Common Ion Effect

With the addition of soluble compounds to the solution leads to the decrease in the solubility of the ionic precipitate due to the ion common with the precipitate, this effect is called the common-ion effect. This effect is due to Le Chatelier’s principle for the ionic association/dissociation reaction for equilibria. Generally, this effect is seen on the solubility of salts and weak electrolytes. The common ion effect is generally used to describe the effect on equilibrium. In heterogeneous equilibria, the solubility products Ksp’s are constants between two different phases. Salts present in the system; they ionize in the solution. Salt containing cation or anion contributes to the concentration of the common ion in the solution.

For example: a solution of acetic acid dissociation equilibrium is represented as:

       CH3COOH (aq) ↔ H+(aq) + CH3COO–(aq)

Or   HAc (aq) ↔ H+ (aq) + Ac– (aq)

Ka = [H+] [Ac–] / [ HAc]

Concentration of H+ ions decreases in addition to CH3COO– ions. On adding H+ ions from an external source, the equilibrium shifts in the undissociated CH3COOH i.e., in the direction of reducing H+ ions. This phenomenon states a common ion effect.

Thus, it is a phenomenon of Le Chatelier’s principle. According to Le Chatelier’s principle, change in any of the factors that determine the equilibrium conditions of a system will cause the system to change so as to reduce the effect of the change. This is applicable to all physical and chemical equilibria.

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Common Ions

When NaCl and KCl dissolve in the same solution, then the Cl– ions are common to both salts. In a solution containing NaCl and KCl, common ions are Cl– ions. 

                           NaCl ↔ Na+ + Cl-

                           KCl ↔ K+ + Cl–

                           CaCl2 ↔ Ca2+ + 2Cl–

                           AlCl3 ↔ Al3+ + 3Cl–

                           AgCl ↔ Ag+ + Cl–

For example- When AlCl3 is dissolved into a solution already containing NaCl, the Cl– ions come from the ionization of both AlCl3 and NaCl. Thus, [Cl–] differs from [Al3+]. 

Example of Common ion effect

1. PbCl2 (s) ↔ Pb2+ (aq) + 2Cl–

    Lead chloride is partially soluble in water giving the above equation. The product contains more chloride ions than lead ions. On adding NaCl, the system will have Cl– ion as a common ion.

NaCl ionizes as-

NaCl (s) ↔ Na+ (aq) + Cl– (aq)

Due to common ion effect, the PbCl2 solubility decreases in the reaction on addition of Cl– ion that shifts the PbCl2 equilibrium to counteract addition of Cl– ion which results into removal of some Cl– ion and forming PbCl2 in the reaction.

Hydrolysis of Salts and the pH of their solutions

Reaction between acid and base results in the formation of salt while undergoing the ionization process in water. On the basis of the nature of salts, the cations/anions formed exist as hydrated ions in aqueous solutions or interact with water to reform acids/bases. The latter process is called hydrolysis. The pH of the solution changes by this interaction. Cations such as Na+, K+, Ca2+, Ba2+, etc. are strong bases whereas anions such as Cl–, Br–, NO3–, ClO4–, etc. constitute strong acid, and they do not get easily hydrated as result of which solutions of salts formed from strong acids and bases having neutral pH 7. Also, other categories of salts do undergo hydrolysis.

Following are the type of hydrolysis of salts: –

1. Salt of weak acid and strong base e.g., CH3COONa
2. Salts of strong acid and weak base e.g., NH4Cl, and
3. Salts of weak acid and weak base, e.g., CH3COONH4

In the 1st type, CH3COONa is a salt of weak acid, CH3COOH and strong base, NaOH.

                              CH3COONa (aq) → CH3COO– (aq) + Na+ (aq)

 Acetate ions thus formed undergo hydrolysis giving acetic acid and OH– ions.

                              CH3COO– (aq) + H2O (l) ↔ CH3COOH (aq) + OH– (aq)

With the increase in concentration of OH– ion, the solution turns into alkaline whereas the acetic acid remains unionized in the solution. The pH of the solution is more than 7.

In the 2nd type, NH4Cl is salt of weak base, NH4OH and strong acid, HCl, hydrolyses completely in the water.

                          NH4Cl (aq) → NH4+ (aq) + Cl– (aq)

NH4+, an ammonium ion undergoes hydrolysis with water to form NH4OH and H+ ions.

                         NH4+ (aq) + H2O (l) ↔ NH4OH (aq) + H+ (aq)

Increased concentration of H+ ion in solution makes the solution acidic while NH4OH, a weak base, remains unionized in the solution.

In the 3rd type, CH3COONH4 salt is formed from weak acid and weak base. The process of hydrolysis is as follows: 

                         CH3COO– + NH4+ + H2O ↔ CH3COO + NH4OH

 CH3COOH and NH4OH remain partially dissociated in the solution. The dissociated ions are as follows:

                         CH3COOH ↔ CH3COO– + H+

                         NH4OH ↔ NH4+ + OH–

                         H2O ↔ H+ + OH–

The degree of hydrolysis is independent of concentration of solution and pH of solution is determined by their pK values.

   pH = 7+ ½ (pKa – pKb)

pH>7, positive difference

pH<7, negative difference

Example 1- The pKa value of acetic acid and pKb of ammonium hydroxide are 4.76 and 4.75 respectively. Calculate the pH of ammonium acetate solution.

Solution-       pH = 7+ ½ [pKa – pKb]

                          = 7+½ [4.76-4.75]

                          = 7+ ½ [0.01]

                          = 7+0.005

                          = 7.005

Buffer Solutions

Body fluids like urine or blood have definite pH and any change in the pH can be seen in the malfunctioning of the body. Optimum pH is required for proper functioning of the body. Buffer solution defines that the solutions that resist change in pH on dilution or with the addition of small amounts of acid or alkali are called buffer solutions. By knowing the value of pKa of acid or pKb of base and desirable ratio of salt and acid or salt and base, buffer solutions of required pH are prepared.

Characteristics of buffer solution:

1. Definite pH
2. pH does not change on dilution
3. Addition of a small amount of a strong acid or a base does not change the pH of the solution.

Types of buffers: – 

1. Acidic Buffer- mixture of weak acid and its salt of a strong base in water is called an acidic buffer. The pH of the acidic buffer is <7.

Example- CH3COOH + CH3COONa and HCOOH + HCOONa

1. Basic Buffer- mixture of weak base and its salt of strong acid in water is called basic buffer. pH of the basic buffer is >7.

Example- NH4OH + NH4Cl and NH4OH + NH4NO3

pH of a Buffer Solution:

pH of a buffer can be calculated by the Henderson-Hasselbalch equation.

Let an acidic buffer consist of weak acid HA and its salt NaA.

                       HA + H2O ↔ H3O+ + A– (slight ionization)

                       NaA → Na+ (aq) + A– (aq) (complete ionization)

                              Ka = [H3O+] [A–]/ [HA]

Salt NaA is completely dissociated. So, [A–] = [NaA] = [Salt] and 

HA is non dissociated [HA] = [Acid] 

Therefore,

                   ∴ Ka = [H3O+] [Salt] / [Acid]

                   ∴ [H3O+] = Ka [Acid] / [Salt]            

                 Taking log of both side to base 10

      log10 [H3O+] = log10Ka  +  log10[Acid] / [Salt]

                 Multiplying both sides by -1

-log10[H3O+] = -log10Ka  –  log10 [Acid] / [Salt]

       ∴pH = pKa  +  log10 [Salt] / [Acid]

  Similarly, for the basic buffer- 

                      pOH = pKa  +  log10 [Salt] / [Base]

Applications of buffer solution:

1. Buffers are used for qualitative and quantitative analysis.
2. Adjust pH of the solutions by the calorimetric method.
3. Used in various processes like electrodeposition of metals, tanning of leather, brewing of alcohols, etc.
4. Used as stabilizers and preservatives e.g. sodium citrate stabilizes penicillin preparations.
5. Buffers maintain pH of the soil.

Common Ion Effect with Weak Acids and Bases

Addition of common ions prevents ionization of weak acid or weak base. Ionization of a weak acid is suppressed by the common ion by adding more ions that are a product of equilibrium.

Example2: The common ion effect of H3O+ on the ionization of acetic acid

Ionization of the weak base is suppressed by the common ion effect by addition of more ions that are a product of equilibrium.

Example3: The common ion effect of OH– on ionization of ammonia.

Reactions shift towards the left side on adding a common ion of hydroxide that leads to decrease the stress according to Le Chatelier’s Principle forming more reactants.

Common Ion Effect on Solubility of Ionic Salts

According to the Le Chatelier’s Principle, any increase in the concentration of any one of ions, then the ion must combine with the opposite charged ion and so salt will precipitate till Ksp=Qsp. Also, if concentration of ions decreases then more salt dissolves to increase the concentration of both ions till Ksp=Qsp.

For example- When HCl is passed through saturated solution of NaCl, then precipitation of NaCl occurs due to increased concentration of Cl– ion obtained on dissociation of HCl. Product obtained is of high purity and the impurities are removed with sodium sulfates and magnesium sulfates.

The common ion effect is used for complete precipitation of an ion as its sparingly soluble salt having low solubility product. Hence, silver precipitated as silver chloride, ferric ion as its hydroxide and barium ion as its sulfates for quantitative estimations.

At lower pH, solubility of weak acids like phosphates increases. Due to protonation, concentration of anion decreases at low pH which in turn increases solubility of the salt so that Ksp=Qsp.

Conclusion 

Thus, the common ion effect states the suppressing effect of ionization of an electrolyte when another electrolyte is added sharing a common ion. At equilibrium constant, two different phases are described on the basis of their solubility products. Salts sharing common cation or anion, then contribute to the concentration of the ion. According to Le Chatelier’s principle, equilibrium shift changes when there is addition of reactant. Buffer solution maintains the pH of the solution.