LeChatelier's+Principle

LeChatelier's Principle
When a chemical reaction reaches a state of equilibrium, it is in a stable state. Sometimes, you hear people talk about their balance, or equilibrium. You can also think of a reaction at equilibrium as being 'in balance'.

Any changes that are then introduced to the equilibrium system will cause it to become unstable, or out of balance.

The unstable reaction will undergo changes to return to equilibrium. These changes are reactions, designed to counter any imbalance. This is the basis of //LeChatelier's Principle//.

The following are ways in which a chemical reaction can get 'out of balance'.

Changes in Amount
The first and most common changes made to an equilibrium system involve adding or removing reactants or products.

Check out this graph of the concentration versus time for the reaction N 2 (g) + 3 H 2 (g) <=> 2 NH 3 (g)



Here is what happened:
 * 1) initially, the reaction is at equilibrium (hence the flat lines - no changes in concentration).
 * 2) When some more H2 is added to the system, it becomes 'out of balance', and the excess H2 must be used up.
 * 3) In the center (tan shaded region), the H2 is consumed, and this creates more NH3. Note also that some N2 is used up.
 * 4) Once the balance is restored, the system reaches a new equilibrium (shaded in blue).

So you can see that adding a reactant caused more product to be formed.

This reaction is analyzed in the Kahn Academy video on LeChatelier's Principle. (YouTube video - might not work at school)

--

To examine this in more detail, look at the generic chemical equation:

**A + B <=> C + D** Imagine that this chemical reaction has occurred in a beaker. At equilibrium, the solution contains some amounts of each: A, B, C, and D are all inside the beaker.



Suppose some more A is added to the beaker. What will this do to B, C, and D?
 * First, it causes an imbalance: there's too much A now and the reaction is out of equilibrium.
 * To compensate, the reaction will consume A. In order to to this, it must use B, so the amount of B will decrease.
 * Because A and B are consumed, the reaction will produce more C and D.
 * Overall, the reaction produces more products, and this is considered a shift to the products side of the equation, or a "right" shift.
 * The same analogy can be applied to the addition of any of the substances in the beaker. Adding some C or D will cause a "left" shift, and produce more A and B, for example.

Now, suppose some A is removed from the beaker. What happens?
 * First, it causes an imbalance: there's too little A now and the reaction is out of equilibrium.
 * To compensate, the reaction will have to make more A. In order to to this, it must use up some C and D, so their amounts will decrease.
 * Because C and D are consumed, the reaction will produce more A and B.
 * Overall, the reaction produces more reactants, and this is considered a shift to the reactant side of the equation, or a "left" shift.
 * The same analogy can be applied to the removal of any of the substances in the beaker. Removing some C or D will cause a "right" shift, and consume more A and B, for example.

Try it out at this interactive demonstration.

Go here for some practice.

Changes in Pressure or Volume
According to Boyle's Law, when the pressure on a gas is increased, the gas will decrease in volume, and vice versa. So how will these changes affect the equilibrium position if a gas is present?

First, the equilibrium system must contain at least one gas.

Also, there must be an uneven number of gas molecules from the reactant side to the product side.

Here is the idea:
 * As pressure on a gas increases, it makes less room for molecules to move around.
 * The molecules will want to re-arrange to give themselves as much space as possible.
 * What this means is that the equilibrium will shift to the side with fewer gas molecules.

For example, look at this chemical equation: 3 O 2 (g) <=> 2 O 3 (g)

Reducing the volume (increasing the pressure) will cause less room for the gases to move around, so the reaction will do what it can (shift right) to give itself more room.

In reality, the molecules get closer together and undergo more collisions; statistically, the side of the equation with more molecules undergoes more collisions (and this more reactions), forcing the equilibrium to shift.

These analogies are opposite for increasing volume (or decreasing pressure) on gases.

Changes in Temperature
Chemical reactions can either produce heat energy or require heat energy to run. Once at equilibrium, any change to temperature will cause an imbalance of heat energy and the system will adapt to a new equilibrium.

For //exothermic reactions//, heat is a product of the reaction.
 * Increasing the temperature (adding heat) will drive the reaction away from the products side (shift left);
 * decreasing the temperature (removing heat) will force the reaction to produce more heat (shift right).

For //endothermic reactions//, heat is a reactant. Increasing or decreasing the temperature will have the opposite effect as those for exothermic reactions.

Catalysts
A catalyst will cause the rate of the reaction to increase, meaning the reaction will go to completion (or reach equilibrium) faster. However, this has no affect on the reaction that has already reached equilibrium, and no change will occur to the reaction.