Organic Chemistry
Review Information for Unit 1
 Atomic
Structure
 MO Theory
 Chemical Bonds
Atomic Structure


Atoms are the smallest representative particle
of an element.
 Three subatomic particles:
 protons
 neutrons
nucleus
 electrons
All atoms of a particular element have the
same number of protons and electrons.
 The number of neutrons can vary.
Atomic Structure


Isotope
 atoms with the same number of protons but
different numbers of neutrons
Carbon has three isotopes:
 12C: 6 protons, 6 neutrons
 most common isotope
 13C: 6 protons, 7 neutrons
 used in structural determinations
 14C: 6 protons, 8 neutrons
 used to determine the age of organic
materials (C-14 dating)
Atomic Structure


The location and energy of the electrons in an
atom are best described using the Quantum
Mechanical Model.
 Electrons have both particle-like and wavelike properties.
Solving the Schroedinger equation leads to a
series of mathematical functions called wave
functions (y)
 describe an allowed energy state for an
electron (orbital)
Atomic Structure


Heisenberg Uncertainty Principle:
 the exact energy and exact location of an
electron in an atom cannot be known
simultaneously
Since solutions to the Schroedinger equation
give the exact energy of the electron, the
exact location of the electron is uncertain.
 Electrons don’t move in well-defined circular
orbits around the nucleus.
Atomic Structure


Although we cannot determine the exact
location of an electron, an orbital describes a
specific distribution of electron density in
space
 the probability of finding an electron in a
particular region of space
The QM model uses 4 quantum numbers to
describe an electron in an orbital:
 n, l, ml are used to describe the orbital
itself
 ms is used to describe the spin of the
electron.
Atomic Structure

Quantum Numbers:
 principal quantum number (n)
 energy of the electron
 relative distance from the nucleus
 azimuthal quantum number (l)
 shape of the orbital
 magnetic quantum number (ml)
 orientation in space of the orbital
 electron spin quantum number (ms)
 direction of electron spin
Atomic Structure

There are four common “types” of orbitals
 s orbital
1s
 spherical
 One per subshell
y

p orbitals
 3-D figure 8
 3 per subshell when n>2
z
 same energy
 different orientation in space
x
px
Atomic Structure
The px, py, and pz orbitals are superimposed
at 90o angles.
y
z
x
x
z
y
The px and py orbitals are in the plane of the
slide while the pz orbital comes out toward
you at 90o from the plane of the slide.
Atomic Structure

d orbitals
 5 per subshell when n>3
 same energy
 different orientations in space
 complex shapes
Note: Most organic
compounds do not utilize
d and f orbitals.

f orbitals
 7 per subshell when n>4
 same energy
 different orientations in space
 complex shapes
Atomic Structure


An orbital diagram or the electron
configuration can be used to describe the
arrangement of the electrons in the orbitals of
an atom.
According to the aufbau principle, the
electrons in an atom in the ground state will be
found in the lowest energy orbital that is
available.
Atomic Structure

Use the diagonal diagram to determine the
relative energies of the orbitals:
1s
2s
2p
3s
3p
3d
4s
4p
4d
4f
5s
5p
5d
5f
Atomic Structure


The Pauli Exclusion Principle and Hund’s rule
also govern the placement of electrons.
Pauli Exclusion Principle:

No two electrons in an atom can have the
same set of four quantum numbers n, l, ml,
and ms.
 Maximum of 2 electrons with opposite
spins per orbital
Atomic Structure

Hund’s Rule:

If more than one orbital with the same
energy is available, electrons will fill empty
orbitals first.
 Keep electrons unpaired as long as an
empty orbital with the same energy is
available.
Atomic Structure
Example: Draw an orbital diagram and write the
electron configuration of N.
# electrons = 7
Orbital diagram:
1s
2s
2p
Atomic Structure
# electrons = 7
Electron configuration: 1s22s22p3
This electron configuration does not clearly
indicate that all three 2p electrons are unpaired.
A better representation is to clearly show where
each p electron is found:
1s22s22px12py12pz1
Atomic Structure

On your exam, you should be able to draw an
orbital diagram or write electron configurations
that clearly indicate the location of each
electron.
 i.e. show whether an electron is in a px, py,
or pz orbital.
MO Theory


Quantum mechanics describes the electrons in
an atom using wave functions called atomic
orbitals.
 Allowed energy states for electrons in an
atom in the QM model
According to Molecular Orbital (MO) Theory, a
chemical bond is formed when 2 atomic orbitals
on different atoms overlap and combine
 Two new molecular orbitals are formed:
 bonding molecular orbital
 antibonding molecular orbital
MO Theory

Molecular orbitals formed when two hydrogen
atoms combine:
Antibonding
molecular orbital
Bonding molecular
orbital
MO Theory


Bonding molecular orbital
 Constructive interference between to atomic
orbitals leads to a build up of e- density
between the nuclei
 lower energy than atomic orbital
Electrons in bonding molecular orbitals stabilize
a chemical bond.
MO Theory


Antibonding molecular orbital
 Destructive interference leads exclusion of
electrons for the region between the nuclei
 Highest electron density is located on
opposite sides of the nuclei
 higher energy than atomic orbital
Electrons in antibonding molecular orbitals
destabilize a chemical bond.
Chemical Bonds


Chemical bond: strong attractive force that
exists between atoms (or ions) in a compound
 ionic bonds
 covalent bonds
 nonpolar covalent bond
 polar covalent bond
Ionic Bond: the electrostatic force of
attraction between oppositely charged ions in
an ionic compound
 metal cation (+)
 non-metal anion (-)
Chemical Bonds


Covalent Bonds: the attractive force between
atoms in a molecule that results from sharing
one or more pairs of electrons
 non-metals
 H2O, O2, CCl4, C6H12O6
In some molecules, electrons are shared
equally.
 nonpolar covalent bonds
 H - H, Cl - Cl, O=O
Chemical Bonds


In some molecules, electrons are not shared
equally due to relatively “large” differences in
electronegativities between atoms in the bond.
 polar covalent bonds
H - O
N - H
 C - Cl
Electronegativity: tendency of an atom in a
compound to draw electrons towards itself
Chemical Bonds

Consider the C - Cl bond:
 Cl is more electronegative than C
 electrons are attracted more strongly to
Cl giving it higher electron density and a
partial negative charge ( d- )
 electrons are drawn away from C giving it
lower electron density and a partial
positive charge ( d+ )
+
d
C
en = 2.5
-
d
Cl
3.2
Polarity


The polarity of a bond is measured by its
dipole moment.
 Amount of charge at either end of the
dipole x bond length
Common dipole moments:
 C - H
0.3 D
 C - O
0.86 D
Increasing
 N - H
1.31 D
polarity
 C - Br
1.48 D
 O - H
1.53 D
D = debye
Chemical Bonds

How do you determine if a bond is polar?
 As a rule of thumb:
D en
Bond Type
< 0.5
nonpolar covalent
0.5 - 2.0
>2.0
polar covalent
ionic
Note: These values are approximate. Bond length
is also important so there are exceptions to these
values!!!
Chemical Bonds

Polarity of bonds can be indicated in a couple
of ways:
+
 partial charges (d and d )
+
 d
on least electronegative atom
 d
on most electronegative atom

a
“+” sign at the positive end of the bond
and an arrow head at the negative end of
the bond
C
Cl
Chemical Bonds
Example: Which of the following contain polar
bonds? Identify all partial charges and indicate
the direction of the dipole moment for each polar
bond.
H2O, F2, HF, CH3CH2OH
Chemical Bonds
Example: Which of the following contain polar
bonds? Identify all partial charges and indicate
O
the direction of the dipole moment for each polar
H
H
bond.
H2O, F2, HF, CH3CH2OH
d-
H
O
d+
H
O
H
H
d+
H
H
O+
d
H
d-
H F
O
H
H F
Chemical Bonds
H
H
d+
H C C
d+
Od
H
H H
H
H
H C C
H H
O
H
Chemical Bonds

Ionic and covalent compounds tend to have
different properties:
 Ionic compounds tend to:
 be water soluble
 be solids at RT
 have higher MP and higher BP
 Covalent compounds tend to:
 be less water soluble or completely
insoluble in water
 be solids, liquids or gases at RT
 have lower MP and lower BP
Chemical Bonds


Valence electrons are involved in the formation
of chemical bonds or ions
 electrons residing in the incomplete outer
shell of an atom
For main group elements, the number of
valence electrons for an element = group
number of the element
 N (group 5A) has 5 valence electrons
Chemical Bonds

Lewis symbols are used to depict the valence
electrons present in an atom (or ion).
 chemical symbol for the element
 dot for each valence electron
 dots are placed on all 4 sides of the
chemical symbol
 up to 2 dots (electrons) per side
Chemical Bonds
Example: Draw the Lewis symbol for oxygen.
Chemical symbol: O
Group number: 6A
# of valence electrons: 6
O