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Precipitation Equilibrium

Precipitation Equilibrium. Precipitation reactions reach a position of equilibrium – even the most insoluble electrolyte dissolves to at least a slight extent, establishing an equilibrium with it’s ions in solution. For example, a solution of lead II chloride. Ion-Product Equilibrium Systems.

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Precipitation Equilibrium

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  1. Precipitation Equilibrium • Precipitation reactions reach a position of equilibrium – even the most insoluble electrolyte dissolves to at least a slight extent, establishing an equilibrium with it’s ions in solution. • For example, a solution of lead II chloride.

  2. Ion-Product Equilibrium Systems • There are 2 types of precipitation equilibria: • between a precipitate and its ions. A precipitate (ppt.) forms when a cation from one solution combines with another forming an insoluble ionic solid. For example: Sr(NO3)2 (aq) + K2CrO4(aq)⇌2KNO3 (aq) + SrCrO4 (s) (2) between a precipitate and the species used to dissolve it. AgCl (s) + 2NH3 (aq) ⇌ Ag(NH3)2+(aq) + Cl-(aq)

  3. The Solubility Product Constant • The Ksp expression – this is the equilibrium constant expression for the dissolving of a solid. • The smaller the Ksp value the less soluble the precipitate. • And like the any equilibrium constant value, the Ksp value is at a fixed temperature.

  4. Writing Ksp Expression: • Write the ion-product expression for each compound: • (a) magnesium carbonate • (b) iron(II) hydroxide • (c) calcium phosphate • (d) silver chromate • (e) silver sulfide *(be careful with sulfide ions, they are so reactive with water that they will produce HS- and OH-)

  5. Calculating Ion Concentration • Calcium phosphate is a water-insoluble mineral, large quantities of which are used to make commercial fertilizers. Taking it’s equilibrium constant value of 1 x 10-33, calculate: • The concentration of the phosphate ion in equilibrium with the solid if the [Ca2+] = 1 x 10-9 M • Ans = 1 x 10-3 M • The concentration of the calcium ions in equilibrium with the solid if the [PO43-] = 1 x 10-5 M • Ans = 2 x 10-8 M

  6. Determining Precipitate Formation: • Ksp values can be used to predict when a precipitate will form or not. To do this we work with the reaction quotient once more: • If Q > Ksp - the reaction will shift to the left so ppt. forms • If Q < Ksp - the reaction will shift to the right so no ppt. forms • If Q = Ksp - then the solution is saturated with ions at the point of ppt.

  7. Determining if a ppt. will form… • Sodium chromate is added to a solution in which the original concentration of Sr2+ is 0.0060 M • (a) assuming the [Sr2+] stays constant, will a precipitate of strontium chromate form when the chromate ion concentration becomes 0.0030 M? • ans: no precipitate will form, Q < K • (b) will a precipitate of strontium chromate form if 0.200 L of 0.0060 M of strontium nitrate solution is mixed with 0.800 L of 0.040 M potassium chromate? • ans: a precipitate will just barely form, Q > K (just slightly)

  8. Ksp & Water Solubility: • One way to establish equilibrium between a slightly soluble solid is to stir the solid with water to form a saturated solution. • Solubility of the solid, in moles per liter, is related to the solubility product constant. • For example, determine the solubility of barium sulfate. • ans: s = 1.0 x 10-5 M

  9. Example: Determine Solubility • Calculate the solubility of barium fluoride in moles/liter and grams/liter. • ans: 3.6 x 10-3 M & 0.63 g/L

  10. Ksp & The Common Ion Effect • The solubility of an insoluble substance decreases when adding a solution with a common ion in comparison to adding water (similar to Le Chatelier’s Principle). • Which will have the higher solubility: • barium sulfate in water OR barium sulfate in a 0.1M solution of sodium sulfate • Why?

  11. Applying the Common-Ion Effect • Taking the equilibrium constant value of barium sulfate into account, estimate its solubility in a 0.10 M solution of sodium sulfate (hint: you will need an equilibrium table). • ans: 1.1 x 10-9 M

  12. Selective Precipitation • A way to separate 2 cations in water solution is to add an ion that precipitates only one of the cations. • For example: • A flask contains a solution 0.10 M Cl- and 0.010 M CrO42-. When AgNO3 is added: • (a) which anion, chloride or chromate, precipitates first? ans: The chloride ion will ppt. first since it requires the lowest silver ion concentration to ppt. • (b) what percentage of the first anion has been precipitated when the second anion starts to precipitate? ans: 99.98% of chloride ions precipitated.

  13. Dissolving Precipitates • Water insoluble ionic solids can be brought into solution by adding a reagent to react with either the anion or the cation. The 2 most useful reagents for this are: • 1. Strong Acids – the H+ will react with the basic anions • 2. NH3 or OH- to react with the metal cations

  14. Dissolving Zinc Hydroxide • Write the equation that represents the dissolving of zinc hydroxide by a strong acid (hint: two-step process). • Determine the equilibrium constant for this reaction (hint: rule of multiple equilibria). The Kw for 1 mole of water is 1 x 10-14.

  15. Strong Acids • Strong acids can be used to dissolve water-insoluble salts in which the anion is a weak base: • Almost all carbonates • Many sulfides • Example: Write balanced equations to explain why each of the following precipitates dissolve in a strong acid (assume the acid is in excess): • Aluminum hydroxide • Calcium carbonate • Cobalt II sulfide

  16. Complex Ion Formation • Ammonia and hydroxides are commonly used to dissolve precipitates containing a cation that forms a stable complex with NH3 or OH- • Write the reaction by which zinc hydroxide dissolves in ammonia and solve for the equilibrium constant of this reaction using: • Step 1: Zn(OH)2(s)⇌ Zn2+ + 2OH- Ksp = 4.0 x 10-17 • Step 2: Zn2+ + 4NH3 (aq) ⇌ Zn(NH3)42+ + 2OH- Kf = 3.6x108 *Kf = equilibrium constant for the formation of complex ions

  17. Dissolving AgCl with Ammonia • Consider the reaction by which silver chloride dissolves in ammonia: • Taking the Ksp AgCl = 1.8 x 10-10 and the Kf Ag(NH3)2+ = 1.7 x 107, calculate the K for this reaction. • ans: 3.1 x 10-3 • Calculate the number of moles of AgCl that dissolves in one liter of 6.0 M ammonia. • ans: 0.30 moles/liter

  18. MC #1 • How many moles of NaF must be dissolved in 1.00 liter of a saturated solution of PbF2 at 25 °C to reduce the [Pb2+] to 1 x 10¯6 molar? (Ksp of PbF2 at 25 °C = 4.0 x 10¯8) (A) 0.020 mole(B) 0.040 mole(C) 0.10 mole(D) 0.20 mole(E) 0.40 mole

  19. MC #2 • What is the net ionic equation for the reaction that occurs when aqueous copper(II) sulfate is added to excess 6-molar ammonia? (A) Cu2+ + SO42¯ + 2 NH4+ + 2 OH¯ ---> (NH4)2SO4 + Cu(OH)2(B) Cu2+ + 4 NH3 + 4 H2O --> Cu(OH)42¯ + 4 NH4+(C) Cu2+ + 2 NH3 + 2 H2O --> Cu(OH)2 + 2 NH4+(D) Cu2+ + 4 NH3 --> Cu(NH3)42+(E) Cu2+ + 2 NH3 + H2O --> CuO + 2 NH4+

  20. MC #3 • The solubility of CuI is 2 x 10¯6 molar. What is the solubility product constant, Ksp, for CuI? (A) 1.4 x 10¯3(B) 2 x 10¯6(C) 4 x 10¯12(D) 2 x 10¯12(E) 8 x 10¯18

  21. MC #4 • A white solid is observed to be insoluble in water, insoluble in excess ammonia solution, and soluble in dilute HCl. Which of the following compounds could the solid be? (A) CaCO3(B) BaSO4(C) Pb(NO3)2(D) AgCl(E) Zn(OH)2

  22. MC #5 • Equal volumes of 0.10-molar H3PO4 and 0.20-molar KOH are mixed. After equilibrium is established, the type of ion an solution in largest concentration, other than the K+ ion, is (A) H2PO4¯(B) HPO42¯(C) PO43¯(D) OH¯(E) H3O+

  23. MC #6 • Barium sulfate is LEAST soluble in a 0.01-molar solution of which of the following? (A) Al2(SO4)3(B) (NH4)2SO4(C) Na2SO4(D) NH3(E) BaCl2

  24. MC #7 • What is the molar solubility in water of Ag2CrO4? (The Ksp for Ag2CrO4 is 8 x 10¯12.) (A) 8 x 10¯12 M(B) 2 x 10¯12 M(C) (4 x 10¯12 M)1/2(D) (4 x 10¯12 M)1/3(E) (2 x 10¯12 M)1/3

  25. MC #8 • The net ionic equation for the reaction between silver carbonate and hydrochloric acid is (A) Ag2CO3(s) + 2 H+ + 2 Cl¯ ---> 2 AgCl(s) + H2O + CO(g)(B) 2 Ag+ + CO32¯ + 2 H+ + 2 Cl¯ ---> 2 AgCl(s) + H2O + CO2(g)(C) CO32¯ + 2 H+ ---> H2O + CO2(g)(D) Ag+ + Cl¯ ---> AgCl(s)(E) Ag2CO3(s) + 2 H+ ---> 2Ag+ + H2CO3

  26. MC #9 • A 20.0-milliliter sample of 0.200-molar K2CO3 solution is added to 30.0 milliliters of 0.400-molar Ba(NO3)2 solution. Barium carbonate precipitates. The concentration of barium ion, Ba2+, in solution after reaction is (A) 0.150 M(B) 0.160 M(C) 0.200 M(D) 0.240 M(E) 0.267 M

  27. MC #10 • A 1.0 sample of an aqueous solution contains 0.10 mol of NaCl and 0.10 mol of CaCl2. What is the minimum number of moles of AgNO3 that must be added to the solution in order to precipitate all of the Cl¯ as AgCl(s) ? (A) 0.10 mol(B) 0.20 mol(C) 0.30 mol(D) 0.40 mol(E) 0.60 mol

  28. FRQ #1 MgF2(s) Mg2+(aq) + 2 F-(aq) In a saturated solution of MgF2 at 18ºC, the concentration of Mg2+ is 1.21´10-3 molar. The equilibrium is represented by the equation above. • (a) Write the expression for the solubility-product constant, Ksp, and calculate its value at 18ºC. • (b) Calculate the equilibrium concentration of Mg2+ in 1.000 liter of saturated MgF2 solution at 18ºC to which 0.100 mole of solid KF has been added. The KF dissolves completely. Assume the volume change is negligible. • (c) Predict whether a precipitate of MgF2 will form when 100.0 milliliters of a 3.00´10-3-molar Mg(NO3)2 solution is mixed with 200.0 milliliters of a 2.00´l0-3-molar NaF solution at 18ºC. Calculations to support your prediction must be shown. • (d) At 27ºC the concentration of Mg2+ in a saturated solution of MgF2 is 1.17´10-3 molar. Is the dissolving of MgF2 in water an endothermic or an exothermic process? Give an explanation to support your conclusion.

  29. FRQ #2 At 25ºC the solubility product constant, Ksp, for strontium sulfate, SrSO4, is 7.6×10-7. The solubility product constant for strontium fluoride, SrF2, is 7.9´10-10. • (a) What is the molar solubility of SrSO4 in pure water at 25ºC? • (b) What is the molar solubility of SrF2 in pure water at 25ºC? • (c) An aqueous solution of Sr(NO3)2 is added slowly to 1.0 litre of a well-stirred solution containing 0.020 mole F- and 0.10 mole SO42- at 25ºC. (You may assume that the added Sr(NO3)2 solution does not materially affect the total volume of the system.) 1. Which salt precipitates first? 2. What is the concentration of strontium ion, Sr2+, in the solution when the first precipitate begins to form? • (d) As more Sr(NO3)2 is added to the mixture in (c) a second precipitate begins to form. At that stage, what percent of the anion of the first precipitate remains in solution?

  30. FRQ #3 • Answer the following questions relating to the solubility of the chlorides of silver and lead. • (a) At 10C, 8.9  10-5 g of AgCl(s) will dissolve in 100. mL of water. • (i) Write the equation for the dissociation of AgCl(s) in water. • (ii) Calculate the solubility, in mol L–1, of AgCl(s) in water at 10C. • (iii) Calculate the value of the solubility-product constant, Ksp for AgCl(s) at 10C. • (b) At 25C, the value of Ksp for PbCl2(s) is 1.6  10-5 and the value of Ksp for AgCl(s) is 1.8  10-10. • (i) If 60.0 mL of 0.0400 MNaCl(aq) is added to 60.0 mL of 0.0300 MPb(NO3)2(aq), will a precipitate form? Assume that volumes are additive. Show calculations to support your answer. • (ii) Calculate the equilibrium value of [Pb2+(aq)] in 1.00 L of saturated PbCl2 solution to which 0.250 mole of NaCl(s) has been added. Assume that no volume change occurs. • (iii) If 0.100 MNaCl(aq) is added slowly to a beaker containing both 0.120 M AgNO3(aq) and 0.150 MPb(NO3)2(aq) at 25C, which will precipitate first, AgCl(s) or PbCl2(s)? Show calculations to support your answer.

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