In my post on chemical equilibrium, I explained that a reaction is at equilibrium when the rates of the forward reaction and the reverse reaction are equal. At equilibrium, certain concentrations of reactants and products remain. Some reactions are product-favored, meaning they produce mostly products, whereas other reactions are reactant-favored, meaning they produce little product in comparison to reactants.
As I said, from a commercial standpoint, it is ideal to run a reaction that produces the maximum amounts of products possible. Thankfully, there is a way to manipulate or “shift” equilibrium so that the reaction produces more (or less) products than it normally would. The principal that governs the shifting of equilibrium by changing reaction conditions is Le Chatelier’s Principle.
Aptly summarized by my chemistry professor, Le Chatelier’s Principle simply states that when you change the conditions of a reaction, the reaction will shift to alleviate the stress on the system. If, for example, you add more energy, the reaction will run to use the extra energy. If you take away reactants, the reverse reaction will run to replace them. If you increase the volume, the reaction will run to maintain pressure.
Le Chatelier’s Principle is interesting (and the most fun part of studying chemical equilibrium, if you ask me) because it’s intuitive and largely qualitative. The reason I described it as being “wibbly-wobbly” in the equilibrium post is because it makes apparent the fact that equilibrium is dynamic (and, apparently, that it doesn’t like change).
There are three basic things you can change to shift equilibrium: concentration, volume/pressure and temperature.
The easiest condition change to grasp intuitively is a change in concentration. Put in simple terms, whenever you remove some of something on one side, equilibrium will shift to that side. Likewise, if you add something on one side, equilibrium will shift to the other side. Take this reaction (taken from my professor’s supplemental homework) for example:
H2 (g) + I2 (g) ⇌ 2HI (g)
Let’s say that, when this reaction is at equilibrium, we remove some of the hydrogen iodide. Equilibrium, then, would shift to accommodate this change by shifting toward products. Likewise, if we were to remove some hydrogen or iodine gas, equilibrium would shift toward reactants.
Pretty simple, eh? I told you it was shiny.
The reason I say volume/pressure is that, for many problems, the effect is the same. Using PV = nRT or our intuition, we know that as volume increases (i.e., the reaction vessel gets larger), pressure decreases and vice versa. I could have technically written volume/concentration for the above heading as well, since adding more solvent to aqueous reactions can both increase the volume and decrease the concentration, but most often when I’ve seen volume changes mentioned they’ve been in the context of gases and pressures.
This is the only application of Le Chatelier’s Principle that seems difficult, and it is mostly because part of it can seem counter-intuitive. Let’s start with the easy part, shall we?
When you decrease the volume of the reaction vessel that a reaction is running in (effectively increasing the pressure), equilibrium will shift in accordance to the number of moles on each side. Look at this reaction:
A (g) + B (g) ⇌ C (g) + D (g)
Here, we have two moles of reactants and two moles of products (the coefficients on each reactant and product are one, and we simply add them together). If we were to change the pressure of this system, equilibrium would not change since the pressure would still be equal on either side.
Now, what if we have this reaction?:
A (g) + B (g) ⇌ C (g)
Here, we have two moles of reactants and one mole of products. At equilibrium, if the volume of the vessel was increased so that the pressures of the reagents decreased, equilibrium would shift to compensate for the pressure decrease. C would be converted into A and B, which would increase pressure (since there are more moles on the reactants side than the products side). The inverse is also true; decreasing volume and increasing pressure would cause the reaction to shift toward products to alleviate the pressure.
Are you still with me? Cool! Two down, one to go.
Changing temperature does interesting things to equilibrium. In my post on rates of reaction, I said that increasing temperature increases energy in a system which, in turn, makes reactions run faster. The shifts of equilibrium resulting from increasing or decreasing temperature of a reaction have to do with energy being added or subtracted from a system in this way.
The most difficult part of working with Le Chatelier’s Principle and temperature is knowing whether a reaction is endothermic or exothermic. Put in simple terms, an endothermic reaction is one that requires energy to run (and, thus, decreases the temperature around it), whereas an exothermic reaction is one that releases energy (and increases the temperature of its surroundings). Combustion is exothermic, whereas ice melting is endothermic.
Why is this significant? Turns out, whether energy is consumed or evolved by a reaction determined whether energy is treated as a product or as a reactant. Take, for example, the following generic reaction:
A (aq) + B (aq) ⇌ C (aq)
Let’s say that this reaction is endothermic. Knowing that energy goes into this reaction, we could write it like this:
A (aq) + B (aq) + energy ⇌ C (aq)
According to Le Chatelier’s Principle, adding reactant causes equilibrium to shift toward products. In this case, increasing temperature would cause more of product C to be produced. Likewise, decreasing the temperature would cause more of reactants A and B to be produced.
The opposite is true of an exothermic reaction. Let’s say the following reaction is exothermic:
D (aq) + E (aq) ⇌ F (aq) + G (aq)
Knowing energy is released, we write it with energy in the right place and get the following:
D (aq) + E (aq) ⇌ F (aq) + G (aq) + energy
In this case, increasing temperature would cause a shift toward reactants, whereas decreasing temperature would cause a shift toward products.
Nifty, huh? Good to know chemicals have problems with change too, eh?
Now that we understand Le Chatelier’s Principle, we know how equilibrium shifts under different conditions. Now we can apply that to acid-base chemistry, where weak and strong acids and bases break up into parts and cause pain.
Questions? Comments? A fan of Doctor Who? Say so in the comments!
Complaints? Corrections? Kindly throw them at me! Flames will be used to reverse exothermic processes. (Sorry, I like making jokes about flames. I’m really not that easy to offend, you guys.)