# Solubility Product Constants

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1 Solubility Product Constants PURPOSE To measure the solubility product constant (K sp ) of copper (II) iodate, Cu(IO 3 ) 2. GOALS 1 To measure the molar solubility of a sparingly soluble salt in water. 2 To prepare a calibration curve based on complex ion formation for absorbance enhancement. 3 To calculate the solubility product constant (K sp ) of a sparingly soluble salt from its molar solubility. 4 To confirm the common ion effect on the molar solubility of a sparingly soluble salt. INTRODUCTION In previous introductory chemistry courses, you learned some basic solubility rules that are useful in determining if an ionic solid will dissolve in water. Solids that dissolve completely, such as NaCl and NH 4 NO 3, were referred to as soluble and others that did not dissolve completely, such as AgCl and BaSO 4, were referred to as insoluble. In fact, very few ionic solids are completely insoluble, meaning that they will not form any ions when placed in aqueous solution. Most solids that are commonly referred to as insoluble are actually slightly soluble and will produce an equilibrium between undissolved solid and ions in solution. For example, when copper (II) iodate (Cu(IO 3 ) 2 ) is placed in water, the following equilibrium is established. Cu(IO 3 ) 2 (s) Cu 2+ (aq) + 2IO 3 (aq) (1) The equilibrium constant associated with this reaction is called the solubility product constant and is given the symbol K sp. K sp = [Cu 2+ ][IO 3 ]2 (2) It is important to emphasize that the equilibrium in equation 1 is only true if some solid is present. If the solid completely dissolves in solution, then the product of the ions as shown in equation 2 is not equal to the K sp. However, as long as some solid is in contact with solution, the solution will become saturated 1 with the ions according to equation 1. The molar solubility 2 of a solid is the maximum number of moles of the solid that will dissolve in one liter of solution. Molar solubility is measured in moles/liter and has units of molarity (M). Molar solubility can be determined by measuring the concentration of the ions formed in a solution 1 %28chemistry% solubility c Advanced Instructional Systems, Inc. and North Carolina State University 1

2 saturated with the solid. For example, if the molarity of Cu 2+ in a solution saturated with Cu(IO 3 ) 2 can be determined, the stoichiometry in equation 1 indicates that this is also equal to the molarity of Cu(IO 3 ) 2 that dissolved in solution or its molar solubility. Alternatively, if the molarity of IO 3 - in a solution saturated with Cu(IO 3 ) 2 can be determined, the stoichiometry in equation 1 indicates that the molar solubility of Cu(IO 3 ) 2 in water is one half the [IO 3 - ] in solution. In a saturated solution, if all the ions came from the solid, then the ratio of cations to anions is known. In the case of Cu(IO 3 ) 2, it is known that 2[Cu 2+ ] = [IO 3 - ]. Thus, if the Cu 2+ concentration can be measured, the IO 3 - concentration can be calculated and vice versa. These concentrations can then be entered into the K sp expression (equation 2) to solve for the solubility product constant of Cu(IO 3 ) 2. The preceding discussion referred to dissolving the sparingly soluble salt in pure water. However, if some of the ions that are to be produced by the solid are present in solution from another source, Le Châtelier s Principle predicts that the equilibrium in equation 1 will shift to the left and less solid will dissolve. This would result in a lower molar solubility. This is referred to as the common ion effect 3. In the example of Cu(IO 3 ) 2, the presence of either Cu 2+ or IO 3 - in solution should result in a lower molar solubility than in pure water. Earlier this semester in the Solutions and Spectroscopy lab, you prepared a calibration curve for Cu 2+ over the range of approximately 0.1 M to 0.4 M. As you may recall, Cu 2+ in solution has a pale blue color and you were able to measure its absorbance at 620 nm. For this experiment, it is necessary to prepare a calibration curve for Cu 2+ over the range of 5 x 10-4 M to 0.01 M. At these concentrations, the absorbance of Cu 2+ is too low to detect directly by spectroscopy. To improve the absorbance of the Cu 2+ solutions, aqueous ammonia (NH 3 ) will be added to each solution to produce the complex ion Cu(NH 3 ) 4 2+ according to the following reaction. Cu 2+ (aq) + 4NH 3 (aq) Cu(NH 3 ) 2+ 4 (aq) (3) Cu(NH 3 ) 4 2+ in solution has a dark blue color and its absorbance at 600 nm can be measured over the range necessary for this experiment. The equilibrium constant associated with equation 3 is called the formation constant and is given the symbol K f. The K f for Cu(NH 3 ) 4 2+ is 4.8 x which indicates that the forward direction of this reaction is greatly favored over the reverse. The result is that with a large excess of NH 3 present, almost all of the Cu 2+ ions are present as the complex ion, Cu(NH 3 ) Thus we will make the assumption that the [Cu(NH 3 ) 4 2+ ] is equal to the [Cu 2+ ] in the solution before the addition of NH 3. In this experiment, you will prepare a calibration curve for Cu(NH 3 ) 4 2+ from four standard Cu(NO 3 ) 2 solutions. Then you will calculate the solubility product constant for Cu(IO 3 ) 2 from measurements of the [Cu 2+ ] in five solutions saturated with Cu(IO 3 ) 2. You will also calculate the molar solubility of copper (II) iodate in pure water and in solutions containing Cu 2+ and IO 3 - and compare your results to predictions of the common ion effect. EQUIPMENT x 100 mm test tubes 3 effect c Advanced Instructional Systems, Inc. and North Carolina State University 2

3 5 stoppers 1 test tube rack 1 spatula 1 glass stir rod 1 MicroLab Spectrophotometer 1 MicroLab Spectrophotometer Instruction Sheet 6 vials ml waste beaker 1 deionized water squirt bottle 1 box of Kimwipes REAGENTS 0.25 g solid Cu(IO 3 ) 2 5 ml 5.0 x 10-4 M Cu(NO 3 ) 2 10 ml 1.0 x 10-3 M Cu(NO 3 ) 2 10 ml 5.0 x 10-3 M Cu(NO 3 ) 2 5 ml 1.0 x 10-2 M Cu(NO 3 ) 2 5 ml 5.0 x 10-3 M KIO 3 5 ml 1.0 x 10-2 M KIO 3 5 ml 3 M NH 3 deionized water SAFETY Cu(NO 3 ) 2 is listed as an oxidizer and corrosive. KIO 3 and Cu(IO 3 ) 2 are listed as strong oxidizing agents. As with all chemicals in the lab, if you come in contact with a solid, you should gently brush off the affected area with a paper towel and then flush the area with water. If you come in contact with a solution, you should flush the affected area with water. WASTE DISPOSAL All solutions generated this week must be placed in the waste bottle in the lab. Designate a waste beaker and set it aside for use during your lab. You can put the small samples of solution you will make into this beaker and empty it into the waste bottle at the end of class, instead of going back and forth to the waste bottle. Always remember not to overfill the waste bottle. If your waste bottle is full, please alert your lab instructor. PRIOR TO CLASS Please read the following sections of the Introductory Material: c Advanced Instructional Systems, Inc. and North Carolina State University 3

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