Learn about Advanced Placement Chemistry, Chemical kinetics 6, in this comprehensive video by bannanaiscool.
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Male Speaker: Now here is another example, if you have a net reaction, ozone breaks down in the upper atmosphere let's say to form oxygen. Now how does that actually occur? Well, here is the possible reaction mechanism, let's see if it actually makes sense. So I've got chlorine from some kind of chlorofluorocarbon, which is been released into the upper atmosphere and it damages the ozone, reacts with the ozone and forms this molecule plus and oxygen atom. Then that free oxygen atom helps to breakdown more ozone into oxygen. Then this reestablishes itself in terms of chlorine, a free atom of chlorine and then makes oxygen. Do these three equations meet the requirements for being a possible reaction mechanism? Let's say that somebody says you the rate law for this reaction is rate equals k times the concentration of O3, times the concentration of O, and you look at this and you say, where do that O come from? Well, it must be that this reaction mechanism has a slow step that has that as a rate law. Okay, let's look. The slow step is this one here. So we look at the molecules that have reacted, does this now -- this rate determining step have a rate law that actually matches the overall rate law? Yes, it does. You would write this rate law for this rate determining step exactly the same way. But here is the second requirement. Do all three of these equations add up to make this one? When you add these three equations together, Cl cancels here and here and CO2 cancels here and here, and this O cancels here and here, 2O3s make 3+1 is 3O2s, it works, it's a possible mechanism. Again because the ClO2 is actually formed and used up and the O is formed and used up, they are called intermediates. But the Cl because it actually is present at the beginning and formed at the end, it's not called an intermediate, but something as present and then formed, it's a catalyst. Catalysts speed up chemical reactions, true? Yeah, in this case, it did, it made this one faster. What else? Catalysts are not used up in chemical reactions, it's used but then made again to go out and do more damage and create another cycle of decomposition of ozone, which is a very bad thing, chlorofluorocarbons are nasty, so catalysts intermediates reaction mechanisms. Activation Energy This is really cool. You can actually calculate an activation energy for a reaction, if you know what its rate constant is and the temperature that that rate constant has been arrived at for, and you know as you change the temperature, what the new rate constant would be for a reaction. If you can do that, take that data. You can then turn it into with calculus, a little y equals MX plus B form and see if this information given will give a straight line when you graph the natural log of the rate constant K versus one over the temperature in kelvins. If you do that and get a straight line, take the slope of that line. When you take the slope of that line, it will equal negative Ea over R, but R is 8.314 joules per kelvin mole, we've dealt with that number before. Then take the slope multiplied by R, divided by negative one, and you're going to get the activation energy, that's how you can calculate that. By the way, I always ask my students this one too. I say to them, from all this data, figure out what order this reaction is overall and you look at and try to figure out and say, I can't do it, I say, that's too bad. Look, the unit for the rate constant is seconds to the negative one, so it's one over seconds, one over seconds is always first order. Yeah that was a giveaway, so be careful.
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