1. Electronic Structure and Covalent Bonding
Essential Organic Chemistry
Paula Yurkanis Bruice
Chapter 1
Electronic Structure and Covalent Bonding

2. Introduction
Reaction illustrating the first synthesis of urea by Friedrich Wöhler in 1828.
“Organic compounds” are now defined as compounds that contain carbon.
What makes carbon so special? Carbon is able to form millions of stable compounds with a wide range of chemical properties simply by sharing electrons.
Figure: UN
Title:
Synthesis of urea
Caption:
Reaction illustrating the first synthesis of urea.
Notes:

3. 1.1 The structure of an Atom
Atoms have an internal structure consisting of one or more subatomic particles: protons, neutrons, and electrons.

4. Summary of Modern Atomic Theory
Most of the mass of an atom is concentrated in the nucleus.
The nucleus contains one or more positively charged protons, and one or more neutrons with no electrical charge.

5. Summary of Modern Atomic Theory
One or more negatively charged electrons are in constant motion somewhere outside the nucleus.
The number of electrons is equal to the number of protons; the atom has no overall electrical charge.

6. Summary of Modern Atomic Theory
Most of the mass of an atom is in its nucleus. However, most of the volume of an atom is occupied by its electrons, and that is where our focus will be because it is the electrons that form chemical bonds.

7. Summary of Modern Atomic Theory
Atomic number:
the number of protons
Mass number:
sum of its protons and neutrons
Isotopes:
have the same atomic number, but different mass numbers
Atomic weight of an element: the average mass of its atoms.
Molecular weight of a compound: the sum of the atomic weights of all the atoms in the molecule.

8. The age of an organic substance can be determined by its 14C content.
For example: Carbon has three isotopes with mass numbers of 12, 13 and 14.
12C: 98.89%, 13C: 1.11%, 14C: trace amount
14C is radioactive, decaying with a half-life of 5730 years.
The age of an organic substance can be determined by its 14C content.

9. 1.2 The Distribution of Electrons in an Atom
The electrons in an atom can be thought of as occupying a set of shells that surround the nucleus.
Each shell contains subshells known as atomic orbitals.
The closer the orbital is to the nucleus, the lower is its energy.
Relative energies of atomic orbitals: 1s＜2s＜2p＜3s＜3p＜3d
Electronic configuration of an atom
Describes what orbitals the electrons occupy
First shell
Second shell
Third shell
Atomic orbitals s
s, p
s, p, d
# of atomic orbitals 1
1, 3
1, 3, 5
Max. # of electrons 2 8 18

10. Three Important Rules Aufbau Principle Pauli Exclusion Principle
An electron always goes into the available orbital with the lowest energy.
Pauli Exclusion Principle
No more than two electrons can occupy each orbital, and the two electrons must be of opposite spin.
Hund’s Rule
When there are two or more orbitals with the same energy, an electron will occupy an empty orbital before it will pair up with another electron.

11. Carbon
Atomic Number = 6

12. Carbon
Atomic Number = 6

13. Oxygen
Atomic Number = 8

14. Oxygen
Atomic Number = 8

15. Figure: UN.T2
Title:
Table 1.2 The Ground State Electronic Configuration of the Smallest Atoms
Caption:
The ground-state electronic configurations of the lightest atoms.
Notes:
Remember that you half fill degenerate orbitals before you pair them up.

16. Core electrons Valence electrons
Electrons in inner shells (those below the outermost shell)
Valence electrons
Electrons in the outermost shell

17. 1.3 Ionic and Covalent Bonds
Chemical Bonds
An attractive force between two atoms in compounds.
Lewis theory (the octet rule):
An atom is most stable if its outer shell is either filled or contains eight electrons and it has no electrons of higher energy.
Lewis Dot Representation of Atoms
Dots around the chemical symbol of an atom represent the valence electrons.

18. Examples

19. Examples

20. Examples

21. Examples

22. Examples

23. Examples

24. Examples

25. Examples

26. Examples

27. Ionic Bonds Transfer of electrons from one atom to another
The electrostatic attraction between oppositely charged ions.
Example:
Sodium chloride

28. Covalent Bonds
Some atoms do not transfer electrons from one atom to another to form ions.
Instead they form a chemical bond by sharing pairs of electrons between them.
A covalent bond consists of a pair of electrons shared between two atoms.

29. Fluorine, F2
Nitrogen, N2
Carbon Dioxide, CO2

30. Electronegativity and Polar Covalent Bonds
Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons.
The polarity of a bond is determined by the difference in electronegativity values.
If the electronegativities are the same the bond is called nonpolar covalent bonds and the electrons are shared equally.
If the atoms have greatly differing electronegativities the bond is called polar covalent bonds.

31. Figure: UN.T3
Title:
Table 1.3 The Electronegativities of Selected Elements
Caption:
Electronegativity is the tendency of an atom to attract bonding electrons toward itself.
Notes:
Note that the electronegativities increase as you go from left to right across a row on the periodic table and decrease as you go down a column.

32. Polar or Nonpolar Covalent BondsF2 HF
A nonpolar covalent bond
A polar covalent bond

33. Polar Covalent Molecules
A polar molecule is a molecule that is electrically asymmetrical, causing it to be oppositely charged at two points.
The molecule possesses a molecular dipole.

34. Figure: UN
Title:
Polar covalent bonds of hydrogen chloride, water, and ammonia
Caption:
A polar covalent bond is result of the differences in electronegativites of the two atoms involved in the bond. A polar covalent bond has a slight positive charge on one atom and a slight negative charge on the other atom. The atom with the greater electronegativity value will have the partial negative charge.
Notes:
The greater the electronegativity difference between the atoms the more polar the bond is.

36. Electrostatic potential maps
Figure: UN
Title:
Electrostatic Potential Maps
Caption:
Electrostatic potential maps for LiH, H2, and HF.
Notes:
The colors indicate the distribution of charge in the molecule. Red is for electron rich areas and blue is for electron deficient areas. Green signifies no charge.

38. 1.4 Representation of Structure
The chemical symbol in which the valence electrons are represented as dots are called Lewis structures.
H atoms are surrounded by no more than two e- and that C, O, N, and halogen (F, Cl, Br, I) atoms are surrounded by no more than eight e-－ they must obey the octet rule.
Valence electrons not used in bonding are called nonbonding electrons or lone-pair electrons.

39. Computing formal charges
Once the atoms and the electrons are in place, each atom must be examined to see whether a charge should be assigned to it.
Compute formal charges and evaluate the structure.
formal charge = (# of valence e-) – (# of lone-pair e- + 1/2 # of bonding e-)

41. A species containing positively charged carbon: carbocation.
Figure: UN
Title:
Lewis dot structures of nitrogen compounds
Caption:
Lewis dot structures of various nitrogen containing compounds.
Notes:
A species containing positively charged carbon: carbocation.
A species containing negatively charged carbon: carbanion.
A species containing an atom with a single unpaired electron: free radical (radical).

43. Writing Lewis Structures
A pair of shared electrons can also be shown as a line between two atoms.
“Normal” bonding: (no formal charge or unpaired e-)
H: one covalent bond
halogens: one covalent bond
O: two covalent bonds
N: three covalent bonds
C: four covalent bonds

44. Normal Covalent Bonding Patterns

45. Writing Lewis Structures- step by step
Find the total number of valence electrons supplied by all of the atoms in the structure.
Write down the skeletal arrangement of the atoms and connect them with a single covalent bond (two electrons).
Distribute pairs of electrons around each atom (except H) to give each atom eight electrons around it (the octet rule).
If there are not enough electrons to give these atoms eight electrons, change single bonds between atoms to double or triple bonds by shifting non-bonded pairs of electrons.
Compute formal charges and evaluate the structure.

46. Examples of Lewis Structures
Phosgene, COCl2
Number of valence electrons:
C = 4 e-, O = 6 e-, 2 Cl = 14 e-; total = 24 e-
Structure:
Build skeleton
Add electrons
Obey octet rule
Used 6 e-

47. Examples of Lewis Structures
Phosgene, COCl2
Number of valence electrons:
C = 4 e-, O = 6 e-, 2 Cl = 14 e-; total = 24 e-
Structure:
Build skeleton
Add electrons
Obey octet rule
18 more e-

48. Examples of Lewis Structures
Phosgene, COCl2
Number of valence electrons:
C = 4 e-, O = 6 e-, 2 Cl = 14 e-; total = 24 e-
Structure:
Build skeleton
Add electrons
Obey octet rule
Make double bond

50. Other structures Kekulé structures
the two bonding electrons are drawn as lines and the lone-pair electrons are usually left out entirely.
Condensed structures
structures simplified by omitting some or all of the covalent bonds and listing atoms bonded to a particular carbon next to it with a subscript to indicate the number of such atoms.

52. Figure: UN.T4
Title:
Table 1.4 Kekule and condensed structures
Caption:
Table contains examples of ways to draw the Kekule and condensed structures of various compounds.
Notes:

54. 1.5 Atomic Orbitals
Electrons are distributed into different atomic orbitals.
An orbital is a three-dimensional region around the nucleus where there is a high probability of finding an electron.
The s orbital is a sphere with the nucleus at its center.

55. 1.5 Atomic Orbitals A p orbital has two lobes.
Each p orbital is perpendicular to the other two p orbitals.
Generally, the lobes are depicted as teardrop-shaped, but computer-generated representations are more like doorknobs.

56. 1.6 Covalent Bond Formation
The covalent bond in H2 is formed when the 1s orbital or one H atom overlaps the 1s orbital of a second H atom.
The covalent bond that is formed when two orbitals overlap is called a sigma (s) bond.

57. The change in energy during bond formation

58. 1.7 Bonding in Methane and Ethane: Single Bonds
Four C－H bonds in methane are identical.
Four different ways to represent a methane molecule are shown.
Methane is a nonpolar molecule.

59. Methane A tetrahedral structure is suggested:
But, the atomic orbitals on carbon do not have the correct geometry to give a tetrahedron.
Generate hybrid orbitals by “mixing” atomic orbitals in order to achieve the correct geometries.

60. Methane
However, mixing the carbon 2s orbital and the 3 carbon 2p orbitals gives 4 sp3 hybrid orbitals:

61. Methane
However, mixing the carbon 2s orbital and the 3 carbon 2p orbitals gives 4 sp3 hybrid orbitals:

62. Figure: 01-04
Title:
Structure of Methane
Caption:
a) Four sp3 orbitals are directed toward the corners of a tetrahedron causing each bond angle to be degrees. b) An orbital picture of methane showing the overlap of each sp3 orbital of the carbon with the s orbital of hydrogen.
Notes:
All molecules with a tetrahedron shape and with no lone pairs of electrons have a bond angle of degrees.

63. Methane
These 4 sp3 hybrid orbitals point toward the corners of a tetrahedron:

64. Methane
Orbital overlap of the carbon sp3 with the hydrogen 1s orbitals gives the four s bonds:
A tetrahedral atom has sp3 hybridization.

65. Ethane
The C－C bond is formed by sp3-sp3 overlap, and each C－H bond is formed by sp3-s overlap.

67. 1.8 Bonding in Ethene: A Double bond
Trigonal planar geometry around each carbon atom.
The molecular orbitals of ethene:
1s orbitals of the four hydrogen atoms
2s, 2px, 2py, and 2pz orbitals of the two carbon atoms

68. Ethene
The 2s, 2px, and 2py orbitals on carbon form three sp2 hybrid orbitals on each carbon atom.
The three sp2 hybrid orbitals point toward the corners of a triangle.
A trigonal planar atom has sp2 hybridization.

69. Ethene
The sp2 hybrids and the 1s orbitals on hydrogen make up the s-framework— s bonds that are relatively low in energy.
A trigonal planar atom has sp2 hybridization.

70. Ethene
Side-to-side overlap of two p orbitals forms a pi (p) bond. Thus, one of the bonds in a double bond is a s bond and the other is a p bond.

71. Ethene
The two p orbitals that overlap to form a p bond must be parallel to each other for maximum overlap to occur.
This forces the two triangles formed by all six atoms to lie in the same plane, and the electrons in the p orbitals occupy a volume of space above and below the plane.

72. 1.9 Bonding in Ethyne: A Triple bond

73. Ethyne A triple bond consists of one s bond and two p bonds.
Because the two unhybridized p orbitals on each carbon are perpendicular to each other, there is a region of high electron density above and below, and in front of and in back of the internuclear axis of the molecule.

76. 1.10 Bonding in the Methyl Cation, the Methyl Radical, and the Methyl Anion

77. 1.11 Bonding in Water
The experimentally observed bond angle is 104.5o.
Oxygen must use hybrid orbitals to form covalent bonds, just as carbon does.
The bond angles in a molecule indicate which orbitals are used in bond formation.

79. 1.12 Bonding in Ammonia and in the Ammonium ion
The experimentally observed bond angles in NH3 are 107.3o.

81. 1.13 Bonding in the Hydrogen Halides
Since hydrogen halides have only one bond, we will assume that the halogens use p orbitals to form bonds.

82. Figure: 01-10
Title:
Orbital Overlap in HF and HCl
Caption:
Overlap of the bonding p orbital of F and Cl with H in HF and HCl.
Notes:
The region where the s orbital of the hydrogen overlaps the p orbital from a halide decreases as the size of the halide increases.

84. 1.14 Summary: Orbital Hybridization, Bond Lengths, Bond Strengths, and Bond Angles
All single bonds are s bonds.
All double bonds are composed of one s bond and one p bond.
All triple bonds are composed of one s bond and two p bonds.
The easiest way to determine the hybridization of a atom is to look at the number of p bonds: no- sp3, one-sp2, two-sp.