ACID,BASE,SALT

Practice Questions
Why does a salt containing a cation from a strong base and an anion from a weak acid form a basic solution?
Why does a salt containing a cation from a weak base and an anion from a strong acid form an acidic solution? 
How do the Ka or Kb values help determine whether a weak acid or weak base will be the dominant driving force of a reaction?
The answers to these questions can be found in the attached files section at the bottom of the page.

Summary of Acidic and Basic Salts
As we have discussed, salts can form acidic or basic solutions if their cations and/or anions are hydrolyzable (able to react in water). Basic salts form from the neutralization of a strong base and a weak acid; for instance, the reaction of sodium hydroxide (a strong base) with acetic acid (a weak acid) will yield water and sodium acetate. Sodium acetate is a basic salt; the acetate ion is capable of deprotonating water, thereby raising the solution’s pH.
Acid salts are the converse of basic salts; they are formed in the neutralization reaction between a strong acid and a weak base. The conjugate acid of the weak base makes the salt acidic. For instance, in the reaction of hydrochloric acid (a strong acid) with ammonia (a weak base), water is formed, along with ammonium chloride. The ammonium ion contains a hydrolyzable proton, which makes it an acid salt.

Practice Questions
Why does a salt containing a cation from a strong base and an anion from a weak acid form a basic solution?
Why does a salt containing a cation from a weak base and an anion from a strong acid form an acidic solution? 
How do the Ka or Kb values help determine whether a weak acid or weak base will be the dominant driving force of a reaction?
The answers to these questions can be found in the attached files section at the bottom of the page.

Summary of Acidic and Basic Salts
As we have discussed, salts can form acidic or basic solutions if their cations and/or anions are hydrolyzable (able to react in water). Basic salts form from the neutralization of a strong base and a weak acid; for instance, the reaction of sodium hydroxide (a strong base) with acetic acid (a weak acid) will yield water and sodium acetate. Sodium acetate is a basic salt; the acetate ion is capable of deprotonating water, thereby raising the solution’s pH.
Acid salts are the converse of basic salts; they are formed in the neutralization reaction between a strong acid and a weak base. The conjugate acid of the weak base makes the salt acidic. For instance, in the reaction of hydrochloric acid (a strong acid) with ammonia (a weak base), water is formed, along with ammonium chloride. The ammonium ion contains a hydrolyzable proton, which makes it an acid salt.

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Salts in Which Both Ions Hydrolyze
The following is a more complicated scenario in which a salt contains a cation and an anion, both of which are capable of participating in hydrolysis. A good example of such a salt is ammonium bicarbonate, NH4HCO3; like all ammonium salts, it is highly soluble, and its dissociation reaction in water is as follows:
[latex]NH_4CO_3(s)\rightarrow NH_4^+(aq)+HCO_3^-(aq)[/latex]
However, as we have already discussed, the ammonium ion acts as a weak acid in solution, while the bicarbonate ion acts as a weak base. The reactions are as follows:

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[latex]NH_4^+(aq)+H_2O(l)\rightleftharpoons H_3O^+(aq)+NH_3(aq)\quad\quad K_a=5.6\times10^{-10}[/latex]
[latex]HCO_3^-(aq)+H_2O(l)\rightleftharpoons H_2CO_3(aq)+OH^-(aq)\quad\quad K_b=2.4\times 10^{-8}[/latex]
Because both ions can hydrolyze, will a solution of ammonium bicarbonate be acidic or basic? We can determine the answer by comparing Ka and Kb values for each ion. In this case, the value of Kb for bicarbonate is greater than the value of Ka for ammonium. Therefore, bicarbonate is a slightly more alkaline than ammonium is acidic, and a solution of ammonium bicarbonate in pure water will be slightly basic (pH > 7.0). In summary, when a salt contains two ions that hydrolyze, compare their Ka and Kb values:
If Ka > Kb, the solution will be slightly acidic.
If Kb > Ka, the solution will be slightly basic.
Hydrolysis of saltsThis video examines the hydrolysis of an acid salt, a basic salt, and a salt in which both ions hydrolyze.

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Salts in Which Both Ions Hydrolyze
The following is a more complicated scenario in which a salt contains a cation and an anion, both of which are capable of participating in hydrolysis. A good example of such a salt is ammonium bicarbonate, NH4HCO3; like all ammonium salts, it is highly soluble, and its dissociation reaction in water is as follows:
[latex]NH_4CO_3(s)\rightarrow NH_4^+(aq)+HCO_3^-(aq)[/latex]
However, as we have already discussed, the ammonium ion acts as a weak acid in solution, while the bicarbonate ion acts as a weak base. The reactions are as follows:

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[latex]NH_4^+(aq)+H_2O(l)\rightleftharpoons H_3O^+(aq)+NH_3(aq)\quad\quad K_a=5.6\times10^{-10}[/latex]
[latex]HCO_3^-(aq)+H_2O(l)\rightleftharpoons H_2CO_3(aq)+OH^-(aq)\quad\quad K_b=2.4\times 10^{-8}[/latex]
Because both ions can hydrolyze, will a solution of ammonium bicarbonate be acidic or basic? We can determine the answer by comparing Ka and Kb values for each ion. In this case, the value of Kb for bicarbonate is greater than the value of Ka for ammonium. Therefore, bicarbonate is a slightly more alkaline than ammonium is acidic, and a solution of ammonium bicarbonate in pure water will be slightly basic (pH > 7.0). In summary, when a salt contains two ions that hydrolyze, compare their Ka and Kb values:
If Ka > Kb, the solution will be slightly acidic.
If Kb > Ka, the solution will be slightly basic.
Hydrolysis of saltsThis video examines the hydrolysis of an acid salt, a basic salt, and a salt in which both ions hydrolyze.

Depending on the acid–base properties of its component ions, however, a salt can dissolve in water to produce a neutral solution, a basic solution, oran acidic solution.When a salt such as NaClNaCl dissolves in water, it produces Na+(aq)Na(aq)+ and Cl−(aq)
Cl(aq)− ions. Using a Lewis approach, the Na+Na+ ion can be viewed as an acid because it is an electron pair acceptor, although its low charge and relatively large radius make it a very weak acid. The Cl−Cl− ion is the conjugate base of the strong acid HCl
HCl, so it has essentially no basic character. Consequently, dissolving NaCl
NaCl in water has no effect on the pHpH of a solution, and the solution remains neutral.
Now let's compare this behavior to the behavior of aqueous solutions of potassium cyanide and sodium acetate. Again, the cations (K+K+ and Na+Na+) have essentially no acidic character, but the anions (CN−CN− and CH3CO−2
CH3CO2−) are weak bases that can react with water because they are the conjugate bases of the weak acids HCN
HCN and acetic acid, respectively.

Depending on the acid–base properties of its component ions, however, a salt can dissolve in water to produce a neutral solution, a basic solution, oran acidic solution.When a salt such as NaClNaCl dissolves in water, it produces Na+(aq)Na(aq)+ and Cl−(aq)
Cl(aq)− ions. Using a Lewis approach, the Na+Na+ ion can be viewed as an acid because it is an electron pair acceptor, although its low charge and relatively large radius make it a very weak acid. The Cl−Cl− ion is the conjugate base of the strong acid HCl
HCl, so it has essentially no basic character. Consequently, dissolving NaCl
NaCl in water has no effect on the pHpH of a solution, and the solution remains neutral.
Now let's compare this behavior to the behavior of aqueous solutions of potassium cyanide and sodium acetate. Again, the cations (K+K+ and Na+Na+) have essentially no acidic character, but the anions (CN−CN− and CH3CO−2
CH3CO2−) are weak bases that can react with water because they are the conjugate bases of the weak acids HCN
HCN and acetic acid, respectively.

Neither reaction proceeds very far to the right as written because the formation of the weaker acid–base pair is favored. Both HCN
HCN and acetic acid are stronger acids than water, and hydroxide is a stronger base than either acetate or cyanide, so in both cases, the equilibrium lies to the left. Nonetheless, each of these reactions generates enough hydroxide ions to produce a basic solution. For example, the pH
pH of a 0.1 M solution of sodium acetate or potassium cyanide at 25°C is 8.8 or 11.1, respectively. From Table 16.9.1
16.9.1 and Figure 16.9.116.9.1, we can see that CN−CN− is a stronger base (pKb=4.79pKb=4.79) than acetate (pKb=9.24pKb=9.24), which is consistent with KCN
KCN producing a more basic solution than sodium acetate at the same concentration.
In contrast, the conjugate acid of a weak base should be a weak acid (Equation 16.9.4
16.9.4). For example, ammonium chloride and pyridinium chloride are salts produced by reacting ammonia and pyridine, respectively, with HCl
HCl. As you already know, the chloride ion is such a weak base that it does not react with water. In contrast, the cations of the two salts are weak acids that react with water as follows:

Neither reaction proceeds very far to the right as written because the formation of the weaker acid–base pair is favored. Both HCN
HCN and acetic acid are stronger acids than water, and hydroxide is a stronger base than either acetate or cyanide, so in both cases, the equilibrium lies to the left. Nonetheless, each of these reactions generates enough hydroxide ions to produce a basic solution. For example, the pH
pH of a 0.1 M solution of sodium acetate or potassium cyanide at 25°C is 8.8 or 11.1, respectively. From Table 16.9.1
16.9.1 and Figure 16.9.116.9.1, we can see that CN−CN− is a stronger base (pKb=4.79pKb=4.79) than acetate (pKb=9.24pKb=9.24), which is consistent with KCN
KCN producing a more basic solution than sodium acetate at the same concentration.
In contrast, the conjugate acid of a weak base should be a weak acid (Equation 16.9.4
16.9.4). For example, ammonium chloride and pyridinium chloride are salts produced by reacting ammonia and pyridine, respectively, with HCl
HCl. As you already know, the chloride ion is such a weak base that it does not react with water. In contrast, the cations of the two salts are weak acids that react with water as follows:

Equation 16.9.4
16.9.4 indicates that H3O+H3O+ is a stronger acid than either NH+4NH4+ or C5H5NH
+C5H5NH+, and conversely, ammonia and pyridine are both stronger bases than water. The equilibrium will therefore lie far to the left in both cases, favoring the weaker acid–base pair. The H3O+H3O+ concentration produced by the reactions is great enough, however, to decrease the pHpH of the solution significantly: the pHpH of a 0.10 M solution of ammonium chloride or pyridinium chloride at 25°C is 5.13 or 3.12, respectively. This is consistent with the information shown in Figure 16.2, indicating that the pyridinium ion is more acidic than the ammonium ionWhat happens with aqueous solutions of a salt such as ammonium acetate, where both the cation and the anion can react separately with water to produce an acid and a base, respectively? According to Figure 16.10, theammonium ion will lower the pHpH, while according to Equation 16.9.516.9.5, the acetate ion will raise the pHpH. This particular case is unusual, in that the cation is as strong an acid as the anion is a base (pKa ≈ pKb). Consequently, the two effects cancel, and the solution remains neutral. With salts in which the cation is a stronger acid than the anion is a base, the final solution has a pH
pH < 7.00. Conversely, if the cation is a weaker acid than the anion is a base, the final solution has a pH
pH > 7.00.