Humboldt State University ® Department of Chemistry

Richard A. Paselk

Chem 328

Brief Organic Chemistry

Summer 2004

Lecture Notes: 20 January

© R. Paselk 2004
 
 

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Introduction

Syllabus

Tentative Schedule

 

Who am I?

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:

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:

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

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.

 

 


Schedule

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Last modified 1 June 2004