Acid–Base Equilibrium for AP Chemistry
Practice problems for these concepts can be found at:
- Equilibrium Multiple Choice Review Questions for AP Chemistry
- Equilibrium Free-Response Questions for AP Chemistry
The reactant quotient can be written at any point during the reaction, but the most useful point is when the reaction has reached equilibrium. At equilibrium, the reaction quotient becomes the equilibrium constant, Kc (or Kp if gas pressures are being used). Usually this equilibrium constant is expressed simply as a number without units, since it is a ratio of concentrations or pressures. In addition, the concentrations of solids or pure liquids (not in solution) that appear in the equilibrium expression are assumed to be 1, since their concentrations do not change.
Consider the Haber process for the production of ammonia:
The equilibrium constant expression would be written as:
If the partial pressures of the gases were used, then Kp would be written in the following form:
There is a relationship between Kc and Kp: Kp = Kc(RT)Δn, where R is the ideal gas constant (0.0821 L atm/mol K) and Δn is the change in the number of moles of gas in the reaction.
Remember: Be sure that your value of R is consistent with the units chosen for the partial pressures of the gases.
For the following equilibrium Calculate Kc for this equilibrium at 25°C.
The numerical value of the equilibrium constant can give an indication of the extent of the reaction after equilibrium has been reached. If the value of Kc is large, that means the numerator is much larger than the denominator and the reaction has produced a relatively large amount of products (reaction lies far to the right). If Kc is small, then the numerator is much smaller than the denominator and not much product has been formed (reaction lies far to the left).
Acid Base Equilibrium
Recall that acids are proton (H+) donors and bases are proton acceptors. Also recall that acids and bases may be strong or weak. Strong acids completely dissociate in water; weak acids only partially dissociate. For example, consider two acids HCl (strong) and CH3COOH (weak). If each is added to water to form aqueous solutions the following reactions take place:
The first reaction essentially goes to completion—there is no HCl left in solution. The second reaction is an equilibrium reaction—there are appreciable amounts of both reactants and products left in solution.
There are generally only two strong bases to consider: the hydroxide and the oxide ion (OH– and O2–, respectively). All other common bases are weak. Weak bases, like weak acids, also establish an equilibrium system, as in aqueous solutions of ammonia:
In the Bronsted–Lowry acid–base theory, there is competition for an H+. Consider the acid–base reaction between acetic acid, a weak acid, and ammonia, a weak base:
Acetic acid donates a proton to ammonia in the forward (left-to-right) reaction of the equilibrium to form the acetate and ammonium ions. But in the reverse (right-to-left) reaction, the ammonium ion donates a proton to the acetate ion to form ammonia and acetic acid. The ammonium ion is acting as an acid, and the acetate ion as a base. Under the Bronsted–Lowry system, acetic acid (CH3COOH) and the acetate ion (CH3COO–) are called a conjugate acid–base pair. Conjugate acid–base pairs differ by only a single H+. Ammonia (NH3) and the ammonium ion (NH4+) are also a conjugate acid–base pair. In this reaction there is a competition for the H+ between acetic acid and the ammonium ion. To predict on which side the equilibrium will lie, this general rule applies: The equilibrium will favor the side in which the weaker acid and base are present. Figure 15.1 shows the relative strengths of the conjugate acid–base pairs.
In Figure 15.1 you can see that acetic acid is a stronger acid than the ammonium ion and ammonia is a stronger base than the acetate ion. Therefore, the equilibrium will lie to the right.
The reasoning above allows us to find good qualitative answers, but in order to be able to do quantitative problems (how much is present, etc.), the extent of the dissociation of the weak acids and bases must be known. That is where a modification of the equilibrium constant is useful.