Lesson video

Hello, my name is Mrs. Collins and I'll be taking you through the learning today.

This lesson forms part of the unit "Industrial chemistry" and is called "Equilibrium: balance between rate, safety, environment and economics." During today's lesson, you will describe how various factors influence the reaction conditions for industrial equilibria, and understand the trade-offs between production rate, equilibrium position, safety, and cost.

Here are the keywords for today's lesson.

Equilibrium position, dynamic equilibrium, reaction rate, Haber process and Contact process.

Here are those words used in context.

What I suggest you do is pause the video here, read through these explanations and write down any notes you feel you need to.

Today's lesson is broken down into two parts, factors affecting industrial equilibria and balancing safety, cost, and efficiency.

So let's start with part one.

Factors affecting industrial equilibria.

Industrial equilibria play a crucial role in modern manufacturing processes.

Some really key examples are the Haber process and the Contact process, which you will have covered before in your learning.

And both of these are important in the production of fertilisers.

The following factors impact reaction rates and dynamic equilibria.

So we've got temperature, pressure, concentration, surface area, and catalysts.

And all of these factors are taken into account when selecting reaction conditions to balance between yield and cost.

So the equilibrium position refers to the relative concentrations of reactants and products in a reversible reaction at equilibrium.

And it's a useful way of indicating whether reactants or products are favoured under given conditions.

Equilibrium position can be affected by changes in temperature, pressure and concentration.

And you will have learned about that already.

Here's a question based on that learning.

Which factors impact both the rate of reaction and the equilibrium position? So you need to think carefully about this.

Pause the video now, answer the question and I'll see you when you're finished.

Welcome back.

The answer's temperature and pressure.

Remember, surface airing catalyst will affect the rate of reaction, but they won't affect the equilibrium position.

So here are the conditions for the Haber process.

450 degrees Celsius, which is a high temperature.

200 atmospheres, which is a high pressure.

An iron catalyst.

And ammonia is condensed and removed from the process and that impacts the concentration.

Here are the conditions for the Contact process.

Again, 450 degrees Celsius.

This time, 1 to 2 atmospheres pressure.

A vanadium oxide catalyst.

And the removal of the sulphur trioxide, again impacting on concentration.

Industrial equilibria involved a trade-off between production rate and equilibrium position.

So this is a classic graph that you may see on examination questions and it shows you the impact of various factors on the yield.

So in this case, the yield of ammonia.

So a high temperature and pressure may not be the most desirable for rate, but it might be for yield and vice versa.

So in this case, we know if we increase the temperature, the rate of reaction will increase.

But if we look at this graph, if you can see when we increase temperature, the actual yield of ammonia is decreasing.

So that is not a desirable at all.

So we need to have some kind of trade-off between the two.

So here's a question based on that learning.

High pressures are only used to increase reaction rates without any other factors considered.

Is that true or false? Justify your answer using the statements below.

So pause the video here and I'll see you when you're finished.

Welcome back.

So the answer to the question is false.

Temperature and catalysts are also considered.

So well done if you got that correct.

We're now gonna move on to task A.

Some pupils are discussing equilibria.

Identify which pupils are correct, and update any incorrect statements.

So pause the video here, answer the question, and I'll see you when you're finished.

Welcome back.

So let's consider each of those statements one at a time, starting with Andeep.

So Andeep says, "Increasing the temperature always increases the reaction rate and favours the formation of products." Well, that's incorrect.

Because increasing the temperature increases the reaction rate, but shifts the equilibrium position depending on whether the reaction is exothermic or endothermic.

So we need to take that into consideration when we are thinking about the equilibrium position.

Lucas is correct.

"Using a catalyst speeds up the reaction rate, but does not change the equilibrium position." Sofia is correct.

"Higher pressure increases the rate of reaction and favours the side with the fewer gas molecules." And then Laura says, "Increasing the surface area of a solid reactant shifts the equilibrium position in favour of the products," which is incorrect.

Increasing the surface area of a solid reactant increases the reaction rate, but does not change the equilibrium position.

So well done if you've got that correct.

We are now moving on to part two of the lesson.

Balancing safety, cost, and efficiency.

So the Haber process involves a reversible reaction as we know, and that reversible reaction will be at dynamic equilibrium.

So we've got nitrogen reacting with hydrogen to form ammonia, one mole of nitrogen reacting with three moles of hydrogen to form two moles of ammonia.

And we know it's a reversible reaction.

We've got the double-headed arrow there.

The forward reaction is exothermic, and the reverse reaction is endothermic.

This means if we increase the temperature, what will happen to the equilibrium position? So let's have a think about that before we pop the answer up.

So if the forward reaction is exothermic and the reverse reaction is endothermic, what will happen if we increase the temperature? So increasing the temperature will speed up the reaction rate because there'll be more frequent, successful collisions.

And that's the rate in both directions, remember.

But it will decrease the yield of the product because it will shift the equilibrium position towards the reactants.

So it favours the reverse reaction.

Increasing temperature will increase reaction rate as we said.

Both the Haber and Contact processes produce a lower yield using higher temperatures, which is a bit difficult.

So we got this balance between a faster rate of reaction and a lower yield by increasing the temperature.

However, using low temperatures results in a slower rate, requiring a longer time to reach equilibrium and that reduces efficiency.

So you can see a balance is needed.

So here's a question based on that learning.

True or false? Increasing temperature with an exothermic forward reaction always results in a higher yield of product.

Is that true or false? And justify your answer using those statements.

So pause the video here, read that question carefully and I'll see you when you're finished.

Welcome back.

So the answer is false.

And that is because higher temperatures can decrease the yield for exothermic forward reactions because the equilibrium shifts to favour the reactants.

So to favour the reverse reaction.

Well done if you got that correct.

So the Contact process involves a reversible reaction at dynamic equilibrium as well.

And we can see that here we've got two moles of sulphur dioxide reacting with one mole of oxygen to form two moles of sulphur trioxide.

The forward reaction produces fewer gas molecules than the reactants.

So increasing the pressure will, and again, we are going to think about this before we see the answers.

So have a look at that.

If we increase the pressure, it'll favour the side with the fewer gas molecules.

So if you look at it, which direction is that going to push the reaction? So it will speed up the reaction rate again because there'll be more frequent successful collisions and it will increase the yield of the product because it will shift the equilibrium position to the right.

It will favour the forward reaction.

There are fewer moles of gas molecules in the products than the reactants.

So we've got three in the reactants and two in the products.

Industry conditions for equilibria have been fine-tuned to balance safety, cost and efficiency.

Both the Haber and Contact processes produce a greater yield using higher pressures.

So you'd think, we'd want as higher pressure as we can possibly get, 'cause we want to increase the yield.

But having high pressure costs money, it also increases the risk of equipment failure and it presents safety concerns 'cause there's a potential for an explosion because you've got these really incredibly high pressures.

So this makes it very hazardous.

So here's a question based on that learning.

Why is it necessary to balance high pressures and temperatures in industrial processes? Pause the video here, answer the question and I'll see you when you're finished.

Welcome back.

So the answers to that question is, to minimise safety risks and reduce costs.

And to shift the equilibrium position favourably.

So well done if got my answer correct.

Both the Haber and the Contact processes use catalysts.

And using a catalyst, remember, speeds up the rate of reaction, but doesn't affect the yield of the product because it doesn't shift the equilibrium position.

So a catalyst speeds up the rate of reaction without being used up in the reaction.

It provides an alternative pathway with a lower activation energy.

The catalyst used will often be a fine powder that gives it a high surface area and that again, increases the rate of the reaction.

Both the Haber and Contact processes remove the final product as the reaction is happening.

And removing this product effectively lowers the concentration of the product that's being produced and therefore, increases the yield of the product because it shifts the equilibrium position towards the product.

Here's a question based on that learning.

Which the following statements best describes the trade-off in industrial equilibria? Pause the video here, answer the question and I'll see you when you're finished.

Welcome back.

So the answer to the question is a high pressure can increase the rate of reaction, but may shift the equilibrium unfavourably.

In industrial reactions, the conditions used are also influenced by the availability of the raw materials.

So it depends where those raw materials come from and how available they are, how much it costs to find and extract those raw materials, the energy supplies and the cost related to those.

For example, the Haber process uses nitrogen from the air, which is a low cost, and hydrogen from sources like natural gas.

And that has a variable cost.

It does depend.

If you extract it using electrolysis, it can be more expensive.

And energy-intensive due to high pressure and temperature, necessitating efficient energy use.

So it's quite energy intensive.

So here's a question based on that learning.

High energy costs can lead to changes in the pressures and temperatures used in industrial reactions.

Is this true or false? And use the statements to justify your answer.

So answer the question and I'll see you when you're finished.

Welcome back.

The answer to that question is true.

To reduce operational costs, industries might use more energy-efficient conditions, even if it means a slightly lower yield.

We're now moving on to task B.

Jun predicts that the reaction between carbon monoxide and hydrogen is exothermic and involves a reduction in gas volume.

So here we've got carbon monoxide and hydrogen forming methanol, and it's a reversible reaction.

We've got one mole of carbon monoxide reacting with two moles of hydrogen to form one mole of methanol.

We've then got a whole series of information in that table there.

So we need to look really carefully.

It is telling us the percentage yield of methanol.

So in that table, in the white section of the table, you can see lots of percentages.

So that's the percentage yield of methanol.

We can see what happens as we increase the pressure.

So it's going from 50 atmospheres, 100 atmospheres, 200, 300, 400 atmospheres, and we can see the impact that's having.

And then down the side, we've got temperature in degrees Celsius, 200 degrees Celsius, 300, 400, 500 degrees Celsius.

So number one, describe and explain whether Jun's prediction are supported by the reaction and the results in the table.

Suggest why the reaction conditions chosen by industry are 200 atmospheres and 300 degrees Celsius, despite not having the highest yield.

So you need to think about this quite carefully.

You need to use the data in the table to support your answer.

So pause the video here, answer the question, and I'll see you when you're finished.

Welcome back.

So let's go through that answer.

So describe and explain whether Jun's predictions are supported by the reaction and the results in the table.

So in endothermic reactions, as the temperature increases, the equilibrium position shifts to the right.

As pressure increases, the equilibrium position shifts to the side with the least number of moles of gas.

So that's the facts.

So the prediction is not supported, the reaction must be exothermic, not endothermic, because equilibrium position moves to the left as the temperature increases.

And that's shown in the table.

The prediction is supported, as pressure increases.

The position of equilibrium moves to the right, fewer moles of gas, and the percentage yield increases.

So well done if you got that correct.

Question two, suggest why the reaction conditions chosen by industry are 200 atmospheres and 300 degrees Celsius, despite not giving the highest yield.

So it balances yield, safety, and costs.

So high pressures, such as 400 atmospheres are costly to maintain and pose safety risks.

Operating at 200 atmospheres reduces these risks and costs while still achieving a relatively high yield.

And you could have quoted the yield values there from the table if you wanted to.

Lower temperatures increases the yield, but slows the rate of reaction.

And a temperature of 300 is a compromise that allows for a reasonable reaction rate and energy efficiency while still producing a decent amount of yield.

So it's a balance between the two.

So well done if you got that correct.

Here is the summary of the lesson.

Industrial equilibria often involve a trade-off between the rate of production and the position of equilibrium.

Temperature, pressure, surface area, concentration, and catalysts all impact reaction rates and equilibrium positions, sometimes in opposing ways.

Maintaining very high pressures is costly and increases the risk of equipment failure.

High pressures and temperatures can improve reaction rates, but must be balanced against safety and economic considerations.

'Cause remember, they can both be expensive to produce.

And low temperatures tend to slow down reactions, impacting production rates, but can shift equilibrium to favour the desired product in some exothermic reactions.

So thank you very much for joining me for today's lesson.