- Principles of Chemistry I: Honors
Fall 2013, Unique 52195
Lecture Summary, 3 September 2013
discussed a variety of experimental measurements that taken
together, form the classical model for the structure of the atom.
Planetary Model of the Atom: We reviewed Rutherford's gold foil experiment which lead to the planetary model of the atom, with a bulky nucleus surrounded by a stable electron in motion.
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 wasabsolutely central to the development of the classical model of the atom. We review periodic trends extensively in this course 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.
The second periodic trend we discussed was electron affinity, which is the energy gained by adding an electron to an atom to make a negatively charged ion, or anion. We saw that the periodic trend for electron affinity is roughly the same as for ionization energy: the magnitude of energy gain increases across a period and decreases down a group. Therefore, the atoms that give up an electron most easily (group I) are also that atoms that are most difficult to make accept an electron, and vice versa. It is very important that you are comfortable with the signs of IE and EA.
Although IE and EA are properties of individual atoms, they influence how an atom will behave when it is involved in a bonding interaction with another atom. The physical properties of the two bonded atoms will give the molecule physical properties that we would like understand and predict. For example, an atom that has a high ionization energy and high electron affinity will tend to sequester the shared electrons in the bond, making the bond polar or even ionic. This is such an important concept that chemists have put the information contained in atomic IE and EA together into the concept of electronegativity. EN is a purely empirical concept; i.e. it is based only on observation and experience, not on logic, and cannot be derived from first principles. It is not a physical quantity that can be measured, which should be clear to you from the discussion in your book on how Mulliken and Pauling established their EN scales. Like IE and EA, EN increases across a period and decreases down a group.