2.7.1. Aqueous Solutions of Various Salts Can Be Neutral, Basic, or Acidic. How Cations and Anions Produced on Dissociation of Salts Interact with Water. On addition of an acid such as HCl or H₂SO₄ to water, the original pH value of 7 goes down because the concentration of the H⁺ goes up. On addition of a base such as NaOH or Ca(OH)₂ to water the pH goes up because the concentration of the H⁺ decreases due to an increase in the concentration of the OH^-. Will the neutral pH of 7 of pure water change on addition of a soluble salt? It might and it might not.

Some salts such as NaCl, Ba(NO₃)₂, and K₂SO₄ have no effect on the pH of water. However, if we dissolve baking soda (NaHCO₃) in water, the resultant solution will be slightly basic (pH ≈ 8). On dissolution of Na₂CO₃ in water, the pH goes to 11 — no kidding! — thereby indicating that a solution of Na₂CO₃ is quite alkaline. On the contrary, aqueous solutions of ZnCl₂ are acidic, the pH being as low as 1 at a high concentration of zinc chloride. Why? Because some ions produced on dissociation of a salt can and do interact with water molecules in an interesting way.

Let us consider an aqueous solution of ZnCl₂ (Figure 2-87). Being a salt, ZnCl₂ readily dissociates in water into the ions Zn²⁺ and Cl^-. Water dissociates too, although to a much smaller degree, to give rise to H⁺ and OH^-. The blue and red ellipses in Figure 2-87 show interactions of the oppositely charged ions derived from ZnCl₂ and water.

Figure 2-87. Interactions between ZnCl₂ and water.


Circled in red in Figure 2-87, are the Cl^- (from ZnCl₂) and H⁺ (from H₂O). When they collide, a molecule of HCl is formed, which is a strong electrolyte that immediately dissociates back into the Cl^- and H⁺. Consequently, the Cl^- anion cannot and does not influence the concentrations of H⁺ and OH^- in water.

Circled in blue in Figure 2-87, are the Zn²⁺ (from ZnCl₂) and OH^- (from H₂O). When they come together, ZnOH⁺ is produced. What is ZnOH⁺ and will this cation dissociate back into the Zn²⁺ and OH^-? The ion ZnOH⁺ could be formed on dissociation of one OH^- from Zn(OH)₂. As Zn(OH)₂ is a weak electrolyte, ZnOH⁺ is also a weak electrolyte, not readily dissociating to Zn²⁺ and OH^-. Consequently, the Zn²⁺ cation in our solution serves as a trap for the OH^- ions. As the concentration of OH^- goes down, the solution becomes less alkaline = more acidic, precisely as can be observed experimentally by a pH measurement.

Next, we consider Na₂CO₃, sodium carbonate, also known as soda ash and washing soda. To understand why solutions of Na₂CO₃ are alkaline, we will use the same type of analysis as in the previous case of ZnCl₂. On dissolution in water, Na₂CO₃ dissociates to Na⁺ and CO₃^2- (Figure 2-88). Water dissociates (to a much lesser extent) to give small quantities of H⁺ and OH^-. A collision of Na⁺ and OH^- produces NaOH that, being a strong electrolyte, dissociates back right away (Figure 2-88, blue ellipse). In contrast, when the H⁺ and CO₃^2- meet, they produce hydrocarbonate, HCO₃^-, a weak electrolyte that does not dissociate readily. We know that HCO₃^- dissociates poorly from the fact that H₂CO₃ is a very weak electrolyte and, consequently, the product of its first dissociation, HCO₃^- is even a weaker acid. Therefore, the CO₃^2- serves as a trap for the H⁺, thereby making the solution less acidic = more basic.

Figure 2-88. Interactions between Na₂CO₃ and water.


Finally, if we apply the same analysis to an aqueous solution of NaCl (Figure 2-89), we will see that neither Na⁺ can trap OH^-, nor Cl^- can sequester H⁺ because both NaOH and HCl are strong electrolytes. Consequently, the pH of an aqueous solution of NaCl is 7, the same as of pure water.

Figure 2-89. The pH of aqueous NaCl is neutral (7) because neither Na⁺ sequesters OH^-, nor Cl^- can trap H⁺.


We say that NaCl is not hydrolyzed in solution (Figure 2-89), whereas Na₂CO₃ (Figure 2-88) and ZnCl₂ (Figure 2-87) are. Hydrolysis is a chemical reaction where water is involved as a reagent. Hydrolysis reactions can be reversible, like those of Na₂CO₃ or ZnCl₂ considered above, but can also be irreversible. Irreversible hydrolysis means that the reaction with water goes to completion, as in the following example.

Al₂S₃ + 6 H₂O = 2 Al(OH)₃↓ + 3 H₂S↑

There is a simple rule for quickly predicting if an aqueous solution of a given salt would be acidic or alkaline. This rule is illustrated in Figure 2-90.

Figure 2-90. A rule of thumb to predict if a solution of a salt would be acidic or alkaline.


First, identify the parent acid and base of a given salt. If the parent acid is stronger an electrolyte than the parent base, an aqueous solution of the salt will be acidic (acid "defeats" base). If the parent base is stronger an electrolyte than the parent acid, the solution will be alkaline (base "defeats" acid). If both the parent acid and the base are strong electrolytes, neither one can "defeat" the other, the salt will not be hydrolyzed and its aqueous solution will be neutral (pH = 7).

As an example, the parent acid and base for KCl are HCl and KOH. As both are strong electrolytes, KCl is not hydrolyzed and aqueous solutions of KCl are neutral. For CuSO₄, the parent acid and base are H₂SO₄ and Cu(OH)₂. Since the acid is strong and the base is weak, solutions of CuSO₄ are acidic. Sodium silicate (Na₂SiO₃) is formed by NaOH, a strong base, and metasilicic acid, H₂SiO₃, a weak acid. Therefore, an aqueous solution of Na₂SiO₃ is alkaline.

Take warning of the often tricky outcome of hydrolysis of hydro salts. A hydro salt is produced if not all of the protons on the molecules of a dibasic or tribasic acid have been displaced by a metal ion(s). Two examples of hydro salts are presented in Figure 2-91.

Figure 2-91. Empirical and structural formulas of NaHSO₄ and NaHCO₃.


Some people make a big mistake thinking that all hydro salts are acidic because of the remaining H atom(s) on the acid remainder. While aqueous solutions of some hydro salts are indeed acidic, those of some other hydro salts are alkaline! Every case should be carefully analyzed to avoid making a mistake.

Aqueous solutions of NaHSO₄ are acidic. This salt is formed by a strong acid and a strong base, so should a solution of NaHSO₄ not be neutral? True, neither OH^- nor H⁺ can be captured by Na⁺ and HSO₄^- produced on dissociation of NaHSO₄ (Figure 1-92). However, the HSO₄^- can and does dissociate to produce H⁺ and SO₄^2-. The dissociation of HSO₄^- is not as efficient as that of H₂SO₄ because polybasic acids always dissociate step-wise and each consecutive step is considerably less facile than the previous one. Nonetheless, HSO₄^- is a medium strength acid, making aqueous solutions of NaHSO₄ quite acidic (Figure 2-92). Note that the acidity of NaHSO₄ solutions does not originate from hydrolysis but rather from dissociation of the HSO₄^- anion into H⁺ and SO₄^2-.

Figure 2-92. NaHSO₄ is not hydrolyzed. The acidity of aqueous NaHSO₄ is due to dissociation of HSO₄^- to give H⁺.


Do not be deceived by the similarity of formulas and structures of NaHSO₄ and NaHCO₃ (Figure 2-91). The case of NaHCO₃ (baking soda) is very different from that of NaHSO₄. Solutions of NaHCO₃ are not acidic but mildly alkaline. (This explains the long-time use of baking soda as a simple remedy for heartburn.) Figure 2-93 accounts for the basicity of a NaHCO₃ solution. Unlike NaHSO₄, NaHCO₃ does undergo hydrolysis because when HCO₃^- and H⁺ collide, H₂CO₃ is formed, which is a weak electrolyte. Serving as a trap for H⁺, the HCO₃^- is responsible for the alkalinity of baking soda solutions. Unlike HSO₄^-, HCO₃^- is an extremely poor electrolyte (even weaker than H₂CO₃) that cannot produce H⁺ in any significant quantities (Figure 2-93).

Figure 2-93. Solutions of NaHCO₃ are mildly alkaline due to hydrolysis. Unlike HSO₄^- (Figure 2-92), HCO₃^- is an extremely weak electrolyte that produces H⁺ only in negligible quantities.


2.7.2. Exercises.

1. By looking at the formulas below determine if an aqueous solution of each salt is acidic, alkaline, or neutral: (a) KBr; (b) K₂CO₃; (c) BaI₂; (d) CuCl₂; (e) BaSO₄; (f) Al₂(SO₄)₃; (g) Ba(NO₃)₂; (h) AgCl; (i) AgNO₃; (j) K₂SiO₃. Answer

2. Using the ion collision analysis (as shown in Figures 2-87, 2-88, and 2-89), support your answers to questions in Exercise 1 above.

3. Would you expect an aqueous solution of Ca(NO₃)₂ to be neutral, acidic, or alkaline? Provide arguments to support your view. Answer