bonding, and how many are non-bonding?

70. Answer
To address this question, we must realize that the nitrogen will contain 8 electrons in its outer shell. Six will be used in the C-N triple bond (shaded box), and two are non-bonding

71. Summary Slides

72. Single Bond: 1 σ + 0 π Double bond: 1 σ + 1 π
Triple Bond: 1 σ + 2 π

73. Bond Strength and Bond Length
The shorter the bond, the stronger it is.

74. A π Bond is Weaker Than a σ Bond