
I just had the idea that I should write an entry in which I truly crystallize what Mount Allison is all about for you. Two problems, though.
1) This is not my last entry, and so I if actually did realize my goal, then any subsequent entry would appear anti-climactic.
2) It's exam time. ERGO, I want to spend as many moments as I can preparing for my exams, in order to pass.
So you know what we're going to talk about today instead? CHEMISTRY!! Yeah, the class you took in grade eleven and were delighted to know you never had to see it again? If you're that person, stop here. This blog is not for you. But if you like chemistry, read on, and read well.
The first chapter that we studied this semester was called "Gases." Well, Raymond Chang's textbook called in "Gases," at any rate. Our professor, Vicki Meli, had another idea. "Understanding physical properties from a microscopic description." I think that she got this title from the kinetic molecular theory of gases. Basically, a gas is composed of point masses that are very far away from each other, that move randomly and undergo completely inelastic collisions, and that do not attract or repel one another. Their average kinetic energy is proportional to their absolute temperature. There's your microscopic description. It's an approximation, of course, but it's useful. Now what about these physical properties?
Let's start with the gas laws. Pressure is proportional to the inverse of volume; volume is proportional to temperature and the number of moles. First, volume is not constant. This is because it doesn't depend on the size of the particles. The kinetic molecular theory lets us treat them as point masses. Rather, volume depends on the motion of the particles, which in turn is a qualitative term for their average kinetic energy. This link between volume and average kinetic energy explains the link between volume and temperature ; kinetic energy is related to temperature. It also explains the link between volume and number of moles. Kinetic energy depends on mass, which depends on number of moles. Hence, volume must be related to the number of moles present. The fact that particles move randomly and collide with each other helps to introduce pressure into the equation. More motion means more "hitting," and more pressure. This is why increasing the temperature of a gas at constant volume will increase its pressure.
Next chapter: "Thermochemistry," or "Energy Accounting 101." The whole chapter is summed up in the first law of thermodynamics. "Energy can be converted from one form to another, but cannot be created or destroyed." It has many forms: kinetic and potential, to name the two that are key to this chapter. In chemistry, we often just call it "chemical energy." In a chemical reaction, energy is transferred either by heat or by work. The change in energy of a system has to equal the sum of the heat given off or taken in by the system and the work done by the system to the surroundings or by the surroundings to the system. The total amount of energy in the universe can't change, so it always has to go somewhere. There are a lot of ways to apply this concept: calorimetry, enthalpy, Hess's Law (that enthalpy is a state function). Basically, you can do whatever you want with energy. You must have to know where it's going.
Speaking of enthalpy, you can't have enthalpy without entropy, and you can't have either one of these lovely quantities without Gibb's Free Energy. Chapter 18: "Entropy and Free Energy," aka "Why some reactions occur spontaneously and others don't." So why do some reactions occur spontaneously while others don't? Simple. "Delta G of the system equals delta H of the system minus T delta S of the system." In other words, the change in the system's free energy is equal to the difference between the change in enthalpy and the product of the temperature and the change in entropy. Whoah. Slow down. ENTROPY???? Don't worry about entropy -- it's just a fancy term for randomness. It's proportional to the natural logarithm of the number of microstates of a system. And entropy is important because in every spontaneous reaction, the entropy of the universe increases. And in every spontaneous reaction, a system either gives off energy or gets more random. So there you have it. Heat and microstates. The two reasons that spontaneous reactions occur.
There's one more chapter that we did with Dr. Meli this semester: "Solutions," or "The whole is not always the sum of its parts." Now why in the world would Dr. Meli have picked a title like that? Basically, in a solution, you take one boring substance, mix it with another boring substance, and you get something really really cool. For example, you take copper and zinc, and you melt them enough to they'll combine, and you get this shiny new substance called brass, which is good for making brass musical instruments and brass faucets. Or you mix carbon dioxide with water and you get soda, known as "pop" in some parts of the world. You can even dissolve hydrogen gas in palladium. I'm not really sure why anybody would want to do that, but if you know why, kindly enlighten me. And weird things happen to these solutions. Like, the volume of the solvent plus the volume of the solute does not equal the volume of the solution. And when you dissolve a solid in a liquid, it looks like the solid disappears. We call this kind of solution a homogeneous solution. You can also get freezing point depression. Ever wonder why we put calcium chloride on the roads in winter? Freezing point depression, that's why. For a definition of this term, consult Wikipedia. While you're at it, look up boiling point elevation. And osmotic pressure. In fact, look up "colligative properties of solutions." Crazy stuff. Then, type "French military victories" into Google and press "I'm feeling lucky." You won't be disappointed. I swear on chemistry.
This concludes the first half of the second semester of Introduction to Chemistry. Was it dull? Be brutal.








