Experimental Procedure and Data
1. For each compound,
a. Follow the directions for and write the correct (best) Lewis structure.
b. Obtain your instructors approval, then build a molecular model from the kits provided.
c. Answer the questions that describe the molecule.
2. Atoms are color coded within each kit. It may be beneficial to evaluate whether you would like to use an “atom” by the type listed or by areas of e density. In some instances (such as with carbon) the areas of e density do not vary by molecule so using the atom from the kit designated as carbon should work when making your molecular model. In other instances, such as with PCl5, it may be more beneficial to look for an atom which can hold 5 areas of electron density for your central atom (there is most likely no “atom” set aside for phosphorus in the kits).
3. You might first want to do your work on a scratch piece of paper. However, if you choose to do this, transfer your answers to the data sheets provided.
Pre-lab Assignment/Questions
- * Note– this pre-lab must be finished before you come to lab.
1. Draw the dot structures for C,H,O, Cl, N, S, and P.
2. If covalent bonding occurs because an atom wants to achieve an octet and therefore fill empty spaces in its orbital, how many covalent bonds would you think are formed by each of the atoms in #1?
3. In some molecules the electron geometry and the molecular shape are the same, but in other molecules they are different. How do the Lewis structures of these two types of molecules differ? Look at Table 1 in the introduction, and consider for example CH4 and NH3.
4. List four elements that always obey the octet rule. List two elements which do not typically obey the octet rule.
5. From your knowledge of its electron structure, why doesn’t hydrogen obey the octet rule?
| Molecule | #Val. e- | Lewis Structure | Areas e Density | Electronic Geometry | #Bonding Areas | # Non-Bonding Areas | Molecular geometry | Bond Angle(s) | Polar Molecule? Yes or No |
|---|---|---|---|---|---|---|---|---|---|
| NH2– | |||||||||
| CH4 | |||||||||
| NO2– | |||||||||
| CO | |||||||||
| CO2 | |||||||||
| H2CO |
| Molecule | #Val. e- | Lewis Structure | Areas e Density | Electronic Geometry | #Bonds | #Lone Pairs | Molecular geometry | Bond Angle(s) | Polar Molecule? Yes or No |
|---|---|---|---|---|---|---|---|---|---|
| SO2 | |||||||||
| XeF4 | |||||||||
| PCl3 | |||||||||
| NH4+ | |||||||||
| SCN– | |||||||||
| SF4 | |||||||||
| NO3– | |||||||||
| SO42- | |||||||||
| SF3+ | |||||||||
| XeF3+ |
Post-lab Questions
1. Without making a model, describe the electron geometry and molecular shape of carbon tetrabromide (CBr4). Would you expect the bonds in this molecule to be polar? Would you expect this molecule to be polar overall? Explain.
2. NH3 and H2CO each have three bonds about the central atom. However, their molecular geometries are not the same. Explain this difference.