Weak base–strong acid reactions | Acids and bases | AP Chemistry | Khan Academy

Ammonia is an example of a weak base, and hydrochloric acid is an example of a strong acid. Ammonia reacts with hydrochloric acid to form an aqueous solution of ammonium chloride. Because this is an acid-base neutralization reaction, there's only a single arrow going to the right, indicating the reaction goes to completion.

Next, let's write the overall or complete ionic equation. Let's start with ammonia. Ammonia is a weak base, and a weak base is only partly ionized in aqueous solution. Therefore, since a weak base is only partly ionized, we're not going to show this as an ion; we're simply going to write NH3 in our equation.

However, for hydrochloric acid, hydrochloric acid is a strong acid, and strong acids ionize 100 percent. Therefore, in aqueous solution, we need to show this as the ions: H⁺ and Cl⁻. Ammonium chloride is a soluble salt; therefore, in aqueous solution, we show it as the ions. Ammonium chloride consists of the ammonium ion NH₄⁺ and the chloride anion Cl⁻.

To save some time, I've drawn in the aqueous subscripts and also put in the reaction arrow and a plus sign. So this represents the overall or the complete ionic equation. We can use the complete ionic equation to find the net ionic equation for this weak base strong acid reaction. To do that, we first need to identify these spectator ions. Remember, these are the ions that do not take part in the chemical reaction.

Since there's a chloride anion on the left side and on the right side, the chloride anion is the spectator ion for this reaction. Once we take out our spectator ion, we're left with our net ionic equation, which is aqueous ammonia plus H⁺ yields NH₄⁺.

So this is one way to write our net ionic equation. However, remember that H⁺ and H₃O⁺ are used interchangeably in chemistry. Therefore, another way to write the net ionic equation is to show aqueous ammonia plus the hydronium ion H₃O⁺ yields the ammonium ion NH₄⁺ plus water.

Now that we have our net ionic equation, we're going to consider three different situations. In the first situation, we have equal moles of our weak base and strong acid. Looking at our net ionic equation, the mole ratio of ammonia to hydronium ion is one to one. Therefore, if we have equal moles of our weak base and strong acid, the weak base and strong acid will completely neutralize each other and produce the ammonium ion NH₄⁺.

So, if our goal is to figure out the pH of the resulting solution, we don't need to consider the weak base or this strong acid. We need to think about the ammonium cation in aqueous solution. In solution, the ammonium cation acts as a weak acid and donates a proton to water to form the hydronium ion H₃O⁺ and aqueous ammonia. The ammonium cation NH₄⁺ is a weak acid; therefore, the Kₐ value is less than one.

Since Kₐ is less than one at equilibrium, there are mostly reactants and not very many products. However, the concentration of hydronium ions in solution is increased, and therefore the resulting solution will be acidic. So, the resulting solution will be slightly acidic, and at 25 degrees Celsius, the pH of the solution will be less than seven. If we wanted to calculate the actual pH, we would treat this like a weak acid equilibrium problem.

Also, it's important to emphasize that the hydronium ions that gave the resulting solution a pH less than seven came from the reaction of the ammonium cation with water. The hydronium ions did not come from the strong acid; all of those hydronium ions were used up in the acid-base neutralization reaction.

For the second situation, we have more of the weak base than the strong acid; therefore, we have the weak base in excess. Because the mole ratio of the weak base to the strong acid is one to one, if we have more of the weak base than the strong acid, all of the strong acid will be used up.

So when the reaction goes to completion, we'll have ammonium cations in solution, and we'll also have some leftover ammonia. After the neutralization reaction is complete and all the H₃O⁺ is used up, we'll have some leftover ammonia. That ammonia will react with water to form hydroxide anions and NH₄⁺.

Because the concentration of hydroxide ions in solution has increased, at 25 degrees Celsius, the resulting solution will be basic, and the pH will be greater than seven. If we wanted to calculate the actual pH, we would treat this like a weak base equilibria problem.

However, we have two sources for the ammonium cation. One source is from ammonia reacting with water to form NH₄⁺, and the other source came from the neutralization reaction. So, actually, this would be a common ion effect problem.

The other way to calculate the pH of this solution is to realize that ammonium NH₄⁺ is a weak acid and ammonia NH₃ is its conjugate base. Therefore, if we have similar amounts of a weak acid and its conjugate base, we have a buffer solution, and we could calculate the pH using the Henderson-Hasselbalch equation.

For our third situation, let's say we have the strong acid in excess. Since the mole ratio of weak base to strong acid is one to one, if we have more of the strong acid than the weak base, all of the weak base will be used up, and we'll have some strong acid in excess.

Therefore, there'll be a concentration of hydronium ions in solution, which would make the resulting solution acidic. So, at 25 degrees Celsius, the pH would be less than 7. We could calculate the actual pH of the resulting solution by doing a strong acid pH calculation problem.

And while it's true that the ammonium cation can function as a weak acid and also increase the concentration of hydronium ions, it's such a small increase compared to the hydronium ions we have in solution from our strong acid that we don't need to worry about the contribution of the ammonium cations. We can just treat this like a strong acid pH calculation problem.