Factors that Effect Rate

Transcription

1 Factors that Effect Rate Now that we know what has to happen for a reaction to take place (a collision between molecules with the correct orientation and minimum activation energy), let s examine what factors we can manipulate in order to change the rate of the reaction. Temperature We ll start by looking at a simple reaction; the dissolving of an Alka-Seltzer tablet in water. When you plop the tablet in water, the water molecules dissolve the citric acid and sodium bicarbonate in the tablet and they react to form sodium citrate, water, and carbon dioxide gas which causes the fizz. We need a baseline experiment so we ll fill a beaker with 500 ml of tap water and put in an Alka-Seltzer tablet and see how long it takes for it to dissolve. Citric acid and sodium bicarbonate react to form carbon dioxide gas which causes the fizz effect. HC 6 H 7 O 7 + NaHCO 3 > NaC 6 H 7 O 7 + H 2 O + CO 2 (g) This can only happen in an aqueous environment, though, so as soon as the water dissolves the tablet, the citric acid and sodium bicarbonate are allowed to react. The faster the water can dissolve the tablet, the faster the reaction goes. financial_hangover/ Temperature ( o C) Time (seconds) seconds. Great. Now let s repeat the experiment except this time we ll use 500 ml of boiling hot water.

2 In this experiment we are using 500 ml of boiling hot water. The time is reduced to 20 seconds. Temperature ( o C) Time (seconds) One last experiment by using the same amount of water that is ice-cold. The time jumps up to 84 seconds. Temperature ( o C) Time (seconds) 2 84 To summarize: hotter temperatures make reactions take place faster. Colder temperatures make reactions take place slower. But why?

3 Number of molecules with a given energy First of all, hotter temperatures make the molecules move faster. The faster the molecules move, the greater the number of collisions will take place. The greater the number of collisions that occur, the greater the chance that the molecules will collide with the correct orientation and thus increase their chances of having the collision be effective. More importantly, though, is the fact that a higher temperature means the molecules will be colliding with each other with more energy because the kinetic energy or speed of the molecules is directly proportional to the temperature. Consequently, a greater number molecules will have the minimum activation energy to make the reaction proceed. Let s look at a graph: 300 K E a, activation energy In this example, we see that none of the molecules have enough energy to get past the activation energy barrier, E a. Thus the reaction is NOT proceeding. Increasing Energy

4 Number of molecules with a given energy Number of molecules with a given energy 300 K 500 K E a, activation energy By increasing the temperature, the distribution of molecular speeds widens. Fewer molecules are at any one speed but there is a wider range of speeds. Consequently, some of the molecules have enough energy to get over the activation energy barrier and thus the reaction occurs slowly. Increasing Energy 300 K E a, activation energy 500 K 700 K Increasing the temperature even more allows for more molecules to have the minimum activation energy necessary so the reaction proceeds even faster. Increasing Energy A faster temperature means the molecules move faster. Thus, there are more collisions and thus a greater chance for a collision with the correct orientation. At the higher temperature, more molecules are past the activation energy line. The reaction proceeds faster.

5 Surface Area Let s stay with the Alka Seltzer experiment for just a minute more. In addition to changing the temperature of the system, we could also grind up the tablet in a mortar. If we compare the time it takes for the ground up tablet to the whole tablet we see a big difference in dissolving time. The ground up tablet reacted a lot faster. Why? Tablet Whole Ground up Temperature ( o C) Time (seconds)

6 Remember that in order for a reaction to proceed, the particles have to collide with each other. When the table is whole, the particles on the inside of the tablet cannot collide with the water molecules until the particles on the outside of the tablet have reacted away. They have to wait. Although they have enough energy to react, they cannot because no contact can be made. When the tablet is ground up, it exposes more surface area of the reactive molecules. This allows for more molecules to collide with each other at the same time as no molecules have to wait until they are exposed. Thus, the reaction proceeds much faster. Water molecules can hit the outer layer of the tablet... but not these here in the center But now that the tablet has been ground up, the water molecules can react with more of the molecules in the tablet at once. Thus, the reaction is faster. More surface area means more collisions can take place at the same time. More collisions means there is a greater chance for an effective collision so the rate increases. Every year there are industrial explosions at factories, farms, and other places that generate a lot of dust. This could be from something simple like flour particles to simple dust kicking up from falling grain, corn, etc. All the exposed surface area from the minute, little pieces can turn a nearby tiny spark into a huge explosion!

7 Concentration What would happen if I put a piece of magnesium ribbon into a beaker of hydrochloric acid? The magnesium would react to form magnesium chloride and hydrogen gas: Mg (s) + 2 HCl (aq) > MgCl 2 (aq) + H 2 (g) Let s do an experiment where we take a piece of magnesium ribbon and put it in two beakers that have the same amount of HCl at the same temperature but have different concentrations. We ll time to see how long it takes for the Mg to dissolve. In the first beaker we have 1 M HCl and in the second beaker we have 6 M HCl. As we can see from the results, it took a lot longer for the 1 M HCl to react away the magnesium. Why? Concentration (M) Time to Dissolve Mg Piece (seconds) As we have stated before, in order for a reaction to proceed, the particles must collide. The more collisions there are, the better chance there is to have an effective collision and so the reactions will speed up. As we can see in the diagrams below, there are more particles in the concentrated solution than in the dilute one. When the magnesium is put into the beakers, there is a much greater chance of a collision occurring in the concentrated one. Thus, the reaction goes faster in the concentrated solution. Dilute Concentrated A greater concentration means there is a greater chance of a collision. The greater the chance of a collision means the reaction will proceed faster.

8 Catalysts Let s examine a car trip that a salesman has to make every week. He starts in his home city and has to drive over a big hill to get to his destination to sell his products. He knows it is going to be profitable but, still, he has to drive his car over that great big hill all the time. Start Finish What can he do about it? By himself, nothing. But if he can exert enough influence, maybe he can get somebody to build a tunnel through the hill so nobody has to go over it. Start Now, it requires much less energy for the car to go from start to finish and it will take the salesman less gas and time to make the trip! Finish

9 The previous page was an analogy of a catalyst a substance that speeds up the rate of a reaction but is not ultimately consumed by the reaction. A catalyst provides the reaction an alternate pathway which lowers the activation energy, making it easier for the reaction to proceed. We can see these effects on the two graphs we have looked at thus far during the unit. The activation energy is much higher for the uncatalyzed pathway here.. than for the catalyzed pathway here. On this graph we see that the original activation energy is here.... but by using a catalyst the activation energy drops to this spot, allowing far more molecules to have the activation energy and thus the reaction goes faster.

10 Remember that catalysts participate in but are not ultimately consumed by the reaction. What that means is that catalysts are often used up in one step and then re-produced in another step. The decomposition of ozone by chlorofluorocarbons (CFCs) is a perfect example of this. We see that Cl is acting as a catalyst. It is a reactant here in step 1 Overall: O 3 + O > 2 O 2 Step 1: Cl + O 3 > ClO + O 2 Step 2: ClO + O > Cl + O 2 And it is a product here in step 2. It does not appear in the overall equation. Step 1: Cl + O 3 > ClO + O 2 Step 2: ClO + O > Cl + O 2 When the 2 steps are added together, the Cl drops out of the equation. It is also good to note here that the opposite of a catalyst is also present in this reaction. This is an intermediate a substance that is produced in one step and then re-used in a subsequent step. Thus neither intermediates or catalysts appear in the overall equation. Step 1: Cl + O 3 > ClO + O 2 ClO is an intermediate. It is made here in step 1... Step 2: ClO + O > Cl + O 2 and re-used here in step 2. It does not appear in the overall equation. Step 1: Cl + O 3 > ClO + O 2 Step 2: ClO + O > Cl + O 2 Combined: Cl + ClO + O 3 + O > Cl + ClO + 2 O 2 Overall: O 3 + O > 2 O 2 Note the catalyst in red and the intermediate in yellow drop out. A catalyst lowers the activation energy by providing a different pathway for the reaction. The reaction speeds up because more molecules have the minimum activation energy. The catalyst is consumed and then re-produced in the reaction so does not appear in the overall equation.