Brief Organic Chemistry
|Lecture Notes: 20 January
© R. Paselk 2004
Who am I?
- Education in Biophysics (NMR of Proteins etc in Grad Sch)
- Also teach Biochemistry, Biochemical Toxicology; Chemical
Instrumentation and all intro Chem courses
- Interests in history of science and technology, particularly
scientific instrumentation and apparatus (Museum in Library,
office and on Web) and Precambrian Eon.
How to study:
Notes are key-nearly everything you
will need to know I will cover in lecture. So how can you get
the most out of your notes?
Don't rely on/be seduced by on-line notes.
Covalent Molecules - Bonding and Shapes
A Quantum Picture of the Atom
We've taken a brief look at the physics underlying atomic structure,
focusing on Schrödinger's Equation and the wave picture of
electron distribution in atoms. Let's flesh this out a bit.
What we need to explain is the energy distribution of electrons
in atoms and how this correlates with atomic properties. First
recall the line spectrum of hydrogen and the Bohr model. We are
going to keep the concepts of ground state and quantized energy
levels from Bohr, after all they worked very well for Hydrogen.
But we will need to build a new structure which will give these
same predictions but with other factors which explain the details
of hydrogen's spectra as well as other atoms. We'll again start
by modelling hydrogen.
Electronic Energy Levels:
- We will designate the primary energy level, corresponding
to the average radial distance of the electron from the nucleus
as a shell, and give it the symbol n. The lowest
possible energy level is then the ground state with n = 1.
- Shells with n > 1 may have subshells which are
different geometrical patterns of electron distribution. Thus:
- The lowest energy pattern is spherical and given the designation
- The next lowest energy distribution is bi-lobed with a planar
symmetry. It is given the designation p.
- The third lowest energy distribution has diagonal planes
of symmetry and is designated d.
- The average energies of the different subshells are the energy
of the shell, thus when subshells are present the energy of the
shell is split. For example, in the n=2 shell the 2s orbital
becomes lower in energy than the shell, while the 2p orbitals
become higher in energy.
- The regions of electron occupancy in subshells are called
- For each shell there is one s orbital.
- For each shell with n = 2 or greater there are three p
orbitals: px, py, and pz.
- For each shell with n = 3 or greater there are five d
orbitals: dxz, dyz, dxy, dx2-
y2, and dz2
- Each orbital can accomodate 1 or 2 electrons
- When there is more than one orbital at an equivalent energy,
electrons will spread out to give one electron per orbital before
pairing up (e.g. a nitrogen atom will have 2 electrons in the
2s orbital, and one electron in each of the 2p orbitals.)
Chemical bonds are the strongest forces that exist between
atoms. They are the forces that hold atoms together in molecules
and atoms or ions together in solids. The two most important and
common strong bond types in chemistry are ionic bonds and
covalent bonds. In this course we will focus mostly on
Covalent Bonds. Covalent Bonds
occur with the sharing of electrons by two atoms with similar
tendencies to gain and lose electrons. Let's look at the formation
of HCl as an example of the creation of a covalent bond:
H2 + Cl2 Æ
In this case can consider that we get two equations each involving
a homo dissociation to give radicals, that is atoms with
- H2 Æ 2 H.
- Cl2 Æ 2 Cl.
These radical then combine to form a bond with these two
electrons shared between the two atoms.
- H. + Cl. Æ
H:Cl (This is not a proper Lewis structure, I have only
shown the bonding pair of electrons.)
One of the simplest models for bonding is the use of Lewis
- In a Lewis structure only the outermost (valence) electrons
- The Lewis structure model then assumes that atoms in the
s- and p- blocks of the Periodic Table will have a strong tendency
to have an octet of electrons in the valence shell (except for
H, Li, and Be which have duets, and B, which is wierd).
- A bond is then represented by a pair of dots between the
symbols of two atoms.
- Lewis Structure:
- Covalent Compound Lewis Structure Examples:
- Ammonium ion
- Carbon dioxide
- Carbon monoxide
Lewis structures are of course quite limited - they work well
only for the representative elements, and even then we have to
stretch the concept to accommodate all covalent structures., Thus
to follow the "Octet rule" we invented resonance
for molecules which don't have enough electrons to give octets
even with multiple bonding. Clark's rules can help you determine
Clark's Method (abbreviated) for determining
bonding in covalent Lewis Structures:
- Add up all of the valence electrons in the structure (remember
to add one electron for each negative charge, or subtract one
for each positive charge)
- If S e- =
6y + 2 where y = # atoms other than H, then octet rule is followed
with single bonds only.
- If S e- <
6y + 2 then probably have multiple bonding with the number of
multiple bonds = D/2 (remember a triple
bond is 2 multiple bonds!). However,
note the exceptions with small atoms (H, Li, Be, and B).
- If S e- >
6y + 2 then have an expanded valence shell. Note that
if D = 2, then pentavalent (10 electrons
in the valence shell) , and if D =
4, then hexavalent (12 electrons in the valence shell).
- If you can draw more than one structure, then chose the most
- If two or more structures are equally symmetrical, then you
probably have resonance and should show all structures connected
by double arrows.
- Resonance example (Using Clark's rules to help determine
- Carbonate ion - CO32-
- valence electrons = 4 + 3 (6) + 2 = 24
- 6y + 2 = 26, but S e-
= 24, therefore expect one multiple bond.
- LS =
- However, other equally symmetrical structures are possible,
Last modified 1 June 2004