How do we treat acids with more than one ionizable hydrogen atom?
In most cases, polyprotic acids can be treated in exactly the same fashion as monoprotic acids.
Treat each ionized hydrogen ion one step at a time.
Successive acid ionization constants are number sequentially: Ka1, Ka2, Ka3, ...
What is the pH of a 0.0050 M solution of sulfuric acid at 25 °C?
Strategy: The first proton of sulfuric acid is completely ionized – it's a strong acid.
Bisulfate ion is a weak acid, so we will need to use the usual weak acid procedures, remembering that the first step of the sulfuric acid ionization generates both hydronium ion and bisulfate ion.
The range of possible pH values is –log(2×0.0050) = –log(0.010) = 2.00, (assuming both hydrogen ions ionized completely) to 7.00 (pure water, assuming sulfuric acid were nonacidic).
What is the pH of a 0.037 M solution of phosphoric acid at 25 °C?
Strategy: H3PO4 has three ionizable hydrogen atoms, so this will be a three-step problem.
This is a weak acid, so treat each ionization reaction as an independent problem, using the results of previous steps, as appropriate.
The range of possible pH values is 0.96 (all three H+ donated, = –log(3×0.037) = –log(0.11)) to 7.00 (pure water, no H+ donated).
Since this is a weak acid, the dihydrogen phosphate and hydrogen phosphate ions should have significantly smaller Ka values, so more realistically, the lower extreme for the pH is probably 1.43 (= –log(0.037)).
The general Brønsted–Lowry reaction for a weak base is:
B(aq) + H2O(l) → ← HB+(aq) + OH–(aq)
The equilibrium constant for this reaction is Kb = [HB+]e[OH–]e [B]e
Kb is the base ionization constant. As with Ka, Kb defines the reaction.
We can also define a log scale:
pKb = –logKb
Better bases give a higher concentration of hydroxide ion: this is indicated by a larger Kb or a smaller pKb.
Metal hydroxides. The strong bases all fit this category. Most metal hydroxides are only sparingly soluble, so are weak bases. We will look at the equilibrium chemistry of salts later.
Most anions. The negative charge on the anion usually can act as a hydrogen ion acceptor in the presence of water. Exceptions: anions of strong acids (Cl–, Br–, I–, ClO4–, NO3–, HSO4–) and anions with ionizable hydrogen ions (HSO4–, H2PO4–, and others).
Organic amines. These have the general structure of R3N: where R is H, C, O, or other main group atoms. The lone pair on the nitrogen is the underlying reason for the ability to accept a hydrogen ion. The conjugate acid of an amine is called an ammonium ion.
R3N:(aq) + H2O(l) → ← R3N–H+(aq) + OH–(aq)
RNH2(aq) + H2O(l) → ← RNH3+(aq) + OH–(aq)
Common amines:
A problem: outside of general chemistry textbooks, Kb (or pKb) is rarely reported. Rather, we find the pKa of the conjugate acid.
What is the relationship between Ka and Kb of a conjugate acid/base pair?
We treat equilibrium problems involving weak bases in exactly the same fashion as weak acids.
Find the pH and percent ionization of a 0.10 M solution of ammonia at 25 °C.
Strategy: Write the reaction, write the mass action expression, find Kb, set up a table of concentrations, plug into the mass action expression, solve for the variable, and then use this information to answer the questions asked.
What are the limits of pH?
If ammonia were a strong base (100% ionized), then pOH = –log(0.10) = 1.00 so the pH = 14.00 – 1.00 = 13.00. If ammonia were completely unreactive, then the pH = 7.00 (pure water).
