Saturday, October 02, 2010

Chapter 6 Preview

As I read and prepare chapter 6 (the students are the not the only people responsible for preparing the material) I am reminded that this chapter is really critical to a fundamental understanding of beginning chemistry. Here is why:

You learn how to balance equations in chapter 5. Fair enough. You learn that the relationship between number of molecules on each side of an equation is also equal to the number of moles of each compound on each side of the equation. Let's review a simple example:  O2(g)  +  2H2(g)  = 2H2O
Let's add atoms on each side of the reaction: 2 atoms O and 4 atoms H on the left while there are 2 atoms O and 4 atoms H on the right side of the equation. It is balanced. What does this mean exactly?
For every one molecule of oxygen gas you must have 2 molecules of hydrogen gas to get 2 molecules of water. Can you get less than two molecules of water? No- because you can't cut a molecule in half. Likewise, you must have one mole of oxygen gas for two moles of hydrogen gas to get two moles of water. The relationships, as you can see, between molecules of compounds and moles of compounds are equal.

This is what enables us to use these relationships in the laboratory in a practical way. From these molecule/molar relationships we can calculate the amount of product we should get in a chemical reaction. Returning to our synthesis of water example:  If we start with two moles of oxygen gas and unlimited hydrogen gas how much water should we get? The mole relationship between oxygen gas and water is needed to figure this out:( 2 moles O2) X( 2 moles H2O/1 moles O2 )= 4 moles H2O. This is the maximum yield. What if you go into the lab and determine that this reaction only yields 3 moles of H2O? Then we get into a calculation of percent yield. The percent yield is the comparison of the actual yield you obtain in the lab with the theoretical yield you calculate based on molar ratios of a balanced equation. Percent yield is as follows: actual/theoretical X 100. In this case it is 3/4 X 100 or 75%. You obtained 75% of the maximum amount of material you could get from this reaction with these amounts of reactants.

This chapter also introduces energy transformations. Just as mass is conserved (atomic theory), energy is conserved in nature. This means that energy is transformed from one form to another. This applies to chemical reactions as well as everyday objects like a ball bouncing on the ground. There are a few key terms here you should really learn well: exothermic vs endothermic.
Exothermic: This means a reaction (or process) gives off energy
Endothermic: This means a reaction (or process) absorbs energy
The diagrams for each of these are in Chapter 6. Familiarize yourself with the energy levels of products and reactants for each case.

And...... my favorite thing in chemistry is introduced. That is Calorimetry. What is calorimetry? It is the measure of the transfer of heat.
There is a very simple demonstration you can do in the classroom to show calorimetry. I haven't figured out how to incorporate it into our class- the ice would melt by the time I get to the topic in class.
You take a styrofoam cup, thermometer, and ice/water. You measure the temperature of the water initially. Then, put a  piece of ice into the water. Let the system come to equilibrium (even temp) and measure the temperature again. This is the temperature change you plug into the equation q= mC(Tf-Ti). This measures the transfer of heat from the surrounding water to the melting ice. In this case, the ice is the system and the water is the surroundings. You are measuring the heat that is being transferred from the water to melt the ice.
In the equation: q (heat transferred) = (mass of ice+cup+water)(C-specific heat of water)(Tf-Ti)
Heat transferred is equal to the mass of the entire system  X 4.184 (constant value) X change in temp.

If you understand this basic experiment it will help you understand the concept of heat transfer.

See you Tuesday.