﻿Hello, everyone. I hope you're all doing well. And welcome back to chemistry. I am Dr. Deng, and I'm going to take you through this final topic in this unit, The Rates and Extent of Chemical Change. In this lesson, we are going to summarise what we've learned about Le Chatelier's principle and look at how it is used in different industrial applications. This topic is very interesting, as we will get to learn about how different industries select optimum conditions when it comes to manufacturing chemical products. So before we begin, please find yourself a quiet space where you can sit and work through this lesson free from any distractions, turn off the notifications on your mobile phone, or switch it off completely if you can, close down all browser tabs that are not being used, so you can make the most of this lesson. When you're ready, let's get started. We are going to start this lesson by recapping Le Chatelier's principle and revisiting the effect of changing conditions on the position of equilibrium in chemical reactions. Then we are going to look at how Le Chatelier's principle is used in different industrial applications. We will go through a couple of different examples of industrial processes and learn why they operate at certain conditions. Before the end of the lesson, you will complete a summary quiz and answer a couple of practise questions independently. In terms of equipment that you need for this lesson, you will only need a writing pen and some paper or something to write on. If you do not have any of these with you right now, please take a moment, pause lesson, and go and collect your equipment, then resume once you're ready. So let's start by going over Le Chatelier's principle and why it's so important in chemical processes. Le Chatelier's principle is a very important principle in chemistry, especially when it comes to chemical processing. It was devised by a French chemist called Henry Louis Le Chatelier, and it is used to predict the effect of changing conditions on a system in equilibrium. These conditions include concentration, temperature and pressure. So before we answer a couple of practise questions, let's go over the effect of changing these conditions on the position of equilibrium using Le Chatelier's principle. So first let's go over the effect of changing concentration on the position of equilibrium for a reaction that involves solutions. So increasing reactants or decreasing products favours the forward reaction and likewise, decreasing reactants or increasing products favours the reverse reaction. Don't forget, in a chemical equation, reactants are displayed on the left hand side of the arrow and products are on the right hand side. So when the concentration of one of these substances is increased it's going to cause the equilibrium to shift away in the opposite direction to counteract this change. Next, let's take a look at how changing temperature affects the position of equilibrium in a reversible reaction. So heating or increasing temperature favours endothermic reaction, whilst cooling or lowering temperature favours exothermic reaction. So it's very important to recognise whether a reaction is endothermic or exothermic and in which direction. And finally, let's take a look at how changing pressure affects the position of equilibrium. in a reversible reaction. Increasing pressure favours the side with fewer moles of gas and likewise, decreasing pressure favours the side with more moles of gas. So don't forget to count the number of gas molecules on each side of the reaction before you determine where the equilibrium is going to shift. Before we move on, and if you haven't done so already in previous lessons, please pause the lesson now and note down these different key points that we've just gone through so you can refer back to this in case you forget. Do this now, then resume the lesson once you're finished. Great. Now let's practise a couple of questions using Le Chatelier's principle. The following reaction shows the decomposition of hydrogen chloride. The forward reaction is endothermic. There are four questions we need to answer. Please read each one of these questions carefully and answer them on your piece of paper. Pause lesson now to attempt this task, then resume once you're finished. Great work. Now let's go through the answers together. So based on this question, the reaction is endothermic in the forward direction. And looking at the symbol equation, I can tell that all the substances involved are gases. So question one says, "What effect does lowering the temperature have on the position of equilibrium?" So do we remember which direction will equilibrium favour when temperature is lowered or when we cool the reaction? That's right, lowering the temperature, or cooling it, will favour the exothermic reaction. So in this case, if the full reaction is endothermic, the reverse reaction must be exothermic. And therefore equilibrium shifts to the left. And don't forget to explain in your answer why this is. Let's take a look at question two. What effect does increasing temperature have on the chlorine gas yield? So once again, let's think about how equilibrium will shift when temperature is increased. So increasing temperature favours endothermic, that is the forward reaction. So therefore, that will cause the equilibrium to shift to the right. And if the equilibrium shifts to the right, it will increase the chlorine gas yield. Therefore a high temperature is going to be favoured if chlorine gas is the desired product. Question three, "What effect does increasing pressure have on equilibrium? " Earlier on, we said that increasing pressure will cause the equilibrium to shift to the side with fewer moles of gas. So first we must count the number of molecules of gas on each side of this reaction. So on the left side, I can see two, because it says 2HCL. And on the right side, that is one H2 and one CL2. So two as well. So if there are two moles of gases on each side, what does that mean in terms of the equilibrium shift? It has no effect. Next question. Explain your answer from question three. So why does increasing pressure have no effect on the equilibrium? And we said this earlier on, because there are equal numbers of gas molecules on each side of the reaction, therefore the equilibrium does not shift. Well done if you answered all of these questions correctly. So now let's take a look at how Le Chatelier's principle is applied in different industrial processes. This part of the lesson is going to be very interesting. So I will need you to put your thinking cap on so we can work out some of these questions together. When it comes to chemical processing in the industry, selecting the right conditions for chemical processes to operate at involves a lot of thinking. These conditions include the ones that we talked about. Concentration, temperature and pressure. It's so important to select the optimum conditions for individual chemical reactions for the following reasons. Firstly, altering reaction conditions can result in the yield of products increasing and the process being more profitable. Secondly, in industry scientists have to balance the yield produced with the rate of reaction. Sometimes increasing the rate of reaction doesn't necessarily mean increasing the yield of the product or the amount of product made. So there needs to be a balance between yield and how quickly the reaction is occurring. And thirdly, other considerations are cost. These include costs to maintain a high temperature, because more energy's needed, and costs to maintain a high pressure, as specialist reaction chambers are needed. So just because increasing temperature or increasing pressure increases yield, industry do not always choose to use these conditions, because they will need to think about the overall cost when it comes to manufacturing products. So now let's take a look at a couple of examples of industrial processes, where we can apply our knowledge of Le Chatelier's principle and determine why certain conditions have been selected by industry for these different processes. I'm going to talk you through how to answer these first couple of questions, and I need you to pay full attention so you can apply this knowledge when answering similar questions later on. The first example we are going to look at is for the production of hydrogen gas. In this process, methane reacts with steam to produce carbon monoxide and hydrogen gas, and hydrogen gas is the desired product. The forward reaction is endothermic. So increasing temperature will favour the forward reaction and increase yield, but using high temperature comes with its disadvantage. So what is the disadvantage of using a high temperature? So, as I said earlier on, this can be very expensive as a lot of energy is needed to heat up this chemical reaction. So what if we use a lower temperature? Lowering the temperature will favour the exothermic reaction. So in this process, the forward reaction is endothermic and the reverse reaction is exothermic. So lowering the temperature is going to shift the equilibrium to the left. What does that mean in terms of the yield? It is going to decrease or lower the yield. So as you can imagine, it's so important to select an optimum temperature for this reaction, so there's a balance between cost and yield. Let's take a look at another example. Let's try and answer this one together. The similar equation in this reaction shows that ethene, C2H4, reacts with steam, H2O, to produce ethanol, C2H5OH. First question is, is the forward reaction exothermic or endothermic? So if we look at the overall energy change, which is given as Delta H, or the change in H, which is the overall change in energy, it is negative 45 kilojoules per mole. So that tells us that this reaction releases energy to the surroundings. So that must mean that this reaction is exothermic. Let's take a look at the next question. So if the forward reaction is exothermic, give a disadvantage of using high temperature. So remember, increasing temperature favours endothermic, and in this reaction, if the foreign reaction is exothermic, that means the reverse reaction must be endothermic. So increasing temperature is going to shift the equilibrium to the left. So what does that mean in terms of the yield? It's going to lower the yield. And the third question. What is the disadvantage of using low temperature then, if this reaction favours low temperature? What do you think? Did you get it? Using low temperatures can cause the reaction to take place too slowly. Once again, selecting an optimum temperature for this reaction in the production of ethanol is extremely important because there needs to be a balance between yield and the rates of reaction. Here's another example of an industrial process. The production of ammonia, also known as the Harber process. In this reaction, nitrogen gas reacts with hydrogen gas to form ammonia. And I'm not going to give any more away because I would like you to answer the following questions independently. So let's look at question one. Is the forward reaction exothermic or endothermic? And in three, two, one. It is exothermic. And hopefully you are able to identify this based on the overall energy change, which is -92.4 kilojoules per mole. And this negative value tells us that this reaction releases energy to the surroundings, which is why the forward reaction is exothermic. Here's your next question. What is the disadvantage of using low pressure in this particular reaction? In three, two, one. It lowers the yield. So remember lowering the pressure favours the side with more molecules of gas. So in this reaction, there are more molecules of gas on the reactant side, and therefore by lowering the pressure it's going to shift the equilibrium to the left and this would decrease the ammonia yield. Next question. What is a disadvantage of using high pressure then? And in three, two, one. It's expensive. As we mentioned earlier on, equipment that is needed to safely contain the reaction can become quite costly. Let's take a look at one more example. So here we have the production of sulfuric acid. In this question sulphur trioxide is going to be produced, which is the precursor used in the production of sulfuric acid. SO2, which is sulphur dioxide reacts with O2, oxygen gas to form sulphur trioxide, SO3. So let's answer the following questions. You are going to answer these independently. Question one, what's the advantage of using a high pressure? And in three, two, one. Increases the yield. So again, when it comes to thinking about pressure, we need to count the number of gas molecules on each side of the reaction. So in this case, the three on the left side and two on the right side, therefore increasing the pressure will shift the equilibrium to the right, increasing yield. Next question. What is a disadvantage of using high pressure then? And a three, two, one, and hopefully you got this. Expensive. Equipment that is needed to safely contain the reaction can become quite expensive. So well done if you have answered all these questions correctly, you're now thinking like a true chemist. So moving on, let's take a look at the advantages of adding a catalyst to chemical reactions in industry. Adding a catalyst or a reversible reaction speeds up the rates of both forward and reverse reactions. Reaction reaches equilibrium more quickly, but adding a catalyst does not affect the position of equilibrium. It's so important that you remember this. And here we have a couple of examples of different catalysts being used in different industrial processes. So in the Harber process, an iron catalyst is used. In the production of sulfuric acid, vanadium pentoxide catalyst is used, and in the production of ethanol, solid silicon dioxide coated with phosphoric acid is used. All right, so we've come to the part of the lesson where you will need to complete a summary quiz. This is going to test your understanding of what we've learned so far. These are going to be quickfire, multiple-choice questions. I will read the questions out and you will get a good couple of seconds to choose an answer. So are you ready? Let's go. The forward reaction of the Harber process is exothermic. What is the effect on yield when temperature is lowered? And in three, two, one. The answer is, A, increases. Next question. The forward reaction of the Harber process is exothermic. What's the disadvantage of using a very low temperature? And in three, two, one. The answer is, D. The reaction is too slow when a low temperature is used. Next question, in this reaction, methane reacts with steam to produce carbon monoxide and hydrogen gas. And the overall energy change is given as positive 65 kilojoules per mole. So the question is, what's the disadvantage of using a very high temperature? And in three, two, one. The answer is, C, it's very expensive. Even though this reaction is endothermic in the forward reaction, meaning increasing temperature will favour the forward reaction, increasing the yield, using a temperature that is very high can get very expensive because a lot of energy is needed. Next question, in this reaction, sulphur dioxide reacts with oxygen to produce sulphur trioxide. What is the disadvantage of using a low pressure? And in three, two, one. The answer is, D, this actually decreases the yield. So don't forget, you need to count the number of gas molecules on each side of the reaction. On the left side, there are three, on the right side, there are two. So lowering the pressure will shift the equilibrium to the left and therefore lowering the yield. Next question, in the same reaction where sulphur dioxide reacts with oxygen to produce sulphur trioxide, what's the disadvantage of using a very high pressure? And in three, two, one. The answer is C. It can get very expensive, even though this reaction favours a high pressure. A massive well done if you have answered all those questions correctly. Okay, so now to consolidate your learning, I would like you to answer a couple of exam style questions independently. So when the questions show up on your screen, please pause the lesson, read the questions carefully, then answer the questions on your piece of paper before resuming so we can check your answers. Here's exam style question one, please pause the lesson to complete this task, then resume once you're finished. Excellent. Let's take a look at exam style question two. Here's your next question. Pause lesson again to complete this then resume once you have finished. Well done. Now let's look at the answers so you can self-assess your work. Here are the answers for exam style question one, please pause the lesson again to check your work, then resume once you're finished. Excellent. Let's move on. Here are the answers for exam style question two, please pause the lesson again to assess your work, then resume once you've finished. Wonderful and well done if you answered most of these questions correctly. Let's summarise what we've learned in this lesson. So at the beginning, we went over Le Chatelier's principle and why it's so important in chemical processes. Next, we looked at how Le Chatelier's principle is applied in many industrial processes, but we also talked about the importance of selecting optimum conditions in order to find a balance between cost, yield and rates of reaction, then you completed a summary quiz, and finally, you had the opportunity to practise a couple of exam style questions. We have now arrived at the end of the lesson. There are downloadable resources available and please don't forget to complete the exit quiz so I know how well you've been listening in this lesson. Thank you so much for joining me as we've studied this unit, The Rates and Extent of Chemical Change. It's one of my favourite topics in chemistry, so I hope you also found it interesting and useful. I look forward to seeing you in our next lesson, where we will look at an overall review of this unit. If you enjoyed learning the topic in this unit, The Rates and Extent of Chemical Change, please don't forget to share your work with Oak National Academy. You've worked so hard and I look forward to seeing the work that you've produced. So if you'd like, please also ask a parent or carer to take a photo of your work and share it on Instagram, Facebook or Twitter, remember to tag @OakNational and hashtag #LearnWithOak. Thank you for joining me and I will see you soon. Bye.

In this lesson, students will identify and explain optimum conditions needed in order to maximise yield in chemical processing using examples in industries.

Chemistry

Y11 Chemistry

Key Stage 4

Year 11

Le Chatelier's principle: Uses in industry

The rate and extent of chemical change

UNIT

Rate of reaction

Rate of reaction using graphs and tangents

Collision theory

Planning an investigation to find rate of reaction

Rate of reaction required practical (Part 1)

Rate of reaction required practical (Part 2)

Effect of changing surface area on rate of reaction

Effect of changing temperature on rate of reaction

Effect of changing pressure on rate of reaction

Catalysts

Reversible reactions

Le Chatelier's principle: Effect of changing concentration and temperature

Le Chatelier's principle: Effect of changing pressure

The Rate and extent of chemical change: Review (Part 1)

The Rate and extent of chemical change: Review (Part 2)

Combined Science

Y11 Combined Science (Higher)

The rate and extent of chemical change (HT)

Lesson overview: Le Chatelier's principle: Uses in industry

Video

Presentation

Worksheet

Exit Quiz

Transcript

16 lessons in The rate and extent of chemical change:

16 lessons in The rate and extent of chemical change (HT):