Le Chatelier's Principle, also known as the Equilibrium Law, states that when a system at equilibrium is stressed, it will shift in response to minimize the applied stress. This sounds a lot more complicated then it actually is.
Consider the reaction `A + B ↔ C` at equilibrium. What would you predict would happen to the reaction if you increased the concentration of A?
We know that the equilibrium constant can be written as:
`K =[[C]]/([A][B])`
If we increase A, then the reaction is no longer at equilibrium! The reaction has to shift in order to somehow get back to equilibrium. Since we increased the concentration of the reactant `A`, the denominator is now greater than it was at equilibrium. Thus, the reaction will shift to the products by forcing more of `B` and `C` in order to minimize the applied stress.
Notice that in response to stress, the system is not moving out of equilibrium, but rather is re-equilibriating back to a new equilibrium. In other words, the equilibrium constant for the reaction remains the same - after all, it is a constant.
When we say stress, what we mean is some change to the system. This change can come in the form of changes in concentration, pressure, volume, temperature, etc. Anytime you change a physical property of a system, the system will have to re-equilibriate. This is Le Chatelier's Principle, informally:
If something changes a property of a system, the system will shift to reduce the effects of that change.
Imagine that you're standing straight up at an "equilibrium"; i.e, you're perfectly content and stable standing straight up. Now imagine that you suddenly find yourself carrying a 20 lb. weight in your right hand. If you wanted to return back to equilibrium, that is standing up straight, you'd have to shift to your left to counteract the 20 lb. weight. That is Le Chatelier's Principle at its core. Notice that the equilibrium position, you standing up straight, doesn't change in response to the stress.
Don't stress out if that analogy didn't make sense. It probably says more about my analogy writing skills than the complexity of the principle. For the rest of this post, we'll be going over common properties that one can apply Le Chatelier's Principle to. By the end, hopefully you'll have developed an intuition with regards to the principle.
For a given reaction `CH_4 + 2O_2 ↔ CO_2 + 2H_2O`, what would happen if we increased the concentration of `CH_4` ?
In general, if we increase the concentration of a reactant, the system will shift to reduce the concentration of the reactant. This means that the reaction will shift towards the products.
What about if we increase the concentration of `H_2O` ? If we increase the concentration of a product, the system will try to reduce the concentration of the product. Therefore, if we increase the concentration of `H_2O` , more of the reactants will be formed.
In summary: if the concentration of a species is increased, the reaction will shift to reduce the concentration of the species.
#1. For the reaction `N_2 + 3H_2 ↔ 2NH_3` , which direction will the reaction shift if:
a) The concentration of `N_2` increases?
a) The concentration of `N_2` decreases?
a) The concentration of `H_2` increases?
a) The concentration of `NH_3` decreases?
a) Right. The system will try to reduce the amount of `N_2` .
a) Left. The opposite of a) should occur.
a) Right. Same reasoning as a)
a) Left. The system will try to reduce the amount of `NH_3` .
If the pressure increases in a gaseous system, the system will shift to reduce the total moles of gas. The opposite also applies for reducing pressure.
For example, the reaction `2NO_(2(g)) ↔ N_2O_(4(g))` has `2 "moles"` of gas on the reactants side and `1 "mole"` of gas on the products side. If we increase the pressure, the reaction will shift towards the products because the products side has less moles of gas.
Intuitively, this makes sense. As we increase the pressure, the gases become more compressed. When the gases are compressed, they're going to try to reduce the amount of volume they take up. The way they do so is by minimizing the number of moles of gas.
The inverse logic applies for gases. Recall from the ideal gas law that pressure and volume are inversely related i.e if pressure increases, volume decreases. Therefore, as we increase the volume of a system, we're decreasing the pressure.
This means that if we increase the volume of a gaseous system, the system will shift towards the side with more moles of gas. The opposite is true for reducing the volume of a gaseous system: the system will shift to minimize the moles of gas.
In summary: if the pressure is increased in a gaseous system, the system will shift to reduce the moles of gas.
#2. For the reaction `2H_(2(g)) + O_(2(g)) ↔ 2H_2O_((g))` , which direction will the reaction shift if:
a) The concentration of `O_2` decreases?
b) The concentration of `H_2` increases?
c) The pressure increases?
d) The volume increases?
a) Left. The system will try to compensate losing `O_2` by producing more of it.
b) Right. The system will try to reduce the amount of `H_2`.
c) Right. The system will shift to favor the side with less moles of gas. Since the products has 2 moles whereas the reactants has 3, the system will shift to the products.
d) Left. The system will shift to increase the number of moles gas.
Recall from the section on thermodynamics that reactions are either exothermic (heat-releasing) or endothermic (heat-absorbing). When writing exothermic/endothermic reactions in the context of Le Chatelier's Principle, we write heat as a reactant or product. This will help to illustrate how heat behaves. For exothermic reactions, heat is a product. For endothermic reactions, heat is a reactant.
For example, the combustion reaction of propane is an exothermic reaction. Since heat is released as a product, we can directly write heat into the equation:
`C_3H_8 + 5O_2 ↔ 3CO_2 + "heat"`
If we increased heat, which direction would the reaction shift? Just like with the previous cases, if we increase the amount of a species, the system will shift to reduce the amount of the species. Therefore, increasing heat will result in a left-ward shift. Decreasing heat will drive the reactants to the products, as the system will try to compensate for the reduced heat.
In summary: for reactions involving heat, treat heat as if it were just another reactant/product.
#3. For the following reaction:
`Na_2SO_(4(s)) ↔ 2Na_((aq))^+ + SO_(4(aq))^(2-) + "heat"`
a) The concentration of `Na^+` decreases?
b) The concentration of `SO_4` increases?
c) The temperature increases?
d) The temperature decreases?
e) The volume increases?
a) Right. The system will shift to produce more `Na^+` .
b) Left. The system will shift to produce more `Na_2SO_4` .
c) Left. The system will shift to reduce the amount of heat.
d) Right. The system will shift to compensate for the loss of heat.
e) No shift. There are no gaseous products, so volume and pressure changes will not stress the system.
1. Le Chatelier's Principle states that a system at equilibrium will shift to minimize stress in response to any stress.
2. When the concentration of a species is changed, the system will shift to produce the opposite of the change. For example, if the concentration of a species is reduced, the system will shift to produce more of the species.
3. When the pressure of a gaseous system is increased, the system will shift to minimize the moles of gas. The same reasoning applies for decreasing pressure and the corresponding volume changes.
4. When heat is involved such as in exothermic or endothermic reactions, write heat as any other species in the reaction and apply the same reasoning to it as you would any other species.
1. The Haber Process
I've written about the Haber Process a lot since it utilizes a lot of the principles we've covered. In case you haven't encountered it, the Haber Process is the following reaction:
`N_(2(g)) + 3H_(2(g)) ↔ 2NH_(3(g)) + "heat"`
The product of the reaction is ammonia `(NH_3)`, which is utilized on an industrial scale for food production. Since large amounts of ammonia are needed to sustain food production, Le Chatelier's Principle is applied to the reaction in order to shift the reaction as far to the right as possible.
This is done by maintaining a pressure of around `200 "atm"` which, as we've learned in this post, will shift the reaction to favor `NH_3` seeing as there are less moles of gas on the products side. Additionally, a "low" temperature of `400-450°C` is used seeing as the reaction is exothermic. In industrial processes, this actually isn't that high of a temperature- certainly a lower temperature could be used, but at lower temperatures the reaction takes a vastly longer time to complete. `400-450°C` is the ideal region for utilizing Le Chatelier's Principle while not having to wait excess times for the reaction to go to completion.
This process is often credited as one of the greatest triumphs of chemistry as it has allowed mankind to produce vast amounts of food. All from a simple principle!