hello everybody and welcome to this a level chemistry video about equilibria which is one of the year 1 topics in this video we will take a look at dynamic equilibrium as applied to reversible reactions we will look at Lycia ateliers principles and how equilibria shift when conditions change and we'll use those ideas to look at industrial examples of how we can maximize yields using our knowledge of a chatelier's and then in a separate video we will take a look at the equilibrium constant KC and associated calculations how you work out the units and how we can affect the value of KC so check that video out after you've watched this one in chemistry we would typically think of a reaction as one where you start with the reactants and the reaction proceeds and we make the products however in reality many reactions are actually reversible and that means that they go both ways they go forwards and they go backwards and to symbolize that we don't have the forwards facing arrow we would have a reversible reactions arrow like this one of the reversible reactions that you're likely to be most familiar with is when hydrated copper sulfate is heated and is converted into and hydrous copper sulfate and water you will probably have seen this reaction in the energetics topic where you were measuring enthalpy changes that can't be measured very easily directly by doing other chemical reactions such as dissolving now in that reaction the forward reaction is endothermic meaning it the temperature drops and in this case it requires being heated to make it work and the reverse reaction the backwards reaction is exothermic meaning that when you add water to the anhydrous copper sulfate the temperature is going to go up and that is a characteristic of any reversible reaction one of the directions will be exothermic and the other will be endothermic by the exact same amount it's just one of them the energy will go up one of them the energy will go down it doesn't have be backwards is 1 and forwards as the other it can totally vary now equilibrium can only be reached in what is known as a closed system and if you remember from the energetics video we discussed that that is where energy can be exchanged between the chemicals and the surroundings but matter cannot and that could be something as simple as just having a lid on a beaker so no gases can go in or out reversible reactions are usually referred to as existing in a state of dynamic equilibrium and in fact that could be a to mark question where you're asked to define what a dynamic equilibrium is and so there's obviously two parts to the answer and so the first part is that the forward and the backward reactions are both occurring at the same rate and that's the origin of the dynamic part of the definition that the forwards reaction is happening and the backwards reaction is happening now the equilibrium part is really sort of like a two-fold part to it the fact that it's at the same rate that's one aspect of the word equilibrium and then the other aspect of the word equilibrium is that the concentrations of all the reactants and products is constant and so you might not have the same amount of both but the amounts that you have of reactants and products is unchanged now the best way of visualizing this is actually to look at a graph if we look at a typical graph of a one direction reaction where reactants turn into products only over time the concentration of reactants will obviously decrease as the reaction goes on it decreases really quickly at first then begins to slow down and then it stops when one of the reactants or both of the reactants or all of the reactants have been used up however in a reversible reaction that is different if we look at the concentration of the reactants that follows the same sort of path of decrease but over time it decreases more slowly until the point where it levels out and it doesn't level out at zero who knows exactly where it levels out but it's definitely not zero there is definitely some reactance left over if we were to map the concentration of products in the same sort of way then obviously at the beginning of the reaction we have zero concentration of products quite quickly we make some products and then the rate of making of products decreases because the amount of reactant that we've got left over has decreased so the rate of the forward reaction is going to slow down so we're going to make less product and then after a time the concentration of products will level out and not change anymore and that coincides with the same place where the reactants doesn't change anymore and this is the point that equilibrium has been reached now to repeat it's a dynamic equilibrium because the forwards and the backwards reactions are both happening but the concentrations of the reactants and products they have no overall change because the product is being made at the same rate that the reactant is being used and because the reaction is reversible that product is being used up at the same rate that it is made as well and reactants are forming again so overall the reactants and products whilst they're both being made and used there is no overall change now the position of the equilibrium is very rarely 50/50 there's no easy way of looking at an equation and working out what the position of equilibrium is you can actually only tell whether there are more products than reactants by doing an experiment to find out how much of one of the reactants or products you've actually got more on that later now a really good analogy for a reversible reaction in dynamic equilibrium is an escalator now escalators the moving stairs that you don't have to walk whilst they are taking you in one direction they will move and you will move with them and so this is moving downwards now obviously you shouldn't probably for health and safety reasons but probably everybody does you can go against the escalator so you can go upwards on a downwards escalator and if you run fast enough you will get to the top and even though the escalator was trying to take you down words the upwards one because the upwards was at a faster rate however you will if you're clever be able to match your upwards walking to the downwards movement of the escalator and here you have reached equilibrium you've reached balance because your position up or down won't be changing you will be effectively still whilst also moving and that's why it's called a dynamic equilibrium because you are moving your legs your muscles are working and the escalator is working so the upwards and the downwards are both happening but your height is not changing because you have reached equilibrium and you could do that at the bottom where we've got here so three steps up or you could do this three steps down you could do this in any position or pour down on the escalator equilibrium that position of not actually moving your overall height your height not changing that is applicable at any of these steps along the escalator just like a dynamic equilibrium might be 90 percent product and 10 percent reactant or any other combination so that's a really good analogy you can reach equilibrium whilst walking up a down escalator I just mentioned that the position of equilibrium could be 40 60 could be any proportions at all and so I wanted to begin this section by looking at what it is that affects that proportion and so the position of equilibrium and that means the proportions or the ratios of reactants or products is at the point where the chemicals are most stable and the fact that they're the most stable at this point means that they have the lowest energy and at that equilibrium there will be a certain set of conditions and these conditions will be to do with an optimum temperature pressure and any other factors that might happen to be affecting the chemical reaction those conditions are such that they are in the most stable temperature for instance as well and so less chatelier's principle states that if a reaction at equilibrium is subjected to a change in concentration pressure or temperature the position of the equilibrium will move to counteract the change which is really fancy sounding but you know he was a professional scientist and he wanted to sound clever however we can reduce that definition down to something that's more usable which is that equilibria shift to oppose change and those few words are really powerful particularly for your understanding evolution at Le A's principle and even in exam situations it can help your approach to any number of types of question about equilibria and we should be absolutely clear here the reason that the equilibrium is shifting to oppose the change it's not through any conscious thought the equilibrium is not something that we can personify it is doing it to maintain stability so if we impose a change upon it that makes it less stable than before and the equilibrium shifts spontaneously to become as stable as possible before we explore the chatelier's a little bit more deeply I just want to clarify two terms about chemical reactions now chemical reactions can be described as homogeneous or heterogeneous now actually heterogeneous are a little bit more common than homogeneous but in terms of equilibria you're only concerned with homogeneous equilibria situations now a homogeneous situation is when all the reactants and products are in the same state and normally they're all going to be gases and so that's derived from homo meaning the same and genius is what links it to state and it would follow that heterogeneous meaning hetero being difference and genius still being States so all the reactants and products could be in a variety of different states so they aren't all in the same state so Litella a tells us that anything that we do to an equilibrium will be opposed by the equilibrium itself shifting and so what that means is if we increase the temperature equal the broom will shift in the direction that lowers the temperature and that is the endothermic reaction if we decrease the temperature equilibrium will shift to oppose that to raise the temperature up again and that will be the exothermic reaction from the point of view of pressure if we decrease the pressure equilibrium will shift to raise the pressure back up again and the higher pressure is the side of the reaction that has got the greatest number of molecules of gas and if we raise the pressure equilibrium will shift to lower the pressure back down again and that will mean that the reaction will shift in the direction where there are fewer molecules of gas and finally if we increase the concentration of a particular chemical in the equilibrium the equilibrium will shift in the other direction to decrease the concentration of that reactant or product and it follows that the exact same thing is true if you were to decrease the concentration of a particular chemical the equilibrium would shift to increase its concentration to make more of it and that means it would shift to the side that has got that particular chemical there by making more of it to restore that most stable of conditions which is what all of these shifts are to do now is probably a good time to point out that the square brackets around a particular formula such as square brackets around X means concentration of chemical X and that is just a shorthand notation for that which definitely makes things a bit faster and is totally fine to do in an exam situation now I haven't mentioned catalysts until this point and the reason that catalysts are less significant is that they have no effect on the position of equilibrium the reason that this is true is obvious if you consider a reaction profile so if we have a look at an exothermic reaction profile in the forward direction you can see why this reaction maybe is reversible because the energy barrier the activation energy in both directions is not particularly big which is what allows both the forwards and the backwards reactions to happen however what catalysts do is they lower the activation energy and if you consider the reaction profile the activation energy gets decreased by the same amount for both the forwards and the backwards reaction because the hill just kind of gets lower so to speak and so the forward reaction and backwards reaction will be sped up by equal amounts and so all that would happen is equilibrium would be reached faster a question that you are probably asking is how does an equilibrium actually shift and the answer is that it shifts by speeding up in one direction so if we consider first of all a system that is at equilibrium the rate of the forward and backwards reactions are the same let's say that they are 1 mole per Desmet meter cubed per second if an equilibrium shift happens say in the forwards direction so it shifts to the right-hand side then the forwards reaction will increase maybe its rates of reaction will increase to 3 moles per decimeter cubed per second and at that point of increase the forward reaction is faster than the backwards reaction so that means that we're making more products than we have been making reactants and so the position of equilibrium which previously was 3 to 7 in terms of the proportions of the reactants to the products has now increased to 2 to 8 but because that we've now got more product and we've got less reactants the rate of the forward reaction decreases down to say 2 moles per decimeter cubed per second and the rate of the backwards reaction increases from 1 mol per decimeter cubed per second up to 2 moles per decimeter cubed per second and so now we have a situation where the forwards and the backwards reactions are equal again do from Reiter before but they're equal again and so now from this point onwards the concentration of the reactants and products will remain constant even though it is still dynamic the reaction is still happening let's take a look at a couple of examples of equilibria and how they would be affected by the changes that we've outlined here so in the first example we've got oxygen and sulfur dioxide producing sulfur trioxide and the Delta H is minus 200 which we need to take a moment to process as meaning it is exothermic and that's really important for this first question if we increase the temperature equilibrium shifts to decrease the temperature which is in the endothermic direction so to the left-hand side here so equilibrium shifts to the left-hand side for Part B if we decrease the pressure equilibrium shifts to increase the pressure and so it shifts suicide with the greater number of molecules of gas and there are three on the left-hand side and only two on the right-hand side so the side on the left hand side is the higher pressure side and so shifting to the left which is what happens has the effect of increasing the pressure and finally adding a catalyst well that's a trick adding a catalyst has no effect on the position of equilibrium at all on to the second example if we decrease the concentration of hydrogen iodide gas then the equilibrium will shift to replace it to make more of that hydrogen iodide gas to try and restore the balance of concentrations to reactants to products and so that means it will shift to the right-hand side and then the final example of increasing the pressure will equilibrium shifts if it can to decrease the pressure now we can't actually do that here because there are two molecules of gas on the left-hand side and two molecules of gas on the right-hand side so the pressure of both sides is equal and so increasing the pressure or decreasing it in fact would have no effect on the position of equal Brinn although it should be pointed out it would affect the rate of reaction we're going to finish this video off by looking at some industrial processes and looking at how we can maximize the yield of a particular target chemical whilst not really putting the rate of reaction under any negative impact unfortunately this isn't always possible and it's really important that you can appreciate that sometimes we have to have a compromise between the amount of product that we make which is the yield remember and how quickly we get that product which is the rate of reaction and something to note from the get-go is 100% yield if it takes a month is not as good as a 10% yield if it takes one day because over a 30 day period in method 1 you get 100% yield which sounds good but in method 2 you get a 10% yield every day for 30 days so that's 300 percent yield for purposes of comparison so method 2 even though it sounds less good on paper it gets us 3 times the amount of product over the same period of time the first industrial process we're going to look at is the synthesis or production of methanol from hydrogen gas and carbon monoxide and so this reaction as you can see from the equation here involves 3 molecules of gas turning into 1 molecule of gas and the forward reaction is exothermic and so we can use our knowledge of leshiy Atelier to work out what the optimum conditions would be to maximize the yield of methanol so we need to think first of all about what we want the equilibrium to do whatever we do to the chemicals however we manipulate the conditions we want to make the equilibrium shift to the right-hand side so the equilibrium needs to be shifting in the forwards direction now the forward direction involves an increase in temperature because it's exothermic when we look at that Delta H value and so it would follow that if we want the equilibrium to increase the temperature before it does it we need to first lower the temperature and have a low-temperature reaction and that would favor the equilibrium shifting to the right also the forward reaction involves a decrease in pressure because the right-hand side is lower pressure than the left-hand side and so to make the equilibrium shift forwards to decrease the pressure we first need to have a higher pressure to maximize the yield of methanol now in practice we use a high pressure about 50 to 100 atmospheres so that means about 50 to 100 times higher pressure than atmospheric pressure and that gives us both a good yield and a fast rate of reaction we use a temperature of 250 degrees C which is not low but from an industrial point of view it's not a high temperature and that's because we want to try and maximize yield remember we'd ideally want a very low temperature but we don't want to compromise on rate of reaction because remember a higher temperature is a faster rate of reaction and so to make up for the fact that we are compromising on the rate of reaction we include a catalyst which is a mixture of copper zinc oxide and aluminium and that helps us to get over our compromise that we're having to do so we're trying to keep the temperature as low as possible without reducing the rate of reaction too much another industrial process you need to know about to see synthesis of ethanol which we explore also in the alcohols video so check that out sometime in the industrial synthesis of ethanol which does not come from renewable materials we use eath in gas and steam and that gets converted into ethanol in an equilibrium reaction that is exothermic now in the same way as for methanol let's look at what conditions would favour this reaction so because we've got two moles of gas on the left-hand side and one on the right-hand side we want equal to shifting the forwards direction which decreases the pressure so ideally we would want to use high pressure and in practice we use quite a high pressure 60 to 70 atmospheres so a little bit lower than for the methanol synthesis and the reason for that is that at high pressures the ething actually tends to polymerize to polythene and also it's really quite expensive to build a power plant that can generate a high pressure and also it costs quite a lot of money to run the high pressure apparatus so 60 to 70 is a slightly different kind of compromise and it's not simply about the rate although because we're compromising on pressure and therefore rate we do use a catalyst which is phosphoric acid most commonly now in terms of temperature once again the forward reaction is exothermic so we want the equilibrium to go to the right hand side in the direction that raises the temperature so in theory we should use a low temperature because that would maximize our yield but because we're compromising already on pressure lowering the rate of reaction that way we actually use a temperature of 300 degrees C so similar to the methanol temperature but a little bit higher once again because we're compromising on the speed of the reaction we have to therefore do something which wouldn't give us the biggest possible yield by having a higher temperature but it's all a compromise where a variety of factors lean into each other safety money rate and yield all leading to these conditions which do give us around 95% of ething eventually converted into ethanol one additional aspect of the practice is that even though the yield is not 100% we get closer to it 95% by recycling the unreacted ethene so that can go into the reactor again and again and that's how we another way that we can push it up to 95% yield one final process that we're going to look at is the manufacture of ammonia now I wouldn't normally ever given a levels to you word equations but the reason I'm doing that is because I wants us to discover some things about the synthesis of ammonia from a graph and so in exams they often give you graphs for you to either interpret in short term sense or give detailed analysis so let's look first of all at general patterns on here so we've got on the graph as we increase the pressure going from left to right the percentage yield of ammonia increases so as we increase the pressure the percentage yield goes up all of these are different forms of sort of positive correlations so because yield increases when pressure increases we know that the forward reaction involves a decrease in the number of molecules of gas because that is the only explanation for a higher pressure making the equilibrium shift to the right-hand side because the forward reaction must lead to a lower pressure the other clue that we've got is we've got four separate lines on our graph at different temperatures and we can see that the highest temperature gives us the lowest yield at the same pressure so if we consider one particular pressure let's say this pressure here the yield for the 770 kelvin is significantly less than for the other temperatures for the same pressure and so because the greatest yield is found for the lowest temperature that must mean that the forward reaction is exothermic because a high yield is favored by a lower temperature so the lower the temperature the more the equilibrium is able to shift at the right-hand sides and maximize that yield so we can look at this graph to deduce at the forward reaction involves a decrease in moles and involves an exothermic process we can also use this graph in a bit greater depth to make some suggestions about the temperature what and the pressure that we should use or to justify the pressure so for instance if we look at the top line after maybe this pressure here there is no point increasing the pressure any further because that will be really expensive there's very little difference in rate of reaction in this entire tracking on here it's not quite flat but it almost is same principle with the 570 line along here you get the greater difference in yield at the higher temperatures and also we could say that perhaps we wouldn't go so higher temperature because the difference in yield at say 670 Kelvin and this particular pressure isn't that much difference at to the 517 certainly 570 and 4/7 see are quite similar so we would probably go for this temperature here 670 Kelvin which by the way is 400 approximately degrees C so again not crazy high and the pressure that we would probably use is about this pressure on here and in practice that is 20,000 Pascal's of pressure and about six hundred and seventy degrees C because that gives us a good yield around about 50% but it gets it as quickly because 670 kelvin is significantly faster than 570 kelvin and that in turn is faster than 470 kelvin so 670 kelvin is the temperature that we use and 20,000 Pascal's okay that's where we're going to end this video don't forget to check out part 2 of the equilibria topic coming up soon bye bye