Acid–Base Equilibrium for AP Chemistry (page 2)
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.
Acidic/Basic Properties of Salts
The behavior of a salt will depend upon the acid–base properties of the ions present in the salt. The ions may lead to solutions of the salt being acidic, basic, or neutral. The pH of a solution depends on hydrolysis, a generic term for a variety of reactions with water. Some ions will undergo hydrolysis and this changes the pH.
The reaction of an acid and a base will produce a salt. The salt will contain the cation from the base and the anion from the acid. In principle, the cation of the base is the conjugate acid of the base, and the anion from the acid is the conjugate base of the acid. Thus, the salt contains a conjugate acid and a conjugate base. This is always true in principle. In some cases, one or the other of these ions is not a true conjugate base or a conjugate acid. Just because the ion is not a true conjugate acid or base does not mean that we cannot use the ion as if it were.
The conjugate base of any strong acid is so weak that it will not undergo any significant hydrolysis; the conjugate acid of any strong base is so weak that it, too, will not undergo any significant hydrolysis. Ions that do not undergo any significant hydrolysis will have no effect upon the pH of a solution and will leave the solution neutral. The presence of the following conjugate bases Cl–, Br–, I–, NO3–, ClO3–, and ClO4– will leave the solution neutral. The cations from the strong bases, Li+, Na+, K+, Rb+, Cs+ Ca2+, Sr2+, and Ba2+, while not true conjugate acids, will also leave the solution neutral. Salts containing a combination of only these cations and anions are neutral.
The conjugate base from any weak acid is a strong base and will undergo hydrolysis in aqueous solution to produce a basic solution. If the conjugate base (anion) of a weak acid is in a salt with the conjugate of a strong base (cation), the solution will be basic, because only the anion will undergo any significant hydrolysis. Salts of this type are basic salts. All salts containing the cation of a strong base and the anion of a weak acid are basic salts.
The conjugate acid of a weak base is a strong acid and it will undergo hydrolysis in an aqueous solution to make the solution acidic. If the conjugate acid (cation) of a weak base is in a salt with the conjugate base of a strong acid (anion), the solution will be acidic, because only the cation will undergo any significant hydrolysis. Salts of this type are acidic salts. All salts containing the cation of a weak base and the anion of a strong acid are acidic salts.
There is a fourth category, consisting of salts that contain the cation of a weak base with the anion of a weak acid. Prediction of the acid–base character of these salts is less obvious, because both ions undergo hydrolysis. The two equilibria not only alter the pH of the solution, but also interfere with each other. Predictions require a comparison of the K values for the two ions. The larger K value predominates. If the larger value is Ka, the solution is acidic. If the larger value is Kb, the solution is basic. In the rare case where the two values are equal, the solution would be neutral.
The following table summarizes this information:
For example, suppose you are asked to determine if a solution of sodium carbonate, Na2CO3, is acidic, basic, or neutral. Sodium carbonate is the salt of a strong base (NaOH) and a weak acid (HCO3–). Salts of strong bases and weak acids are basic salts. As a basic salt, we know the final answer must be basic (pH above 7).
Practice problems for these concepts can be found at: