Salts

An acid and base react to form a salt. Therefore when an acid or a base is "neutralized" a salt is formed.

Depending on the acids and bases the salt that is formed can be neutral, acidic, or basic.   This is all just a different language for what you have already learned.   For example, if formic acid is combined with sodium hydroxide, it generates a salt, sodium formate and water

\[\rm{HCOOH(aq) + NaOH(aq) \rightleftharpoons Na(HCOO)(aq) + H_2O(l)}\]

In this case, the salt is a basic salt since it contains the weak base, formate (HCOO^-) [and the spectator ion Na⁺].

In contrast, if a strong acid and a strong base are combined, like hydrochloric acid and potassium hydroxide you get a neutral salt, potassium chloride

\[\rm{HCl(aq) + KOH(aq) \rightleftharpoons KCl(aq) + H_2O(l)}\]

This is because both the strong acid and the strong base result in ions that are merely spectators.

Finally, it is possible to make acidic salts by neutralizing a weak base such as ammonia, NH₃ with a strong acid like HCl

\[\rm{NH_3(aq) + HCl(aq) \rightleftharpoons NH_4Cl(aq) + H_2O(l)}\]

Here the neutralization of NH₃ forms the ammonium ion, NH₄⁺ which is a weak acid.  Thus the ammonium chloride salt is acidic.

These salts can be isolated from solution by removing the water. This will leave behind the solid ionic compound.

What are some examples of basic salts? KCN, potassium cyanide. CN^- is the conjugate base of HCN. Na(HCOO), sodium formate. The formate ion, HCOO^- is the conjugate base of formic acid. What are some acidic salts? CH₃NH₃Cl, methylammonium chloride. Methylammonium is the conjugate acid of methylamine, CH₃NH_2.

There is a worksheet on identifying acid/base compounds on the worksheet page

Neutralization Reactions

It is critical in acid/base chemistry to first determine the majority of the chemical species that are in the solution.  This is particularly true when mixing two solutions together. Once you know the dominate species, you can then worry about solving the equilibrium problem to determine any small concentrations of interest (such as the pH).

As such, when mixing two solutions together, you need to first look at any neutralization reaction to figure out what will (for the most part) remain in solution. A neutralization reaction is the reaction of an acid and base.  In particular strong acids will always react in the presence of any base. Similarly strong bases will always react ion the presence of any acid.

  Let's look at an example of a reaction of formic acid and hydroxide.

\[\rm{HCOOH(aq) + NaOH(aq) \rightleftharpoons Na(HCOO)(aq) + H_2O(l)}\]

Which side does this equilibrium favor?  Note: This is the reverse reaction for the reaction of putting acetate (as weak base) into water. Therefore, this reaction strongly favors the righthand side of the reaction. We can assume this reaction goes 100% to the right. This reaction forms the salt sodium formate, Na(HCOO). We will see later that this salt is basic (since it forms a basic solution when placed in water). If we wanted to know the concentrations in a solution formed by mixing equal parts of formic acid and sodium hydroxide it would be the same as solving for the concentrations in a solution of sodium formate. This is because neutralizing formic acid with sodium hydroxide creates a solution of sodium formate.

To determine what is present after mixing any two acid/base solutions, we must realize that it is not possible to simultaneously have high concentrations of certain species.

We cannot have high concentrations of both H₃O⁺ and any base.

We cannot have high concentrations of both OH^- and any acid.

The simplest case is the "neutralization" reaction when you have exactly the same amount of acid and base. That is neither the acid nor the base is in excess. They will react until one or the other of them is gone from the solution. In the case of perfect "neutralization" they will both be gone and you'll end up with 100% products. Then you can work the equilibrium problem. Note: for weak acids and weak bases neutralization does not end up forming a solution with a neutral pH

This is the procedure you want to use for all neutralization reactions. First react the H₃O⁺ and any base (weak or strong). Alternatively you would react OH^- and any acid (weak or strong). Next use the limiting reagent to determine what reactants (if any) will remain in solution. When you are finished, you should have either no remaining H₃O⁺ or no remaining base .  Alternatively you should have no remaining OH^- or no remaining acid (or neither of either one). Then you can look at the solution and decide what type of solution you have. The remaining solution will fit into one of five categories:

1. It will either still have H₃O⁺, this is a strong acid solution. This is what happens when the strong acid is in excess.
2. It will still have OH^-, this is a strong base solution. This is what happens if the strong base is in excess.
3. It will have only the protonated base, this is a weak acid solution. This is what happens when a weak base and a strong acid are mixed in exact proportions.
4. It will have only the deprotonated form of the acid, this is a weak base solution. This is what happens when a weak acid and a strong base are mixed in exact proportions.
5. You will have both the protonated and deprotonated form of a conjugate pair.  This is a buffer solution. These solutions form by partially neutralizing either a weak acid or a weak base.

You already know how to solve for the equilibrium concentrations of the first four types of solution. We will soon cover the buffer situation.

Let's look at the neutralization reactions for a generic weak acid HA (BH⁺).  This would occur by mixing a weak acid solution with that of a strong base. This is the reaction we can assume will go 100% until either all of the HA is reacted or all of the OH^- is reacted.

\[\rm{HA(aq) + OH^-(aq) \rightleftharpoons A^-(aq) + H_2O(l)}\]

\[\rm{BH^+(aq) + OH^-(aq) \rightleftharpoons B(aq) + H_2O(l)}\]

The neutralization of a weak base, B (A^-),  with H₃O⁺ can also be assumed to go 100%.

\[\rm{B(aq) + H_3O^+(aq) \rightleftharpoons BH^+(aq) + H_2O(l)}\]

\[\rm{A^-(aq) + H_3O^+(aq) \rightleftharpoons HA(aq) + H_2O(l)}\]

Thinking About it

With all neutralization problems, it is important to think about the problems systematically. This is what is meant by "thinking like a chemist". Before leaping to a formula, you need to know what you have in solution and what reactions are taking place. Which concentrations are dominant and which ones are very small.

1.  Identify all the compounds (acids, bases, strong, weak, spectator ions, ...).

2. Figure out what is in solution.

3. Neutralize any strong acids or bases (if there are other bases/acids in solution).  This will require looking for the limiting reagent, reacting the compounds to completion, and identifying what remains in solution.

4.  After figuring out what is left in the solution, solve the equilibrium.  The remaining solution will either be a strong acid, weak acid, buffer, weak base, or strong base solution.

Remember, if you have any H₃O⁺ after neutralization you have a strong acid solution.  If you have any OH^- after neutralization you have a strong base solution.  If you have substantial amounts of both the protonated and deprotonated forms of a conjugate pair then you have a buffer.