Hydrolysis of Ions Conjugate to Weak Acids and Bases — Examples

Previously, we explored neutralization reactions. We saw that reacting $\underline{strong}$ acids and bases produced water. So, our pH was neutral. But what happens with weak acids and bases? We know that the conjugates of weak acids and bases are strong. So, how do they affect pH?

What is Hydrolysis?

Conjugates affect pH through hydrolysis reactions. Hydrolysis is any chemical reaction where water molecules break a bond. If you study chemistry in college, you'll see hydrolysis a lot in organic chemistry. But, for acids and bases, hydrolysis happens because of the self-ionization of water. On occasion, two water molecules collide and generate hydronium and hydroxide ions. What happens if we add a conjugate of a weak acid or base?

Let's say we add sodium acetate to water. Technically, sodium is a Lewis acid because it can accept electrons. But, it's unbelievably weak. Acetate is the conjugate base of a weak acid. So, it's a strong base. Let's look at the balanced chemical equations.

$(1)\, CH_{3}COONa + H_{2}O \rightarrow CH_{3}COO^{-} + Na^{+}$

$(2)\, 2H_{2}O \rightleftharpoons H_{3}O^{+} + OH^{-}$

$(3)\, CH_{3}COO^{-} + H_{3}O^{+} \rightleftharpoons CH_{3}COOH + H_{2}O$

The first reaction is the dissociation of the ionic salt. Reaction two is the self-ionization of water. Let's focus on the third reaction. The acetate anion reacts with hydronium. Where did hydronium come from? Reaction two! So, by Le' Chatelier's principle, we know that will shift the equilibrium of the second reaction. But, only the hydronium is reacting. So, we have excess hydroxide. Our solution's pOH is decreasing. By adding acetate, our solution became basic. What would happen if we used the conjugate of a weak base? The opposite. Hydroxide gets consumed. So, our solution becomes more acidic.

Stong Acid with Weak Base

What happens when we react a strong acid and a weak base? We'll assume the acid and base neutralize. But, we still produce a conjugate acid. So our solution will have a pH below 7.0. Let's look at an example.

$HCl + NH_{3} \rightarrow NH_{4}^{+} + Cl^{-}$

Here we have a reaction between hydrochloric acid and ammonia. Ammonia is a Lewis base because of nitrogen's lone pair of electrons. It is also the simplest amine. Hydrochloric acid is a strong Arrhenius acid. Our products are ammonium and chloride ions. Again, chloride is an unbelievably weak base. But, ammonium is a strong acid.

$2H_{2}O \rightleftharpoons H_{3}O^{+} + OH^{-}$

$NH_{4}^{+} + OH^{-} \rightleftharpoons H_{2}O + NH_{3}$

We see that the ammonium ion consumes hydroxide. So, we have excess hydronium. Our solution is acidic!

Strong Base with Weak Acid

Let's look at a reaction between a strong base and weak acid.

$NaOH + HCO_{3}^{-} \rightarrow H_{2}O + Na^{+} + CO_{3}^{2-}$

Here our reaction is between sodium hydroxide and bicarbonate. Bicarbonate is amphoteric. It can be either an acid or a base. Since it's reacting with a base, it's an acid. The products are sodium and carbonate ions.

$2H_{2}O \rightleftharpoons H_{3}O^{+} + OH^{-}$

$CO_{3}^{2-} + H_{3}O^{+} \rightleftharpoons HCO_{3}^{-} + H_{2}O$

Our carbonate anion consumes hydronium. So, we have excess hydroxide. Our solution is basic!

Weak Acid with Weak Base

Sadly, there is one challenging scenario. What happens when we react weak acids and bases? We will produce a conjugate acid and conjugate base. So, we have to determine the strengths of the conjugates. We have a couple of options. First, we can compare either the ionization constants or the dissociation constants. Second, we can evaluate either the starting acids and bases or their conjugates. Why do we have these options? Because for any acid/base and it's conjugate, $K_{a} + K_{b}=K_{w}$ and $pK_{a} + pK_{b}=pK_{w}$.

• If $K_{a}=K_{b}$ then the solution is neutral. In this case, each ion is reacting with water. The conjugate acid generates hydronium. The conjugate base generates hydroxide. But, both occur at equal rates. So, they balance each other.

• If $K_{a}>K_{b}$ then the solution is acidic. Here, the conjugate acid is stronger than the conjugate base. So, there's excess hydronium. Our pH will be less than 7.0.

• If $K_{a}<K_{b}$ then the solution is basic. The conjugate base is stronger than the conjugate acid. So, there's more hydroxide. Our pH will be greater than 7.0.

Equal Values Example

Let's look at an example of when $K_{a}$ and $K_{b}$ are equal. Acetic acid and ammonia have the same values. For ease, we'll compare the dissociation constants of the conjugates. For acetic acid and ammonia, $pK_{a/b}=4.74$. For the ammonium and acetate ions, $pK_{a/b}=9.3$. Let's look at their reactions.

$CH_{3}COOH + NH_{3} \rightleftharpoons CH_{3}COO^{-} + NH_{4}^{+}$

$CH_{3}COO^{-} + H_{3}O^{+} \rightarrow CH_{3}COOH + H_{2}O$

$NH_{4}^{+} + OH^{-} \rightarrow NH_{3} + H_{2}O$

So, we see that acetate reacts with hydronium. Meanwhile, ammonium reacts with hydroxide. Again, hydronium and hydroxide come from the self-ionization of water. But, the $pK_{a/b}$'s are equal. Remember, dissociation constants are just equilibrium constants. Way back, we derived equilibrium constants from rate laws. So, we know both reactions occur at the same rate. They counteract one another!

(Video) Salt Hydrolysis: How to deduce Nature of Salt in Salt Hydrolysis

By Chemistry Guru

In this video, the Chemistry Guru explains salt hydrolysis. You'll learn why the conjugate ions of weak acids and bases change the pH. He'll also show you many examples. After watching this video, it will be easy to determine how a conjugate ion affects pH!