General Chemistry Spring 2009

This is a series of flash Cards and animations of molecules to help memorize the necessary information that relates number of bonding groups and lone pair electrons on the central atom with the bond angle and molecular geometry. Accompanying this information is an animation showing what the molecular geometry looks like.

In this clip I discuss several multiple choice questions. The first multiple question we need to understand the concept of isoelectronic. In the second questions (3’ 35”) we use the concept of ionization energy. The third question (7’ 55”) the electron configuration of a transition metal cation is determined. The fourth question (11’ 38”) we use the energy level diagram for the hydrogen atom to reach the correct answer. The fifth question (14’ 00”) we must match an electron configuration with a particular group in the periodic table. The sixth question (18’ 6”) is a Hess’ Law problem that has some fun features. The seventh question (26’ 22”) we must calculate the enthalpy of combustion.

In this clip I discuss predicting the products of four reactions (Question 1), and how to write the ionic and net ionic equation for one of the reactions (Q2 @7’ 51”). Question 3 (10’ 57”) asks for the electron configuration and the number of unpaired electrons for three elements. Using the same theme, an orbital diagram is focus of Q4 (18’ 58”). Question 5 (22’ 35”) is a multipart question beginning with the calculation of the frequency of an emission line in the hydrogen atom spectrum, then a calculation of the value of n that the electron falls to, a drawing of the energy level diagram for the hydrogen atom, and finally a calculation of the amount of energy required to ionize the electron. In Q6 (37’ 35”) effective nuclear charge is used to explain periodic trends.

In this lecture we continued our discussion of hybrid orbitals on central atoms that are used to form covalent bonds. We briefly reviewed sp^3 hybrid orbitals, then we talked about sp^2 and sp hybrid orbitals. Next we looked at the bonding in ethene (ethylene) to show how we view a double bond. We discussed sigma and pi bonds. We also looked at ethyne (acetylene) to see the nature of the bonding in that molecule. We discussed nitrite as an example of a polyatomic anion that exhibits delocalize electrons. Finally we started talking about gases. We did a demonstration of atmospheric pressure and looked at the PLE experiment that everyone did for class.

In this lecture we began with a review of VSEPR models. Then we discussed polar bonds, and polar and nonpolar molecules. Then we discussed the types of orbitals that are involved in foming covalent bonds. In today's lecture we focused on examples of covalent bonds in substances that only used atomic orbitals. Right near the end of class we looked at methane, which is a case where atomic orbitals can NOT be used to described the orbitals on carbon that form covalent bonds. We talked about hybrid orbitals.

In this lecture we continued our discussion over Lewis structures by looking at ethane, ethene and ethyne. These compounds are more complicated compared to the compounds discussed in the previous class due to the presence of mulitple central atoms. Bond order and bond lengths were introduced with these three compounds. Ethane has a carbon-carbon bond order of one, ethene has a carbon-carbon bond order of 2 and ethyne a carbon-carbon bond order of three. The next substance discussed was ozone. Ozone was interesting because the concept of resonance and resonance structures was introduced to completely represent the Lewis structure of the allotropic form of oxygen. The oxygen-oxygen bond order was determined. Next we drew the ressonance structures for nitrite (NO2^-) and nitrate (NO3^-). The next grop of compounds considered were the oxyacids such as sulfuric acid, nitric acid, carbonic acid and many others. At the end of the lecture formal charge was introduced and the formal charge for the atoms in a compound

In this lecture we began discussed Lewis structures for simple molecules. There are a specific series of steps that must be followed to drraw a complete Lewis structure. Several examples were discussed. The examples focused on how to draw a skeletal structure based on deciding from a chemical formula which atom(s) are terminal and which atom(s) are central. Examples with multiple bonds were also discussed.

In this lecture we began by reviewing the effective nuclear charge. The ACA for today's class had some questions on ENC that many students had difficulty with, so we did a brief review. Next we discussed the two anomilies in the first ionization energies for the 2nd period elements. Then we defined electron affinity and looked at the experimental values for the elements in the 2nd period and discussed the trend. Next we began discussing Chapter 8 Chemical Bonding. We began witha review of ionic and covalent compounds and their properties. We looked at the Lewis structure of atoms, and the symbolic representation of the valence electrons. We discussed ionic bonding and covalent bonding and looked a several examples. We discussed lattice energy as a measure of the strength of an ionic bond. Right an the end of class we looked at an animation depicting the overlap of two hydrogen 1s orbitals to form a covalent bond.

In this lecture we began by reviewing the electron configurations of chromium and copper. Then we discussed the electron configuration of ions. We discussed electron configurations for cations and anions and how to determine which shells and subshells electrons would be removed from (cations) or added to (anions). Next we discussed effective nuclear charge, we learned how to calculate the effective nuclear charge for any electron in an atom. We then discussed how our electronic structure of the atom can be applied to experimental data for atomic radii and ionization energy. We looked at atomic radii for the elements and discussed how effective nuclear charge explains the trend in atomic radii going across a period. effective nuclear charge can not be used to explain the trend in atomic radii going down a group. That trend is explained in terms of the shell the valence electrons are located. We discussed first ionization energy trends going across a period in terms of ionization energy. We also looked at 1st

In this lecture we began by reviewing responses from the ACA and the PLE. We then covered two more examples of electron configurations for Europium and for rhenium. We talked about f-orbitals and how to manage the d-orbitals for elements just before the f-block, in the f-block and following the filling of the 4f-block. We looked at the orbital shapes for 's', 'p', 'd' and 'f' orbital. Then we looked at orbital diagrams for the element hydrogen through copper. Check the podcast on Orbital Diagrams for Multielectron Atoms for a brief review of this material.