UNIT 03 CHEMICAL EQUILIBRIUM –II (IONIC EQUILIBRIUM)

EQUILIBRIA INVOLVING IONS

Electrolytes when dissolved in water splits up into charged particles called ions.. The process is called ionisation or dissociation. Certain electrolytes such as NaCl, KCl, HCl are almost completely ionised in solutions, whereas electrolytes such as NH₄OH, CH₃COOH etc. are weakly ionised. The electrolytes which are almost completely ionised in their solutions are called strong electrolytes. On the other hand, electrolytes which are weakly ionised in their solutions are called weak electrolytes. In case of solutions of weak electrolytes, the ions produced by dissociation of electrolyte are in equilibrium with undissociated molecules of the electrolyte. The equation for dissociation of strong electrolytes are written with only a single arrow directed to the right.

KCl(aq)        ®     K⁺(aq)    +  Cl^-(aq)

NH₄Cl(aq)     ®    NH₄⁺(aq)  +  Cl^-(aq)

On the other hand, equations for the dissociation of weak

VARIOUS CONCEPTS OF ACIDS AND BASES

ARRHENIUS CONCEPT OF ACIDS AND BASES

According to Arrhenius concept :

An acid is a substance which can furnish hydrogen ions in its aqueous solution.  A base is a substance which can furnish hydroxyl ions in its aqueous solution.

For example, substances such as HNO₃, HCl, CH₃COOH etc are acids, whereas substances such as NaOH , KOH , NH₄OH etc. are bases, according to this concept.

Acids such as HCl and HNO₃ which are almost completely ionised in aqueous solution are termed as strong acids whereas acids such as CH₃COOH which are weakly ionised are called weak acids.

          Similarly, bases which are almost completely ionised in aqueous solution are called Strong bases , for example, NaOH and KOH. The bases such as NH₄OH are only  slightly ionised are called weak bases.

          According to Arrhenius theory , neutralisation of acids and bases is basically a reaction between H⁺ and OH^- ions in solutions.

Nature of Hydrogen ion in aqueous solutions

          Hydrogen atom contain one proton and one electron. H⁺ ion is formed by loss of this electron. Therefore H⁺ ion is simply a proton. Charge density of this unshielded proton is very high. Therefore, it is not  likely to exist as H⁺ ion. In an aqueous solution, H⁺ ion is considered to be present in hydrated form in combination with a water molecule as H₃O⁺.

H⁺  +   H₂O   ®   H₃O⁺

H₃O⁺ is called hydronium ion.

BRONSTED-LOWRY CONCEPT OF ACIDS AND BASES

 In 1923, Bronsted and Lowry independently proposed new definitions for acids and bases. They proposed that :

   An acid is a substance that can donate a proton.

A base is a substance that can accept a proton.

These definitions  are more general  because according to these definitions even ions can behave as acids or bases. Moreover, these definitions are not restricted to reactions taking place in aqueous solutions only. In order to understand this concept of acids and bases , consider some specific example :

From the above equations, it is obvious that acid base reactions according to Bronsted-Lowry concept involve transfer of proton from the acid to a base.  A substance can act as an acid only if another substance capable of accepting a proton is present.

Conjugate acid-base pairs

 An acid after losing a proton becomes a base  whereas a base after accepting the proton becomes an acid.  For example, consider the reaction between water and ammonia as represented by the following equilibrium:

In the forward reaction, water donates a proton to ammonia(base) and acts as acid. In the reverse reaction NH₄⁺ ions donate a proton to OH^- ions (base) and acts as acid. A base formed   by the loss of proton by an acid is called conjugate base of the acid, whereas an acid formed by gain of a proton by the base is called conjugate acid of the base. In the above example, OH^-  ion is the conjugate base of H₂O and NH₄⁺ ion is the conjugate acid of NH₃.  Acid- base pairs such as H₂O - OH^-  and NH₄⁺- NH₃ which are formed by loss or gain of proton are called conjugate acid-base pairs.

          A strong acid would have large tendency to donate a proton. Thus, conjugate base of a strong acid would be a weak base. Similarly, a conjugate base of a weak acid would be a strong base.

Some more conjugate acid-base pairs has been given in the following equations :

It may be noted that in equation (1) H₂O is behaving as a base whereas in equation (2) it s behaving as an acid. Similarly HCO₃^- ion in equation (3) acts as an acid and in equation (4) it acts as a base. Such substances which act as acids as well as bases are called amphoteric substances.

          In both Arrhenius and Bronsted concepts, acids are sources of protons. Hence all Arrhenius acids are also Bronsted acids. However, there is a difference in the definition of bases. Arrhenius theory requires base to the source of OH^- ions in aqueous medium, but Bronsted theory requires base to be a proton acceptor. Hence Arrhenius bases may not be Bronsted bases. For example, NaOH is a base according to Arrhenius theory because it gives OH^- ions in aqueous solution, but NaOH does not accept proton as such. Hence it may not be classified as a base according to Bronsted theory.

Problems

1.        Write  the conjugate acids for the following Bronsted bases :

F^- , HSO₄^- , CO₃^{2 -} 

2.        Write the conjugate acid base of :

(i)      HCO₃^-  (ii)   H₂O    (iii)   HS^-    (iv)    NH₃

3.        In the reaction :

        which reactant would be considered as a Bronsted

        acid.

4.        What will be the the conjugate bases for the following acid : HF, H₂SO₄ and HCO₃^-  ?

5.        The species H₂O, HCO₃^-, HSO₄^- and NH₃ can  both act as Bronsted acids and bases. For each give the corresponding conjugated acid and base

6.        Which of the following are Lewis bases ?

H₂O, BF₃, H⁺, and NH₃

7.        What is the conjugate base of HCN  ?

8.        Which conjugate base is stronger CN^-  or F^-   ?

Strengths of acids and bases

  Strength of an acid is measured in terms of its tendency to lose proton whereas strength of a base is measured in terms of its tendency to accept proton. The conjugate base of a strong acid is a weak base.

On the other hand, conjugate base of a weak acid is a strong base.

A base is considered to be strong if it has great tendency to accept a proton. Therefore, conjugate acid of a strong base has little tendency to lose proton and hence is a weak acid.

On the other hand, conjugate acid of a weak base is a strong acid.

The strength of acids or bases is experimentally measured by determining its ionisation or dissociation constants.

THE LEWIS ACIDS AND BASES

Although Bronsted-Lowry theory was more general than Arrhenius theory of acids and bases , but failed to explain the acid base reactions which do not involve transfer of protons. For example it fails to explain how acidic oxides such as anhydrous CO₂, SO₂, SO₃ etc. can neutralise basic oxides such as CaO, BaO etc. even in absence of solvent.

          Lewis proposed a more general definition for acids and bases, which do not require the presence of protons to explain the acid-base behaviour.

 Accoding to Lewis concept :

An acid is a substance which can accept a pair of electrons.

A base is a substance which can donate a pair of electrons.

Acid-base reactions according to this concept involve the donation of electron pair by a base to an acid to form a co-ordinate bond. Lewis bases can be neutral molecules such as

having one or more unshared pairs of electrons. , or anions such as -CN^- , -OH^- , Cl^-, etc.

          Lewis acids are the species having vacant orbitals in the valence shell of one of its atoms. The following species can act as Lewis acids.

(a)     Molecules having an atom with incomplete octet.

       For example , BF₃ and AlCl₃.

       

       

(b)     Simple cations.

        For example H⁺, Ag⁺ etc.

       

        

( c)  Molecules in which central atom has vacant orbitals

       and may acquire more than octets of valence electrons.

       Forexample, SiF₄.

      

(d)     Molecules containing multiple bonds. For example CO₂, SO₂, etc.     

          It may be noted that all Bronsted bases are also Lewis bases but all Bronsted acids are not Lewis acids. Lewis bases generally contain one or more lone pairs of electrons and therefore , they can also accept a proton (Bronsted base). Thus, all Lewis bases are also Bronsted bases. On the other hand, Bronsted acids are those which can give a proton, for example , HCl, H₂SO₄ . But they are not capable of accepting a pair of electrons . Hence , all Bronsted acids are not Lewis acids.

Problems

9.        Classify the following into acids and bases according to Lewis concept:

        SO₃,   CaO,   OH^-,   BF₃,   RNH₂,   S^2-,   Ag⁺.

10.     In the reaction:

       SnCl₄   +  2 Cl^- ®   [SnCl₆]^2-

     Which is Lewis acid and which one is the Lewis base ?

11.     Classify the following species into Lewis acids and bases and show how these act as such :

(a)  HO^-   (b)  F^-       (c)  H⁺      (d)   BCl₃

IONISATION OF ACIDS AND BASES

          Acids like HCl, HNO₃ and H₂SO₄ when dissolved in water dissociate completely and thus producing a large number of H⁺ ions. Hence these acids are called strong acids.  On the other hand, acids like CH₃COOH, HF, H₂CO₃, H₃PO₃ etc,. dissociate only  to a small extent in the aqueous solution giving small amount of  H⁺ ions and hence are called weak acids. Similarly bases like NaOH and KOH dissociate almost completely in the aqueous solution producing a large number of OH^- ions and are called strong bases, whereas bases like NH₄OH, Ca(OH)₂, Al(OH)₃ etc  dissociate only to a small extent in aqueous solution and are therefore , called weak electrolytes.

THE DISSOCIATION CONSTANTS OF ACIDS (K_a)

Strong acids dissociate almost completely in water and therefore the molar concentrations of H₃O⁺ions in the solution is same as that of acid itself. But weak acids are not completely dissociated and relative strengths of weak acids can be compared in terms of their dissociation constants. For example, the dissociation equilibrium of an acid HA may be represented as :

Applying the law of Chemical equilibrium:

Since the concentration of water is very large and remains almost constant in solution, it can be combined to give another constant K_a.

Here K_a is called Dissociation constant of the acid.

          The value of dissociation constant gives an idea about the relative strength of the acid. Larger the value of K_a ,greater is the concentration of H₃O⁺ions and stronger is the acid. If dissociation constants of two acids are known, their relative strength can be compared. For example, consider the following examples:

             

                     = 1.80 x 10^-5

        

                                        

Since K_a for CH₃COOH is larger than K_a for HCN, acetic acid is a stronger acid than HCN.

          It may be noted that the above expression of K_a is applicable to monobasic acids. The polybasic acids like H₂S, H₂CO₃,  H₃PO₄  etc. dissociate in stages and different stages of dissociation have different values of dissociation constants. For example, the dissociation stages of H₂S , a dibasic acid are given as follows :

Calculation of [H₃O⁺] and degree of Dissociation

From the knowledge of Ka , it is possible to calculate hydronium ion concentration and degree of ionisation of a weak acid.

Degree of dissociation or ionisation may be defined as  the fraction of the total number of molecules an electrolyte (acid or base) which dissociates into ions. Thus ,

          

 As an example, consider acetic acid. The following equation represents the ionisation of acetic acid in aqueous solution.

Suppose C moles of CH₃COOH are dissolved per litre of the solution and a the degree of ionisation of CH₃COOH, then at equilibrium the concentration of various species would be as follows :

Since for weak acids a is very small as compared to 1, a in the denominator can be neglected.  The expression for Ka then  becomes :

Knowing the value of Ka , it is possible to calculate the degree of ionisation of weak acid at any particular concentration C.

From the degree of ionisation, hydronium ion concentration can be calculated as :

THE DISSOCIATION CONSTANTS OF BASE (K_b)

The ionisation constant K_b for weak base BOH can be represented as follows :

(Concentration of water remains constant)

Smaller the value of ionisation constant for a base, weaker is the base.

          If C is the molar concentration of base and a is its degree of dissociation, then the concentrations of each species at equilibrium are:

But for weak base , a is very small so that ( 1 - a ) »  1

Therefore.

The relative strengths of  acids and bases can be  compared in terms of their dissociation or ionisation constants. The acid (or  a base ) having a higher value of dissociation constant is stronger. But for a quantitative comparison, the concentration of H⁺ ions  should be determined for both acids in solution of equal molarity. In general, strengths of two acids HA(1) and HA(2)

 

Where K_a(1) and K_a(2) are dissociation constants.

Similarly for two bases,

             

Ionisation constants of some acids

Acid	Ka	Acid	Ka
Hydrofluoric acid	3.5 x 10-4	Benzoic acid	6.5 x 10-5
Nitrous acid	4.5 x 10-4	Formic acid	1.8 x 10-4
Niacin	1.5 x 10-5	Acetic acid	1.85 x 10-5
Hypochlorous acid	3.0 x 10-8	Hydrocyanic acid	4.9 x 10-10
Phenol	1.3 x 10-10

Ionisation constants of some bases

Base	Kb	Base	Kb
Dimethyl amine	5.4 x 10-4	Triethyl amine	6.45 x 10-5
Ammonia	1.77 x 10-5	Quinine	1.10 x 10-6
Pyridine	1.77 x 10-9	Aniline	4.27 x 10-10
Urea	1.3 x 10-14

Problems

12.     Calculate [H₃O⁺] ion concentration of 0.01 M solution of an acid HA. The dissociation constant of the acid is 1.6 x 10^-5.

13.     At 298 K a 0.1 M solution of acetic acid is 4 % ionised. What is the ionisation constant Ka for acetic acid.

14.     Calculate the degree of ionisation of 0.01 M solution of HCN , if its Ka is 4.8 x 10^-10 . Also calculate the hydronium ion concentration.

15.     Calculate the concentration of H₃O⁺ ion in a mixture of 0.02 M acetic acid and 0.2 M sodium acetate. Ka of acetic acid is 1.8 x 10^-5.

16.     The acid dissociation constants of HCN, CH₃COOH and HF are 7.2 x 10^{- 10} , 1.80 x 10^{- 5} and 6.7 x 10^{- 4} respectively. Write their names in the increasing order of their strengths.

SELF IONISATION OF WATER

        The conductivity measurements of water indicate that water is a weak electrolyte. Even in pure state water is weakly ionised to give H₃O⁺(aq) and OH^-(aq) ions as shown under :

Applying the law of equilibrium :

Since concentration of water is very high ( 55.5 M ) and only a small fraction of it undergo ionisation, therefore [H₂O] may be taken as constant and may be combined with K to have another constant Kw .

The constant Kw is called ionic product of water. Its value is       1.008 x 10^-14 mol² L^-2.

          In pure water [H₃O⁺] and [OH^-] are equal. Therefore ,

Thus in pure water,

[H₃O⁺] = [OH^-]  = 1.0 x 10^-7  mol L^-1   at 298 K.

Since with increase in temperature dissociation of water increases, therefore, value of Kw increases as the temperature is increased. However, at all temperatures [H₃O⁺] remains equal to [OH^-] in pure water. The values of Kw at different temperatures are given in TABLE.

Temperature (K)	Ionic Product (Kw)
273	0.113 x 10-14
283	0.292 x10-14
298	1.008 x 10-14
313	2.917 x 10-14
323	5.474 x 10-14
373	7.50 x 10-14

CONCENTRATIONS OF [H₃O⁺] AND [OH^-] IN AQUEOUS SOLUTIONS OF ACIDS AND BASES

          In pure water, the [H₃O⁺]equal to [OH^-]. But on addition of some acid or base to water, these concentrations no longer remain equal. However, the value of ionic product of water Kw at a particular temperature always remain constant irrespective of the fact that whether water is pure or some acid or base has been added to it. For example, if an acid (say HCl) is added to water, the concentration of hydronium ion becomes quite high. Consequently, dissociation equilibrium of water shifts in the reverse direction, i.e., H₃O⁺ ions would combine with OH^- ions to form undissociated  water molecules, so that the value of Kw in the solution may remain the same as that in pure water. Thus, addition of an acid in water decreases the [OH^-] according to the relation,

Similarly, the addition of a base such as NaOH increases the [OH^-] and decreases the [H₃O⁺] according to the relation :

From the above discussion, it is clear that H₃O⁺ and OH^- ions are always present in aqueous solution whether it is acidic or basic. However, the relative concentrations of these ions  vary in different solutions.

In general,

      In neutral solution [H₃O⁺]   = [OH^-]

      In acidic solutions  [H₃O⁺]   > [OH^-]

      In basic solutions  [H₃O⁺]    < [OH^-]

EXPRESSING HYDRONIUM ION CONCENTRATION  - pH SCALE

Sorensen (1909) suggested a new term for expressing the concentration of hydrogen ion known as pH scale. The symbol pH is derived from the Danish word Potenz meaning power. Thus pH means power of hydrogen ion.

          The pH is defined as the negative logarithm of H₃O⁺ ion concentration  in moles per litre. Mathematically, it may be expressed as :

For acidic solutions, [H₃O⁺] concentration is more than                    1 x 10^-7 mol L^-1. Therefore, pH of acidic solution is less than 7. For basic solutions, the pH value is greater than 7. Corresponding range of pH is from 0 to 14. The solution having pH between 0 and 2 are strongly acidic, those with pH between 2 to 4 are moderately acidic, while others having pH between 4 to 7 are weakly acidic.  Similarly, the solutions having pH value between 7 to 10 are weakly basic, those having pH 10 to 12 are moderately basic whereas others which have pH range between 12 to 14 are strongly basic.

pOH OF A SOLUTION

     pOH of a solution is the negative logarithm to the base 10  of OH^- ion concentration expressed in moles per litre.

Thus :

                     pOH = - log[OH^-]

Relation between pH and pOH

In any solution :

            [H₃O⁺] [OH^-] = Kw = 10^-14

Taking  logarithms on both sides:

log[H₃O⁺] + log  [OH^-] = log Kw = log [10^-14]

Reversing the signs throughout we get :

-log[H₃O⁺] - log  [OH^-] = -log Kw = -log [10^-14]

                    pH   +   pOH   =   pKw   =  14

pH of Some common substances

The name of fluid	pH	Name of fluid	pH
Saturated soln of NaOH	» 15	Blank coffee	5
0.1 M NaOH solution	13	Tomato juice	» 4.2
Lime water	10.5	Soft drinks & vinegar	» 3.0
Milk of magnesia	10	Lemon juice	» 2.2
Egg white, sea water	7.8	Gastric juice	» 2.2
Human blood	7.4	1 M HCl	» 0
Milk	6.8	Concentrated HCl	» -1.0
Human saliva	6.4

Ionisation of Weak Electrolytes –

Degree of ionisation

Weak electrolytes are ionised partially. The ions produced as a result of dissociation of weak electrolytes are present in dynamic equilibrium with the undissociated molecules. The fraction of total number of molecules of the electrolyte dissolved, that ionises at equilibrium is called  degree of dissociation. Let us consider the ionisation of some weak electrolyte AB in water. Let C be the concentration of the electrolyte in the solution and a the degree of ionisation. The concentration of various  species at equilibrium would be as given under:

According to Law of chemical equilibrium:

For weak electrolytes under  normal concentrations, a is very small as compared with unity and  hence ( 1 - a ) can be taken as  1 in the denominator.

So the degree of ionisation  is inversely  proportional to the square root of concentration. Thus, as the concentration decreases, the degree of ionisation increases.  It can also be followed from the above relationship that,  as C approaches zero or dilution approaches infinity, the degree of dissociation approaches unity, i.e., its maximum value. In other words, in this dilution , the whole of electrolyte is present in the dissociated form as ions. This important generalisation is known as Ostwald dilution Law.

SOLUBILITY EQUILIBRIA OF SPARINGLY SOLUBLE SALTS

          The solubility of ionic solids in water varies a great deal. Some are so soluble that they are hygroscopic in nature and even absorb water vapours from atmosphere. Others have so little solubility  and are called insoluble. The solubility depends on a number of factors, important among  which are the lattice energy of the salt and how well the ions are hydrated in aqueous solution. Each salt has its characteristic solubility and its dependence on temperature. We classify salts on the basis of their solubility in the following three categories.

·         Category I    Soluble           Solubility > 0.1 M

·         Category II   Slightly soluble     Solubility < 0.1 M

·         Category III  Sparingly soluble   Soubility  < 0.01 M

SOLUBILITY PRODUCT CONSTANT

Certain electrolytes such as BaSO₄ and AgCl are sparingly soluble in water. Even in their saturated solutions, the concentration of the electrolytes is very low. So , whatever little of electrolyte goes into solution, undergoes complete dissociation (due to low concentration). Therefore , in saturated solutions of such electrolytes solid electrolyte is in equilibrium with the ions as represented below :

Consider a saturated solution of a salt containing the solid salt. There are two equilibria, one between solid salt and dissolved salt and second between the dissolved salt and its ions.

Applying the Law of mass action  to the second equilibrium,

where K is the equilibrium constant and [AB] is the concentration of the dissolved salt.  Cross multiplying we get

Since the solution is saturated , the concentration of the dissolved salt remains constant at a fixed temperature. 

Hence . [A⁺][B^-] = K x Constant = K_sp where K_sp is another constant. This constant K_sp is known as the solubility product of the electrolyte. It is the maximum value of product of concentrations of the ions of the electrolyte.

          In the case of silver chloride, we have :

K_sp = [Ag⁺]   [Cl^-]

For lead chloride :

K_sp = [Pb²⁺]   [Cl^-]²

          In general , for any sparingly  soluble salt A_xB_y

which dissociates to set up the equilibrium :

where A^Y+ and B^X- denote the positive and negative ions , x and y represent the number of these ions in the formula of the electrolyte. The  solubility  product constant may be expressed as :

K_sp = [A^Y+]^x   [B^{X -}]^y

Thus solubility product of a sparingly soluble salt at a given temperature may be defined as the product of the concentrations of its ions in the saturated solution, with each concentration term raised to the power equal to the number of  times the ion occurs in the equation representing the dissociation of the electrolyte.

Problem

17.   Express the solubility products of :

i)        PbCl₂   ii)   Mg(OH) ₂      iii)   Ca₃(PO₄) ₂   iv)    ZnS

Calculation of Solubility product

 Knowing the solubility of the salt, its solubility product can be calculated. Consider the salt AB. Suppose at a particular temperature its solubility is S mol L^-1. S moles of salt on ionisation give S moles of A⁺ and S moles of  B^- ions.

Knowing the solubility S, K_sp can  be calculated.

          In general , for any sparingly  soluble salt A_xB_y which dissociates to set up the equilibrium :

Problems

18.     The solubility of AgCl is 1.06 x 10^-5 mol L^-1at 298 K. Find out the solubility product of AgCl at this temperature.

19.     The solubility of PbCl₂ at 298 K is 2 x 10^-2 mol L^-1. Find out the solubility product at this temperature.

20.     Which of the following is more soluble ?

(a)      AgCl (Ksp = 1.7 x 10^-10)

(b)     Ag₂CrO₄ (Ksp = 4 x 10^-12 )

21.     The ionisation constant of phenol is 1.0 x 10^-10 . What is the concentration of phenate ion in 0.05 M solution of phenol ? What will be its degree of ionisation if  the solution is also 0.01 M  in sodium phenate.

Applications of Solubility Products

1.       Calculation of Solubility

 If solubility product of a sparingly soluble salt at a particular temperature is known, its solubility at that temperature can be calculated.

2. Predicting the precipitation of a salt

  The concept of solubility product principle helps us to predict whether  a salt will precipitate or not, on mixing the solutions containing its ions under particular conditions. In the solution, the ionic product, i.e., product of concentration of the ions of the salt(raised to appropriate power) cannot exceed the value of its solubility product because solubility product is the highest limit of ionic product at a particular temperature. Thus, if ionic product exceeds solubility product, excess ions combine each other to form precipitate of the salt. So in order to predict whether a salt will precipitate or not, ionic product of the salt is calculated. If it exceeds the value of solubility product, the salt will be precipitated , otherwise not. Thus, it can be concluded that :

Precipitation occurs : 

                          if calculated ionic product > K_sp

No precipitation    :

                         if calculated ionic product <  K_sp.

Problems

22.     The solubility product of Ag₂CrO₄ at 298 K is 4 x 10^-12. Find out the solubility at this temperature.

23.     The solubility product of AgBr at a certain temperature is      2.5 x 10^-13. Find out the solubility of  AgBr in gram per litre at this temperature.

24.     How many moles of AgBr (Ksp = 5 x 10^-13 mol²L^-2) will dissolve in 0.01M NaBr solution.

25.     What [H₃O⁺] must be maintained  in a saturated solution to precipitate Pb²⁺ , but not Zn²⁺ from a solution in which each ion is present at a concentration of 0.01 M ? [Ksp (H₂S) = 1.1 x 10^-22, Ksp (ZnS) = 1.0 x 10^-21 )

26.     Equal volumes of 0.02 M sodium sulphate and 0.02 M Barium chloride solutions are mixed together. Predict the precipitation will occur or not (K_sp of BaSO₄ is 1.5 x 10^-9)

27.     50 ml of 0.01 M solution of calcium nitrate is added to 150 ml of 0.08 M solution of ammonium sulphate. Predict whether CaSO₄ will be precipitated or not. (Ksp of CasO₄ = 4 x 10^-5.

3.  Precipitation of Soluble Salts

The principle of solubility product is also applied in the precipitation of soluble salts from their saturated solution, in pure state. The phenomenon known as salting out, is used in the purification of sodium chloride. This is done by preparing a saturated solution of impure sodium chloride in water when the following equilibrium exists :

HCl gas is passed through this solution. The concentration of chloride ions increases considerably. The ionic product exceeds the solubility product of NaCl and therefore , it precipitates out from the solution in pure state. The impurities remain in solution.

4. Inorganic Qualitative Analysis

The classification of basic radicals into different groups in the inorganic qualitative analysis is based upon the knowledge of solubility  products of salts of these basic radicals. For example , chlorides of Hg₂²⁺, Pb²⁺and Ag⁺ have very low solubility products. On the basis of this knowledge these radicals are grouped together in Group-I and are precipitated as their chlorides by adding dil. HCl  to their solutions.  For adjusting the condition for precipitation, another concept called Common Ion Effect  plays a very important role.