CHEM 3013
ORGANIC CHEMISTRY I
LECTURE NOTES
CHAPTER 1
1. Atomic orbitals
a.
Heisenberg Uncertainty Principle The exact position of an electron cannot be
specified; only the probability that it occupies a certain position of space.
b.
The electron can only exist in certain regions of space called atomic orbitals.
c.
The energy of an electron in a particular orbital has a very precise vale.
i.
ii.
iii.
Orbitals are described by two quantum numbers:
The first descriptor is called the principal quantum numbers and is an
interger value from one to seven. Electrons having higher numbers have
higher energy.
The second quantum number is called the angular(azimuthal)
quantum numbers. It describes the shape of the orbital. The four letter
descriptors of these shapes are s,p,d,f.
d.
Orbitals of higher principal quantum number have more nodes.
e.
Higher energy orbitals have the electron density located further from the nucleus.
node
1s orbital
2s orbital
three 2p orbitals
Orbitals
2. Electronic Configuration
a.
Aufbau Rules Building the complex atoms of the periodic table.
3p
3s
ENERGY
2p
2s
1s
Energy of different orbitals in a many electron atom
i.
ii.
iii.
Electrons are added to orbitals starting with those of lowest energy
Only two orbitals per atomic orbital and these must differ in their
spin quantum number (Pauli Exclusion Principal).
For orbitals of equal energy, single electrons with identical spin are
added untill the orbital is half-full, then the electrons of opposite
spin are added (Hunds Rule).
1
2
3p
3p
3s
3s
ENERGY
ENERGY
2p
2p
2s
2s
1s
1s
carbon 1s2 2s2 2p2
oxygen 1s2 2s2 2p4
Filling the orbitals of carbon and oxygen
3.
Ionic Bonding
a.
Inert Gas Configuration; Octet Rule. Certain elements in the periodic table are
especially stable. These are the inert gases, helium, neon, argon,krypton, xenon
and radon found in group 8A. The common features of these elements is that they
all have a completely filled valence shell of electrons. Except for helium which has
a filled shell of two electrons, all the remaining inert gases have 8 electrons in their
outer shell. This special configuration can also be attained in other elements; it is
the major driving force in the formation of ions (by gain or loss of electrons).
Elements or ions which have a filled shell of electrons are said to obey the octet
rule.
b.
Formation of ions. Much of chemistry can be understood in terms of the various
elements' wish to gain the filled shell configuration. In the reaction of lithium with
fluorine gas, both elements can satisfy the octet rule by transfer of one electron
from the lithium to the fluorine. During this process the fluroine becomes a fluoride
anion, the lithium a lithium cation. Both ions have the same configuration as neon.
Same electronic
configuration as Ne
The lithium half-reaction):
Li (1s2 2s1)
1 e- + Li+ (1s2)
Elemental lithium,
a source of one
electron
available
electron
The fluorine half-reaction
F (1s2 2s2 2p5)
+
1 e-
Same electronic
configuration as Ne
F- (.....2p6)
Elemental fluorine, electron
an acceptor of one needed
electron
Filled shell configuration by loss or gain of electrons
Elements at the left of the periodic table (like groups 1A and 2A) can most easily
attain a filled shell configuration by donating electrons. Conversely, elements on
the right side of the periodic table are converted to the octet configuration by
accepting electrons.
c.
4.
Isoelectronic Chemical Species have the same electronic configuration. Thus the
ions of O-2,F -, Ne, Li+ and Mg+2 are isoelectronic, having the numbers of
protons which makes them distinct chemical entities.
d.
Electrostatic Attraction The simple attraction of oppositely charged particles which
holds the ions together in an ionic solid.
Covalent Bonding.
a.
Not all compounds are comprised of ions. In particular, many carbon compounds
are not ionic. The reason for this difference can be found in the electronic
configuration of carbon (1s2,2s 2,2p x1,2p y1). To gain a filled shell configuration,
carbon would have to either gain or lose four electrons. This represents an
energetically unfavorable situation.
The solution to the bonding problem posed by carbon is to share electrons
in a covalent bond. This type of bonding was proposed by G. N. Lewis in 1916. The
formalism depicts bonding resulting from havinga pair of nuclei being shared between two
nuclei. The models resulting from this formalism are called Lewis Structures. This model
is the cornerstone of understanding organic reaction mechanisms and is exceptionally
useful for electronic bookkeeping.
b.
Writing Lewis Structures
i.
ii.
iii.
iv.
v.
vi.
Start with a central atom (usually the heaviest, but make it carbon if carbon
is present).
Place all the outer shell electrons on four imaginary sides. Don't pair
electrons unless more than four are in the outer shell of an atom.
Bring the other atoms of the molecule in to the central atom so that single
electrons on the central atom can be paired with single electrons on the other
atoms.
If, after all atoms have been added, two adjacent atoms have unshared
electrons, they share the unshared electrons to make multiple bonds.
If the species is an ion, then add or subtract the appropriate number of
electrons in order to obtain the charge of the ion.
The final Lewis-dot structure should have eight electrons around all atoms
except hydrogen,which has two. (There are a few exceptions to the octet
rule but these usually don't involve compounds containing carbon.)
Covalent Bonding in methane
H
H C H or H C H
H
H
H
C
+
4H
Note that the hydrogen of methane have a helium closed shell and the
carbon has a neon closed shell
LEWIS STRUCTURES FOR COVALENT MOLECULES
c.
Unshared pairs. Not all electrons in a covalent compound are necessarliy involved
in bonding; some structures have orbitals that contain unshared pairs of electrons.
While these electrons are not used in determining the primary bonding arrangement
of the covalent compound, they are of basic importance in determining how the
compoundwill react in later reactions. In drawing Lewis Structures, the lone pairs
are often not shown, but we shall draw them explicitly at this stage of the course.
Lewis Structures are also useful in understanding the bonding of covalent
compounds containing multiple bonds (double and triple).
3
4
H
H
C
H
H
O
C
H3C
H
H
CH3
dimethyl
ether
unshared pairs
H3C
H
C
H
O
C O
O
formaldehyde
H3C
Lewis structures with unshared pairs and/or multiple bonds
5.
VSEPR Theory
a.
The shape of a molecule can be predicted by allowing all the bonded and
nonbonded electron pairs to move as far apart as possible.
b.
For the given number of different bonds and nonbonded pairs of electrons,
the expected shapes would be as shown below.
Tetrahedral. WXYZ can be bonds or non-bonding electrons
H
Z
W
X
A
O
H
H
Y
H
H
HOH = 104.5
bent
N
H
HNH = 107.3
pyramidal
H
H
C
H
HCH = 109.5
tetrahedral
Trigonal. XYZ can be bonds or non-bondingelectrons.
X
A
Z
F
B
Y
F
F
FBF = 120
planar
Digonal (Linear). YZ can bebonds or non-bonding electrons
X
A
Y
H
C
C
H
Bonds and geometry
6. Molecular Orbitals
a.
Quantum Mechanics tells us that combining two atomic orbitals leadsto two
molecular orbitals. The positive combination is a bonding molecular orbital and the
negative combination is an antibonding molecular orbital. The formation of a bond
between two hydrogen atoms is a very favorable (104 Kcal/mole) process. This is
because we put two electrons in a bonding MO downhill in energy. The
combination of two atomic orbitals with spherical symmetry (e.g. 1s) generates a
MO with cylindrical symmetry (symmetric about the internuclear axis). This
called a sigma (σ) bond.
b.
An orbital is described as a probability function that defines the distribution of
electron density in space, the overlap region of a MO, between the two nuclei has a
considerably higher electron density. The covalent bond arises from the
electrostatic attraction between the positive nuclei and the region of increased
electron density.
+
Antibonding
Anti-bonding
MO
Combination
E
N
E
R
G
Y
Region of decreased
electron density.
Energy of
isolated H
atoms
sigma bond
+
Bonding
Combination
Bonding MO
Overlap region of increased
electron density
Molecular orbitals for the hydrogen molecule
7.
Hybrid Orbitals
a.
As a carbon atom begins to form bonds with up to four other atoms, the ground
state electronic configuration of 1s2 2s2 2p2 is purturbed and begins to change. An
electron in a 2s orbital is promoted into a 2p orbital and the new electronic
configuration is 1s2 2s2 2px1 2py1 2pz1.
b.
Two-electron bonds are like small magnets, the bonds tend to repel each other.
Forming four covalent bonds to carbon by using three p orbitals and one s orbital is
not the best energetic situation. This situation would place the sigma bond
formed using the s orbital quite close to the bonds formed from the three p
orbitals. By undergoing hybridization, nature creates four new orbitals of lower
energy called sp3 hybrid orbitals which now have a tetrahedral structure.
E
N
E
R
G
Y
2p
2p
2s
2s
1s
1s
Elemental
carbon
Hybridization in carbon
Electron promoted to
2p orbital
2sp3
1s
Completely hybridized
carbon in organic
compounds
5
6
8.
Examples of sp 3 Compounds
The value of this bonding model is that it gives the chemist the ability to predict the
geometry of chemical structures. This can be seen in the case of water. The oxygen in
water has a ground state electronic configuration of 1s2, 2s2, 2p x2, 2p y1, 2p x1.This
would suggest that oxygen would form covalent bonds with hydrogens using the half-filled
1s orbitals of each hydrogen and the half-filled 2py and 2pz orbitals of oxygen. Thus, only
two sigma bonds from oxygen to hydrogen, which would seem to suggest that water has a
bent structure, with a HOH bond angle of 90 degrees (the angle of the 2py and 2pz orbitals
with respect to each other). But does water have a 90 degree bent structure? NO!
Hybridization results in the formation of four sp3 orbitals, two filled and two half-filled.
The orientation of the four new orbitals is tetrahedral (109 degrees). When hydrogen
covalently bonds to the two half-filled sp3 orbitals, the HOH bond angle is 109
degrees.
E
N
E
R
G
Y
2p
2s
mix 2s and
1s
all three 2p
Elemental
oxygen
2sp
3
new
hybrid
orbitals
1s
Completely hybridized
oxygen in water
H
H
Hybrid Orbitals of Oxygen in Water
9.
Other Hybridizations
Other hybridization states exist for the carbon atom (and others) when the number
of 2p orbitals mixed in with the 2s is less than three. Hybrid orbitals of sp2 are formed
when only two of the three available 2p orbitals are mixed in. This now results in the
formation of three "new" sp2 orbitals, which are oriented in a trigonal planar orientation
( they look like a three-bladed propeller). There is one un-hybridized 2p orbital left over,
which is oriented perpendicular to the plane of the propeller.
7
E
N
E
R
G
Y
2p
2p
2p
2s
2s
mix 2s and two
1s
1s
2p orbitals
1s
sp2 hybridized
carbon
Electron promoted to
2p orbital
Elemental
carbon
new
2sp2 hybrid
orbitals
2p
sp2
sp2
sp2
sp2 Hybridization in carbon
Two carbon atoms of sp2 hybridization will combine to form a ∑ (sigma) and ∏
(pi) bond. It should be noted that a double bond is comprised of two two-electron bonds
(∑ and ∏); this is why double bonds are shorter and stronger than single bonds. The ∑
bond is formed by end-on overlap of the carbon sp2 orbitals. It has cylindrical symmetry
about the carbon-carbon bond axis. The ∏ bond is formed by sideways overlap of the two
2p orbitals. It is not cylindrically symmetric. A ∏ bond is about 1/2 as strong as a ∑ bond.
2 C(sp2) +
4 H(1s) ======>
H2C=CH2
H
π bond (two p orbitals)
H
+
H
H
σ bond (two sp2 orbitals
from the two carbons).
Bonding
in sp2 Hybridized Carbons
Hybrid orbitals of sp hybridization are formed when one 2s and one 2p orbital are
mixed together. A carbon atom which has sp hybridization has two sp orbitals (oriented at
180˚ from each other) and two orthogonal 2p orbitals.
8
E
N
E
R
G
Y
2p
2s
mix 2s and one
1s
1s
2p orbital
Elemental
carbon
2p
2p
2s
Electron promoted to
2p orbital
sp
1s
sp hybridized
carbon
2p orbitals
sp hybrid
orbital
sp Hybridization in Carbon
new
hybrid
orbitals
sp hybrid
orbital