|CH301H - Principles
of Chemistry I: Honors
Fall 2011, Unique 51040
Lecture Summary, 30 August 2011
|Potential Energy Diagrams:
To begin studying the classical model of chemical bonding, we
thought about electrostatic forces contributing to the structure and
stability of the atom. We reviewed Coulomb's law to calculate
electrostatic force (F) and electrostatic potential energy. We used a calculation of potential energy (V) to construct a potential energy diagram (V(r) vs. r) for several examples of interacting charges (proton + electron, proton + proton, etc.). We found that when F < 0 and V < 0, the system is in an attractive or stabilized state, and when F > 0 and V
> 0, the system is in a repulsive or destabilized state. We
then used this information to reconstruct Rutherford's gold foil
experiment, and used his data to estimate the size of the Au nucleus.
We found that under the correct experimental conditions, we could
calculate a very accurate estimate of the Au nuclei radius, although
our estimate must necessarily be slightly higher than the actual value.
Periodic Trends: The summary of atomic characteristics into periodic trends was absolutely central to the development of the classical model of the atom. We still review periodic trends extensively because they are a fantastic way to develop chemical intuition, or the ability to predict the properties and reactivity of any atom simply based on its location in the periodic table. Today we talked about ionization energy (IE), which is the energy needed to remove one electron from an atom. By examining data of IE's, we found that IE increases across a period, peaks at the noble gas, then drops sharply in the next period. We found that IE > 0 always. We also found that the second ionization energy, the energy required to remove an electron from a species that is already carrying a charge of +1, is always higher than the first ionization energy; i.e IE(2) > IE (1) always. By examining trends in increasing IE(n), we found that electrons are not removed equally from the atom, but appear to come in groups. This observation led to the shell model of the atom.