Equilibrium is a term that will be used extensively in this module. Some reactions that take place are reversible. Equilibrium, then is a general term that expresses a balance between opposing reactions.   We express this balance as follows:

The concepts that observation and microscopic behavior of systems often fails us in looking at equilibrium.  Open the kitchen cabinet and look, just look at the sealed bottle of clear, colorless cider vinegar there.  Nothing appears to be happening, that is our sense anyway.  But the appearance of lack of change masks the physical and chemical changes that are actually proceeding.  But these changes are taking place so that the forward and reverse processes take place at the same rate.  Let's look at three of these changes.

Vinegar is a water solution of acetic acid, CH3COOH.  First, let's think about the water in the bottle, in the closet at 25 o C.   As we recall, we will find water vapor in the space above the liquid from the following physical transformation: H2O(l)  -->   H2O(g) And the water vapor can return to liquid by the reverse of the above equation: H2O(g)  -->   H2O(l) Each of these processes takes place at a rate effected by the concentration of liquid and vapor and by the temperature.  But what is certain is that on the shelf, in the closed container, after many years, if the temperature remains constant,  the amount of vapor in the space above the liquid remains constant.  Not because nothing is happening, but because the two reactions, vaporization and condensation occur at the same rate. The system is in equilibrium.

But if the closet warms, the rate of vaporization increases - remember vapor pressure increases with temperature so more vaporization will occur.   So the system is no longer in equilibrium since the forward and reverse rates are no longer the same.  Only when the return rate of vapor to liquid increases as result of increased vapor concentration, will the system return to equilibrium.  Equilibrium is a hard state to maintain!

The second consideration of our vinegar container is the chemistry of the acetic acid dissolved in the water in the bottle.  This solution is in equilibrium also if we do not disturb it.

We are misled by the benign and static condition of the container.  Two reactions are taking place, the ionization of acetic acid into a hydronium ion ( H3O⁺) and the acetate anion (CH3COO^- ) and the reverse reaction restoring the charged species to water and unchanged acetic acid.   In the unopened container at constant temperature, these processes reach the same rate and we achieve an equilibrium concentration of each of these species at equilibrium.  These concentrations will remain constant unless something disturbs the system.

Acetic acid is a liquid with a vapor pressure, too.   And this leads us to the third consideration of equilibrium.  As with water, the free, unionized acetic acid will go into the vapor state.

Suppose the vapor finds its way behind the seal and into the metal of the cap of a vintage bottle of vinegar.  There corrosion reactions will take place.  Let's speculate that acetic acid will react with iron in the cap:

Fe  +  2CH3COOH   -->  Fe⁺²  +2CH3COO^-   +  H2

This process removes acetic acid from the vapor.   The reverse reaction that allows acetic acid vapor to return to solution slows because of the now lowered concentration of vaporous acid.  Therefore more acetic acid goes into the vapor than returns.  The system is no longer in equilibrium but it is trying to get there by increasing vapor concentration so that vaporization and condensation rates will be the same.   In this old bottle, it will likely be a futile quest, since corrosion of the cap will continue to pull away acetic acid vapor.  Eventually the cap will be eaten through and the contants of the bottle will be free to escape to the outside.  Any chance for an equilibrium condition is lost.

It is important to understand that acid-base reactions are all reversible reactions. Reactions are proceeding in a forward reaction and at the same time, the reaction is being reversed.  We can seek to describe the equilibrium conditions but we must be wary of truly reaching that state.

In acid/base chemistry we express the reaction of an acid and a base as proceeding to produce two new products, the conjugate acid of the base formed by the addition of a proton and the conjugate base of the acid formed by loss of a proton:

At equilibrium, these reactions (acids to conjugate bases and back again, and bases to conjugates acids)  are constantly taking place, but at equal velocities.

Some reactions are irreversible, such as drug absorption. In this case, the reactant (drug) is being removed, and thus the reaction will proceed in a forward fashion according to LeChatelier's Principle.

Let’s take a look at a simple acid-base reaction. Hydrochloric acid dissociates completely in water, and is thus classified as a strong acid. This reaction is not an equilibrium reaction, since the reaction proceeds completely until all of the HCl is dissociated. Notice the arrow points only to the right.

According to the Bronsted-Lowry theory, HCl is an acid because it is capable of donating a proton.

Water may either donate or accept a proton, so it is termed amphiprotic. In this case, water acts as a base, since it accepts a proton to form the hydronium ion. The hydronium ion on the right side of the equation then becomes an acid,  called the conjugate acid, and the anionic chloride ion is a base, albeit a very weak base, called the conjugate base, capable now of accepting a proton to then become an acid again! This illustration shows how the species are "paired"; capable of being an acid or a base, and these reactions are taking place simultaneously. So at any given time within a solution of HCl in water, there are both acids and bases, existing in equal amounts.

Let’s take a look at a corresponding example of a weak base and water:

NH₃ is the base, and NH₄⁺ its conjugate acid. (Likewise, H₂0 can be considered an acid and OH^- its conjugate base.) In this example, NH₃ is a weak base with little tendency to react with water to form the ammonium ion.  The reaction proceeds in both forward and reverse directions.   Implicit in this explanation is the fact that at equilibrium the forward reaction proceeds at the same velocity as the reverse reaction, while the concentrations of the reactants and products remains the same.

Written generically, these acid-base reactions are as follows:

According to the law of mass action, the equilibrium constant for this reaction may be expressed as the product of the ionized species divided by the product of the uninoized species:

Take a look at each reaction individually.

The rate constant for the forward reaction may be expressed as:

k_r=[HA] [H₂0]

and the rate constant for the reverse reaction as:

k_f= [A^-] [H₃0⁺]

Please remember that the [] indicates molar concentration of the reactants. Since the reaction takes place in a one-to-one ratio of reactants, the exponent for each of the concentrations in THIS case is 1.

Also please note that the k values represent a velocity of the reactions. In an equilibrium reaction, the reactions in both directions are taking place at the same velocity, so k is equal for both the forward and reverse reactions.

This brings us to the important topic of the ionization of or autoprotolysis of water.

 Something to remember:

• Many processes can be at equilibrium.  But with changes in condition - concentration, temperature -  the system will no longer be at equilibrium and will adjust to try to get there again.