Reversible Reactions

A piece of iron rusting in the atmosphere may not appear to change very rapidly, but it is not an equilibrium. Since the oxygen required for rusting comes from a virtually unlimited supply, the metal will completely rust away, given enough time. This is an open, not a closed system.

By now you should realize that chemical reactions can go in two directions:

from reactants to products (the forward reaction)
from products to reactants (the reverse reaction).

All chemical reactions are reversible. Reactions in closed systems will reach equilibrium, given enough time, though that may be a very long time indeed!

If the system is not closed, materials can enter or leave. The reaction will not reach an unchanging state. Equilibrium can only occur in a closed system.

When a system is at equilibrium it has constant observable properties.

However, every reaction that does not appear to change is not necessarily at equilibrium:

the reaction may react so much towards the products that it appears to have gone to completion. Such reactions are actually equilibria, but since almost all the reactants turn into products, we normally don't apply equilibrium principles to them. When a reaction goes to completion, the reaction reaches an equilibrium that is almost entirely products.

the reaction is so slow that it has not reached equilibrium yet, though it will given enough time.

if the system is open, and materials are being removed or added in such a way that there is no net change, a steady state has been reached. This is not an equilibrium, but it is a fairly common state for a reaction to reach.

One of the characteristics of equilibrium is that the same position is reached, whether you start the reaction on the product or reactant side. One way to identify if a reaction is really at equilibrium is to:

start with some reactants, and no products. Let the reaction continue until it appears to stop.

start with some products, and no reactants. Let the reaction continue until it appears to stop.

If you arrive at the same end condition in both trials, the reaction is at equilibrium.

dichromatestart.jpg (7935 bytes)
Dichromate solutions at the start

dichromateend.jpg (8370 bytes)
Dichromate solutions at the end

Time-lapse video of this experiment.

Consider the solubility equilibrium of potassium dichromate (K₂Cr₂O₇) in water. This time-lapse video shows two beakers of solutions of K₂Cr₂O₇. The beaker on the left is a saturated solution at room temperature. It was prepared by putting solid K₂Cr₂O₇ into water at room temperature, and stirring, until the solution became saturated. In other words, this reaction was started from the reactant side of the solubility equation:

K₂Cr₂O₇ (s)		2 K⁺ (aq) + Cr₂O₇^2- (aq)
Reactants		Products

The beaker on the right contains an unsaturated solution of K⁺ and Cr₂O₇^2- at an elevated temperature. At first there is no solid potassium dichromate present, so it is not yet at equilibrium.. This beaker starts on the products side of the solubility equation. As the solution cools, the solubility of the potassium dichromate decreases, and so it begins to crystallize. After approximately 30 minutes, the two beakers reach the same temperature, and they have the same unchanging color intensity. This indicates they have reached the same equilibrium concentration.

Reaching the same concentration level from either the starting point of solid K₂Cr₂O₇ or the dissolved ions K⁺ (aq) and Cr₂O₇^2-(aq) shows that a saturated solution is a true equilibrium.