Captain Nemo's Guide to the Chemically Impaired: Chapter 6 By: Captain Nemo November 30, 1998 Greetings and Salutations. If you are reading you: a). Have too much free time b). Want to run out your print cartridge c). Actually learned from my last guide (comments can be sent to the webmaster of the site you got this from). Anyway, we are moving on from the Land of Huge Equations to the Land of Gases. Just a warning: Le' Professeur Harkins may combine 6-7 in one. This guide will be updated accordingly. With that out of the way: Onward Ho! For starters, there are three things (the Key Three) you can measure gases in: Volume (cm3), Temperature (degrees Celsius, or Kelvin -->to be explained), and Pressure. Since going in that order would probably be the easiest, the book goes pressure-volume, which we will follow (regretfully). Now, some might be asking: "What about sight, smell, and all that other nifty stuff we do in labs? Can't we measure gases in that?" Well, as far as I know, there is no SI measurement for smell (odor? milliordor? kiloodor? The nasal scale?). But while they may be important in labs, they are mostly meaningless in equations. On to pressure. Pressure is the first of the Key Three. But fortunately for you, the Captain has discovered that pressure is so easy, you won't need to break a sweat. No sweat, I promise. The book describes pressure as force over a given area (gm/cm3). While this definition is true, gas is not measured in that unit for the rest of the unit. Instead the book jumps to my favorite part: The ping pong ball analogy. Read it. Then ask yourself: How did the ping pong balls get there? How do they move around? Why isn't it that when you shake it they don't all move in the same direction? And who is the crack addict who though of this? Well, to sum it up: Higher temperature=more particle movement=more pressure. Another analogy: If two women weigh the same, but women A has high heels and women B is in sneakers, who will exert more force if they step on rapist C's foot? Get it? Well, it was a trick question. They both exert the same force, but the high heels have less area, so they exert more pressure. So Women A will break his foot while women B will waste her energy until....never mind. Let's end that rant. Anyway, now that you understand pressure (hopefully) you can work on measuring pressure. On the subject of measuring pressure, the book suggests a barometer (a device that measures atmospheric pressure). Of course we won't have a barometer on the test, but it's helpful to know it, since the other contraptions in this chapter are just more complex versions. It's filled with mercury, which is how one describes measurements on a barometer. Which brings us to the most useful part of the chapter: The Table of Gas Measurement(p.199). The Captain suggests a). Memorize the table, b). Put it on a note card, or c). Ask yourself again why you are reading this. All three are very productive. Actually, this stuff is nothing very difficult. As a matter- o- fact, it's nowhere near difficult. Really, it is no where near the area where difficult might be in......never mind. You get the point. p. 200 Practice (spoiler warning!!!) 1. 74 cm Hg * 10mm Hg * 1kPa * 0.009869 atm = 0.980 atm 1cm Hg 7.509mm Hg 1 kPa Remember: When converting-----> What you want What you have 2. 97.5 kPa * 0.009689 atm = 0.962 atm 1 kPa Doesn't get much easier than that. Know the table and remember you conversion. You'll have no problem. As a matter-o-fact you can stick your face in your book and shout "Ha ha you stupid book! I understand this! And I can do this quicker than I can turn a page on your sorry butt." Or not. Broken a sweat yet? I didn't think so. Alright, now we're rolling. And guess what ladies and gentlemen: It get's even easier. Let's look at Dalton's Law of Partial Pressure (DLPP). Intimidating name, yes; hard concept, no. As a matter of fact, this law actually helps more than it hurts. It says that the pressures of gasses do not effect each other, so the total pressure of the container (whatever is holding the gases you are measuring) is equal to the sum of the pressures of each individual gas. Wow! Ladies and Gents, this is fantastic. All chemistry in this part of the chapter has been reduced to addition and subtraction problems. Yes, you heard me. Let's examine one. Example 6-2, p. 203 (yes, I know they show how to do this, but bear with me.) First thing we need is some translation: "...collected over water" - In addition to the gases given, their will be water vapor in the container "...water levels inside and outside the collection bottle are equal at the end of the experiment" - The pressure inside the bottle = the pressure outside the bottle. So let's see what we have: 99.4 kPa Atmospheric pressure 3.0 kPa Water vapor pressure Because of our translation, we know that the pressure of the atmosphere (not measured in atm) is equal to the pressure in the bottle. So we can add that we know: 99.4 kPa Total pressure So: 99.4 kPa (total) - 3.0 kPa (H2O vapor) = 91.4 kPa of Hydrogen gas Now you must be shaking your head. No way can honors chemistry be subtraction? Well guess what, it is. I told you, you won't break a sweat. But for all of you busy bodies, here's another (spoiler alert!!!). #3 p. 204 "...described in the example..." - Collected over water, with equal pressure inside and out. 99.4 kPa - 4.5 kPa = 94.9 kPa That's right folks: Subtraction. And #4 is the same without water vapor. Hey! We're done half the chapter. No sweat right?:^)