1. L e c t u r e 2L e c t u r e 2L e c t u r e 2L e c t u r e 2 Precipitation equilibrium Associate prof. L.V. Vronska Associate prof. M.M. Mykhalkiv

2. Outline 1.Precipitation equilibrium as heterogeneousequilibrium 1.Precipitation equilibrium as heterogeneous equilibrium 2.Calculation of solubility and solubility product K sp 3.Influence of chemical factors is on solubility of precipitate 4.Completeness of precipitation and factors which influence on its 5.Conditions of dissolution of precipitation

3. 1. PRECIPITATION EQUILIBRIUM AS HETEROGENEOUSEQUILIBRIUM Precipitation equilibrium 1. PRECIPITATION EQUILIBRIUM AS HETEROGENEOUS EQUILIBRIUM Precipitation equilibrium   precipitate An insoluble solid that forms when two or more soluble reagents are combined.

4.   The most common precipitation reaction is a metathesis reaction, in which two soluble ionic compounds exchange parts. Thus, the precipitation of PbCl 2 is written as Pb 2+ (aq) + 2Cl – (aq) = PbCl 2 (s)   In the equilibrium treatment of precipitation, however, the reverse reaction describing the dissolution of the precipitate is more frequently encountered. PbCl 2 (s) = Pb 2+ (aq) + 2Cl – (aq)

5. Precipitation equilibrium The equilibrium constant for this reaction is called the solubility product, K sp, and is given as K sp = [Pb 2+ ][Cl – ] 2 = 1.7 ·10 –5  and for all electrolytes A m B n K sp = [A] m [B] n   solubility product K sp - the equilibrium constant for a reaction in which a solid dissociates into its ions

6. Concentrational (real) constant for all electrolytes A m B n (use, when we have real conditions (influence of ionic strength)) Concentrational (real) constant solubility product K R sp for all electrolytes A m B n (use, when we have real conditions (influence of ionic strength)) K R sp = [A] m [B] n for all electrolytes A m B n Thermodynamic constant solubility product K T sp for all electrolytes A m B n

7.  depends on:  Thermodynamic constant solubility product K T sp depends on:  Temperature  Pressure  Nature of solvent  Nature of precipitate

8. Thermodynamic constants solubility product K T sp are adduction in reference books

9.  Conditional constant for all electrolytes A m B n  Conditional constant solubility product K C sp for all electrolytes A m B n

10.  We use, when we have the following real conditions:  Temperature  Pressure  Influence of ionic strength  Influence of competitive reactions

11. 2. CALCULATION OF SOLUBILITY AND SOLUBILITY PRODUCT K sp   Solubility is a property of matter to give/ to form a solution with a certain solvent at certain conditions We determine Solubility as:   Coefficient of Solubility (k s )   Molar Solubility (S)

12. Coefficient of Solubility (k s )  It is mass of matter which dissolves at this temperature in 100 g or 100 mL of solvent

13. Coefficient of Solubility (k s )

14. Molar Solubility (S)  It is a molar concentration of matter in the saturated solution А m В n  mА + nВ K R sp = [A] m [B] n [A] = m[A m B n ] = mS [B] = n[A m B n ] = nS K R sp= (mS) m (nS) n = m m n n  S m+n

15. A rule of solubility product: in saturated solution above sediment product of ions concentrations is permanent at a stationary temperature  In unsaturated solution [A] m [B] n  K sp  In saturated solution [A] m [B] n =K sp In supersaturated solutionIn supersaturated solution [A] m [B] n > K sp

17.  If ionic strength can be adopted even a zero and to neglect of competitive reactions, solubility of precipitate is expected on the size of K T sp (if μ→0 then f→1; α=1)  If to take into account ionic strength, but to neglect of competitive reactions, solubility is expected after the size of K R sp (if μ≠0 then f≠1; α=1)  If we cannot neglect by competitive reactions, then solubility is expected on the size of K C sp (if μ≠0 then f≠1; α ≠ 1)

18. 3. INFLUENCE OF CHEMICAL FACTORS IS ON SOLUBILITY OF PRECIPITATE  The Common-Ion effect  The activity effect  Acid-base reactions  Complexation reactions  Red-ox reactions

19. The Common-Ion Effect

20. The common-ion effect. The solubility of at 25°C decreases markedly on addition of ions. Note that the calculated solubility is plotted on a logarithmic scale.

21.   The solubility of precipitate decreases in the presence of a solution that already contains one of its ions. This is known as the common ion effect.

22. The activity effect   Clearly the equilibrium position for the reaction AgIO 3 (s)  Ag + (aq) + IO 3 – (aq)   depends on the composition of the solution. When the solubility product for AgIO 3 is calculated using the equilibrium concentrations of Ag + and IO 3 – K sp = [Ag + ][IO 3 – ] its apparent value increases when an inert electrolyte such as KNO 3 is added.

23.   The true thermodynamic equilibrium constant, K sp, for the solubility of AgIO 3, therefore, is K T sp = a(Ag + ) a(IO 3 – ) K T sp =[Ag + ][IO 3 - ]f(Ag + )f(IO 3 – ) K T sp = K R sp f(Ag + )f(IO 3 – ) K R sp = K T sp / f(Ag + )f(IO 3 – )   To accurately calculate the solubility of AgIO 3, we must know the activity coefficients for Ag + and IO 3 –.

24. Mention !!!   First, as the ionic strength approaches zero, the activity coefficient approaches a value of one. Thus, in a solution where the ionic strength is zero, an ion’s activity and concentration are identical. We can take advantage of this fact to determine a reaction’s thermodynamic equilibrium constant. The equilibrium constant based on concentrations is measured for several increasingly smaller ionic strengths and the results extrapolated back to zero ionic strength to give the thermodynamic equilibrium constant.   Second, activity coefficients are smaller, and thus activity effects are more important, for ions with higher charges and smaller effective diameters. Finally, the extended Debye–Hückel equation provides reasonable activity coefficients for ionic strengths of less than 0.1. Modifications to the extended Debye–Hückel equation, which extend the calculation of activity coefficients to higher ionic strength, have been proposed.

25. The pH of the Solution   An ionic compound that contains a basic anion becomes more soluble as the acidity of the solution increases. The solubility of CaCO 3, for example, increases with decreasing pH because the CO 3 2- ions combine with protons to give HCO 3 - ions. As CO 3 2- ions are removed from the solution, the solubility equilibrium shifts to the right, as predicted by Le Châtelier’s principle. The net reaction is dissolution of CaCO 3, in acidic solution to give Ca 2+ ions and HCO 3 - ions:

27. Formation of Complex Ions The solubility of an ionic compound increases dramatically if the solution contains a Lewis base that can form a coordinate covalent bond to the metal cation. Silver chloride, for example, is insoluble in water and in acid, but it dissolves in an excess of aqueous ammonia, forming the complex ion [Ag(NH 3 ) 2 ] +. A complex ion is an ion that contains a metal cation bonded to one or more small molecules or ions, such as NH 3, CN - or OH -. In accord with Le Châtelier’s principle, ammonia shifts the solubility equilibrium to the right by tying up the Ag + ion in the form of the complex ion:

28. Formation of Complex Ions Silver chloride is insoluble in water (left) but dissolves on addition of an excess of aqueous ammonia (right).

29. The solubility of AgCl in aqueous ammonia at 25°C increases with increasing ammonia concentration owing to formation of the complex ion [Ag(NH 3 ) 2 ] +. Note that the solubility is plotted on a logarithmic scale.

30. 4. COMPLETENESS OF PRECIPITATION AND FACTORS WHICH INFLUENCE ON ITS The precipitation is considered practically complete, if the concentration of the precipitate’s ions in solution above precipitate does not exceed a 10 -6 mol/L

31. Factors which influence on completeness of precipitation 1.Excess of precipitation reagent (50 %) 2.Strength of electrolyte-precipitator 3.pH of solution 4.Fractional precipitation

32. Separation of Ions by Selective Precipitation   A convenient method for separating a mixture of ions is to add a solution that will precipitate some of the ions but not others. The anions SO 4 2- and Cl - for example, can be separated by addition of a solution of Ba(NO 3 ) 2. Insoluble BaSO 4 precipitates, but Cl - remains in solution because BaCl 2 is soluble.   Similarly, the cations Ag + and Zn 2+ can be separated by addition of dilute HCl. Silver chloride, AgCl, precipitates, but Zn 2+ stays in solution because ZnCl 2 is soluble.

33. Separation of Ions by Fractional Precipitation  Ions Ba 2+ and Ca 2+ can be separates if concentration of SO 4 2- ions is controlled.  BaSO 4 has K sp = 1  10 –10 and CaSO 4 has K sp = 2,3  10 –5

34.  That precipitate of BaSO 4 have been removed out, his ionic product must be greater K sp, but, that precipitate of CaSO 4 did not removed out it is simultaneously necessary, that ionic product [Ca 2+ ][SO 4 2- ]  K sp CaSO 4, but [Ba 2+ ][SO 4 2- ]  K sp BaSO 4  Therefore, if concentrations both ions are 10 -2 mol/L, concentration SO 4 2- ions must be between and

35. 5. CONDITIONS OF DISSOLUTION OF PRECIPITATION It is necessary for dissolution of sediment, that its ionic product became more small constants of solubility product: [Kat + ][An – ]  K sp_KatAn

36. Decrease of ions concentration it can be carried out the followings methods: 1. strong dilution of solution Descriptive term Descriptive term Approximate volume of solvent in milliliters per gram of solute Approximate volume of solvent in milliliters per gram of solute very soluble less than 1 freely soluble freely soluble from 1 to 10 from 1 to 10 soluble soluble from 10 to 30 from 10 to 30 sparingly soluble sparingly soluble from 30 to 100 slightly soluble slightly soluble from 100 to 1000 very slightly soluble very slightly soluble from 1000 to 10 000 from 1000 to 10 000 practically insoluble more than 10 000 more than 10 000

37.  The term 'partly soluble' is used to describe a mixture of which only some of the components dissolve.

38. Decrease of ions concentration it can be carried out the followings methods: 2. The ions of precipitate are connected in compounds which well water-soluble Co(OH)Cl  + HCl = CoCl 2 + H 2 O 3. The ions of precipitate are connected in compounds which give gas ZnS  + 2HCl = ZnCl 2 + H 2 S 

39. Formation and dissolution of Cr(OH) 3 Formation and dissolution of Cr(OH) 3

40. Decrease of ions concentration it can be carried out the followings methods: 3. The ions of precipitate are connected in compounds which are complex AgCl  + 2NH 3 = [Ag(NH 3 ) 2 ]Cl 4. Oxidation and reduction of ions of precipitate in others compounds MnO(OH) 2  + H 2 C 2 O 4 + H 2 SO 4 = MnSO 4 + 2CO 2 + 3H 2 O MnSO 4 + 2CO 2 + 3H 2 O

42. Dissolution of sulfatic-precipitate sodium carbonate extraction is a translation of sulfates of second analytical group in carbonates  BaSO 4 K sp =1,1  10 -10  BaCO 3 K sp =5,1  10 -9

43. Aplication of Aplication of Precipitation equilibrium  Gravimertic analysis: - Particulate gravimetry - Precipitation gravimetry

44. Thanks for your attention!