Equilibrium-class 12 chemistry

Equilibrium in physical and chemical processes, dynamic nature of equilibrium, law of mass
action, 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 poly basic acids, acid strength, concept of pH, buffer solution, solubility product,
common ion effect , Arrhenius concept, brownsted lowry concept, acidic buffer, basic buffer, dissociation constant, equilibrium constant, see saw game, equilibrium in solid, liquid and gaseous state, solubility product.

1. CHEMICAL EQUILIBRIUM SHORT NOTES BY S.ARUNKUMAR

2. Equilibrium 12 Periods Equilibrium in physical and chemical processes, dynamic nature of equilibrium, law of mass action, 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 poly basic acids, acid strength, concept of pH, buffer solution, solubility product, common ion effect (with illustrative examples ). Deleted :Hydrolysis of salts (elementary idea), Henderson Equation.

3. Equilibrium   is actually a state, when forces from both the side become equal. According to chemistry: It is a point in a chemical reaction, when rate of forward reaction becomes equal to rate of backward reaction.  Or we can say, it is the state when concentration of reactants becomes equal to concentration of products .

4. Equilibrium state : H 2 O(l) H 2 O( vap ) Equilibrium in which the number of molecules leaving the liquid equals the number returning to liquid from the vapour . The system has reached equilibrium state at this stage. The mixture of reactants and products in the equilibrium state is called an equilibrium mixture. A dynamic equilibrium for a chemical reaction is the condition in which the rate of the forward reaction equals the rate of the reverse reaction.

5. Chemical equilibrium in a chemical reaction may be classified in three groups. (i)The reactions that proceed nearly to completion and only negligible concentrations of the reactants are left. In some cases, it may not be even possible to detect these experimentally. (ii) The reactions in which only small amounts of products are formed and most of the reactants remain unchanged at equilibrium stage. (iii) The reactions in which the concentrations of the reactants and products are comparable, when the system is in equilibrium.

6. Equilibrium involving ions in aqueous solutions which is called as ionic equilibrium. Equilibrium in Physical Processes : The most common equilibria involving physical processes are those which involve phase transformation : Solid    Liquid Liquid    Gas Solid     Gas Solid     Solution Gas     solution

7. Solid − liquid equilibrium : If some ice cubes along with some water at 0°C and normal atmospheric pressure are placed in a thermo flask so that no heat can enter or leave the system, the mass of ice and water is found to remain constant . At equilibrium Rate of melting of ice      Rate of freezing of water The temperature at which the solid and liquid form of a pure substance are in equilibrium at the atmospheric pressure is called the  normal freezing point on melting point  of that substance .

8. L iquid − vapour equilibrium : Consider a closed vessel connected to a manometer and having arrangement for evacuation and addition of liquid into it . Suppose the vessel is first evacuated. The level of mercury in both the limbs of the manometer will be same. Now suppose water is added into the vessel and the whole apparatus is allowed to stay at room temperature. It is observed that the level of mercury in the left limb of the manometer begins to fall and that in the right limb begins to rise. After sometime, the levels become constant. The system is then said to have attained equilibrium. The difference in the levels of mercury in the two limbs gives the equilibrium vapour pressure of the water at room temperature.

9. In the beginning, more and more of the water is changing into vapours i.e. number of water molecules in the vapour phase increases. Some of these molecules strike back on the surface of the water and are captured into it .The amount of water vapour becomes constant i.e.  now as much water changes into vapours the same amount of water vapour change back into the liquid. Rate of evaporation      Rate of condensation The liquid with higher vapour pressure is more volatile and has lower boiling point.

10. If we expose three watch glasses containing separately 1ml each of acetone, ethyl alcohol and water to atmosphere and repeat the experiment with different volumes of the liquids in a warmer room, it is observed that in all cases the liquid disappears and time taken to complete evaporation depends on (i) Nature of the liquid (ii) The amount of the liquid (iii)The temperature.

11. Solid     Vapour Equilibrium This type of equilibrium is attained for solids which undergo sublimation . If solid iodine is placed in a closed vessel ,  violet vapours start appearing in the vessel whose intensity increases with time and ultimately it becomes constant. At this stage ,equilibrium is attained i.e. Rate of sublimation  of solid I 2  to form vapour    Rate of condensation of I 2  vapour to give solid I 2 I 2  ( s )     I 2  (vapour)

12. Solid – Solution Equilibrium : Suppose more and more of sugar is added into a fixed volume of water at room temperature and stirred thoroughly with a glass rod. The sugar will keep on dissolving but then a stage will come when no more sugar dissolves. Instead it settles down at the bottom. The solution is now said to be saturated and the state of equilibrium. At this stage, as many molecules of sugar from the surface of the undissolved sugar go into the solution, the same number of molecules of sugar from the solution are deposited back on the surface of the undissolved sugar .

13. As a result ,the amount of an dissolved sugar and the concentration of sugar in solution remains constant. Rate of dissolution    Rate of precipitation Rate of dissolution of sugar=Rate of Crystallisation of sugar The amount of the solid in grams that dissolves in 100 gram of the solvent to form a saturated solution at a particular temperature is called the  solubility  of that solid in the given solvent at that temperature.

14. Gas − solution equilibrium: This type of equilibrium is found in soda water bottle. The equilibrium that exist within the bottle is: CO 2  ( g )     CO 2  ( solution) The amount of gas dissolved is governed by  Henry’s law  which states that : The mass of a gas dissolved in a given mass of a solvent at any temperature is directly proportional to the pressure of the gas above the solvent, i.e. m ∝ p or m= kp

15. In a sealed soda water bottle ,the pressure of the gas is very high above the liquid ,so the mass of the gas dissolved is also high. As soon as the bottle is opened ,the pressure tends to decrease to atmospheric pressure ,so the solubility decreases, i.e. the dissolved gas escapes out.

16. For solid  liquid equilibrium, there is only one temperature(melting point) at 1 atm at which two phases can coexist. For liquid Vapour equilibrium, the vapour pressure is constant at a given temperature. For dissolution of solids in liquids, the solubility is constant at a given temperature. For dissolution of gases in liquids, the concentration of a gas in liquid is proportional to the pressure of the gas over the liquid. Equilibrium will not reach in the open system.

17. General characteristics of physical equilibrium The measurable properties become constant. It can be established only in closed vessel. At equilibrium, the opposing forces become equal. The equilibrium is dynamic in nature .That is the reaction keeps on going only the rate becomes constant. At equilibrium, the concentration becomes constant. The magnitude of equilibrium value, gives indication about the extent of reaction.

18. A reaction in which not only the reactants react to form the products under certain conditions but also the products react to form reactants under the same conditions is called a reversible reaction .

19. Irreversible reactions If a reaction cannot takes place in the reverse direction i.e. the products formed do not react to give back the reactants under the same conditions, it is called  irreversible reaction. It is represented by putting a single arrow between the reactants and products, pointing from reactants towards products. AgNO 3  ( aq ) + NaCl ( aq ) ——> AgCl ( s ) + NaNO 3  ( aq ) BaCl 2  ( aq ) + Na 2 SO 4  ( aq )—–> BaSO 4  ( s ) + NaCl ( aq ) 2Mg ( g ) + O 2  ( g ) —-> 2MgO ( s )

20. Concept of Chemical Equilibrium :

21. 1 ) In the beginning the concentration of A and B are maximum and the concentration of C and D are minimum. 2) As the reaction proceeds, the concentration of A and B are decreasing with passage of time whereas the concentration of C and D are increasing. 3) The rate of forward reaction is decreasing while the rate of backward reaction goes on increasing . 4 ) A stage comes , when the rate of forward reaction becomes equal to rate of backward reaction. The reaction is then said to be in a  state of chemical equilibrium.

22. In the Haber’s process ,starting with definite amount of N 2  and H 2  and carrying out the reaction at a particular temperature, when equilibrium is attained ,the concentration of N 2 , H 2  and NH 3  become constant. If the experiment is repeated by taking deuterium in place of H 2  but with the same amounts and exactly similar as before equilibrium is attained containing D 2  and ND 3  in place of H 2  and NH 3  but in the same amount. If the two reaction mixtures are mixed ,then after sometime, it is found that the concentration of ammonia and hydrogen are same except that now all forms of ammonia and all forms of hydrogen are present. This shows that an equilibrium , the reaction is still going on i.e. equilibrium is dynamic in nature.

23. When the equilibrium is reached, the concentration of each of the reactants and products become constant. In the reaction between H 2  and I 2  to form HI , the colour becomes constant because the concentration of H 2  , I 2  and HI becomes constant. As much of the reactants react to form the products ,the same amount of products react to give back the reactants in the same time. Hence the equilibrium is dynamic in nature and not static.

24. Law of mass action : It states that: Rate of a reaction is directly proportional to concentration of reactants raised to their respective moles. Consider a reaction: aA + bB --> cC + dD According to this law : Where ‘K’ is rate constant or velocity constant. We can attain equilibrium from both sides : For forward reaction : Therefore,

25. For backward reaction : Therefore At equilibrium: Rate of forward reaction = rate of backward reaction K f [A] a [B] b  = K b [C] c [D] d But, K c = K f /K b  (where K c  is equilibrium constant) K c  = [A] a [B] b /[ C] c [D] d At particular instant of time: The equilibrium constant is called as reaction quotient (Q). That is: Q=[C] c [D] d  /  [A] a [B] b At equilibrium Q= K c For example: In case of gases, we can also express it in terms of their partial pressures :

26. At equilibrium, The rate of forward reaction = Rate of backward reaction.

27. At equilibrium, The rate of forward reaction = Rate of backward reaction

28. Law of chemical equilibrium : At a given temperature, the product of concentrations of the reaction products raised to the respective stoichiometric coefficient in the balanced chemical equation divided by the product of concentration of the reactants raised to their individual stoichiometric coefficients has a constant value.

29. The equlibrium constant for a general reaction, aA+bB  cC+dD is expressed as, K c = Where [A], [B], [C], [D] are the equilibrium concentrations of the reactants and products. Equilibrium constant for the reaction . 4NH 3 (g)+5O 2  4NO(g)+6H 2 O(g ) is written as  

30. Molar concentrations of different species is indicated by enclosing these in square bracket and as mentioned above, it is implied that these equilibrium concentrations. While writing expression for equilibrium constant, symbol for phases( s,l,g ) are generally ignored. Let us write the equilibrium constant for the reaction, =x

31. The equilibrium constant for the reverse reaction is the inverse of the equilibrium constant for the reaction in the forward direction. K 1 C = Equilibrium constant for the reverse reaction is the inverse of the equilibrium constant for the reaction in the forward direction.  

32. As K c and K 1 C have different numerical values, it is important to specify the form of the balanced chemical equation when quoting the value of an equilibrium constant.

33. Characteristics of Equilibrium constant : 1.The value of the equilibrium constant for a particular reaction is always constant depending only upon the temperature of the reaction and is independent of the concentration of the reactant with which we start or the directions from which the equilibrium is approached .

34. If the reaction is reversed, the value of equilibrium constant is reversed .

35. If the equation having equilibrium constant K is divided by 2 ,the equilibrium constant for the new equation is the square root of K .

36. If the equation having equilibrium constant is multiplied by 2 ,the equilibrium constant for the new equation is a square of equilibrium constant .

37. If the equation having equilibrium constant is written in 2 steps then equilibrium then K= K 1  × K 2

38. The value of equilibrium constant is not affected by the addition of a catalyst. This is because the catalyst increases the speed of forward reaction and backward reaction to same extent.

39. Types Of Chemical Equilibria 1) Homogeneous Equilibria When in an equilibrium reaction, all the reactants and the products are present in the same phase , it is called homogeneous equilibrium. Type 1 : In which the number of moles of products is equal to the number of moles of reactants. H 2  + I 2     2HI N 2  + O 2     2NO CO + H 2 O    CO 2  + H 2

40. Type 2 : In which the number of moles of products is not equal to the number of moles of reactants N 2  + 3H 2       2NH 3 2SO 2  + O 2     2SO 3 PCl 5     PCl 3  + Cl 2 The reactions in the liquid phase are: CH 3 COOH + C 2 H 5 OH    CH 3 COOC 2 H 5  + H 2 O

41. Equilibrium constant in gaseous systems :

42. Heterogeneous Equilibria : When in an equilibrium reaction, all the reactants and the products are present in two or more than two phases , it is called heterogeneous equilibrium. CaCO 3  ( s )     CaO ( s ) + CO 2  ( g ) 3Fe ( s ) + 4H 2 O ( g )    Fe 3 O 4  (s ) + 4H 2  ( g )

43. Expression and Units Of Equilibrium Constant: For Homogeneous Equilibrium: CH 3 COOH ( l ) + C 2 H 5 OH ( l )     CH 3 COOC 2 H 5  ( l ) + H 2 O ( l ) NH 3  ( g ) + 5O 2  ( g )    4NO ( g ) + 6H 2 O ( g )

44. For Heterogeneous Equilibrium: Units of Equilibrium Constant: Keq = ( mol L-1 ) c+d / ( mol L-1 ) a+b Keq =( mol L-1 ) ( c+d ) – (a + b ) Keq = ( mol L-1 ) Δ n

45. Kp = ( atm ) ( c+d ) / ( atm ) a+b Kp = (bar) ( c+d ) / (bar) a+b Kp = ( atm or bar )( c+d ) – ( a+b ) Kp = ( atm ) Δn or (bar) Δn If Δn =0 i.e. number of moles of products = number of moles of reactants , Keq or Kp will have no units. Units Of Equilibrium Constant

46. For Ex: 1) H 2  ( g ) + I 2  ( g )  2HI ( g ) Δn =0 , K c  or K p  will have no units. 2) N 2  ( g ) + 3H 2  ( g )  2NH 3  ( g ) Δn = 2 – ( 1+ 3 ) = -2 The units will be ( mol L -1  ) -2  or atm -2  or bar -2

47. Important features of equilibrium constant : Expression for equilibrium constant is applicable only when concentration of the reactants and products have attained constant value at equilibrium state. The value of equilibrium constant is independent of initial concentrations of the reactants and products. Equilibrium constant is temperature dependent having one unique value for a particular reaction represented by a balanced equation at given temperature. The equilibrium constant for the reverse reaction is equal to the inverse of the equilibrium constant for the forward reaction.

48. Predicting the extent of reaction: The magnitude of the equilibrium constant gives an idea of the relative amount of the reactants and the products. a)  larger value of the equilibrium constant  ( > 10 3  ) shows that  forward reaction is favoured  i.e. concentration of products is much larger than that of the reactants at equilibrium. For Ex: 1) H 2  ( g ) + Br 2  ( g )      2HBr ( g)   Kc = 5.4 × 10 18 2) H 2  ( g ) + Cl 2  (g )      2HCl ( g )   K c  = 4 × 10 31

49. 3) H 2  ( g ) + ½ O 2  ( g )    H 2 O ( g ) Kc = 2.4 × 10 47 This shows that at equilibrium , concentration of the products  is very high , i.e. reaction go almost to completion. b)  Intermediate value of K  ( 10 -3  to 10 3  ) show that the  concentration of the reactants and products are comparable . For Ex: Fe 3 +  ( aq ) + SCN – ( aq )      (Fe(SCN)) 2 + ( aq )   Kc = 138 at 298 K 2) H 2  ( g) + I 2  ( g )   2HI ( g )  Kc = 57 at 700 K

50. c)  low value of K  ( < 10  -3 ) shows that  backward reaction is favoured  i.e. concentration of reactants is much larger than that of products i.e. the reaction proceeds to a very small extent in the forward direction . For Ex: N 2  ( g ) + O 2  ( g )     2NO ( g ) Kc =4.8 × 10  -31  at 298 K 2) H 2 O ( g )    H 2  ( g ) + ½ O 2  ( g ) Kc = 4.1 × 10  -48

51. Predicting the direction of a reaction: At any stage of the reaction, other than the stage of chemical equilibrium, concentration ratio, as given by the expression for the law of chemical equilibrium , is called  concentration quotient or reaction quotient. It is usually represented by  Q c  or Q . Thus,

52. 1) If Q= K c  , the reaction is in equilibrium 2) If Q > K c  , Q will tend to decrease so as to become equal to K. As a result, the reaction will proceed in the backward direction. 3) If Q < K c , Q will tend to increase so as to become equal to K. As a result, the reaction will proceed in the forward direction. Consider the gaseous reaction of H 2 with I 2 , H 2 (g) + I 2 (g) ⇌ 2HI(g ); Kc = 57.0 at 700 K Suppose we have molar concentrations [H 2 ] t=0.10M, [I 2 ] t = 0.20 M and [HI] t = 0.40 M

53. Thus, the reaction quotient, Qc at this stage of the reaction is given by, Qc = [HI]t 2 / [H 2 ] t [I 2 ] t = ( 0.40) 2/ (0.10)×(0.20) = =8.0 Now, in this case, Qc (8.0) does not equal Kc (57.0), so the mixture of H2 (g), I2 (g) and HI(g) is not at equilibrium. If Qc < Kc , net reaction goes from left to right • If Qc > Kc , net reaction goes from right to left. • If Qc = Kc , no net reaction occurs.

54. Calculating equilibrium concentration: Knowing the initial concentration of reactants, equilibrium concentration of all the reactants and products can be calculated as: 1) Write the balanced equation for the reaction. 2) Assume x as the amount of the reactants reacted or a product formed. 3) Calculate the equilibrium concentration of each reactant and product from its stoichiometry of the equation. 4) Write expression for K c  or K p . 5)Substitute equilibrium concentration and calculate x. 6) Check the result by substituting calculated values of equilibrium concentration to get the value of  K c  or K p

55. Relationship between equilibrium constant reaction quotient Q and Gibbs energy G.

56. Factors affecting Equilibria : Le- Chatelier’s principle : It states that a change in any of the factors that determine the equilibrium conditions of a system will cause the system to change in such a manner so as to reduce or to counteract the effect of the change.

57. Effect of concentration change : When the concentration of any of the reactants or products in a reaction at equilibrium is changed, the composition of the equilibrium mixture changes so as to minimize the effect of concentration changes. Let us take the reaction H 2 (g)+I 2 (g)  2HI(g) If H 2 is added to the reaction mixture at equilibrium, then the equilibrium of the reaction is disturbed. In order to restore it, the reaction proceeds in a direction where in H 2 is consumed, more H 2 and I 2 react to form HI and equilibrium shifts in forward direction.

58. Effect of concentration-An experiment : According to Le Chatelier's Principle, the system will react to minimize the stress. Since Fe 3+  is on the reactant side of this reaction, the rate of the forward reaction will increase in order to "use up" the additional reactant. This will cause the equilibrium to  shift to the right , producing more FeSCN 2+ . For this particular reaction we will be able to see that this has happened, as the solution will become a darker red color . There are a few different ways we can say what happens here when we add more Fe 3+ ; these all mean the same thing: equilibrium shifts to the right equilibrium shifts to the product side the forward reaction is favored

59. I f we add more FeSCN 2 + Again, equilibrium will shift to use up the added substance. In this case, equilibrium will shift to favor the  reverse  reaction, since the reverse reaction will use up the additional FeSCN 2+ . equilibrium shifts to the left equilibrium shifts to the reactant side the reverse reaction is favoured

60. Effect of pressure change: When pressure increases, equilibrium position shifts in the direction that decreases the total moles of gas When pressure decreases, equilibrium position shifts in the direction that increases the total moles of gas. Consider the reaction, Here, 4 mol of gaseous reactant(CO+3H2) become 2 mol of gaseous products (CH4 +H2O).

61. Suppose equilibrium mixture kept in a cylinder fitted with a piston at constant temperature is compressed to one half of its original volume. Then, total pressure will be doubled . The reaction is no longer at equilibrium. The direction in which the reaction goes to re-establish equilibrium can be predicted by applying the Le- Chatelier’s principle. Since pressure has doubled, the equilibrium now shifts in the forward direction , a direction in which the number of moles of the gas or pressure decreases .

62. Effect of inert gas addition: If the volume is kept constant and an inert gas such as argon is added which does not take part in the reaction, the equilibrium remains undisturbed. It is because the addition of an inert gas at constant volume does not change the partial pressures or the molar concentrations of the substances involved in the reaction. The reaction quotient changes only if the added gas is a reactant or product involved in the reaction .

63. Effect of temperature change: Whenever an equilibrium is disturbed by a change in the concentration, pressure or volume, the composition of the equilibrium mixture changes because the reaction quotient Qc no longer equals to Kc. When a change in temperature occurs, the value of equilibrium constant Kc is changed. The equilibrium constant for an exothermic reaction (negative H) decreases as the temperature increases.

64. The equilibrium constant for an endothermic reaction (positive H) increases as the temperature increases. Ex: Production of Ammonia Is an exothermic process. According to Le chatelier’s principle, raising the temperature shifts the equilibrium to left and decreases the equilibrium concentration of ammonia.Low temperature is favourable for high yield of ammonia, but at very low temperature slow down the reaction so we use the catalyst.

65. Effect of Temperature-An experiment :

66. The first tube is placed in a hot water bath. The second is placed in ice-water. The third is placed in a flask filled with dry-ice/acetone slush. The fourth is placed in liquid nitrogen. NO 2  is brown and N 2 O 4  is colourless . The intensity of the brown colour decreases as the temperature decreases. Therefore, a decrease in temperature yields and increase in N 2 O 4 . Equilibrium is shifted to the N 2 O 4  side upon a decrease in temperature . 2 NO 2  (g) <=> N 2 O 4  (g)

67. Effect of a catalyst: A catalyst increase the rate of the chemical reaction by making available a new low energy pathway for the conversion of reactants to products. It increase the rate of forward and reverse reactions that pass through the same transition state and does not affect equilibrium. Catalyst lowers the activation energy for forward and reverse reactions by exactly the same amount. Catalyst does not affect the equilibrium composition of a reaction mixture.

68. In the manufacture of sulphuric acid by contact process. Platinum or V2O5 is used as a catalyst to increase the rate of the reaction. 2SO 2 (g)+O 2 (g)  2SO 3 (g)

69. We classify substances in to two types: Electrolytes Non electrolytes Electrolytes : The substances which dissociate into ions in solution on passing current. For example: AB  A +  + B - Non electrolytes : The substances that do not dissociate into ions, when current is passed through them. That is AB à on passing current , nothing happens.

70. Further we can classify electrolytes as: Strong electrolyte Weak electrolyte Strong electrolyte : The substances which completely dissociate into ions, when current is passed through them. It means if take an example of AB then on passing current it dissociate completely without leaving AB anywhere in solution . Example: HCl,H 2 SO 4   etc

71. A strong electrolyte is defined as a substance which dissociates almost completely into ions in aqueous solution and hence is a very good conductor of electricity. For Ex:  NaOH , KOH, HCl , H 2 SO 4  , NaCl HCl + H 2 O ———> H 3 O +  + Cl − NaOH + aq ———> Na +  ( aq ) + OH –   ( aq )ionisation constant .

72. A weak electrolyte is defined as a substance which dissociates to small extent in aqueous solution and hence conduct electricity also to a small extent. For Ex: NH 4 OH , CH 3 COOH CH 3 COOH + H 2 O  CH 3 COO ‾ + H 3 O + NH 4 OH + aq     NH 4 +  ( aq ) + OH‾ ( aq )

73. The ionisation of a weak electrolyte , AB , is represented as : AB ( s ) + aq     A +  ( aq ) + B‾ ( aq ) Such and equilibrium is called Ionic Equilibrium between the ions and the undissociated electrolyte. Applying the law of chemical equilibrium to the above equation :

74. where K is the  ionisation constant. When an ionic compound is dissolved in water ,the ions which are already present in the solid compound separates out. The process is called  dissociation. When a neutral molecule which does not contain ions but when dissolved in water splits to produce ions in the solution, the process is called  ionization. An ionic compound is a cluster of positively and negatively charged ions held together by electrostatic force of attraction. When such an ionic compound is put into water ,the high dielectric constant of water reduces the electrostatic forces of attraction. Thus, ions become free to move in the solution.

75. A cids and bases and Salts : Acid  is a substance whose aqueous solution possessed the following characteristic properties 1)conduct electricity 2) reacts with active metals like zinc ,magnesium to give hydrogen 3) turns blue litmus to red 4) has a sour taste 5)whose acidic properties disappear on reaction with a base Base  is a substance whose aqueous solution possessed the following characteristic properties: 1)turns red litmus blue 2) has a bitter taste 3) has a soapy touch

76. Arrhenius concept of acids and bases : Arrhenius in 1884 put forward a theory popularly known as  Arrhenius theory of ionization When an electrolyte is dissolved in water ,it dissociates into positively and negatively charged ions. An  acid  is defined as a substance which contains hydrogen and which when dissolved into water gives hydrogen ions (H + ) Substances like HCl , HNO 3  , H 2 SO 4  containing hydrogen, when dissolved in water dissociates completely into H +  ions and negative ions.

77. HCl ———-> H +  + Cl ‾ H 2 SO 4  ————–> 2H +  + SO 4 2- HNO 3  ———> H +  + NO 3 ‾ Such acids are called  strong acids. Substance like acetic acid ,carbonic acid, phosphoric acid when dissolved in water dissociate into ions to a small extent . CH 3 COOH    CH 3 COO‾+ H + H 2 CO 3     2H +  + CO 3 2‾ H 3 PO 4      3H +  + PO 4 3‾ Such acids are called  weak acids .

78. A base is defined as a substance which contain hydroxyl group, when dissolved in water gives hydroxide ion (OH‾). Substances like NaOH and KOH containing hydroxyl group when dissolved into water, dissociate completely to give OH‾ ions as: NaOH ————> Na +  + OH‾ KOH —————> K +  + OH‾ These are called strong bases. Substances like NH 4 OH , Ca (OH) 2  , Mg(OH) 2  , Al(OH) 3   dissociates to a small extent as :

79. NH 4 OH    NH 4   +  + OH‾ Ca (OH) 2    Ca 2 +  + 2OH¯ These are called  weak bases . H +  ion is simply a proton which is very small in size. It has a strong electric field. It takes up a lone pair of electrons from water molecule to form called  hydronium ion, H 3 O + . The  dissociation of an acid  in water may be expressed as : HCl + water    H +  ( aq ) + Cl ‾ ( aq ) CH 3 COOH + water    H +  ( aq ) + CH 3 COO‾ ( aq )

80. The  dissociation of base  may be expressed as NaOH + water  Na +  ( aq ) + OH‾ ( aq ) NH 4 OH  NH 4 +( aq )  + OH‾( aq ) When an acid and base are mixed together, they react to form salt and water. The reaction is called  neutralization . NaOH + HCl ———–> NaCl + H 2 O Neutralization  is defined as the process in which hydrogen ions given by the acid and hydroxyl Ion given by the base combine to form unionized molecules of water.

81. Advantage of Arrhenius concept   : The Arrhenius concept of acids and bases was able to explain a number of phenomena like neutralization, salt hydrolysis ,strength of acids and bases . Limitation of   Arrhenius concept  1) Nature of H +  and OH‾ ion According to Arrhenius concept ,acids and bases were defined as substances which gave H +  ions and OH‾ ions in aqueous solution. But these ions cannot exist as such in aqueous solution but exist as hydrated ion.

82. 2) Inability to explain acidic and basic character of certain substances Arrhenius concept says that an acid must contain hydrogen and a base must contain hydroxyl group. However a number of substances like NH 3  , Na 2 CO 3 , CaO are known to be basic but do not contain any hydroxyl group. Similarly a number of substances like CO 2 , SO 2  ,SO 3   etc are known to be acidic but do not contain any hydrogen. NH 3  (g) + H 2 O     NH 4 +  (aq) + OH‾ (aq) CaO+ H 2 O    Ca 2 +  (aq) + 2OH‾ (aq)

83. 3)Inability to explain the reaction between an acid and base in the absence of water. NH 3  (g) + HCl (g) ——> NH 4 Cl (s) CaO (s) + SO 3 (g) ——–>CaSO 4  (s )

84. Bronsted - lowry concept of acid and bases: Bronsted and lowry in 1923 proposed concept of acids and bases. Acid is defined as a substance which has the tendency to give a proton and base is defined as a substance which has a tendency to accept a proton. An  acid  is a  proton donor  whereas  base  is a  Proton acceptor . HCl + H 2 O     H 3 O+ + Cl‾ CH 3 COOH + H 2 O   H 3 O +  (aq ) + CH 3 COO‾ (aq)

85. 1.HCl and CH 3 COOH are acids because they donate a proton to H 2 O. 2) NH 3  and CO 3 2‾   are bases because they accept a proton from water. 3)Not only molecules but even the ions can act as acids or bases. 4)Water is accepting a proton and hence is base. Water acts both as an acid as well as base and hence is called  amphoteric .

86. 5) Bronsted – lowry definition of acid and bases are not restricted to aqueous solution. HCl is acid because it gives a protons and NH 3  is a base because it accepts the proton. 6) The presence of hydroxyl group is not essential for a substance to act as a base. The only requirement is that it should have tendency to accept a proton. 7) A substance acts as an acid i.e. gives a proton only when another substance to accept the proton i.e. a base is present.

87. 8) The reverse reaction are also acid – base reaction. In reaction  HCl +H 2 O    H 3 O+ + Cl ‾ ,the reverse process H 3 O +  can give a proton and hence is an acid while Cl ‾ can accept the proton and hence is a base. There are two acid base pair in reaction.These are HCl-Cl ‾ and H 3 O +  -H 2 O .These acid base pair are called  conjugate acid – base pair.

88. Advantages of Bronsted-lowry concept 1) Bronsted-lowry concept is not limited to molecules but includes even the ionic species to act as acids or bases. 2) It can explain the basic character of the substances like Na 2 CO 3  , NH 3   etc on the basis that they are proton acceptor. 3) It can explain the acid base reaction in the non aqueous medium or even in the absence of solvent.

89. Disadvantage of Bronsted-lowry concept 1) It cannot explain the reaction between acidic oxides like CO 2  , SO 2  , SO 3  and the basic oxides like CaO , BaO , MgO which takes place even in the absence of solvent. 2) Substances like BF 3  , AlCl 3   etc do not have any hydrogen and hence cannot give a proton but are known to behave as Lewis acid.

90. Lewis concept of acids and bases: G.N. Lewis in the same year i.e. 1923 ,proposed: An  acid  is defined as substance which is capable of accepting a pair of electrons and  base  is defined as a substance which is capable of donating an unshared pair of electrons. An acid is an electron pair acceptor while a base is an electron pair donor. The reaction between an acid and base amount to the formation of a co-ordinate Bond or dative bond between them:

91. For Ex: Reaction between BF 3  and NH 3  Since NH 3  can donate a lone pair of electron while BF 3  can accept a pair of electron , NH 3  is a base and BF 3  is an acid.

92. Advantage of Lewis concept It could explain even those acids and base reactions which could not be explained by others concepts . Drawback   of Lewis concept : 1) Lewis concept is so general that it considers every reaction forming a co-ordinate bond to be acid base reaction. This however, may not be always true. 2) The requirement in Lewis concept is the formation of a co-ordinate bond between the acid and base. However acids like HCl and H 2 SO 4  do not form any co-ordinate Bond and therefore should not be acids according to this concept.

93. IONIZATION OF ACIDS AND BASES : H ydrochloric acid ( HCl ), hydrobromic acid ( HBr ), hyrdoiodic acid (HI), nitric acid (HNO3) and sulphuric acid (H2SO4) are termed strong because they are almost completely dissociated into their constituent ions in an aqueous medium, thereby acting as proton (H+) donors. Similarly , strong bases like lithium hydroxide ( LiOH ), sodium hydroxide (NaOH ), potassium hydroxide (KOH), caesium hydroxide ( CsOH ) and barium hydroxide Ba(OH)2 are almost completely dissociated into ions in an aqueous medium giving hydroxyl ions , OH–.

94. According to Arrhenius concept T hey are strong acids and bases as they are able to completely dissociate and produce H3O+ and OH– ions respectively in the medium. Brönsted - Lowry concept of acids and bases, wherein a strong acid means a good proton donor and a strong base implies a good proton acceptor. Consider , the acid-base dissociation equilibrium of a weak acid HA. HA(aq) + H2O(l)  H3O +(aq) + A–(aq) conjugate conjugate acid base

95. Consider the two acids HA and H3O+ present in the above mentioned acid-dissociation equilibrium. We have to see which amongst them is a stronger proton donor. Whichever exceeds in its tendency of donating a proton over the other shall be termed as the stronger acid and the equilibrium will shift in the direction of weaker acid. Say , if HA is a stronger acid than H3O+, then HA will donate protons and not H3O+, and the solution will mainly contain A– and H3O+ ions. The equilibrium moves in the direction of formation of weaker acid and weaker base.

96. because the stronger acid donates a proton to the stronger base. It follows that as a strong acid dissociates completely in water, the resulting base formed would be very weak i.e., strong acids have very weak conjugate bases. Strong acids like perchloric acid (HClO4), hydrochloric acid ( HCl ), hydrobromic acid ( HBr ), hydroiodic acid (HI ), nitric acid (HNO3) and sulphuric acid (H2SO4 ) will give conjugate base ions ClO4 –, Cl , Br –, I–, NO3 – and HSO4 – , which are much weaker bases than H2O . Similarly a very strong base would give a very weak conjugate acid.

97. On the other hand, a weak acid say HA is only partially dissociated in aqueous medium and thus, the solution mainly contains undissociated HA molecules . Typical weak acids are nitrous acid (HNO 2 ), hydrofluoric acid (HF) and acetic acid (CH 3 COOH ). It should be noted that the weak acids have very strong conjugate bases . For example , NH 2 –, O 2 – and H– are very good proton acceptors and thus, much stronger bases than H 2 O .

98. The Ionization Constant of Water and its Ionic Product: Some substances like water are unique in their ability of acting both as an acid and a base. In presence of an acid, HA it accepts a proton and acts as the base while in the presence of a base, B– it acts as an acid by donating a proton. In pure water, one H2O molecule donates proton and acts as an acid and another water molecules accepts a proton and acts as a base at the same time. The following equilibrium exists:

99. H 2 O(l) + H 2 O(l )  H 3 O+(aq) + OH–(aq) acid base conjugate .A conjugate.B The dissociation constant is represented by, K = [H 3 O+] [ OH–] / [H 2 O] The concentration of water is omitted from the denominator as water is a pure liquid and its concentration remains constant. [H 2 O] is incorporated within the equilibrium constant to give a new constant, K w, which is called the ionic product of water . K w = [H+][OH–]

100. The concentration of H+ has been found out experimentally as 1.0 × 10–7 M at 298 K. And, as dissociation of water produces equal number of H+ and OH– ions, the concentration of hydroxyl ions, [OH–] = [H+] = 1.0 × 10–7 M. Thus, the value of K w at 298K, K w = [H 3 O+][ OH–] = (1 × 10–7)2 = 1 × 10 –14 M 2

101. The value of K w is temperature dependent as it is an equilibrium constant. The density of pure water is 1000 g / L and its molar mass is 18.0 g /mol. From this the molarity of pure water can be given as , [ H2O] = (1000 g /L)(1 mol/18.0 g) = 55.55 M. Therefore , the ratio of dissociated water to that of undissociated water can be given as: 10 –7 / (55.55) = 1.8 × 10–9 or ~ 2 x 10 –9 ( thus, equilibrium lies mainly towards undissociated water ) We can distinguish acidic, neutral and basic aqueous solutions by the relative values of the H 3 O+ and OH– concentrations: Acidic: [H 3 O+] > [ OH–] Neutral: [H 3 O+] = [ OH–] Basic : [H 3 O+] < [ OH–]

102. The pH Scale : The pH of a solution is defined as the negative logarithm to base 10 of a (H+ ) of hydrogen ion . Relationship between pH and pOH :

103. Ionization of weak acid : When acetic acid is dissolved in water, it dissociates partly into H +  and H 3 O +  and CH 3 COO‾ ions as: CH 3 COOH + H 2 O    CH 3 COO‾ + H 3 O + In dilute solution , concentration of water is constant. The product of K and constant

104. The product of K and     is denoted by  K a , the  ionization constant or dissociation constant of the acid . If C represents the initial concentration of the acid in moles L -1 and  α , the  degree of dissociation  , then equilibrium concentration of the ions ( H 3 O +  and CH 3 COO‾ ) is equal to Cα and that of the undissociated acetic acid = C ( 1- α ) i.e. we have

105. In case of weak electrolyte, The value of α is very small and can be neglected in comparison to 1 i.e. 1-α =1.Hence we get α = √ Ka / C If V is the volume of the solution in litres containing 1 mole of the electrolyte , C = 1/ V . Hence , we have, α = √ Ka  × V

106. For a weak electrolyte , the degree of ionisation is inversely proportional to the square root of molar concentration  or directly proportional to the square root of volume containing one mole of solute . This is called  Ostwald’s dilution law .

107. Relate K a , K w  and K b

108. Di-polybasic Acids and Di- and Polyacidic bases :

109. Factors affecting acid strength:

110. Common Ion Effect: If to an ionic equilibrium, AB    A +  + B‾ , a salt containing a common ion is added, the equilibrium shifts in the backward direction. This is called  common Ion effect . Acetic acid being a weak acid, ionizes to a small extent as: CH 3 COOH     CH 3 COO‾ + H +   

111. To this solution , suppose the salt of this weak acid with a strong base is added. It ionizes almost completely in the solution as follows and provides the common acetate ions: CH 3 COONa <————> CH 3 COO‾ + Na + the concentration of CH 3 COO‾  ions increases and by Le Chatelier’s principle, the dissociation equilibrium shift backwards i.e. dissociation of acetic acid is further suppressed. 2) To the solution of the weak base , NH 4 OH , NH 4 Cl is added which provides the common NH 4 +  ions ,dissociation of NH 4 OH is suppressed:

112. NH 4 OH      NH 4 +  +OH‾ NH 4 Cl ———->NH 4 +  + Cl ‾ 3) To the solution of silver chloride in water, if NaCl is added which provide the common Cl ‾ ions, the solubility AgCl decreases : AgCl   Ag +  + Cl ‾ NaCl ——–> Na +  + Cl ‾ Increase in the concentration of Cl ‾ ions shifts the equilibrium in the backward direction i.e. some solid AgCl separates out. If to the solution of a weak electrolyte, which ionises to a small extent ,a strong electrolyte having a common ion is added which ionizes almost completely, the ionization of the weak electrolyte is further suppressed . If to the solution of a sparingly soluble salt, if a soluble salt having a common ion is added ,the solubility of the sparingly soluble salt further decreases.

113. Salt Hydrolysis Salt hydrolysis  is defined as the process in which a salt reacts with water to give back the acid and the base. Salt +water ———-> Acid + Base BA + H 2 O ———> HA + BOH All salts are strong electrolytes and thus ionize completely in the aqueous solution. (1) If the  acid  produced is  strong  and the  base  produced is  weak. B +  + A‾ + H 2 O ——-> H +  + A‾ + BOH or B +  + H 2 O ——-> H +  + BOH In this case the cation reacts with water to give an acidic solution. This is called  cationic hydrolysis. (2) If the  acid  produced is  weak  and the  base  produced is  strong. B +  + A‾ + H 2 O ——-> HA + B +  + OH‾

114. A‾ + H 2 O ——-> HA + OH‾ In this case the anion reacts with water to give basic solution. This is called  acidic hydrolysis. Salt hydrolysis  may be defined as the reaction of the cation or the anion of the salt with water to produce acidic or basic solution.

115. 1) Salts of strong acid and strong base NaCl , NaNO 3 , Na 2 SO 4 , KCl , KNO 3  , K 2 SO 4 NaCl + H 2 O ——-> NaOH + HCl Na +  + Cl ‾ + H 2 O ——-> Na +  + OH‾ + H +  + Cl ‾ H 2 O ——->OH‾ + H + It involves only ionization of water and no hydrolysis . So the solution is neutral . The salts of strong acids and strong bases do not undergo hydrolysis and the resulting solution is neutral .

116. Salts of weak acid and strong bases CH 3 COONa, Na 2 CO 3 , K 2 CO 3 , Na 3 PO 4 CH 3 COONa + H 2 O  CH 3 COOH + NaOH CH 3 COO‾ + Na +  + H 2 O  CH 3 COOH + Na +  + OH‾ As it produces OH‾ ions, the solution of such a salt is alkaline in nature . Salts of strong acid and weak base: NH 4 Cl , CuSO 4 , NH 4 NO 3  , AlCl 3 , CaCl 2 NH 4 Cl +  H 2 O      NH 4 OH + HCl NH 4 +  + Cl‾ + H 2 O     NH 4 OH + H +  + Cl‾ NH 4 +  + H 2 O ——->     NH 4 OH + H +

117. As it produces H +  ions, the solution of such a salt is acidic in nature. (4) Salts of weak acid and weak base: CH 3 COONH 4   +  H 2 O     CH 3 COOH + NH 4 OH It involves both anionic and cationic hydrolysis. The p K a of acetic acid and p Kb of ammonium hydroxide are 4.76 and 4.75 respectively . Calculate the pH of ammonium acetate solution . Solution pH = 7 + ½ [ p K a – p K b ] = 7 + ½ [4.76 – 4.75] = 7 + ½ [0.01] = 7 + 0.005 = 7.005

118. BUFFER SOLUTIONS: A  buffer solution is defined as a solution which resist any change in its pH value even when small amount of acid or base are added to it . Types of the buffer solution 1) Solution of single substance The solution of the salt of weak acid and weak base eg : ammonium acetate ( CH 3 COOH) act as a buffer.

119. Solution of Mixture: These are of 2 types: Acidic buffer  : It is the solution of mixture of the weak acid and a salt of this weak acid with a strong base. For Ex: CH 3 COOH + CH 3 COONa or HCOOH + HCOONa Basic buffer :  It is the solution of a mixture of weak base and a salt of this  weak base with a strong acid. For Ex: NH 4 OH + NH 4 Cl  or NH 4 OH + NH 4 NO 3

120. Buffer Action The property of a buffer solution to resist any change in its pH value even when small amount of the acid or the base are added to it is called  Buffer action . Preparation of Acidic Buffer To prepare a buffer of acidic pH we use weak acid and its salt formed with strong base. We develop the equation relating the pH, the equilibrium constant, Ka of weak acid and ratio of concentration of weak acid and its conjugate base. For the general case where the weak acid HA ionises in water,

121. Preparation of Acidic Buffer To prepare a buffer of acidic pH we use weak acid and its salt formed with strong base. We develop the equation relating the pH, the equilibrium constant, Ka of weak acid and ratio of concentration of weak acid and its conjugate base. For the general case where the weak acid HA ionises in water .

122. AgCl (s )  Ag +  ( aq ) + Cl ‾ ( aq ) Ksp = [Ag + ][ Cl ‾ ] where K sp  is the solubility product and is equal to ionic product  for a saturated solution. BaSO 4  (s )  Ba 2 +  + SO 4 2‾ AgSCN ( aq )   Ag +  ( aq )+ SCN‾ ( aq )