- Collision theory
- The Arrhenius equation
- Forms of the Arrhenius equation
- Using the Arrhenius equation
- Collision theory and the Maxwell–Boltzmann distribution
- Elementary reactions
- Reaction mechanism and rate law
- Reaction mechanism and rate law
- The pre-equilibrium approximation
- Multistep reaction energy profiles
- Types of catalysts
- Types of catalysts
A catalyst is a substance that increases the rate of a chemical reaction without being consumed in the reaction. A catalyst works by providing a different pathway for the reaction, one that has a lower activation energy than the uncatalyzed pathway. This lower activation energy means that a larger fraction of collisions are successful at a given temperature, leading to an increased reaction rate. Created by Jay.
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- At5:09how come [I-] can be included in the rate law? I thought the experimental rate law would only involve H2O2 since it is the only reactant shown in the overall reaction? (I- is a only a catalyst and thus not included?)(12 votes)
- I know this comment is 4 years ago, and You probably passed your high school, but it will help others
The rate of the reaction depends on the reactants of the slowest step whih is the rate-determining step, and their molecularity in that step is the power of their concentrations in the rate expression
Hope it helps(10 votes)
- At6:56, you showed that the second activation energy is the difference of energy between the initial
reactant and the second transition state. I think it should be the energy difference between the second stable state to second transition state. Please help!!(6 votes)
- You are correct. The activation energy for the second step is the energy difference between the intermediate and the second transition state.(9 votes)
- Could i more explanation about the role of catalysts in reaction? like homogenous catalysis and heterogenous catalysis ,how does it take place ?(7 votes)
- A catalyst will simply speed up a reaction with the net result of not reacting itself. They actually stabilize a reaction in order for it to reach a lower transitional energy state in the reaction. Homogenous/Heterogenous catalysts simply refer to the medium (solid/liquid/gas) but otherwise function identically.(3 votes)
- Does the valley at6:39always have to be at the same x-location as the activation energy of the non-catalyst reaction as shown in the diagram? Or can the valley be anywhere?(3 votes)
- No, it can be anywhere. In fact, those two lines should not strictly speaking be in the same graph, since the x axis is different in both cases. In other words, the catalyst actually does not lower the barrier itself, it actually takes the reactants through a completely different path.
It's like you were lost in some mountains and then some sherpa would show you a secret path to get to the other side.(7 votes)
- Catalysts don't appear in overall reaction and we usually say that amount of catalyst doesn't affect reaction rate. However here you put the catalyst in the rate equation so at least the initial concentration of catalyst affects the rate. Then can we say that the amount of the catalyst affect the rate or we just say catalyst increase the rate constant k so ignore the amount of catalyst?
- what concept do you use to determine fast and slow rate determining step?(3 votes)
- That's determined experimentally. Not even deltaG or the spontaneity can tell you how fast the reaction goes forward.(4 votes)
- Is there something called negative catalysts? Something that reduce the rate of reaction?(2 votes)
- Yes, there are negative catalysts which slow down chemical reactions. These don't work by raising the activation energy (because, if they did, the reaction would simply proceed by the uncatalysed route which has a lower activation energy). Instead, they either work by reacting catalytically with intermediates, thus breaking the reaction chain, or else by interfering in some way with a "positive" catalyst that may be present.
Importantly, negative catalysts are still catalysts in that they are not consumed in the chemical reaction and can, at least in theory, be recovered afterwards to be used again.(5 votes)
- Does the catalyst actually decrease the activation energy, or does it just split it up between 2 "hills".(2 votes)
- Can catalyst also be used to slow down a reaction?
As in previous videos i heard that 'positive' catalyst increases rate. What about 'negative' catalyst do they even exist?
I am confused about this, if we add a negative catalyst suppose it increased activation energy but why would it the reactants will follow higher energy mechanism. That is too confusing. Help me.(2 votes)
- Catalysts don't slow down a reaction per ce, they only decrease the energy required to start the reaction.
If you want to increase the energy required to start a reaction, you can use an inhibitor, which is the opposite of a catalyst.(1 vote)
- A catalyst is a substance that increases the rate of a reaction, but it itself is not consumed in the overall reaction. So let's look at the decomposition of hydrogen peroxide, so H2O2 is hydrogen peroxide, and when it decomposes you get water and also oxygen. This reaction occurs at room temperature. However, it's very slow, so this is a slow reaction. So to speed it up you need to add a catalyst and if you're doing a demonstration like the famous elephant's toothpaste demonstration in general chemistry, you need to add a source of iodide ions. So this is one of the catalysts that you could use. You could use potassium iodide or sodium iodide. So you add a source of iodide anions, and that makes this reaction fast. So your iodide anion is your catalyst; it increases the rate of a reaction. Let's take a look at the mechanism for the reaction when we add our iodide anion as our catalyst. So in the first step of the mechanism, you can see we have H2O2 and our iodide catalyst, and this forms the hypoiodite ions. So this is our intermediate, so the hypoiodite anion is our intermediate and we also are given the information that this first step of the mechanism is the slow step. And the second step of the mechanism, alright we have another molecule of hydrogen peroxide reacts with our intermediate, our hypoiodite ion and we get our oxygen, and this step is fast. Remember, for a mechanism, a possible mechanism must have elementary steps that add up to the overall reaction. So if we add our two steps together, we should get our overall reaction. So we're gonna add all of our reactants together, so that would be H2O2 plus I- plus another H2O2 plus IO- and that should give us our products. So our products, circle all of them over here, we have H2O plus IO- plus H2O plus O2 plus I-. So we have a lot going on there. Let's see what we can cancel out. So what do we have on both sides? Well we can cancel out the iodide, that's on the left and that's on the right. That's our catalyst. It increases the rate of the reaction but it's not consumed. You can see we're using it in the first step, but the iodide anion is regenerated in the second step. So overall, our catalyst is not consumed. And then we have our intermediate. The hypoiodite ion is on the left side and on the right side, so we can cancel that out. Our intermediate is created in the first step, but then it's consumed in the second step. So what are we left with here? We'll be left with two H2O2, so we have two H2O2 for our reactants, and then on the right, we would have two H2Os, so two waters and also oxygen, so plus O2. So we get back, we get back our original reaction, our overall reaction. Also, a possible mechanism must be consistent with the experimental rate law for the overall reaction. And we've seen how to do that in an earlier video. To write your rate law, you need to first recognize the rate determining step in your mechanism. And the rate determining step is the slow step in a mechanism. So step one is our rate determining step. And we can write the rate law for the reaction from the rate determining step, which we know is an elementary reaction. Alright, this is an elementary reaction, it's bimolecular so let's go ahead and write down our rate law. Alright, so the rate of the reaction should be equal to our rate constant, k, so whatever the rate constant happens to be for this, times the concentration of H2O2. So we have times the concentration of H2O2 and since we have a coefficient of one here, remember for an elementary reaction we can turn that coefficient into an exponent. So we have to the first power, and then we also have the concentration of iodide anions, so the concentration of I- and once again our coefficient is a one here, so since this is an elementary reaction, we take our coefficient and turn it into our exponent. And this is the rate law that we would predict using the rate determining step, and this is in agreement with the experimental rate law for our overall reaction. So this is a possible mechanism for our overall reaction, the decomposition of hydrogen peroxide. So how does our catalyst actually increase the rate of our reaction? Well let's look at an energy profile for our uncatalyzed reaction, so the conversion of hydrogen peroxide into water and oxygen. And here's the uncatalyzed version. We start with a certain energy for our reactants and we know at the top here, that represents the energy of the transition state. And the difference between those two would be our activation energy, alright. So that represents our activation energy for our uncatalyzed reaction. When we add the source of iodide ion, when we add our catalyst, this actually provides a different mechanism, a lower energy mechanism, and we know that mechanism occurred in two steps, so let me go ahead and sketch this. The energy of our reactants is the same, right, It's the same level, but we're going to decrease the activation energy, so let me go ahead and draw this in here. So it might look something like this, I'm sure I'm not drawing it perfectly. So let's say that's what our energy profile looks like with the addition of our catalyst. So this, this would be the transition state for the first step of our mechanism, and you can see the activation energy has decreased. So here we have a lower activation energy. So remember from an earlier video if you decrease the activation energy you increase the rate of your reaction, which you can see with the Arrhenius equation. So that's what the catalyst does, and then we reach this point right here, this valley if you will would represent the energy of the intermediate, the hypoiodite ion, and then, we have a second activation energy for the second step of our mechanism. So this would be the first step of our mechanism, so I'm gonna write Ea1, and this would be the activation energy for the second step of our mechanism which is Ea2. So the activation energy for the first step is higher because that's our rate determining step. So again, the catalyst does not affect the energy of your reactants or of your products. It's still the same energy for both of those. What the catalyst does is decrease the necessary activation energy which increases the rate of your reaction.