UNIT 4: Heterogeneous EQUILIBRIUM (Chap 17-pg 759)

Transcription

1 UNIT 4: Heterogeneous EQUILIBRIUM (Chap 17-pg 759) 4:1. Review your understanding of the Solution Process by reading CHAPTER 12 in particular on pages Then answer the following given the H values for dissolving 1 mol of each of the following gases in water: GAS H (dissolve) a) Based on the H, which gas is most likely to dissolve NH kj/mol b) Based on the H, which gas is least likely to dissolve SO kj/mol c) What is the effect of entropy on the solubility of these gases CH kj/mol d) What is the effect of temperature on the solubility of these gases e) Suggest a reason why the value of H sol n for NH 3 is negative 4:2. Attempt questions a) 12.4 b) 12.5 c) 12.7 d) 12.8 e) 12.9 f) g) h) i) j) on page 566 then k) on page 568 then * l) on page 570 * Determining Ksp values (17.1 pg 760) 4.3. Attempt questions a) b) c) d) e) on page 787/88 * Using Ksp values to determine solubilities ( pg 765 ) 4.4 Calculate a) the molar solubility & b) the solubility in µg/l for Al(OH) 25 C 4:5. Attempt questions # a) b) c) on page What is the conc'n of OH - ions in a saturated sol'n of zinc hydroxide? 4:7. A litre of sol'n which is in eq'm with a precipitate of Mn(OH) 2 contains 5 times as many moles of OH - as Mn 2+. How many moles of Mn 2+ are present? (Ksp = 4.0 x for Mn(OH) 2 ) * Common Ion Effect ( pg 767 ) 4:8. Attempt questions a) b) c) on page :9. Calculate the sol'y in ng/500 ml of Ag 2 CrO 4 in a 0.03 M AgNO 3 sol'n 4:10. If 2.0 x 10-5 mol of sparingly soluble Cu(IO 3 ) 2 can dissolve in 2 L of a very soluble NaIO 3 sol'n, what is the molar concentration of the NaIO 3 sol'n? (Ksp = 1.4 x 10-5 for Cu(IO 3 ) 2 ) * Will a Precipitate form? ( pg 770 ) 4:11. Attempt question a) on page 788 4:12. Determine whether a ppte forms when the following sol'ns are mixed a) Equal volumes of 4 x 10-4 M Ba(NO 3 ) 2 and 8 x 10-6 M Na 2 SO 4 b) One drop (0.05 ml) of 5.0 M Pb 2+ is added to 100 ml of 3x10-3 M I -. c) 2.0 g of Mg(ClO 3 ) 2 are added to 100 ml of 0.05 M KNO 3 d) 10.0 ml of 2x10-2 M AlCl 3 is mixed with 25.0 ml of 0.4 M Pb(NO 3 ) 2 & then diluted to 10 L 4:13. When 100 ml of 2.5 x10-5 M ferrous chloride is added to 150 ml of 6.7 x 10-5 M sodium hydroxide a precipitate just starts to form. What is the Ksp for Fe(OH) 2? * Complex Ion Equilibria (17.4 pg 780) Use LeChateliers Principle and all relevant equations to illustrate why oxalic acid ( sparingly soluble in water ) dissolves readily in NaOH solutions 4:15. Attempt question a) 17.7 on page 787 and b) on pg 788 and c) on pg 789

2 LAB 4-1: Determination of the Solubility Product for PbCl 2 - Evaporation Technique (** The saturated PbCl 2 sol'n has been made up by shaking excess PbCl 2 with water) 1. Determine the mass of a clean, dry glass evaporating dish - record. 2. Carefully pipet 25.0 ml of saturated PbCl 2 sol'n into the evap dish, taking care that no solid gets into the dish. Record the temperature of the sol'n. 3. Heat gently to evaporate the water. Take particular care to avoid splattering especially near the end. Ideally you could slowly evaporate about two-thirds of the water then leave the rest to dry naturally (in oven) 4. Determine the mass of the evap dish which now contains the PbCl 2 residue. From this result find the mass of solid lead chloride that was dissolved in the original 25 ml of solution. 5. CALCULATE: a) the concentration of the original saturated PbCl 2 sol'n in mol/l ie (molar solubility ) then b) the Ksp value for lead chloride LAB 4-2: Determining the Solubility Product for PbCl 2 - Precipitation Technique 1) Mass two pieces of regular filter paper and record. Set up for filtration. 2) Pipet 25 ml of the sat PbCl 2 sol'n into a 150 ml beaker, be sure no solid is transferred. HEAT until close to boiling. 3) Obtain approximately 50 ml of 0.1 M sodium iodide sol'n provided. HEAT this sol'n close to boiling and pour it into the PbCl 2 sol'n. Cool in the sink and pour the mixture into the filter funnel. (digesting a precipitate) *Rinse the beaker out with several aliquots of distilled water and filter these as well. 4) Carefully remove the filter paper and leave it in an evap dish in the drying oven overnight. Weigh the next day. 5) CALCULATE: a) the quantity of PbI 2 precipitate in g/25 ml then mol/l b) the concentration of "lead" that was in the original PbCl 2 solution in mol/l (explain your reasoning/logic behind this answer) c) the Ksp value for lead chloride PREPARE AND HAND-IN A COMPLETED LAB WRITE-UP FOR BOTH EXPERIMENTS ** Include a comparison of the 2 results and a comparison of your result with the value on your data sheet with a discussion/comment on these comparisons ( hint: look ar page ) L-1. A silver chloride ppte weighing g was obtained from g of a soluble salt mixture. Calculate the % of chloride in the soluble salt mixture. L-2. What is the % of iron in a g sample of an acid soluble iron ore which yields a ferric oxide precipitate weighing g? L-3. A sample of iron chloride weighing 0.3 g was dissolved and treated with excess silver nitrate to precipitate out the chloride as silver chloride. The silver chloride precipitate weighed g. Calculate the moles of chloride present in this solution then the moles of iron and finally the correct formula/name for this iron chloride. L-4 Attempt questions 4.77 & 4.78 on page 190

3 The Thermodynamics of Solubility Equilibrium Typically when we speak about Ksp values we are interested in very slightly soluble salts--the kinds of compounds you learned as "insoluble" in your solubility rules. But the quantity can be calculated for any salt as long as you have solubility data. However, determining Ksp values from simple experiments in which solubility is measured almost never yields accurate results, even for very insoluble compounds. The reasons for this are complex and have to do with ion interactions which occur even in rather dilute solutions These ions can clump together in groups, or establish quite large hydration spheres with water, etc. In other words, sometimes the actual concentration of " ions" in a solution is difficult to determine by simple experimental means. The most accurate determinations of Ksp are done by examining successively more dilute solutions. When this is done and the ion concentration is extrapolated to zero, the " official" Ksp or what is sometimes call the thermodynamic Ksp is found. This value should come close to the value you would get if you found G o for the process and then used the expression G o = -RTlnKsp to determine the Ksp. This often results in disagreements between ordinary experimental values and calculated values related by factors of 10 or 100!!!! So. there are "ideal" solutions (the kind we talk about) and "real" solutions (the kind we deal with) just as there are ideal and real gases. For the purposes of this experiment we will ignore the effects of ion interaction. In fact, we will be using a very soluble salt, KNO3, because it has characteristics that make it easy to work with, it's inexpensive, relatively harmless as used, and easily recovered for reuse. The equilibrium established in a sat solution of KNO 3 may be represented as: KNO 3 (s) K + - (aq) + NO 3 (aq) As with all equilibria, the relative amounts of products and reactants at a given temperature are determined by the "tension" between the tendency toward maximum randomness ( S) and the tendency toward minimum energy ( H). In this experiment you will first determine the solubility of potassium nitrate at a variety of temperatures and then use the solubility to calculate the "Ksp" of the salt. Recall the relationship between the equilibrium constant (Ksp) and the free energy change ( G ) : G =-RT(lnK sp ) where R=8.314 J/mol K and T is the temperature in Kelvin. So you can determine the free energy change( G) at each temperature from the Ksp at each temperature. The sign of the enthalpy change ( H) should be clear from the solubility behavior (i.e., whether more or less dissolves at higher temperatures) but the value of H can also be determined from the value of the Ksp ( G= H-T S). It is known that over small temperature ranges, K and H are related as follows: H 1 ln K sp = - R T + C y = m x + b So a plot of ln Ksp vs. 1/T (in Kelvin) will yield a nearly straight line with slope =- H/R. And, of course, once you know G and H, you can determine S at each temperature. Again, it should be clear what the sign of the entropy change will be considering the process. To summarize: You will determine the solubility of KNO 3 in water at different temperatures and use this information to calculate the "Ksp" of the salt. From the Ksp you will determine the free energy change ( G), enthalpy change ( H) and entropy change( S) for this process and compare them to the values calculated from tables in your text. Procedure 1. Add approximately 3.0 g of KNO 3 to a 10 ml graduated cylinder ( remove the plastic base before heating) and add distilled water to the 4 ml mark (including the solid). 2. Place the cylinder in a hot water bath and stir until the solid has completely dissolved. Remove from the bath and patiently wait for crystallization to occur (saturation point). Try to minimize the amount of time the cylinder spends in the hot water, i.e.,don't leave it there any longer than needed to dissolve the solid. Otherwise you will spend a lot of time waiting for the solution to cool before crystals form. Always keep the solution stirred and record the first temperature that crystals become visible. Experience seems to indicate that the falling temperature levels off a little as crystallization occurs. Be careful not to lose solid as you remove the thermometer from the cylinder to add more water. 3. Repeat, diluting the sol n by adding 1 ml of water each time 4. Recycling the KNO 3! After you have your four readings, pour your solution into the beaker indicated by your instructor. NOTE:. In order to determine the solubility in moles/litre you need to know the moles of solute and the volume of the solution. The initial volume can be read directly from the cylinder once the solid has dissolved. Each 1mL addition thereafter will increase the volume, but again you should record the volume only after redissolving the solid in the hot water bath.

4 Calculations 1. Use the mass of KNO 3 and the measured volumes of the various mixtures to calculate the Molarities (solubilities) for each of the four trials. Then use the calculated solubilities to determine the "Ksp" for potassium nitrate at the four temperatures measured. Show all your work in the table below! 2. Now plot a graph as described in the introduction to determine the value of H for this process. Look up the H for values for the process in your text and calculate H from them. Explain any disagreement. 3. Calculate the Gibbs free energy change( G) for this process at each of the four temperatures measured, using your "Ksp" values 4. Use the experimental values of G and H to calculate S at each of the four temperatures. How does the sign on S "fit" this process? Is this process entropy or enthalpy driven--or both? Explain. Solubility KNO 3 (mol/l) Ksp ln(ksp) Temp ( C) 1/Temp( 1/K) H (kj) G (kj) S (J)

5 Heterogeneous Equilibria - Assignment 1. a) Sea water on our planet is saturated with silver chloride. The chloride ion concentration in sea water is 0.53M. Calculate the number of nanograms of silver in 1 L of sea water. b) How many litres of sea water would you need to obtain 1 mole of silver ions? 2. In a saturated solution of MgF 2 at 18 C, the concentration of Mg 2+ is 1.21x10-3 M. The equilibrium is represented by this equation.. MgF 2 (s) Mg 2+ (aq) + 2 F - (aq) (a) Calculate the Ksp value for MgF 2 at 18 C. (b) Calculate the equilibrium concentration of Mg 2+ in 1.00 liter of saturated MgF 2 solution at 18 C to which mole of solid KF has been added. The KF dissolves completely. (c) Predict whether a precipitate of MgF 2 will form when 100 ml of a 3.00x10-3 M Mg(NO 3 ) 2 solution is mixed with 200 ml of a 2.00xl0 3 M NaF solution at 18 C. (d) At 27 C the concentration of Mg 2+ in a saturated solution of MgF 2 is 1.17x10-3 M. Is the dissolving of MgF 2 in water an endothermic or an exothermic process? Give an explanation(lechat) to support your conclusion. 3. At 25 C the solubility product constant, K sp, for strontium sulfate, SrSO 4, is 7.6x10-7. The K sp for strontium fluoride, SrF 2, is 7.9x (a) What is the solubility of SrSO 4 in mg/100ml at 25 C? (b) What is the solubility of SrF 2 in mg/100ml at 25 C? (c) A solution of Sr(NO 3 ) 2 is added slowly to 500 ml of a well-stirred solution containing mole F - and 0.05 mole SO 2-4 at 25 C. Which salt precipitates first? (You may assume that the added Sr(NO 3 ) 2 solution does not materially affect the total volume of the system.) 4. How many milliliters of ammonia gas at 20 C and 130 kpa would have to be dissolved in 1.0 L of water to dissolve 1.0 x10-4 moles of the very insoluble salt cobalt(ii) carbonate ( CoCO 3 Ksp = 1.0 x ) in the form Co(NH 3 ) 6 2+? 5. a) Explain why a solution of copper(ii) nitrate is light blue yet turns a darker blue-purple when ammonia is added to it, include a 3-dimensional sketch of the molecules involved b) Attempt questions a) b) & c) on page 942