Chemical Equilibrium - LEVEL I- Part - I 1. Concentration of reactant & product does not become equal but constant at constant temperature is equilibrium state. Justify the statement. 2. Define reversible & Irreversible reactions. 3. Prove dynamic nature of equilibrium, using CaCo3 as starting substance. 4. Give example of reversible & irreversible reactions. 5. Classify equilibrium based on extend of reaction. 6. What the types of equilibrium. Define them & give 3 example of each. 7. Write a notes on : (i) Solid Liquid - equilibrium (ii)liquid gas equilibrium (iii)equilibrium involving dissolution of solid (v) Solid Gas equilibrium 8. Give General Charactarics of equilibrium in involved is physical process [Hint - Topic 4.4 Pg 96].***GMP*** 9. Give graphical representation of chemical equilibrium. 10. Give dynamic nature of chemical equilibrium taking NH3 as example.****gmp*** 11. Define and derive equilibrium constant from law of mass action or prove that ratio of forward & reverse rate constant is called equilibrium.[gmp] 12. Show equilibrium constant for following reactions - H2(g) + I2(g) 2HI(g) - 4NH3(g) + 5O2(g) 4NO(g) + 6H2O(g) - N2(g) + 3H2(g) 2NH3(g) 13. Give relation between following equilibrium constant. (a) (i) H2 +I2 2HI (b) (i) H2 + I2 2HI (c) (i) H2 + I2 2HI (ii) 2HI H2 + I2 (ii) H2 + I2 HI (ii)nh2 + ni2 2nHI 14. Give units of Kp & Kc.Derive relation between Kp & Kc (****GMP****) 15. Give types of chemical equilibrium; define each of them & give 2 example of each. 16. In heterogeneous equilibrium const.of pure solid or pure liquid is taken constant.justify. 17. Derive kp & kc for decomposition of CaCO3 & for following reaction - Ag (3) + 2HNO3(aq) 2AgNo3(aq) + H2O(g) 18. Write characteristics of equilibrium constants. 19. Relate K & G.Also explains spontaneity using K. 20. Write Le Chatelier s principal & explain effect of (i) Concentration (ii) Principal (c) Temperature (iv) Inert Gas (v) Catalyst with proper example 21. What happens when - (i) Fe +3 is added 142
(ii) KSCN added or (iii)[fe(scn)] +2 added to the following equilibrium Fe +3 + SCN - Fe(SCN)] +2 Chemical Equilibrium Level I Part - II 22. Define the following with at least one example. (i) Reversible reaction (ii) Irreversible reaction (iii) Homogeneous (iv) Heterogeneous reversible reaction (v) Active mass (vi) Equilibrium state (vii) Equilibrium constant (viii) Law of mass action (ix) Le-Chatelier s principle 23. The equilibrium constants of : (i) N2 + 3H2 2NH3 and (ii) N2 + H2 NH3 are K1 and K2. Prove that K2 =. 24. Assuming the reaction, A2(g) + 2B2(g) C(g) to be exothermic, how will the yield of the product C change with : (a) increase of temperature (b) increase of pressure (c) addition of catalyst (d) addition of inert gas (e) removal of C 25. Consider the equilibrium, PCl3 (g) + Cl2(g) PCl5(g) How would the following affect the position of equilibrium? (a) Addition of PCl3 (b) Addition of Cl2 (c) Removal of PCl5 (d) Decrease in volume of container (e) Addition of He at constant volume 26. Manufacture of NH3 is represented by the equation : N2(g) + 3H2(g) 2NH2(g) + Heat On the basis of this equation, deduce the optimum conditions for best result. 27. Discuss the application of law of chemical equilibrium to heterogeneous equlibrium to heterogeneous equilibria involving : (i) dissociation of CaCO3 (ii) dissociatgion of ammonium hydrogen sulphide. (iii) dissociation of ammonium carbonate 28. Explain why it is so : (a) in the synthesis of NO, high temperature is maintained. (b) in the synthesis of NH3, low temperature is maintained (c) high pressure is applied in the synthesis of ammonia (d) A catalyst is used in the synthesis of ammonia (e) In presence of Cl2, dissociation of PCl5 decreases (f) The dissociation of HI is independent of pressure. (g) Kp is equal to KC when n is zero. (h) for the completion of decomposition of CaCO3, CO2 is allowed to escape. 143
Chemical Equilibrium - Level II 1. What is change in Gibbs free energy at equilibrium? 2. What happens to the equilibrium constant if temperature is increased? 3. The equilibrium constants for the reactions, N2 + O2 2NO..(i) 2NO + Cl2 2NOCl..(ii) are K1 and K2 respectively, what will be the equilibrium constants for the reaction? N2 + O2 + Cl2 2NOCl 4. Using following data, explain whether the reaction is exothermic or endothermic. K = 1.24 10-3 at 25 C K = 2.34 10-3 at 50 C 5. What conclusions can be drawn for the following? (i) Q = K; (ii) Q > K ; (iii) Q < K 6. Let us consider following two reversible reactions (i) CH3COH(l) + C2H5OH(l) CH3COOC2H5(l) + H2O(l) (ii) CaCO3 (s) CaO(s) + CO2(g) Explain, why the equilibrium (i) can be attained in an open vessel but the equilibrium (ii) Is heterogeneous involving a gas hence it is not possible in an open vessel. CO2 gas will escape out in the open vessel CaCO3(s) CaO(s) + CO2(g) 7. What is the ratio of KP to KC for the reaction? 2SO2(g) + O2(g) 2SO3(g) at 25 C Consider the unit of concentration in mol per litre and that of pressure in atmospheres. 8. Explain, why the addition of inert gas des not affect the equilibrium in a rigid vessel? 9. Give the units of equilibrium constant Kp of the following reactions. (i) N2(g) + 3H2(g) 2NH3(g) (ii) P4 (s) + 6H2(g) 4PH3(g) Ionic equilibrium 1. (a) Mention conjugate base of each of the following : HS -, H3O +, H2PO4 -, HSO4 -, HF, CH3COOH, C6H5OH, HClO4, NH4 + (b) Mention the conjugate acid of each of the following : OH -, CH3COO -, Cl -, CO3 2-, H2PO4 -, CH2NH3, CH3COOH, NH2 - (c) Which of the following behave both as Bronsted acids as well as Bronsted bases? H2O, HCO3 -, H2SO4, H3PO4, HS -, NH3 (d) Which is stronger acid in each of the following pairs? HCl, HI; H2CO3, H2SO4; H2O, H2S; C6H5OH, C2H5OH; Na +, K + (e) Which is stronger base in each of the following pairs? OH -, Cl - ; OH -, NH2 - ; OH -, CH3COO - ; CH3COO -, Cl - ; NH3, CH3NH2 (f) Classify the following into acids and bases according to lewis concept. S 2-, H +, OH -, BF3, NI 2+, NF3, AlCl3, SnCl4, NH3, (CH3)2 O 2. What is the difference between a conjugate acid and a conjugate base? 3. What is the concentration of hydroxyl ions in pure water? 144
4. The product of molar concentrations of H + and OH - ions is a constant at a constant temperature. The constant is known as. 5. Define ph. 6. The solution with constant ph is called... 7. Write Henderson s equation for ph of an acidic buffer. 8. Diethyl ether, C2H5OC2H5 acts a s a Lewis base. Which of its atoms acts as electron donor? 9. Two acids HA and HX have values of dissociation constants as 2 10-5 and 3.0 10-9 respectively. Which acid is more stronger? 10. How can it be predicted whether a salt will precipitate or not on mixing the solutions containing its ions? 11. What is the approximate ph of a solution containing 10-8 mol HCl per litre? 12. What happens when HCl gas is passed through saturated NaCl solution? 13. What happens to the ph if few drops of an acid are added to CH3COONH4 solution >? 14. What are the units of ionic product of water? 15. In a given solution, the [H + ] > 10-7 mol L -1 at 25 C. What is the nature of the solution? 16. The pka for acid A is higher than the pka for acid B. Which is a strong acid? 17. Predict whether the following substances will give acidic, basic or neutral solution? (i) K2CO3 (ii) NaCl (iii) FeCl3 (iv) CuSO4 (v) AlCl3 (vi) NH4Cl (vii) CH3COONH4 (viii) NaCN (ix) Na2S (x) Na2B4O7 18. Explain the following : (i) Anhydrous HCl is a bad conductor of electricity while aqueous solution of HCl is a good conductor. Anhydrous HCl is a covalent compound. However, it combines with water in aqueous solution and forms H3O + and Cl - ions, HCl + H2O H3O + + Cl - (ii) Silver nitrate solution does not give white precipitate of silver chloride with pure chloroform but gives white precipitate with sodium chloride. A white precipitate of AgCl is formed when Ag + ions react with Cl - ions. Chloroform being a covalent compound does not give Cl - ions while sodium chloride being an electrolyte gives Cl - ions in solution. (iii) A mixture of sodium acetate and acetic acid resists changes in ph value on adding small amounts of acids or bases. Hint : The mixture is buffer solution. (iv) ph of an aqueous 10-2 M acetic acid is not 2. Acetic acid is a weak electrolyte. The comcentration of hydrogen ions in 10-2 M acetic acid solution is not 10-2 where is degree of ionization which is less than 1)] (v) Ammonia is a Lewis as well as Bronsted base. NH3 combines with proton (NH3 + H + NH4 + ) Thus, it is Bronsted base also] (vi) Water acts both as an acid and as a base. 145
Hint : H2O + H2O H 3 O + + OH - Acid Base Acid Base Conjugate Conjugate 19. Account the following : (a) Solubility of AgCl is more in water than in a solution of common salt. NaCl is a strong electrolyte. It furnishes a large concentration of Cl - ions. Due to this common ion, the solubility of AgCl is less in NaCl solution. (b) A white precipitate appears when HCl is passed through barium chloride solution. HCl gives a large concentration of Cl ions. On account of this, the ionic product of barium chloride exceeds its solubility product and the precipitation of barium chloride occurs. (c) The sulphides of IVth group cannot be precipitated in presence of HCl. The values of solubility products of the sulphides of IVth group are high, i.e., require higher concentration of sulphide ions. This concentration is not obtained in presence of HCl as it suppresses ionization of weak electrolyte H2S (d) A weak base becomes weaker in presence of its salt. The salt is a strong electrolyte and furnishes a large concentration of cations. This suppresses ionization of the weak base due to common ion effect. Thus, OH - ion concentration is quite less in presence of the salt. (e) (f) Sodium chloride is precipitated from the saturated solution when HCl gas is passed through it. Same explanation as in (ii) An excess of H2SO4 is added in gravimetric estimation of barium ions as BaSO4. See text for answer. 20. Which salt out of NaCl, FeCl3 and KNO3 will get hydrolysed? what will be the nature of the solution obtained? 21. What is buffer solution? Explain its properties. 22. What do you understand by ionic product of water? what is its value at 25 C? 23. If the solubility of the sparingly soluble salt Ag2CrO4 is s, write the expression relating the solubility product Ksp and s. 24. What do you understand by hydrolysis? explain giving suitable example. 25. Derive a relation between ionization constant and degree of ionization of a weak electrolyte. How can the strengths of two monobasic acids of same molar concentrations be compared? 26. What are solubility product and common ion effect? Discuss their applications in qualitative analysis. 27. Explain the terms ionic product of water and ph value. What is ph scale? how does it tell us whether a solution is acidic or basic? Can a solution have ph either zero or less than zero? 28. What is buffer solution? what are the different types of buffers? Discuss their buffer action. Give henderso s equation. What is the importance of buffer in biological processes? 146
29. The species : H2O, HCO3 -, HSO4 -, NH3 can act both as Bronsted acids and bases. For each case give the corresponding conjugate acid and conjugate base. Conceptual Questions 1. What is the difference between dissociation and ionization? - Dissociation is used to refer the process of separation of ions in water which already exist as such in the solid state of the solute. NaCl (Na + Cl - ) Na + (aq) + Cl - (aq) (Ionic compound) - The term ionization on the other hand is used for the process in which a neutral molecule (polar covalent molecule) splits into ions in the solution. HCl H + + Cl - polar covalent compound (neutral molecule) 2. Indicate whether the following compounds are Arrhenius, Bronsted or Lewis acids and bases. NH3(g), HCl(aq), Na2CO3(aq), CH3COOH(aq), CO2(g), BF3, Ag +, HCO3 - - NH3(g) bronsted and Lewis base HCl(aq) Arrhenius and Bronsted acid Na2CO3(aq) Bronsted base CH3COOH(aq) Arrhenius and Bronsted acid CO2(g) Bronsted and Lewis acid BF3 Lewis acid Ag + Lewis acid HCO3 - Both acid and base according to Bronsted concept. 3. What is the effect of dilution on degree of ionization of weak electrolytes? According to Ostwald s dilution law, degree of ionisation of weak electrolytes increases on dilution. = Where C is concentration On dilution, Concentration C decreases which results into increase in degree of ionization. 4. Prove that : pka + pkb = pkw for weak acid and its conjugate base. let us consider the weak acid, HCN HCN H + + CN - ; Ka = [ ] [ ] (i) [ ] 147
CN + H2O HNC + OH ; Kb = [ ] [ ] (ii) [ ] multiplying equations (i) and (ii), we get, (a)ka Kb = [H + ] {OH ] = Kw; Ka Kb = Kw. (iii) Taking logarithm of above equation we get, log10ka + log10kb = log10kw log10ka log10kb = log10kw pka + pkb = pkw 5. What are the PH of boiling water and neutral solution at human body temperature? Ans: (i) ph of boiling water = 6.5625. (ii) ph of neutral solution at human body temperature = 6.8 6. What is the effect of temperature on the ionic product of water? Ans: 2H2O H3O + + OH kw = [H3O + ] [OH ] On increasing the temperature, ionic product of water increases because with rise in temperature the dissociation of H2O increases. 7. What is the ph of blood? Why does it remain almost constant? ph of blood is 7.4. its ph remains almost constant due to presence of a buffer. 8. Calculate : (i) pkw at 25 C (ii) ph and poh for a neutral solution with Kw = 10 12 (i) Kw at 25 C = 10 14 pkw = logkw = log10 14 = 14 (ii) Kw = 10 12 [H + ] = [OH ] = 10 6 ph = poh = 6 9. Predict whether Cu 2+ or Ni 2+ will precipitate first on passing H2S gas in acid medium. Cu +2 is II group radical whereas Ni 2+ is of IV group. Ksp CuS < Ksp NiS Thus, on passing H2S gas, CuS will precipitate first. 10. Why sufficient amount of Nh4OH is added before (NH4)2CO3 in Vth group qualitative analysis? Ammonium carbonate contains ammonium bicarbonate with it. Thus, ammonium hydroxide is added to convert ammonium bicarbonate into ammonium carbonate. NH4HCO3 + NH4OH (NH4)2CO3 + H2O 11. Give conjugate acid and base of the following compounds : (i) NH3 (ii) H2o (iii) HSO4, (iv) HCO3 Compound Conjugate baseconjugate acid NH3 NH2 NH4 + H2O OH H3O + HSO4 SO4 2 H2SO4 148