Week 3

This past week in AP Chemistry, we began with some extended applications of stoichiometry to real world situations. We were asked to calculate the empirical formula of a compound consisting of only carbon, hydrogen and oxygen with a known mass, given only that it had reacted with an unknown amount of oxygen gas in a combustion reaction forming certain masses of carbon dioxide and water. Up to this point we had only learned how to calculate empirical formulas given the mass percentage of each element directly, through the use of molar masses and mole ratios.

This new challenge required no additional knowledge of stoichiometry, but significantly more thought and work. One had to retrieve previous information from other units such as the law of chemical equations stating that the number of atoms of each element as well as the total mass must be equivalent on the reactant and product sides of every balanced chemical reaction. Using what we had learned in this lesson, we could solve for the mass of each of the three elements through the mass percentages of the elements in the products. Subtracting the total mass of the reactants by the given compound gave the mass of the oxygen gas that had been burned. You were left with the mass of each element in the compound being examined; the same setup as all other problems this unit.

Challenging problems such as these are greatly beneficial to our learning process in chemistry. They help students connect multiple ideas, often from different content areas together to answer a more complex problem than those they have encountered in the past. These are the best way to assess and strengthen a students true understanding of the material at hand because they must pull the information together by feel, navigating on their own, rather than simply following a step-by-step process they were given for one specific set of problems.

After being assessed on stoichiometry, we moved on to the Lewis structures of atoms (as depicted in the images below). Developed by scientist G. N. Lewis, this notation is commonly used to show the structure of molecules through the covalent bonds they form. Covalent bonds form between atoms that share electrons in order to complete their octet and become more stable like the noble gases nearest them, aside from hydrogen which only forms a duet with a single bonding pair. The octet rule is mainly a result of the eight electron spaces in their valence electron shell. The valence electrons of any atom are those in the s and p orbitals of their highest energy level. These particular electrons are the only reason that the nonmetal elements form compounds, because nonmetal elements will not form positive ions (aside from ammonium). For the best practice drawing Lewis structures, see the link below.

Lewis structure practice:
http://www.chem.purdue.edu/vsepr/practice.html

Lewis structure of HCN (left) and other common
molecules (right)