WEBVTT 00:00:04.420 --> 00:00:08.850 If you're driving down the highway and want to see how fast you're going, you look down 00:00:08.850 --> 00:00:13.450 at your speedometer. If you're in a lab and want to see how fast a chemical reaction is 00:00:13.450 --> 00:00:18.380 going, it's a little more complicated. In this video we'll look at a few factors that 00:00:18.380 --> 00:00:21.730 influence the speed of chemical reactions. 00:00:21.730 --> 00:00:26.039 This video is part of the Equilibrium video series. It is often important to determine 00:00:26.039 --> 00:00:30.680 whether or not a system is at equilibrium, to do this we must understand how a system's 00:00:30.680 --> 00:00:34.720 equilibrium state is constrained by its boundary and surroundings. 00:00:34.720 --> 00:00:39.970 Hi, my name is George Zaidan and I'm a graduate of the chemistry department here at MIT. 00:00:39.970 --> 00:00:45.070 Before watching this video, you should have a basic understanding of chemical equilibrium. 00:00:45.070 --> 00:00:48.390 After watching this video, you will be able to: 00:00:48.390 --> 00:00:52.820 Understand how reaction rate is influenced by reactant concentration 00:00:52.820 --> 00:00:57.309 Explain how reaction rates change as a system establishes equilibrium, AND 00:00:57.309 --> 00:01:02.129 Predict relative equilibrium concentrations of reactant and product, based on rates of 00:01:02.129 --> 00:01:09.129 forward and reverse processes. 00:01:09.880 --> 00:01:16.880 Suppose you have this general reaction where A, B, C, and D are molecules and lowercase 00:01:16.920 --> 00:01:23.920 a, b, c, d are their molar coefficients. As the reaction progresses, A and B will be 00:01:24.259 --> 00:01:27.119 consumed, and C and D will be formed. 00:01:27.119 --> 00:01:32.158 The speed at which any of these four processes happen multiplied by the reciprocal of the 00:01:32.158 --> 00:01:37.158 appropriate molar coefficient is called the rate of this reaction. 00:01:37.158 --> 00:01:42.149 If you want your car to go faster, you step on the gas, but speeding up a chemical reaction 00:01:42.149 --> 00:01:47.499 is not that simple. There are lots of different factors that affect reaction rate. Three common 00:01:47.499 --> 00:01:52.609 ones are: the concentration of reactants, the temperature of the reaction mixture and 00:01:52.609 --> 00:01:57.679 the presence of a catalyst. In this video we'll focus on the effect of concentration 00:01:57.679 --> 00:02:00.460 on homogenous reactions. 00:02:00.460 --> 00:02:04.670 Most chemical reactions are the result of two or more molecules colliding with each 00:02:04.670 --> 00:02:10.330 other... but not every collision leads to a reaction. The molecules must collide in 00:02:10.330 --> 00:02:15.510 the proper physical orientation, and they must do so with enough energy to break their 00:02:15.510 --> 00:02:17.040 bonds. 00:02:17.040 --> 00:02:22.200 Imagine a beaker with 100 mL of water, in which you dissolve 106 molecules of reactant 00:02:22.200 --> 00:02:29.200 A and 106 molecules of reactant B. How likely is the reaction to form C and D? Do a back 00:02:29.430 --> 00:02:36.430 of the envelope calculation to support your answer. Pause the video. 00:02:39.439 --> 00:02:46.439 So let's see. 100 mL of water is 100 grams of water, which is 100 over 18.02, about 5.5 00:02:46.610 --> 00:02:53.610 moles of water. That means there are about 3.3x1024 molecules of water in the beaker, 00:02:53.629 --> 00:03:00.629 and only 106 molecules of each reactant...So in any given volume of this solution, 3.3x1024 00:03:01.450 --> 00:03:08.450 over (3.3x1024 + 106 + 106), or 99.99999999999999994 %, of molecules are nonreactive water! 00:03:22.720 --> 00:03:28.370 Clearly, successful collisions between the reactants would be very rare in this situation. 00:03:28.370 --> 00:03:32.980 Intuitively, you would expect that the more concentrated the reactants, the faster the 00:03:32.980 --> 00:03:35.680 reaction, and this is generally true. 00:03:35.680 --> 00:03:42.680 Let's draw a reaction coordinate diagram to better understand our hypothetical reaction. 00:03:45.390 --> 00:03:52.390 Let's assume that the reaction is exothermic, in other words, that energy is given off as 00:03:55.709 --> 00:03:58.629 the reaction progresses, so we draw the curve like this. 00:03:58.629 --> 00:03:59.189 The x-axis is the progress of the reaction, and the y-axis is potential energy. 00:03:59.189 --> 00:03:59.560 Here are the reactants, and here are the products. 00:03:59.560 --> 00:04:05.870 This hump is the activation energy, the point of highest potential energy of the reaction. 00:04:05.870 --> 00:04:11.319 Two molecules must collide with at least this amount of energy in order to successfully 00:04:11.319 --> 00:04:12.120 react. 00:04:12.120 --> 00:04:17.798 Do you notice anything about this diagram? There's no indication of directionality; in 00:04:17.798 --> 00:04:22.650 other words, there's no reason that the reaction couldn't proceed backwards just as well as 00:04:22.650 --> 00:04:24.260 forwards. 00:04:24.260 --> 00:04:28.290 You're probably used to thinking of chemical reactions as processes that happen in only 00:04:28.290 --> 00:04:35.060 one direction (forward), but most reactions are actually reversible. There are a few exceptions, 00:04:35.060 --> 00:04:42.060 for example combustion. *You can't un-burn a match. But most reactions that chemists 00:04:42.620 --> 00:04:47.350 carry out in the lab or that happen in our bodies are reversible. 00:04:47.350 --> 00:04:51.250 Draw your own reaction coordinate diagram and label the forward and reverse reaction 00:04:51.250 --> 00:04:56.650 paths. Suggest what the relationship might be between the activation energy and the relative 00:04:56.650 --> 00:05:03.650 rates of the forward and reverse reactions. Pause the video. 00:05:07.360 --> 00:05:11.720 This is the activation energy for the forward reaction. The higher this activation energy, 00:05:11.720 --> 00:05:16.160 the slower the forward reaction, because the number of reactant molecules with sufficient 00:05:16.160 --> 00:05:20.360 energy to react when they collide will decrease. 00:05:20.360 --> 00:05:25.600 But now let's look at this picture in reverse. The reverse activation energy is not the same 00:05:25.600 --> 00:05:31.100 as the forward activation energy! In this case, it's higher. So that means that the 00:05:31.100 --> 00:05:35.140 reverse reaction will be slower than the forward reaction. 00:05:35.140 --> 00:05:40.700 If we were to somehow increase the activation energy, both the forward and the reverse rates 00:05:40.700 --> 00:05:46.500 would slow down, but the relationship between them -- namely that the forward rate is faster 00:05:46.500 --> 00:05:49.650 than the reverse rate -- would be preserved. 00:05:49.650 --> 00:05:54.850 Here are four other diagrams, each for a different hypothetical reaction. Predict the relative 00:05:54.850 --> 00:06:01.850 rates of the forward and reverse reactions in each of these cases. Pause the video. 00:06:06.520 --> 00:06:13.520 In 00:06:29.720 --> 00:06:33.930 real life chemical processes, the rates of the forward and reverse reactions are often 00:06:33.930 --> 00:06:39.520 very different. Many of the reactions that you may have previously thought of as "irreversible"actually 00:06:39.520 --> 00:06:44.750 just have wildly different forward and reverse rates. For example, you've probably seen the 00:06:44.750 --> 00:06:48.720 dissociation of a strong acid in water written like this. 00:06:48.720 --> 00:06:55.180 But really, it's this. In this case, the forward reaction is many orders of magnitude faster 00:06:55.180 --> 00:07:02.180 than the reverse reaction, so we write it as just the forward reaction. 00:07:04.940 --> 00:07:10.300 So far we've seen that concentration and activation energy can each independently affect the rate 00:07:10.300 --> 00:07:15.150 of a chemical reaction. The higher the concentration of a reactant, the faster the reaction; and 00:07:15.150 --> 00:07:20.490 the lower the activation energy, the faster the reaction. But there's a twist... Thinking 00:07:20.490 --> 00:07:27.490 about the hypothetical reaction A and B yields C and D. Suggest what that twist might be. 00:07:29.810 --> 00:07:36.810 Pause the video. 00:07:37.620 --> 00:07:42.050 There's no reason that the activation energy of this reaction would change as the reaction 00:07:42.050 --> 00:07:49.050 progresses. But that's not true for the concentrations: as A and B are converted to C and D, the concentration 00:07:49.300 --> 00:07:56.300 of all four species change! And as the concentrations change, the reaction rate also changes. 00:07:56.810 --> 00:08:01.710 To see if we can understand how the reaction rate changes as the reaction progresses, let's 00:08:01.710 --> 00:08:06.930 go back to the reaction coordinate diagram. Let's start by considering both the forward 00:08:06.930 --> 00:08:11.210 and reverse reactions as if they're completely separate. 00:08:11.210 --> 00:08:18.210 Here are A and B reacting to form C and D. Let's assume this forward reaction is relatively 00:08:18.270 --> 00:08:24.490 fast. As more and more A and B get converted to products, their concentrations decrease, 00:08:24.490 --> 00:08:30.950 and the initially fast reaction rate slows over time. Here are C and D reacting to form 00:08:30.950 --> 00:08:37.950 A and B. This reaction would be slower than this one, but as C and D react, their concentrations 00:08:38.649 --> 00:08:43.099 decrease with time and the reaction becomes even slower. 00:08:43.099 --> 00:08:50.099 Except that in reality, these aren't two separate reactions. The products of one reaction are 00:08:50.970 --> 00:08:55.649 the reactants for the other. Let's now consider the forward and reverse 00:08:55.649 --> 00:09:01.600 reactions at the same time. At the very very beginning, only A and B are present -- no 00:09:01.600 --> 00:09:06.680 C or D. So the initial rate of the forward reaction will be relatively high, since the 00:09:06.680 --> 00:09:12.740 concentrations of A and B are high. The reverse reaction can't happen at all yet, because 00:09:12.740 --> 00:09:17.279 there is no C and D present, so its initial rate is zero. 00:09:17.279 --> 00:09:23.629 As the reaction proceeds, two things happen: the concentrations of A and B decrease, and 00:09:23.629 --> 00:09:29.920 the concentrations of C and D increase. So the forward reaction rate slows, and the reverse 00:09:29.920 --> 00:09:34.319 reaction rate speeds up. Eventually, we reach a point where the rates 00:09:34.319 --> 00:09:39.839 of the forward and reverse reactions are the same: this is the definition of a dynamic 00:09:39.839 --> 00:09:46.670 chemical equilibrium. When this happens, the concentrations of A, B, C, and D stop changing 00:09:46.670 --> 00:09:53.670 with time. Be careful not to confuse equilibrium with "no reaction."A and B are still reacting 00:09:54.500 --> 00:10:00.290 to form C and D; and C and D are still reacting to form A and B. But the rate of formation 00:10:00.290 --> 00:10:05.100 of products equals the rate of disappearance of reactants and vice versa. That means that 00:10:05.100 --> 00:10:10.310 the concentrations of A, B, C, and D don't change with time. 00:10:10.310 --> 00:10:15.180 Let's go back to our reaction coordinate diagrams for our four hypothetical reactions. We've 00:10:15.180 --> 00:10:20.459 already worked out the relative rates of the forward and reverse reactions in each case. 00:10:20.459 --> 00:10:24.730 Using these relative rates, see if you can predict the relative concentrations of the 00:10:24.730 --> 00:10:31.730 reactants and products at equilibrium. Pause the video. 00:10:36.279 --> 00:10:43.220 -Endo, high Ea: lots of reactant, not much product 00:10:43.220 --> 00:10:47.379 -Exo, high Ea: lots of product, not much reactant 00:10:47.379 --> 00:10:54.149 -(In both of these cases, the forward rate is the same as the reverse rate...)Same energy, 00:10:54.149 --> 00:11:01.149 high Ea: ratio of product/reactant will depend on specific reaction Same energy, low Ea: 00:11:08.569 --> 00:11:11.680 ratio of product/reactant will depend on specific reaction 00:11:11.680 --> 00:11:16.819 Now what's the difference between the last two scenarios? In each case, the forward reaction 00:11:16.819 --> 00:11:21.050 rate will be the same as the reverse reaction rate. But the rates for this scenario will 00:11:21.050 --> 00:11:27.480 be much faster than for this one... so equilibrium will be reached much more quickly here than 00:11:27.480 --> 00:11:31.180 here! The most important thing to understand about 00:11:31.180 --> 00:11:35.769 reaction rates and equilibrium is that just because the forward and reverse rates are 00:11:35.769 --> 00:11:42.629 the same DOES NOT mean that the concentrations of the reactants and products are the same. 00:11:42.629 --> 00:11:47.129 Some equilibria, like the dissociation of a strong acid in water, strongly favor the 00:11:47.129 --> 00:11:53.740 products. Others strongly favor the reactants. Many of the reactions that keep us alive are 00:11:53.740 --> 00:11:57.449 equilibria, and our body goes to a great deal of effort to make sure that the position of 00:11:57.449 --> 00:12:04.449 equilibrium heavily favors one side or the other. 00:12:07.089 --> 00:12:10.749 We hope this video has helped you understand the relationship between reaction rates and 00:12:10.749 --> 00:12:17.009 equilibrium. We saw that, in general, the rate of a reaction decreases as reactant concentration 00:12:17.009 --> 00:12:18.040 decreases. 00:12:18.040 --> 00:12:23.079 And while we may think of reactions as irreversible, most are actually reversible, it's just that 00:12:23.079 --> 00:12:25.620 we "see"the faster of the two. 00:12:25.620 --> 00:12:30.899 As the forward reaction rate decreases, the reverse reaction rate increases until equilibrium 00:12:30.899 --> 00:12:35.949 is established. At equilibrium, the rate of the forward reaction equals the rate of the 00:12:35.949 --> 00:12:41.059 reverse reaction, but this does not imply anything about the equilibrium concentrations 00:12:41.059 --> 00:12:44.399 of the products and reactants.