ABC - Hydrolysis of Ions and Salts (Lesson)
Hydrolysis of Ions and Salts
Thus far, all of the substances that have been referred to, for the most part, have been neutral substances. What about charged ions? Can these substances act as acids? bases? The short answer is yes; however, in order to determine whether these substances behave as Bronsted-Lowry acids or bases, a few guidelines must be set forth to assist with the identification. Furthermore, ions cannot exist alone without a counterion present. For example, a solution cannot contain a carbonate ion alone without a corresponding positively charged cation. In the context of acids and bases, such compounds are referred to as salts. This lesson will focus on categorizing ions as either weak acids or weak bases as well as determining if salts themselves will raise or lower the pH of a solution.
Generally speaking, anions tend to form solutions that are either basic or neutral whereas cations typically form solutions that are either acidic or neutral. The example of ammonium chloride below can be used as an example. Notice first that because ammonium chloride is soluble in water it dissociates completely:
NH4Cl (s) ⟶ NH4+ (aq) + Cl- (aq)
As can be seen in the equilibrium below, NH4+ acts as a weak Bronsted-Lowry acid due to its ability to donate a proton to water:
NH4+ (aq) + H2O (l) ⇌ NH3 (aq) + H3O+ (aq)
The table below illustrates several examples of acid/base conjugate pairs and lists them in order of strength.
Anions as Weak Bases
Notice in the table above that many of the substances categorized as bases are anions. Furthermore, these anions are the conjugate bases of the corresponding acid in the left column. Generally speaking, the following two observations can be made:
- Anions that are conjugate bases of strong acids are pH neutral (i.e. form solutions with a pH = 7.00).
- e.g. Br- is the conjugate base of HBr and forms neutral solutions.
- Anions that are conjugate bases of weak acids are themselves weak base.
- e.g. Acetate (CH3CO2-) is the conjugate base of acetic aid (CH3CO2H) and forms basic solutions as evidenced by the production of hydroxide in the equilibrium shown below.
CH3CO2- (aq) + H2O (l) ⇌ CH3CO2H (aq) + OH- (aq)
Cations as Weak Bases
Generally speaking, cations can be subdivided into two categories for the purposes of discussing their acid properties. The first category is cations of strong bases such as KOH, NaOH, Mg(OH)2, etc. These cations interact with water, but their interaction is not strong enough to remove a proton from water. As a result, this means that solutions of all alkali and alkaline earth metal cations result in neutral solutions with a pH of 7.00.
The second category is the conjugate acids of weak bases which are themselves weak acids. This can be illustrated by examining ammonia (NH3) from the table above. The conjugate acid of NH3 would be NH4+ and results in an acidic solution when placed in water due to the production of hydronium ions as shown in the equilibrium below:
NH4+ (aq) + H2O (l) ⇌ NH3 (aq) + H3O+ (aq)
You Try It!
In the following self-assessment activity, identify the solutions as either neutral, basic, or acidic. Click on the plus sign to check your answer!
Determining the pH for a Solution of an Anion Acting Base
Since ions and salts can alter the pH of pure water, it is possible to determine the pH of these solutions. As an example, consider a 0.115 M solution of NaClO. What would be the pH of this solution? The first step is to understand what ions are present in water and whether either of those ions can alter the pH of pure water:
NaClO (s) ⟶ Na+ (aq) + ClO- (aq)
Analysis of the ions present after dissociation reveals the following:
- Na+ is an alkali metal and will not alter the pH of water.
- ClO- is the conjugate base of a weak acid and will therefore increase the pH of water (acts as a weak base).
The ClO- ion in water will establish the equilibrium shown below:
ClO- (aq) + H2O (l) ⇌ HClO (aq) + OH- (aq)
The next step would be to write the equilibrium expression for this equilibrium; however, a unique problem is often encountered. A table of Ka values is readily available for a wide variety of weak acids, but a corresponding table for weak bases is not nearly as exhaustive. Since ClO- is a weak base, a Kb value for the equilibrium is necessary; how can this value be determined?
The relationship between Ka, Kb, and Kw can be utilized to determine this value. After determining that the Ka for HClO is 2.9 x 10-8 the following calculations are used to determine the corresponding Kb value for ClO-:
Ka x Kb = Kw = 1.0 x 10-14
(2.9 x 10-8)Kb = 1.0 x 10-14
Kb = 3.4 x 10-7
Returning to the example above, it can be seen that all the required information is now available. An equilibrium consisting of 0.115 M of the weak conjugate base hypochlorite (ClO-) is present in water and the Kb value for this weak base is 3.4 x 10-7. From this point, an ICE table can be utilized and the problem begins to look very similar to previously worked examples:
NH3 (aq) | H2O (l) | NH4+ (aq) | OH- (aq) | |
---|---|---|---|---|
Initial | 0.115 M | -- | 0 | 0 |
Change | - x | -- | + x | + x |
Equilibrium | 0.115 - x | -- | x | x |
The equilibrium expression for this scenario would look like the following:
Kb=[HOCl][OH−][OCl−]=3.4×10−7
Substituting the values from the Equilibrium line of the table into this expression yields the following:
Kb=[HOCl][OH−][OCl−]=(x)(x)0.115−x=3.4×10−7
As before, the x in the denominator is ignored to avoid a quadratic equation:
(x)(x)0.115=3.4×10−7
x2=(0.115)(3.4×10−7)=3.9×10−8
x=√3.9×10−8=2.0×10−4
As is the case for all other weak base equilibria, the value for x corresponds to the [OH-] at equilibrium and requires the calculation of pOH before pH:
pOH = -log[OH-]
pOH = -log(2.0 x 10-4)
pOH = 3.70
pH + pOH = 14.00
pH = 10.30
You Try It!
In the following self-assessment activity, calculate the pH value. Click on the plus sign to check your answer!
[CC BY 4.0] UNLESS OTHERWISE NOTED | IMAGES: LICENSED AND USED ACCORDING TO TERMS OF SUBSCRIPTION