AP Chemistry: Salt Hydrolysis

When you dissolve table salt in water, you expect a neutral solution. But what happens when you dissolve ammonium chloride or sodium acetate? The pH can change dramatically because the ions themselves react with water in a process called salt hydrolysis. Mastering this concept is essential for predicting solution behavior in chemical synthesis, buffer design, and even understanding physiological processes, from blood pH regulation to drug absorption.

The Foundation: Acids, Bases, and Their Salts

To predict hydrolysis, you must first classify the parent acid and base that formed the salt. A strong acid (like HCl, HNO3, H2SO4, HI, HBr, HClO4) completely dissociates in water. Its conjugate base is so weak it shows negligible basicity. A weak acid (like CH3COOH, HF, H2CO3) only partially dissociates; its conjugate base is relatively strong and will accept a proton. The same logic applies to bases: strong bases (Group 1 and 2 hydroxides, like NaOH and KOH) fully dissociate, leaving a conjugate acid too weak to donate a proton. A weak base (like NH3, CH3NH2) has a conjugate acid that is relatively strong and will donate a proton.

A salt is the product of an acid-base neutralization. Hydrolysis occurs when a cation or anion from the dissolved salt reacts with water, altering the concentration of $H_3{O}^{+}$ or $O{H}^{-}$ ions. Whether this happens depends entirely on the strength of the parent acid and base.

The Four Categories of Salts and pH Prediction

By combining the strength of the parent acid and base, we can categorize any salt into one of four types and predict the resulting solution pH.

1. Salt of a Strong Base and a Strong Acid Example: NaCl (from NaOH and HCl). Here, the cation ($N{a}^{+}$) is the conjugate acid of a strong base (NaOH) and is pH-neutral. The anion ($C{l}^{-}$) is the conjugate base of a strong acid (HCl) and is also pH-neutral. Neither ion undergoes hydrolysis. The solution contains only neutral ions and water, resulting in a pH = 7 (at 25°C).

2. Salt of a Strong Base and a Weak Acid Example: NaCH3COO (Sodium acetate, from NaOH and CH3COOH). The cation ($N{a}^{+}$) is neutral. However, the anion ($C{H}_{3}CO{O}^{-}$) is the conjugate base of the weak acid acetic acid. It is a relatively strong base that will hydrolyze: $$C{H}_{3}CO{O}^{-}(aq)+H_{2}O(l)\rightleftharpoons C{H}_{3}COOH(aq)+O{H}^{-}(aq)$$ This reaction produces $O{H}^{-}$ ions. Result: pH > 7 (basic).

3. Salt of a Weak Base and a Strong Acid Example: NH4Cl (Ammonium chloride, from NH3 and HCl). The anion ($C{l}^{-}$) is neutral. The cation ($N{H}_4^{+}$) is the conjugate acid of the weak base ammonia. It is a relatively strong acid that will hydrolyze: $$N{H}_4^{+}(aq)+H_{2}O(l)\rightleftharpoons N{H}_3(aq)+H_3{O}^{+}(aq)$$ This reaction produces $H_3{O}^{+}$ ions. Result: pH < 7 (acidic). This is why ammonium salts are used in some fertilizers to create slightly acidic soil conditions.

4. Salt of a Weak Base and a Weak Acid Example: NH4CH3COO (Ammonium acetate, from NH3 and CH3COOH). Both ions can hydrolyze. The $N{H}_4^{+}$ cation tends to make the solution acidic, while the $C{H}_{3}CO{O}^{-}$ anion tends to make it basic. The final pH depends on the relative strengths of the parent acid and base, specifically the comparison of their ionization constants ($K_a$ of the weak acid vs. $K_b$ of the weak base).

• If $K_a>K_b$, the weak acid is relatively stronger, meaning its conjugate base ($A^{-}$) is weaker. The cation's hydrolysis wins, and the solution is acidic.
• If $K_b>K_a$, the weak base is relatively stronger, meaning its conjugate acid ($B{H}^{+}$) is weaker. The anion's hydrolysis wins, and the solution is basic.
• If $K_a\approx K_b$, the effects nearly cancel, and the solution is approximately neutral.

Quantitative pH Calculations for Hydrolyzing Salts

For salts where only one ion hydrolyzes (Categories 2 & 3), we can calculate the pH directly using an equilibrium constant for hydrolysis.

For a Salt of a Weak Acid and Strong Base (Basic Solution): The anion ($A^{-}$) is the conjugate base of weak acid HA. Its hydrolysis constant, $K_b$ for $A^{-}$, is related to $K_a$ for HA by the water autoionization constant, $K_w$. The relationship is: $K_a\times K_b=K_w=1.0\times 10^{-14}$ at 25°C. Therefore, $K_b=K_w/K_a$.

Let’s calculate the pH of a 0.10 M sodium acetate ($NaC{H}_{3}COO$) solution. $K_a$ for $C{H}_{3}COOH$ is $1.8\times 10^{-5}$.

1. Find $K_b$ for $C{H}_{3}CO{O}^{-}$: $K_b=K_w/K_a=(1.0\times 10^{-14})/(1.8\times 10^{-5})=5.6\times 10^{-10}$.
2. Set up an ICE table for the hydrolysis: $C{H}_{3}CO{O}^{-}+H_{2}O\rightleftharpoons C{H}_{3}COOH+O{H}^{-}$.

Initial: [$C{H}_{3}CO{O}^{-}$] = 0.10 M, [$O{H}^{-}$] = 0. Change: -x, +x, +x. Equilibrium: 0.10 - x, x, x.

1. $K_b=[C{H}_{3}COOH][O{H}^{-}]/[C{H}_{3}CO{O}^{-}]=(x)(x)/(0.10-x)\approx x^2/0.10=5.6\times 10^{-10}$.
2. Solve for x: $x=[O{H}^{-}]=\sqrt{(5.6\times 10^{-10})(0.10)}=\sqrt{5.6\times 10^{-11}}=7.5\times 10^{-6}M$.
3. Find pOH: $pOH=-\log (7.5\times 10^{-6})=5.12$.
4. Find pH: $pH=14.00-5.12=8.88$.

For a Salt of a Weak Base and Strong Acid (Acidic Solution): The cation ($B{H}^{+}$) is the conjugate acid of weak base B. Its hydrolysis constant, $K_a$ for $B{H}^{+}$, is: $K_a=K_w/K_b$.

To find the pH of a 0.10 M $N{H}_{4}Cl$ solution, given $K_b$ for $N{H}_3$ is $1.8\times 10^{-5}$:

1. Find $K_a$ for $N{H}_4^{+}$: $K_a=K_w/K_b=(1.0\times 10^{-14})/(1.8\times 10^{-5})=5.6\times 10^{-10}$.
2. Hydrolysis: $N{H}_4^{+}+H_{2}O\rightleftharpoons N{H}_3+H_3{O}^{+}$.
3. $K_a=[N{H}_3][H_3{O}^{+}]/[N{H}_4^{+}]=x^2/0.10=5.6\times 10^{-10}$.
4. $x=[H_3{O}^{+}]=\sqrt{(5.6\times 10^{-10})(0.10)}=7.5\times 10^{-6}M$.
5. $pH=-\log (7.5\times 10^{-6})=5.12$.

Notice the symmetry: salts of weak acid/strong base and weak base/strong acid with equal $K_a$/$K_b$ values yield pH values that are symmetric around 7.

Common Pitfalls

1. Misidentifying "Spectator" Ions: The most common error is forgetting that the conjugate acid of a strong base (e.g., $N{a}^{+}$, $K^{+}$, $C{a}^{2+}$) and the conjugate base of a strong acid (e.g., $C{l}^{-}$, $N{O}_3^{-}$, $Cl{O}_4^{-}$) do not hydrolyze. They are true spectators. Only assume hydrolysis for ions derived from weak parents.

1. Incorrectly Applying $K_a$ and $K_b$: When calculating pH for a salt of a weak acid, you must use the *$K_b$ of the anion*, which you find from $K_w/K_a$. Students often mistakenly try to use the $K_a$ of the parent weak acid directly in the calculation for a basic salt solution. Remember: you are dealing with the base behavior of the anion, so you need a base ionization constant ($K_b$).

1. Overcomplicating Salts of Weak Acid and Weak Base: For a salt like $N{H}_{4}C{H}_{3}COO$, you don't need to perform two simultaneous equilibrium calculations at the AP level. Instead, compare $K_a$ and $K_b$. If they are equal ($K_a$ of $C{H}_{3}COOH\approx K_b$ of $N{H}_3$), the solution is neutral. If they are not, you need more advanced math beyond the AP scope; simply predict if the solution is acidic or basic based on which constant is larger.

1. Ignoring Concentration: The pH of a hydrolyzing salt solution does depend on its concentration, as shown in the quantitative calculations. A more dilute solution of $N{H}_{4}Cl$ will be less acidic (pH closer to 7) than a more concentrated one. Always consider the initial salt concentration when setting up your ICE table.

Summary

• Salt hydrolysis is the reaction of a salt's cation or anion with water, changing the solution's pH. The driving force is the formation of a weak acid or weak base.
• Predict pH by analyzing the parent acid and base: Strong Acid + Strong Base = Neutral. Strong Base + Weak Acid = Basic. Weak Base + Strong Acid = Acidic. Weak Acid + Weak Base = pH depends on relative $K_a$/$K_b$.
• For quantitative pH calculations of salts with a single hydrolyzing ion, use the relationships $K_a\times K_b=K_w$ and standard ICE table procedures to find $[H_3{O}^{+}]$ or $[O{H}^{-}]$.
• Avoid pitfalls by correctly identifying non-hydrolyzing spectator ions, using the correct hydrolysis constant ($K_a$ for a cationic acid, $K_b$ for an anionic base), and accounting for the salt's initial concentration.