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IONIC EQUILIBRIUM IN SOLUTION

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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. The electrolytes which are almost completely ionised in their solutions are called strong electrolytes .

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Strong electrolytes are:1.All water soluble salts

(KCl,Na2SO4,Ca(NO3)2 ,etc.2.Alkalines (NaOH, KOH, Ca(OH)2, Ba(OH)2),

etc.3.Mineral acids (H2SO4, HNO3, HCl, HBr,

HI),etc.The equation for dissociation of strong electrolytes are written with only a single arrow directed to the right.

KCl(aq) → K+ (aq) + Cl −(aq)

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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.    

NH4OH(aq) NH4+ (aq) + OH−(aq)

Equations for the dissociation of weak electrolytes are written with double arrows( ).

 CH3COOH(aq) CH3COO−(aq) + H+ (aq)

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VARIOUS CONCEPTS OF ACIDS AND BASES 1.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.

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HNO3 → H+(aq) + NO3−(aq)

HCl(aq) → H+(aq) + Cl−(aq)

CH3COOH(aq) H+(aq) + CH3COO−(aq)

NaOH(aq) → Na+(aq) + OH−(aq)

KOH(aq) → K+(aq) + OH−(aq)

NH4OH(aq) NH4+(aq) + OH−(aq)

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According to Arrhenius theory , neutralisation of acids and bases is basically a reaction between H+ and OH− ions in solutions. H + + OH− H2O

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2.BRONSTED-LOWRY CONCEPT OF ACIDS AND BASES. The Brønsted-Lowry definition, formulated in 1923, independently by Johannes Nicolaus Brønsted in Denmark and Martin Lowry in England It is based upon the idea of protonation of bases through the de-protonation of acids

Johannes Nicolaus Brønsted Martin Lowry

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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.

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It is a reversible reactions that involve proton transfer from the acid to the base HA + B  HB+ + A−

Acid Base

Acid is known as Proton Donor.Base is known as Proton Acceptor.

HCl → H+ + Cl−

Acid ( Proton Donor, donate H+ ) Base ( Proton Acceptor, accept H+)

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

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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.

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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.

                 HCl(aq)       H +  +   Cl−(aq)  strong acid                                       weak base

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

  CH3COOH(aq) H+ (aq) + CH3COO−(aq) weak acid strong baseThe strength of acids or bases is experimentally measured by determining its ionisation or dissociation constants.

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3. 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.

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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 .

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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.

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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.Molecules having an atom with incomplete octet.

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For example , BF3 and AlCl3 .

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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.

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THE DISSOCIATION CONSTANTS OF ACIDS (Ka ) Strong acids dissociate almost completely in water and therefore the molar concentrations of H+ 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 :

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    HA(aq)        H+ (aq) + A −(aq)Applying the law of Chemical equilibrium:

Here Ka is called dissociation constant of the acid.

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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 H+ 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: CH3COOH(aq) H+ (aq) +CH3COO−(aq)

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Factors affecting acid strength The extent of dissociation of an acid depends on the strength and polarity of the H−A bond. In general , when strength of H−A bond decreases , that is , the energy required to break the bond decreases. HA becomes a stronger acid. Also, when the H−A bond becomes a stronger acid. Also, when the H−A bond becomes more polar i.e., the electronegativity difference between the atoms H and A increases and there is marked charge separation, cleavage of bond becomes thereby increasing the acidity. But it should be noted that while comparing elements in the same group of the periodic table, H−A bond strength is a more important factor in determining acidity than its polar nature.

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As the size of A increases down the group, H−A bond strength decreases and so the acid strength increases. For example,

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Degree of ionisation (α) = (Number of ions (n)) ÷ (Total number of ions and molecules (N)).

α =

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According to Arrhenius theory of electrolyte dissociation, the molecules of an electrolyte in solution are constantly splitting up into ions and the ions are constantly reuniting to form unionized molecules. Therefore, a dynamic equilibrium exists between ions and unionized molecules of the electrolyte in solution. It was pointed out by Ostwald that like chemical equilibrium, law of mass action can be applied to such systems also.

Ostwald’s Delution Law.

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   H3CCOOH(aq) H+ (aq) + CH3COO−(aq)

where:Ka: constant of dissociationα: degree of dissociation C(CH3COO-): concentrations of anionsC(H+): concentration of cationsC(CH3COOH): concentration of associated electrolyte.

C(1-α) Cα Cα

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For very weak electrolytes, α <<< 1, (1 - α ) = 1 K = Cα2

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

Thus, degree of dissociation of a weak electrolyte is proportional to the square root of dilution.

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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.                    AB     A+  +  B−    AB

         

                    (solid salt)        (dissolved salt)                (ions)

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                                       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                      K[AB] = [A+] [B−]

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

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Hence . [A+] [B−]= K × Constant = KSp where KSp 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 :                      AgCl Ag+  + Cl−

KSp = [Ag+] [Cl−]

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

                      Ax By x Ay+ y Bx−

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where Ay+ and Bx− 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 :

                      KSp = [Ay+]x [Bx− ]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.

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KSp = [A+] [B−] = S × S = S2

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.

AB       A+(aq) +  B− (aq)

In general , for any sparingly soluble salt A x B y which dissociates to set up the equilibrium :                      Ax By In general , for any sparingly soluble salt A x B y which dissociates to set up the equilibrium :                      Ax By x Ay+ y Bx−

                      [Ay+] = x S and [Bx− ] = y S

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KSp   =[x S]x   [y S ]y = xx yy S(x+y)

The concept of solubility product principle helps us to predict whether a salt will precipitate or not. Precipitation occurs : if calculated ionic product > K sp

No precipitation : if calculated ionic product < KSp.

Thank You!

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