Notes and Study Materials -Chemical and Ionic Equilibrium

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About this unit

Equilibrium in physical and chemical processes, dynamic nature of equilibrium, law of chemical equilibrium, equilibrium constant, factors affecting equilibrium-Le Chatelier’s principle; ionic equilibrium- ionization of acids and bases, strong and weak electrolytes, degree of ionization, ionization of polybasic acids, acid strength, concept of PH., Hydrolysis of salts (elementary idea), buffer solutions, Henderson equation, solubility product, common ion effect (with illustrative examples).

 

CHEMICAL EQUILIBRIUM

REVERSIBLE REACTIONS

Reactions which do not always proceed to completion and may be made to proceed in the opposite direction under suitable conditions are called reversible reactions e.g.
3Fe + 4H2O Fe3O4 + 4H2
H2 + I2 2HI
N2O4 2NO2
N2 + 3H2  2NH3

IRREVERSIBLE REACTIONS

Reactions which always proceed to completion in one direction only are called irreversible reactions.
 
       
         

CHEMICAL EQUILIBRIUM

When a reversible reaction is carried out in a closed vessel a stage reached when the speed of the forward reaction equals the speed of the backward reaction and chemical equilibrium is said to be established.

CHARACTERISTICS OF CHEMICAL EQUILIBRIUM

  1. Equilibrium can be attained from either side.
  2. Equilibrium is dynamic in nature i.e. at equilibrium, reaction does not stop.
  3. At equilibrium there is no change in the concentration of various species.
  4. The equilibrium state remains unaffected by the presence of catalyst. Catalyst helps to attain the equilibrium state rapidly.
  5. It can be achieved in a closed container.
  6. The observable properties of the process become constant and remain unchanged.

EQUILIBRIUM STATE AND FREE ENERGY CHANGE

At equilibrium G is equal to zero and we have
G = H – TS
H = TS

LAW OF MASS ACTION

It was put forward by Guldberg and Waage. The law states that the rate at which a substance reacts is directly proportional to its active mass and the rate of a chemical reaction is directly proportional to the product of the active masses of the reacting substances. For a general reaction
aA + bB   cC + dD
Rate of forward reaction
Rate of backward reaction
where Kf and Kb are velocity constants for forward and backward reactions respectively. At equilibrium point,
Rate of forward reaction = Rate of backward reaction
Kc is called the equilibrium constant.

FACTORS INFLUENCING EQUILIBRIUM CONSTANT

EQUILIBRIUM CONSTANT IS NOT INFLUENCED BY

  1. Concentration of reactants and products.
  2. Presence of a catalyst.
  3. Pressure.
  4. Presence of inert materials.
  5. The direction from which the equilibrium state is reached.

EQUILIBRIUM CONSTANT IS INFLUENCED BY

  1. Temperature : The variation of equilibrium constant is given by Van’t Hoff  equation.
where ,   (Kp)1 and (Kp)2 = Equilibrium Constant at temperature T1 & T2
R = Universal Gas Constant
For exothermic reaction : Kp decreases with increase of temperature since Kf decreases.
For endothermic reaction : Kp increases with increase of temperature since Kf increases.
  1. The mode of representing the reaction :
The reaction A + B C + D may be written as
   C + D  A + B
            
 
  1. Stoichiometric representation of equation :
N2 + 3H2   2NH3
 NH3
          

USE OF PARTIAL PRESSURE INSTEAD OF CONCENTRATIONS

For gaseous reacting substances partial pressures are conveniently used since at any fixed temperature partial pressure is directly proportional to concentration. For a general reaction
aA + bB   cC + dD

RELATION BETWEEN KC AND KP

where
n = [moles of products – moles of reactants] gaseous only.

RELATION BETWEEN KC AND KP FOR DIFFERENT TYPES OF REACTIONS

  1. When n = 0, Kp = Kc e.g. for reaction A B.
  1. When n = +ve, Kp > Kc e.g. for reaction A  2B.  
  1. When n = –ve, Kp < Kc e.g. for reaction 2A  B.  

UNITS OF Kp AND Kc

  • Unit of Kp = (atm) n
  • Unit of Kc = (mol lit–1) n

CHARACTERISTICS OF EQUILIBRIUM CONSTANT

  1. It has definite value for every chemical reaction at a particular temperature
  2. The more is the value of Kc or Kp, the more is the completion of reaction or the more is the concentration of products.
  3. When the reaction can be expressed as sum of two other reactions, the Kc of overall reaction is equal to the product of equilibrium constants of individual reactions.

HOMOGENEOUS EQUILIBRIUM

In homogeneous equilibrium the reactants and products are present in the same phase (gaseous or liquid).
2SO2 (g) + O2 (g)  2SO3 (g)

HETEROGENEOUS EQUILIBRIUM

In heterogeneous equilibrium the reactants and products are present in two or more phases.
3Fe(s) + 4H2O(g)  Fe3O4 (s) + 4H2 (g)

CHEMICAL EQUILIBRIUM APPLIED TO HOMOGENEOUS SYSTEM

GASEOUS SYSTEM

They are of two types
  1. Gaseous reactions in which the number of moles of products remain the same as that of reactants
 
Hydrogen – iodine equilibrium : Suppose, a moles of H2 and b moles of I2 are present in a container of V litres. At equilibrium x moles of each have combined to form HI.
           
Initial molar conc.             0
Eqb. molar conc.         
Applying the law of chemical equilibrium
 ……….(i)
The equilibrium constant written as Kc indicates that active masses are expressed in terms of molar concentrations.
The eq. (i) does not contain the volume term. Thus equilibrium is independent of volume and therefore of pressure.

 

  1. Gaseous reactions in which the number of moles of products and reactants are different.
        
Dissociation of PCl5 : Suppose  ‘a’ moles of PCl5 are present in a container of V litres. At equilibrium x moles have dissociated.
              
Initial molar Conc.         0         0
Eqb. molar Conc.                        
Applying the law of chemical equilibrium
 …….(i)
The eq.(i) contains the V term in denominator . If volume increases, the dissociation of PCl5 must also increase to keep Kc constant. The decrease of pressure will cause increase in volume and so the dissociation.
If the value of x is small then   ; 

LIQUID SYSTEM

Examples are :
  1. Esterification of acetic acid
At equilibrium 2/3rd of acetic acid is converted into ester. Hence alcohol consumed will also be 2/3rd.
  1. Reaction between amylene and tricholoroacetic acid

CHEMICAL EQUILIBRIUM APPLIED TO HETEROGENEOUS SYSTEM

  1. Dissociation of calcium carbonate
Applying the law of chemical equilibrium
The active mass of a solid reactant and product is assumed to have a constant value and is taken as unity. The equilibrium  constant is determined by gaseous substances only. Therefore,
  1. Reaction of steam on heated iron
Partial pressures of solid is taken unity.
  1. Water gas reaction
Since partial pressure of carbon (solid) is taken as unity, the equilibrium constant is given by

VAN’T HOFF ISOCHORE

A relationship between the equilibrium constant Kp, at  any temperature T and constant pressure P, and heat of reaction H°.
The enthalpy change DH does not vary appreciably with change in partial pressures of reactants and products. Therefore DHº can be taken as DH whatever may be the partial pressures of reactants and products
The integrated form of the equation is
Three important conditions may arise
  1. when H = 0 no heat is evolved or absorbed
Equilibrium constant does not change when no change in temperature.
  1. when H = +ve i.e. heat is absorbed
Equilibrium constant increases with increase of temperature.
  1. when H = –ve i.e. heat is evolved
  ;   Equilibrium constant
decreases with increase of temperature
  1. when n = 0 i.e. there is no change in volume during a reaction KP = Kc. The variation of equilibrium constant with temperature is given by
, E heat of reaction at constant volume.

VAN’T HOFF REACTION ISOTHERM

It gives the free energy change of a reaction at any given temperature, pressure and composition of the reacting system.
G = G° + RT ln J
At equilibrium G = 0 then G° = –RT ln Jeq
J stands for reaction quotient of partial pressure of products and reactants.
viz.
Jeq means the partial pressure of the products and the reactants at the equilibrium. Hence Jeq can be replaced by Kp.
G° = –RT ln

HENRY’S LAW

The mass of a gas dissolved per unit volume of solvent is proportional to the pressure of the gas in equilibrium with the solution at constant temperature.
  • The volume of the gas dissolved remains the same inspite of increase in pressure.
  • The dissolution of a gas in a liquid is spontaneous process (G = 0), accompanied by decrease in entropy (S = –ve). Since G = H – TS, G can only be negative if H is –ve. Therefore dissolution of a gas in a liquid is always exothermic in nature.

FACTORS ALTERING THE STATE OF EQUILIBRIUM – LE CHATELIER’S PRINCIPLE

There are three main factors which alter the state of equilibrium. They are (I) Concentration, (II) Temperature, and (III) Pressure.
Le Chatelier’s principle states that if a system at equilibrium is subjected to a change of concentration, pressure or temperature, the equilibrium shifts in the direction that tends to undo the effect of the change.
  1. Effect of change of concentration :
If at equilibrium the concentration of one of the reactants is increased, the equilibrium will shift in the forward direction and vice versa. Consider the following equilibrium
 Fe3+ (aq)      + SCN (aq)         [Fe(SCN)]2+ (aq)
Pale yellow           Colourless       Dark brown
If ferric salt is added the colour of the solution darkens immediately i.e. Fe3+ ions are consumed and more [Fe(SCN)]2+ are formed. If some sulphocyanide salt is added the colour also darkens. If Potassium ferrisulphocyanide capable of giving complex ion [Fe(SCN)]2+ is added the colour lightens to pale yellow.
  1. Effect of change in pressure :
    1. No effect of pressure on equilibria having same moles of reactants and products e.g. N2 + O2 2NO
      H2 + I2  2HI
    2. When there is change in the number of moles the equilibrium will shift in the direction having smaller number of moles when the pressure is increased and vice versa e.g.
N2 + 3H2 2NH3
More pressure more ammonia
PCl5 PCl3 + Cl2
The more the pressure, the lesser the dissociation of PCl5.
  1. Effect of temperature :
    1. When process is exothermic – Low temperature favours the formation of products.
    2. When process is endothermic – High temperature favours the formation of products
e.g. N2 + 3H2  2NH3 + 24.0 kcal.
Since the production of NH3 is exothermic low temperature favours its formation.
  1. Effect of addition of inert gas :
    1. Addition of Inert gas at constant volume : The total pressure of the system is increased, but the partial pressure of each reactant and product remains the same. Hence no effect on the state of equilibrium.
    2. Addition of Inert gas at constant pressure : The total volume is increased, the number of moles per unit volume of each reactant and product is decreased. Hence equilibrium will shift to the side where number of moles are increased e.g.
PCl5 (g)    PCl3 (g) + Cl2 (g)
Introduction of inert gas at constant pressure will shift the equilibrium to right hand side.
  1. Effect of catalyst : The presence of catalyst does not change the position of equilibrium. It simply fastens the attainment of equilibrium.

LE CHATELIER’S PRINCIPLE APPLICABLE TO PHYSICAL EQUILIBRIUM

  1. Effect of pressure on solubility : The increased pressure, will increase the solubility of a gas and vice versa.
  2. Effect of temperature on solubility : The substances which dissolve with the absorption of heat, their solubility will increase with increase of temperature and vice versa e.g. dissolution of NH4Cl, KCl, KNO3 is endothermic which increases with increase of temperature. The dissolution of calcium acetate and Calcium hydroxide is exothermic, their solubility is lowered at higher temperature.
  3. Effect of pressure on the melting point of ice :
Ice  liquid water
The ice occupy the more volume than liquid water, so increased pressure will result in melting of ice according to Le Chatelier’s principle.

FAVOURABLE CONDITIONS FOR SOME IMPORTANT REACTIONS

Synthesis of ammonia (Haber’s process)
N2 (g) + 3H2 (g)  2NH3 (g) + 22.4 kcal
  • Low temperature (500°C)
  • High pressure (200 – 1000 atm.)
  • Excess of N2 and H2
Synthesis of NO (nitric acid birkland eyde process)
N2 (g) + O2 (g)  2NO(g)– 43.2 kcal
  • High temperature
  • Excess of N2 and O2
  • No effect of pressure
Formation of SO3 (sulphuric acid contact process)
2SO2 (g) + O2 (g)  2SO3 + 42.0 kcal
  • Low temperature
  • High pressure
  • Excess of SO2 and O2
Formation of nitrogen dioxide
2NO + O2 2NO2 + 27.8 kcal
  • Low temperature
  • High pressure
  • Excess of NO and O2
Dissociation of nitrogen tetraoxide
N2O4  2NO2 – 14 kcal
  • High temperature
  • Low pressure
  • Excess of N2O4
Oxidation of CO by steam (Bosch process)
CO + H2O  CO2 + H2 + x kcal
  • Low temperature
  • Excess of steam and CO
  • No effect of pressure
Dissociation of PCl5
PCl5 PCl3 + Cl2 – 15 kcal
  • High temperature
  • Low pressure
  • Excess of PCl5

TRIPLE POINT

The temperature and pressure at which the three states of a substance can exist in equilibrium is known as triple point e.g.
Ice (s) water (l)  vapour (g) can exist at 0.0098°C and 4.58 mm.

DEGREE OF DISSOCIATION FROM DENSITY MEASUREMENT

The density of one mole of gas is given by
D = where M = Mol. wt of gas; P = Total pressure.
The volume of the gas increases on dissociation in proportion to increase in the total number of moles, but total weight remains constant. Hence density decreases in the same proportion. Consider dissociation of PCl5. Let x be degree of dissociation
PCl5        PCl3   + Cl2
1 – x  x             x Total moles (1 + x)

 

where D is the theoretical vapour density and
D= Molecular mass
d is observed vapour density at temperature tºC.
If nx moles of products are formed, then total number of moles after dissociation
1 – x + nx = 1 + x (n – 1)

IONIC EQUILIBRIUM

 

ARRHENIUS THEORY OF IONISATION

On the basis of colligative properties of solutions of salts, acids and bases, Arrhenius proposed the theory of ionisation i.e. splitting of these substances into ions in solution. It is reversible process, effects electrical conductivity, colligative properties like depression in freezing point, elevation in boiling point, lowering of vapour pressure, osmotic pressure.

EVIDENCES IN FAVOUR OF IONISATION

  1. X-ray diffraction studies
  2. Ionic reactions
  3. Heat  of neutralisation
  4. Colour of compounds and their solutions
  5. Colligative properties
  6. Conductance of electrolytes in solution

DEGREE OF IONISATION OR DISSOCIATION ()

The fraction of the total number of molecules which is ionised at the equilibrium state is known as degree of ionisation or dissociation.

FACTORS AFFECTING IONISATION OR DISSOCIATION

  1. Nature of electrolytes : The stronger the electrolyte, the more is the ionisation and vice versa.
  2. Nature of solvent : The more the dielectric constant of solvent, the more is the ionisation.
  3. Concentration : The lesser the concentration, the more is the ionisation.
  4. Temperature : The higher the temperature, the more is the ionisation.
  5. Solvation : The more the solvation, the more is the ionisation.
  6. Presence of the ions in the solution : Ionisation decreases in presence of common ions.

ELECTROLYTE

A substance which splits into ions in solution is called electrolyte. It can be an acid, base or salt.
  1. Strong electrolyte :  Which dissociates almost completely into ions even in concentrated solution eg. NaOH, KOH, HCl, H2SO4, NaCl, CaCl2.
  2. Weak electrolyte :  Which dissociates to a small extent into ions in solution eg CH3COOH, NH4OH etc.
    Note : Salts are always strong electrolytes.

OSTWALD’S DILUTION LAW

The degree of ionisation or dissociation (a) of weak electrolytes increases with dilution and law of mass action can be applied to them.
AB             A+    + B–
C  0 0 initial conc.
C(1-) C C equilibrium conc.
Ionisation constant K = C2
Concentration of A+ or B = C
 approaches unity with dilution.

ACIDS AND BASES

  1. Arrhenius concept : An acid is a substance that dissociates to give hydrogen ions when dissolved in water eg. HCl, CH3COOH, H3PO4.
A base is a substance that dissociates to give hydroxyl ions when dissolved in water eg. NaOH, Ca(OH)2.
  1. Lowry and bronsted concept : An acid is a substance which has a tendency to donate a proton (H+) to any other substance.
A base is a substance which has a tendency to accept a proton (H+) from any other substance.
Acid H+ +  + Base
Acid and base differing by a proton are known as conjugate pair. The weaker the acid, the stronger the base in conjugate pair and vice versa.
  1. Lewis concept :  An acid is a substance which can accept a pair of electrons from any other substance e.g. BF3, AlCl3 (incomplete octet), SnCl4, SF4 (central atom has vacant d-orbital) or cations Fe3+, Cu2+ etc.
A base is a substance capable to donating a pair of electrons to any other substance eg. anions X, OH, CNor neutral molecules having lone pair(s) of electrons on one or more atom   , , etc.
Lewis acid may be any of the following types of substances:-
  • Molecules having an atom with incomplete octet
  • Simple cations
  • Molecules with central atom having empty d-orbitals
  • Molecules with a multiple bond between atoms of different electronegativities
    Strength of some Lewis acids
BX3 > AlCl3 > FeX3 > GaX3 > SbX5 > InX3 > SnX4 > AsX5 > ZnX2 > HgX2
  1. Extended Lewis concept : When the central atom is bonded to atoms of different electronegativities by multiple bonds, the substance  is known as extended Lewis acid e.g. CO2, CS2 etc.
Extended Lewis base e.g. CO and unsaturated hydrocarbons like alkenes, alkynes etc. are also known as border line Lewis bases.
  1. Hard acids : Cations of lighter elements, smaller size, higher charge not easily polarisable e.g. light alkali and alkaline  metal ions of B, Al, Si, Ti4+, Cr3+, Co2+, Fe3+ (lighter transition elements).
Soft acids : Cations of heavier elements, larger size, lower charge and easily polarisable e.g. heavy transition metal ions (second and third row) e.g. Hg2+, Pd2+, Cd2+, Cu+, Ag+, Hg+ etc.
  1. Hard bases : Species having donor atoms of higher electronegativity and low polarisability e.g. N, O, F, Cl etc. Examples H2O, NH3, ROH.
Soft bases : Species having donor atom of lower electronegativity and higher polarisability e.g. P, As, S, Se etc. Examples R3P, R2S, I.
  1. Lux-flood concept of acids and bases : An oxide ion donor is a base and an oxide ion acceptor is an acid.
  1. Ingold concept : All electrophiles are acids and nucleophiles are bases.

STRENGTH OF ACIDS AND BASES

The greater the value of Ka or Kb the stronger is the acid or base. the smaller the value of pKa the stronger is the acid.
  1. Relative strengths of acids: For weak acid Ka = C 2. For two acids with dissociation constants and at the same concentration C,
  1. Relative strengths of bases : For weak base
Kb = C 2.. For two bases with dissociation constants and at the same concentration C.

LEVELLING EFFECT

All the strong acids in aqueous, solution appear almost equally strong since water acts as strong base. For example HClO4, HBr, H2SO4, HCl and HNO3 appear equally strong.
HA + H2 O   H3O+ + +A
Hence relative strengths in aqueous solution cannot be compared. This phenomenon is known as levelling effect.

EFFECT OF SOLVENT ON ACID STRENGTH

  1. In acetic acid :  
HA + CH3COOH CH3COOH2+ + A
As acetic acid has a little tendency to accept proton, even strong acids are feebly ionised in acetic acid. For example
  1. In liquid NH3 :  
        NH4+ + A
As ammonia has a great tendency to accept proton, even weak acids appear strong in liquid ammonia. For example HCl, HNO3 and CH3COOH appear  equally strong in liquid ammonia.
  1. In HF : Since HF is a strong acid, the other acids act as a base when dissolved in HF eg.
HNO3 + HF

RELATION BETWEEN Ka AND Kb

Ka × Kb = Kw  or pKa + pKb = pKw = 14 at 25°C.
Ionisation of polybasic acids : Polybasic acids ionise in various steps e.g. Orthophosphoric acid H3PO4.
H3PO4  H+ + H2PO4
H2PO4 H+ + HPO4– –
HPO4– –    H+ + PO4– –
K1 > K2 > K3 and overall dissociation const. K = K1 × K2 × K3

AMPHOTERIC OR AMPHIPROTIC SUBSTANCE OR AMPHOLYTES

A substance acting as an acid as well as a base, eg. water acts as an acid with ammonia  and as base with acetic acid. A substance acting as proton donor and proton acceptor.

COMMON ION EFFECT

The degree of ionisation of an electrolyte is suppressed by the addition of another electrolyte having a common ion. This is known as common ion effect e.g. ionisation of CH3COOH is suppressed by the addition of HCl or CH3COONa.
  • It helps in controlling the concentration of ions furnished by weak electrolytes.
  • It effects the solubility of salts.

MIXTURE OF WEAK ACID AND ITS SALT WITH A STRONG BASE

The hydrogen ion (H+) concentration of a mixture of a weak acid HA and its highly dissociated salt say NaA is given by
(I) HA   H+ + A–
(II) NaA    Na+ + A–
Ka =
   [H+] =
HA being weak acid and secondly due to common ion (A) remains almost unionized. Salts are almost 100% ionised.

MIXTURE OF WEAK BASE AND ITS SALT WITH A STRONG ACID

The hydroxyl ion (OH) concentration of mixture of weak base BOH and its highly dissociated salt say BCl is given by
(I) ionisation of BOH B+ + OH–       (negligible)
(II) ionisation of BCl   B+ + Cl (100%)

SOLUBILITY PRODUCT (KSP)

At constant temperature and pressure the saturated solution of a sparingly soluble salt has an equilibrium between the excess of the solute and the ions furnished by it. e.g.
AgCl                    AgCl             Ag+ + Cl
Solid  dissolved but               ions in sol.
undissolved   not ionised
Applying law of mass action,
or K[AgCl] = [Ag+] [Cl–], Ksp = [Ag+] [Cl–]

 

The constant Ksp is known as solubility product. It is equal to the product of the concentration of ions in saturated solution.
When Ksp > [Ag+] [Cl] Solution is not saturated
When Ksp < [Ag+] [Cl]  Solution is supersaturated and precipitation takes place
When Ksp = [Ag+] [Cl] Solution is saturated
For general electrolyte AxBy.
Ksp = [A+y]x [B–x]y
  • Ksp is independent of the source of ions.
  • Helps to know the solubility of electrolytes.
  • Predicting ionic reactions.
  • Qualitative analysis.
  • Purification of common salt, salting out of soap and Solvay ammonia soda process.
Relation between solubility product (Ksp) and solubility (S).
  1. For binary electrolyte e.g. AgCl, BaSO4
S =
  1. For ternary electrolyte e.g. CaF2, PbI2
S =
Representation of Ksp for various electrolytes.
Mg(OH)2 Ksp = [Mg++] [OH]2
Ag2S Ksp = [Ag+]2
Sb2S3 Ksp = [Sb3+]2

DISSOCIATION CONSTANT OF WATER / IONIC PRODUCT OF WATER (Kw)

Water being weak electrolyte is slightly ionised as follows :
   
 or  
or  

 

Kw is known as ionic product of water or dissociation constant of water. It is equal to the product of concentration of [H3O+] and [OH] ions in water. At constant temperature of 25°C, the value of Kw is 1.0 × 10–14.

 

In pure water [H3O+] = [OH] = 1.0 × 10–7  at 298 K
Molar concentration of water is 55.55 mol/lit
K[H2O] = [H+] [OH]
K × 55.5 = Kw.

 

Hence ionic product of water is 55.5 times greater than K. Kw increases with temperature.
The addition of salt, acid or base does not change value of KW. Its value changes with temperature only.

 

Hydrogen and hydroxyl ion concentration in aqueous solution of Acids and Bases
When an acid is added to water H+ (aq.) ion combine with OH (aq.) ions so that Kw remains  constant. Thus addition of an acid decreases the conc. of OH(aq.) ions and addition of base decreases the conc. of H+ ions. In both cases the self ionisation of water is suppressed due to extra supply of H+ or OH ions.

EXPRESSING HYDROGEN ION CONCENTRATION : pH SCALE

Any aqueous solution of some electrolyte or nonelectrolyte having equal concentrations of H+ and OHions, is neutral. It has been observed that H+ ion concentrations can usually vary from 0 to  moles/l. Sorensen represented the acidic or basic character of an aqueous solution in terms of pH. The pH of a solution is the numerical value of the negative power  to which ten must be raised to express the H+ ion concentration.
[H+] = 10–pH

 

In pure water [H+] = 10–7 = 10–pH at 25ºC or 298 K
Hence pH of pure water is 7.
log [H+] = – pH log 10
pH = – log (H+) =
Thus pH of a solution is the negative logarithm of hydrogen ion concentration. Similarly, negative logarithm of hydroxyl ion concentration is known as pOH.
pOH = –log[OH]

 

Aqueous solution having pH value less than 7, is acidic and more than 7 at 298 K is basic.
pH + pOH = 14
The pH changes with temperature. It decreases with rise in temperature.

 

pH range of some important substances:

SALT HYDROLYSIS

Salts are strong electrolytes and on dissolution in water split into ions which react with H+ or OH ions furnished by water yielding acidic or basic solution. The process is known as salt hydrolysis.

HYDROLYSIS OF SALTS OF A STRONG ACID AND STRONG BASE

e.g. KCl, NaCl, Na2SO4 etc.
KOH + HCl
Representing strong electrolytes as ions
K+ + OH + H+ + Cl
we have overall reaction H2O   H+ + OH
The solution is neutral since [H+] = [OH], pH = 7, thus salts of strong acids and strong bases are not hydrolysed.

HYDROLYSIS OF SALTS OF WEAK ACIDS AND STRONG BASES

e.g. CH3COONa, KCN
CH3COONa + H2O  CH3COOH + NaOH
or CH3COO+ Na+ + H2O CH3COOH + Na+ + OH
CH3COO + H2O  CH3COOH + OH
  1. Due to OH ions, the solution becomes basic, hence salt is hydrolysed.
  2. Hydrolysis is known as anionic hydrolysis because CH3COO is an anion.
  3. Salt hydrolysis constant Kh =
  4. pH of solution more than 7
  5. Degree of hydrolysis
  6. pH = ,
where C is  conc. of anion i.e. CH3COO

HYDROLYSIS OF SALTS OF WEAK BASES AND STRONG ACIDS

e.g. NH4Cl
      
or       
   
  1. Due to H+ ions, the solution becomes acidic, hence salt is hydrolysed
  2. Hydrolysis is known as cationic hydrolysis because  is a cation
  3. Salt hydrolysis const. Kh =
  4. pH of solution less than 7
  5. Degree of hydrolysis
  6. pH =

HYDROLYSIS OF SALTS OF WEAK BASES AND WEAK ACIDS

e.g. CH3COONH4
CH3COONH4 + H2O   CH3COOH + NH4OH
or CH3COO+ NH4+ + H2O   CH3COOH + NH4OH
  1. The solution may be neutral, acidic or basic depending upon the relative strength of weak acid  CH3COOH and weak base NH4OH formed. The salt is said to be hydrolysed.
  2. Hydrolysis is known as cationic as well as anionic hydrolysis.
  3. Salt hydrolysis Const., Kh =
  4. pH of the solution may be 7, > 7 or <7
  5. Degree of hydrolysis,
  6. pH =

BUFFERS

Solutions which resist the change in the value of pH when small amount of acid or base is added to them are known as buffers. Types
  1. Simple buffers : Solution of salt of weak acid and weak base CH3COONH4, NH4CN.
  2. Mixed buffers :
    1. Acidic buffers : Solution of equimolar mixture of a weak acid and its salt with a strong base e.g. (CH3COOH + CH3COONa); (H2CO3 + NaHCO3); (Boric acid + borax); (H3PO4 + NaH2PO4)
    2. Basic buffers : Solution of equimolar mixture of a weak base and its salt with a strong acid e.g. (NH4OH + NH4Cl)

 

Buffer action of simple buffer : CH3COONH4 exist as ions CH3COO and  in solution. Added acid (H+) combine with CH3COOto give feebly ionised CH3COOH and added base(OH) combine with NH4+ to give feebly ionised NH4OH. Thus pH remains unchanged

 

Buffer action of mixed buffers :
  • Acidic buffer :
CH3COONa     CH3COOH
When small amount of an acid is added to the buffer the added H+ ions combine with CH3COO to form feebly ionised CH3COOH and when small amount of a base is added to the buffer the added OH ions combine with H+ to form feebly ionised H2O. In the latter case more CH3COOH ionises to keep Ka of acid constant and hence constant concentration of H+ ions. Thus whether we add acid or a base, the H+ concentration remain constant and pH of solution does not change
  • Basic buffer :
NH4Cl        NH4OH
When small amount of an acid is added to the buffer the added H+ ions combine with OH to form feebly ionised H2O, equilibrium is disturbed. More NH4OH ionises to keep [OH] fixed.
When small amount of a base is added to the buffer the added OHions combine with NH4+ to form feebly ionised NH4OH to keep [OH] fixed. Hence there is no change of pH in both cases.
Buffer capacity =
Calculation of pH value of buffers : Acidic buffer contains weak acid and its salt with strong base.
The H+ ion concentration is given by
pH = pKa + log
This is known  as Henderson’s equation.
Basic buffer contains a weak base and its salt with strong acid. The OH– ion concentration is given by
pOH = pKb + log
Note : pH + pOH = 14 = pKw
A buffer has maximum buffer capacity when
= 1
In such case pH = pKa for acid buffer and pOH = pKb for basic buffer. For a buffer the ratio of concentration of salt to acid or base must lie between 0.1 and 10. Thus pH range for acid buffer is from pKa – 1 to pKa + 1. For basic buffer the pOH range is pKb – 1 to pKb + 1.

NEUTRALISATION

The combination of H+(aq) and OH(aq) furnished by an acid and a base in aqueous solution to produce undissociated water to maintain the value of KW at  1 × 10–14, is known as neutralisation.
HA + Water   H+(aq) + A(aq)
BOH + Water    OH (aq) + B+(aq)
H+(aq) + OH(aq) H2O
  1. Neutralisation of a strong acid with strong base : The strong acids and strong bases are almost completely ionised in aqueous solutions. When their solutions are mixed, the only effective change is the formation of unionised water.
Net reaction H+ (aq) + OH (aq)   H2O,
H = –13.7 kcal/mole
Neutralisation is complete. pH of resulting solution is 7.
  1. Neutralisation of a weak acid with a strong base : While sodium hydroxide is completely ionised in aqueous solution, the acid is only weakly ionised.
CH3COOH + NaOH    CH3COONa + H2O
Net reaction CH3COOH + OH  CH3COO + H2O,
H = –13.20 kcal/mole.
Solution remains basic (pH > 7) due to presence of OHions. Heat of neutralisation is less by 0.5 kcal/mol which is utilised for the dissociation of acetic acid.
  1. Neutralisation of a strong acid and weak base : The strong acid is completely ionised in aqueous solution and weak base is feebly ionised.
NH4OH + HCl    NH4Cl + H2O
Net reaction  NH4OH + H+    NH4+ + H2O,
H = –12.3 kcal/mole.
The solution remains acidic due to presence of H+ ions(pH < 7) 1.4 kcal/mole is utilised for dissociation of NH4OH.
  1. Neutralisation of a weak acid and a weak base : The acid and base are weakly ionised in aqueous solution.
CH3COOH + Water
     
Net reaction   H+  + OH   H2O,
H = –11.8 kcal/mol
The combination of H+ and OH take place to a limited extent only. This combination is sufficient to disturb., the equilibrium in (1) and (2) and dissociation is complete.
When neutralisation is completed, pH is almost 7. Heat evolved during neutralisation is even less, due to effecting the ionisation of weak acid and weak base.

HYDROGEN ION INDICATORS OR ACID BASE INDICATORS

Such indicators change their colour with the change of the pH of solution. They are either weakly acidic or basic in nature e.g. phenolphthalein which is colourless below pH 8 and pink at 9.8.

THEORY OF ACID-BASE INDICATORS

  1. Ostwald’s theory : The unionised molecule of the indicator has one colour while the ionised form has another colour.
HIn              H+      +      In
Colour A Colour B
Consider the ionisation of phenolphthalein which is weakly acidic in nature
HPh                  H+  +  Ph
Colourless                   Pink
Addition of a strong base will disturb the equilibrium (OH + H+  H2O) and more phenolphthalein will ionise giving pink solution. Addition of a strong acid will suppress the ionisation of phenolphthalein by common ion effect, solution will be colourless. The indicator is not suitable for titrating weak base like NH4OH against strong acid. The OH– ions furnished by weak base are insufficient to shift the equilibrium and pink colour does not appear just at the end point. Excess of weak base is required.

 

Action of methyl orange : Methyl orange is a weak base and is ionised as follows :
MeOH         Me+ + OH
Yellow            Red
Addition of a strong acid will disturb the equilibrium (H+ + OH  H2O) and more methyl orange will ionise giving red solution. Addition of a strong base will suppress the ionisation of methyl orange by common ion effect hence solution will be yellow in colour. The indicator is not suitable for titrating weak acid like CH3COOH against strong base. The H+ ions furnished by weak acid are not sufficient to shift the equilibrium and red colour does not appear just at the end point. Excess of weak acid is required.
  1. Quinonoid theory :
In the unionised form the indicator is generally in the benzenoid form which is less intense in colour and in the ionised form it is in the Quinonoid form which is more intense in colour.

 

Phenolphthalein :
Methyl Orange :

TYPES OF INDICATORS

  1. Internal : Which can be added to reacting substances further they can be acid base indicator (phenolphthalein, methyl orange etc.) redox  indicators, used for redox titrations (N-phenyl anthranilic acid) adsorption indicators (starch).
  2. External : Which can not be added to reacting substances e.g. potassium ferricyanide for titration of K2Cr2O7 vs. FeSO4(NH4)2SO4.6H2O.
  3. Radioactive : 8O18 and D2 for studying reaction mechanism.

 

Choice of indicator
  • For strong acid and strong base – methyl orange, phenolphthalein and Litmus.
  • For weak acid and strong base – phenolphthalein.
  • For strong acid and weak base – methyl orange.
  • For weak acid and weak base – Phenol red.

SALTS

The compounds containing positive and negative ions are known as salts. Their solutions may be acidic, basic or neutral. They are classified as
  1. Simple salts : Formed by neutralisation of an acid and a base. They may be further classified as follows.
    1. Normal salts : Salts not containing replaceable hydrogen or a hydroxyl group e.g. Na2SO4, KNO3,Al2 (PO4)3, CaCl2.
    2. Acidic salts : Salts containing replaceable hydrogen atoms NaHSO4, Na2HPO4, NaHCO3.
    3. Basic salts : Salts containing replaceable hydroxyl groups Zn(OH)Cl, Mg(OH)Cl, Fe(OH)2Cl.
  2. Double Salts : Formed by the combination of two simple salts
FeSO4.(NH4)2SO4.6H2O,  K2SO4. Al2(SO4)3. 24H2O.
  1. Complex Salts : Salts containing complex ion e.g. K4[Fe(CN)6],   [Cu(NH3)4]SO4.
  2. Mixed Salts : Salts containing more than one type of cation or anion e.g. ; NaKSO4.

 

All salts are strong electrolytes.

SOLVENTS

The substances which provide medium to carry out chemical reactions. They are classified as
  1. Protophilic : Such solvents have a tendency to accept proton e.g. H2O, liquor ammonia and alcohol etc.
  2. Protogenic : Such solvents have a tendency to donate proton e.g. HCl (l), glacial CH3COOH and H2O etc.
  3. Amphiprotic : Solvents capable of donating and accepting a proton e.g. H2O, alcohol etc.
  4. Aprotic : Solvents which neither donate nor accept a proton e.g. benzene, CS2, CCl4 etc.

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