Le Chatelier's Principle - Online Tutor, Practice Problems & Exam Prep

Le Chatelier's principle states that a system at equilibrium will shift to counteract disturbances, such as changes in concentration, pressure, or temperature. Increasing reactants or decreasing products shifts the equilibrium to the right, producing more products. Conversely, increasing pressure favors the side with fewer moles of gas. Temperature changes affect equilibrium based on the reaction's enthalpy; increasing temperature shifts away from heat, while decreasing temperature shifts towards it. Catalysts do not shift equilibrium but speed up reactions. Understanding these concepts is crucial for predicting reaction behavior in chemical systems.

Understanding Le Chatelier's Principle

Now Le Chatelier's principle says that if a chemical reaction is at equilibrium and it's disturbed, it adjusts itself by shifting in a certain direction. Now the direction of the shift depends on minimizing or canceling the disturbance and reestablishing equilibrium. Remember, a chemical reaction in these cases wants to be at equilibrium, so it's just moving to the left or to the right in order to re-obtain that equilibrium status.

Now we're going to say here that the following disturbances are carried out at constant temperature. Later on we'll talk about what if temperature is changing? What effect would that have? But for now, let's focus on the factors that have nothing to do with temperature. So here we're talking about the disturbances of chemical equilibrium. So we're looking at different factors. We're talking about changing the concentrations of your gaseous or aqueous compounds, adjusting pressure or volume, and then the introduction of inert gases, what we call mostly noble gases.

We'll talk about the type of changes. Are we increasing or decreasing these effects? We'll look at an example of a chemical reaction and then we'll explain what's going on. All right. So first let's talk about concentrations of gaseous and aqueous compounds. So here, let's say that I want to increase my reactants. Now realizing here that if you're increasing reactants, what's the opposite of that? Decreasing products? OK. So whether you're increasing reactants or decreasing products, the same effect will happen. So think of it as a balancing act.

Let's say that I have added additional reactants to my chemical reaction. It was at equilibrium, was fine. But then I decided to add more reactants. I got to remove that excess reactants that I've just added. I can't physically take my hand and dunk it into the solution because maybe it's dissolved in there. So what is the course of action that needs to be taken by the chemical reaction? Well, if I put reactants, I got to move in the forward direction to get rid of it. Moving in the forward direction will convert those excess reactant molecules into products.

In another way, let's say I decrease my products, OK, I've just depleted my products. How do I make more? Again, we would move in the forward direction. Moving in the forward direction towards my products would cause an increase in the amount of products. So here we're going to say the arrow moves this way, OK. So moving this way can cause a decrease in the amount of reactants that I just added or it can replenish the products that I've lost. So you're going to have more products will be made.

Next, pressure and volume. If you looked at my videos dealing with the gas laws, we know that pressure and volume have an inverse relationship. If pressure goes up, volume goes down. So let's say that I'm increasing pressure. If you're increasing pressure, that's equivalent to saying that I'm decreasing volume. Here we're going to say we have to look at the gaseous compounds on both sides of the arrow to understand which way it will shift. On the reactant side, here's a gas and I have two moles of it. This is aqueous, so it doesn't count. On the product side, I have one mole of gas and then another 4 moles of gas. That's five moles of gas total.

If I increase my pressure or if I decrease my volume, then my reaction shifts to the side with less moles of gas. So here it would shift this way, it shifts towards the reactants because the reactant side has less moles of gas. Next we have inert gas. Now when I say inert gas, we generally are talking about your noble gases. Of course there can be other types of inert gas that exist, but for this level of chemistry, we're restricting into the noble gas. All right, so adding an inert gas can happen under two conditions other under constant volume or under constant pressure.

And for adding an inert gas under constant volume. So here it will cause no shift. The inert gas we're adding is neon gas. We're going to say here partial pressures of gases are not changed. Now, if volume or pressure not mentioned, assume no shift. So let's say the question says you're adding some nitric neon gas to your chemical reaction. They don't mention that it's under constant volume or constant pressure. You just assume that it's being done under constant volume and no shift, no reaction will take place.

Now, if we're adding under constant pressure, what effect happens? Well, if you're adding under constant pressure, then something can happen again. Remember, this side has two moles of gas total and this side here has five. If you're adding an inert gas under constant pressure, realize that the addition of more gas into your reaction means that the volume has to expand in order to accommodate it. So your volume would increase and this causes a shift to the side with more moles of gas. In this example, the product side has more moles of gas, so it shifts to the right.

Now these are things that we need to remember. Also remember that the side that you're shifting to, wherever you're shifting to will be increasing in amount, and if you're increasing in that direction, the other side will be decreasing in amount. So here we're shifting towards reactants. So reactants will be increasing in amount and this ought to be decreasing in amount. If there's no shift, there's no increasing or decreasing on either side. Again, here we're shifting towards the product side. So the product side would be increasing in amount and this ought to be decreasing in amount.

OK, so keep in mind these are the different changes that you can do playing around with the concentrations of gaseous or aqueous compounds, playing around with your volume or pressure or the introduction of inert gas. Now let's say you're adding a catalyst. Adding a catalyst does not cause a shift in equilibrium position, it simply changes the reaction rates. Remember, catalyst lowers your energy of activation, thereby speeding up the reaction. It has no control of which way your reaction will shift, to the left or to the right.

OK, so here recall these different changes that we can do and remember if this is true under these conditions. If you reverse them, then the reverse would be true. Meaning that if I decrease my reactants amounts or I increase my product amounts, then we would have more reactants being formed. If I decreased my pressure or increased my volume, then I'd shift to the side with more moles of gas. OK here. I didn't put the other scenario, just memorize this version. Here you see it, you have the rule for it. If you're seeing the opposite on your test, then just remember the conditions have been flipped then what I see here, so the outcome or explanation will be flipped.

OK, so keep this in mind when looking at Le Chatelier's principle and the different types of changes that can take effect.

Le Chatelier's Principle Example

Here it says consider endothermic reaction at equilibrium. So here this chemical reaction is endothermic. It's as predictive, the reaction will shift to the right, left or no shift after each disturbance all right. So for the first one it says we're removing some O₂. If we look at our equation, O₂ is a product. Think about keeping things balanced. You want to maintain your equilibrium status. I've just removed some product from my chemical reaction. I need to make more. So which way would we need to shift in order to produce more product? We'd have to move in the forward direction or to the right.

Next, I'm adding some glucose right here I'm adding some product got to maintain that equilibrium status. I got a shift in order to reestablish it. How do I get rid of the product that I just added? I have to deplete it, so I have to move in the reverse direction. Moving in the reverse direction will cause a decrease in the product because it's going to be converted into some reactant.

Next, volume of container decreased. So remember if your volume decreases, that's equivalent to the pressure increasing. Remember, if volume is decreasing or pressure is increasing, we shift to the side with less moles of gas. So if we look, we have 12 moles of gas total on this reacted side and then we have 6 moles of gas over here on the product side. We're shifting to the side with less moles of gas. So we have to shift to the right.

Finally, xenon gas added to reaction mixture. Here we're not told that it's under constant volume or constant pressure. So we're assuming it's being done under constant volume. Under constant volume, there will be no shift because you have not disrupted your equilibrium status. Again, they'd have to explicitly say under constant pressure for you to take into account that the changes happen and what which way do I need to go?

All right, So just remember, keep reviewing those Le Chatelier Principle rules and you'll be able to answer questions like this that much easier.

Temperature Changes

Now let's take a look at the effect that temperature has on our equilibrium status, based on the Schotteler's principle. Here we're going to say changing the temperature of the reaction at equilibrium will disrupt the equilibrium and cause a shift. But how exactly does it work? Well, we're going to have to say here that enthalpy, which remembers ΔH of a reaction, plays a big role in the direction of the change.

Here we're going to say, recall that the equilibrium constant is temperature dependent. So changing the temperature not only can affect which way we shift, but it can also affect the scale or value of your equilibrium constant K. So here, let's take a look. We have enthalpy type. We have the type of change that can occur whether we're increasing or decreasing. We have examples and then we have explanations.

So in the first row, we have the explanation that we're shifting away from heat. Now we shift away from heat if we're increasing the temperature. Now this is true whether it's exothermic or endothermic. Increase in the temperature causes the shift away from the heat source. If we're talking about an exothermic process, ΔH is negatively charged. Exothermic means we're releasing heat. So heat has to be a product.

All right. So we increase the temperature, we're going to shift away from the heat, which means I'd have to move away from it. So I'm going to move towards the reactant side. We're going to go to the left. Remember the direction you shift in that side would be increasing an amount, so our reactants would increase in amount and conversely the product should be decreasing out. Because remember if one side is increasing, the other side has to be decreasing.

Now here for the 2nd row, we're shifting towards the heat. Now this is true whether it's endothermic or exothermic. If we decrease our overall temperature of the chemical reaction, the reaction have to shift towards the heat. So here, if you're an endothermic process, your enthalpy is pop. Your enthalpy is positive. If it's positive, you're absorbing heat, so he has to be a reactant. We are decreasing our temperature, so we're going to shift towards it. In this case, we're moving towards the left, again towards the heat. So again, wherever you're moving to, we'll be increasing an amount. If this side is increasing, the other side is decreasing.

So when it comes to this idea of temperature, Le Chatelas principle, first it's important to understand is your reaction endothermic or exothermic? This will tell you what side heat will be found. Once you've established a location of heat, then you look to see OK, is my heat, is my temperature increasing or decreasing? If temperature is increasing, I shift away from the heat. If temperature is decreasing, I shift towards the heat. So keep that in mind anytime you're dealing with a question that ties together temperature and Le Chatelier's principle.

Le Chatelier's Principle Example

Here it says consider the reaction. We have one mole of nitrogen gas reacting with one mole of oxygen gas and one mole of bromine gas to produce 2 moles of Nobr gas. Here we'll told the enthalpy value was -32.5 kilojoules. So here we're saying that the enthalpy is negative, so it's exothermic, which means he does released. So heat would have to be a product all right.

So now the phone changes will shift equilibrium to the left except all right. So removing N₂, if you remove some reactor, you have to replenish the reactor you lost, so you would have to move in the left direction to remake it. So this does move to the left, increasing the partial pressure of NOBr. So when we say increase the partial pressure of a gas, that just means adding more moles of that gas. Because remember, the quantity of a gas is directly connected to its partial pressure. The more of it in it that we have, the higher the partial pressure.

So in essence, this part C is saying that I'm increasing the amount of product to get rid of it. I'd have to move to the left again to get rid of it. So this is out. Increase the pressure in container. Now this is different. We're not increasing the partial purchase of any gases. We're increasing the overall pressure of the container. Increasing pressure is equivalent to decreasing volume in either instance. All that means that we're shifting to the side width. Less moles gas on this side here, see this side here, we have 3 moles of gas. On this side here, we have two moles of gas. So again, we're shifting to the side that has less moles of gas. South, we're shifting to the right.

So right now this is an example of a ship that's not to the left. Next, adding some NOBr is equivalent to option C, where we're increasing the partial pressure of NOBr. They're both saying the same thing. We shift to the right, to the left to get rid of it, decreasing the temperature. If we're decreasing the temperature, we have to shift towards the heat. The heat is on the product side, so we'd be shifting towards the right. So this is also an example of the equilibrium shifting to the right and to the left, sorry, to the right to the right.

Now lastly, F decrease the container volume, which is the same thing as saying E word, decreasing container volume, which is the same thing as increasing pressure of the container. So we're going to shift to the side width less moles of gas. So this would also shift to the right. So in this one, EDNF would all shift to the right, so they would be the answer. They're the ones that are not shifting to the left. So again, DE and F are our final answers.

Select correct answer(s) that would yield more products in the following reaction.

CH₄(g) + 4 F₂(g) ⇌ CF₄(g) + 4 HF(g)     ΔH = 38.2 kJ/mol

a) increase pressure                          c) add 0.31 moles of F₂                   e) cool down reaction vessel
b) increase temperature                    d) add some Xenon gas                  f) decrease volume

• A. a) and d)
• B. a) and e)
• C. b) and c)
• D. b) and e)
• E. a), d) and e)
• F. b), c), and f)

At 26°C K_p = 2.3 x 10^–12 and at 56°C K_p = 3.7 x 10^–5 for the hypothetical reaction. Determine if the reaction is endothermic or exothermic.

AB(s) + 2 B(g) ⇌ 3 BB (g) + C (g)

• A. endothermic
• B. exothermic
• C. enthalpy change equal to zero

Here’s what students ask on this topic:

What is Le Chatelier's Principle and how does it apply to chemical equilibrium?

Le Chatelier's Principle states that if a chemical reaction at equilibrium is disturbed, the system will adjust itself to counteract the disturbance and reestablish equilibrium. This principle applies to changes in concentration, pressure, volume, and temperature. For example, if the concentration of reactants is increased, the system will shift towards the products to reduce the added reactants. Similarly, if the pressure is increased by decreasing the volume, the system will shift towards the side with fewer moles of gas. Temperature changes also affect equilibrium; increasing temperature shifts the equilibrium away from the heat source in exothermic reactions and towards the heat source in endothermic reactions.

How does changing the concentration of reactants or products affect the equilibrium position according to Le Chatelier's Principle?

According to Le Chatelier's Principle, changing the concentration of reactants or products will cause the equilibrium to shift in a direction that counteracts the change. If the concentration of reactants is increased, the system will shift towards the products to reduce the excess reactants. Conversely, if the concentration of products is increased, the system will shift towards the reactants to reduce the excess products. This shift helps to reestablish equilibrium by balancing the concentrations of reactants and products.

What is the effect of pressure and volume changes on chemical equilibrium?

Pressure and volume changes affect chemical equilibrium in systems involving gases. According to Le Chatelier's Principle, increasing the pressure (by decreasing the volume) will shift the equilibrium towards the side with fewer moles of gas. Conversely, decreasing the pressure (by increasing the volume) will shift the equilibrium towards the side with more moles of gas. This shift occurs because the system adjusts to minimize the disturbance caused by the change in pressure or volume.

How does the addition of an inert gas affect the equilibrium position of a reaction?

The addition of an inert gas, such as a noble gas, does not affect the equilibrium position of a reaction if it is added at constant volume. This is because inert gases do not react with the components of the reaction and do not change the partial pressures of the reactants or products. However, if an inert gas is added at constant pressure, the volume of the system increases, causing the equilibrium to shift towards the side with more moles of gas to accommodate the increased volume.

How does temperature change affect the equilibrium constant and the direction of the shift in a chemical reaction?

Temperature changes affect both the equilibrium constant (K) and the direction of the shift in a chemical reaction. For exothermic reactions (ΔH < 0), increasing the temperature shifts the equilibrium towards the reactants, decreasing the value of K. For endothermic reactions (ΔH > 0), increasing the temperature shifts the equilibrium towards the products, increasing the value of K. Conversely, decreasing the temperature shifts the equilibrium towards the heat source: towards the products in exothermic reactions and towards the reactants in endothermic reactions. This shift helps to reestablish equilibrium by balancing the heat absorbed or released.