# Hydrolysis of Salt Solutions

### Dissociated Salts Form Acidic, Basic, or Neutral Solutions

The acidity of a salt solution can be predicted by determining how the dissociated cations and anions interact with water.

A salt is an ionic compound produced by a neutralization reaction of an acid and a base. According to the Brønsted-Lowry theory of acids and bases, a conjugate acid is a substance that accepts a proton from a base. A conjugate base is what remains after the acid donates a proton. Conjugate acids and bases are the products of acid-base reactions.

The pH scale shows how acidic or basic a solution is. The value of pH is calculated by taking the negative logarithm of the hydrogen ion concentration of a solution. The range of pH is typically between 1 and 14, with 1 being strongly acidic, 7 being neutral, and 14 being strongly basic. Values less than 1 and greater than 14 are also possible.

A strong acid completely dissociates into protons and anions in an aqueous solution. An example of a strong acid is hydrochloric acid (HCl). Weak acids partially dissociate and reach an equilibrium. Weak acids release fewer protons in an aqueous solution as a result. An example of a weak acid is acetic acid (CH3COOH).

Each salt consists of a cation and an anion. The cation is bound to the conjugate base of a strong acid, and the anion is bound to the conjugate acid of a strong base. Many salts dissociate in water. To predict whether the resulting ions of the dissociated salt will produce an acidic, basic, or neutral solution, consider the following:

• the effect of the cation on the pH of water
• the effect of the anion on the pH of water

These two effects are then compared, and the stronger effect determines the acidity/basicity of the solution.

For example, dissolving trisodium phosphate (Na3PO4) in water results in sodium ions (Na+) and phosphate ions (PO43–) in solution. Sodium is in group 1 of the periodic table, and ions of elements from groups 1 and 2 (the alkali and alkaline earth groups) and group 17 of the periodic table do not react with water in general. As a result, they do not produce either protons or hydroxide ions in solution and therefore do not affect pH. Phosphate (PO43–), however, is the conjugate base of the weak acid hydrogen phosphate (HPO42–). This means it will react with protons to form HPO42. Therefore, the phosphate ions will lower the proton (H+) concentration, making the solution slightly more basic. Conjugate bases of weak acids make solutions more basic.

If both the cation and the anion affect the pH, it is possible to determine which will have a greater effect by comparing their equilibrium constants, the constant that relates the amount of reactants and products in an equilibrium. For example, in ammonium phosphate ((NH4)3PO4), the cation is the ammonium ion (NH4+) rather than the sodium ion (Na+). Ammonium phosphate ((NH4)3PO4) dissociates in water:
$({\rm{NH}}_4)_3{\rm{PO}}_4\rightarrow3{\rm{NH}_4}^++{\rm{PO}_4}^{3-}$
Ammonium ion (NH4+) will donate a proton to water because it's the conjugate acid of a weak base, and phosphate (PO43–) will take a proton from water because it's the conjugate base of a weak acid. These two reactions are in competition with each other. To determine which reaction will have a greater effect, consider the acid equilibrium constant Ka and the base equilibrium constant Kb of the ions:
\begin{aligned}K_{\rm{a}}\;\text{of }\rm{NH}{_4}{^+}&=5.6\times10^{-10}\\K_{\rm{b}}\;{\text{of }\rm{PO}_4}^{3-}&=4.2\times10^{-13}\end{aligned}
In general the greater of the two numbers will determine the direction at which equilibrium will settle. In this case the Ka of NH4+ is greater than the Kb of PO43–, so the salt solution will be acidic.

### Metal Cations Create Weak Acid Solutions

Water molecules that are part of a metal-hydration complex can donate protons and form weak acid solutions.
Many common salts contain metal ions that are not in groups 1 or 2 of the periodic table. Transition metals are examples of such metals. When soluble salts containing these metals are dissolved in water, the solution becomes slightly acidic. The metal ions themselves do not donate protons. Instead, the polar water molecules form a hydration complex around these metals, and the molecules in this hydration complex can donate protons to the surrounding solution.

#### Hexaaquairon(III) Complex

Consider the coordination complex, a structure with a metal atom at the center and surrounded by other structures (ligands), that forms around an iron(III) ion. This iron cation has a 3+ charge and attracts lone pairs on the oxygen atoms around it. This shifting of electrons toward the large metal cation in the center of the complex results in less electron density around the hydrogen atoms at the edges of the complex. In turn, single hydrogen nuclei and protons will more readily leave the water molecules to form hydronium ions:
$[{\rm{Fe}}({\rm{H}}_2{\rm{O})_6}]^{3+}(aq)+{{\rm{H}}}_2{\rm{O}}(l)\rightleftarrows[{\rm{Fe}}({\rm{H}}_2{\rm{O}})_5({\rm{OH}})]^{2+}(aq)+{\rm{H}_3}{\rm{O}}^+(aq)$
The complex can lose up to two additional protons to the solution, each via its own equilibrium reaction.
$[{\rm{Fe(H}}_2{\rm{O)}_5}]^{2+}(aq)+{\rm{H}}_2{\rm{O}}(l)\rightleftarrows[{\rm{Fe(H}}_2{\rm{O)}}_4{\rm{(OH)}_2}]^{+}(aq)+3{\rm{H}}_3{\rm{O}}^{+}(aq)$
$[{\rm{Fe(H}}_2{\rm{O)}_4}]^{+}(aq)+{\rm{H}}_2{\rm{O}}(l)\rightleftarrows[{\rm{Fe(H}}_2{\rm{O)}}_3{\rm{(OH)}_3}](s)+3{\rm{H}}_3{\rm{O}}^{+}(aq)$
If the complex formed around the iron(II) cation instead, it would have less effect on the acidity because it would donate fewer protons. In general, cations of transition metals with 2+ or higher charges will contribute the greatest acidity to the solution.