### Richard A. Paselk

Chem 328

Brief Organic Chemistry

Summer 2004

Lecture Notes: 20 January

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# Introduction

### Tentative Schedule

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.

# 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 s.
• 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 orbitals.
• 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.)

Atomic Orbitals Supplement

# Bonding Review

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. 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 2 HCl

In this case can consider that we get two equations each involving a homo dissociation to give radicals, that is atoms with unpaired electrons:

• 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.)

Lewis Stuctures: One of the simplest models for bonding is the use of Lewis Structures.

• In a Lewis structure only the outermost (valence) electrons are shown.
• 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
• Methane
• 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 octet violations:

 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 symmetrical. 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 the structure):
• 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, so:

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