Lab 10 - Solubility Product Constants

Purpose

To measure the solubility product constant (K_sp) of copper (II) iodate, Cu(IO₃)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 NH₄NO₃, were referred to as "soluble" and others that did not dissolve completely, such as AgCl and BaSO₄, 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(IO₃)₂) is placed in water, the following equilibrium is established. The equilibrium constant associated with this reaction is called the solubility product constant and is given the symbol K_sp. 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 with the ions according to equation 1. The molar solubility 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 saturated with the solid. For example, if the molarity of Cu²⁺ in a solution saturated with Cu(IO₃)₂ can be determined, the stoichiometry in equation 1 indicates that this is also equal to the molarity of Cu(IO₃)₂ that dissolved in solution or its molar solubility. Alternatively, if the molarity of IO₃^- in a solution saturated with Cu(IO₃)₂ can be determined, the stoichiometry in equation 1 indicates that the molar solubility of Cu(IO₃)₂ in water is one half the [IO₃^-] 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(IO₃)₂, it is known that 2[Cu²⁺] = [IO₃^-]. Thus, if the Cu²⁺ concentration can be measured, the IO₃^- 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(IO₃)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. In the example of Cu(IO₃)₂, the presence of either Cu²⁺ or IO₃^- in solution should result in a lower molar solubility than in pure water. Earlier this semester in Lab 2—Solutions and Spectroscopy, you prepared a calibration curve for Cu²⁺ over the range of approximately 0.1 M to 0.4 M. As you may recall, Cu²⁺ in solution has a pale blue color and you were able to measure its absorbance around 630 nm. For this experiment, it is necessary to prepare a calibration curve for Cu²⁺ over the range of 5 x 10^-4 M to 0.01 M. At these concentrations, the absorbance of Cu²⁺ is too low to detect directly by spectroscopy. To improve the absorbance of the Cu²⁺ solutions, aqueous ammonia (NH₃) will be added to each solution to produce the complex ion Cu(NH₃)₄²⁺ according to the following reaction. Cu(NH₃)₄²⁺ 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(NH₃)₄²⁺ is 4.8 x 10¹². which indicates that the forward direction of this reaction is greatly favored over the reverse. The result is that with a large excess of NH₃ present, essentially all of the Cu²⁺ ions are present as the complex ion, Cu(NH₃)₄²⁺. Thus, we will make the assumption that the [Cu(NH₃)₄²⁺] that forms after NH₃ is added to a Cu²⁺ solution is equal to the [Cu²⁺] in the solution before the addition of NH₃. In this experiment, you will prepare a calibration curve for Cu(NH₃)₄²⁺ from four standard Cu(NO₃)₂ solutions. Then you will calculate the solubility product constant for Cu(IO₃)₂ from measurements of the [Cu²⁺] in five solutions saturated with Cu(IO₃)₂. You will also calculate the molar solubility of copper (II) iodate in pure water and in solutions containing Cu²⁺ and IO₃^- and compare your results to predictions of the common ion effect.

Equipment

Reagents

Safety

Cu(NO₃)₂ is listed as an oxidizer and corrosive. KIO₃ and Cu(IO₃)₂ are listed as strong oxidizing agents. NH₃ is corrosive. These chemicals can attack the skin and cause permanent damage to the eyes. If one of these solutions splashes into your eyes, use the eyewash immediately. Hold your eyes open and flush with water for at least 15 minutes. If contact with skin or clothing occurs, rinse the affected area with water for at least 15 minutes. Have your lab partner notify your teaching assistant about the spill and your exposure. The NH₃ solution also has irritating vapors. It should be used in a fume hood to avoid inhaling the vapors. 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 rinse the area with water. If you come in contact with a solution, you should rinse the affected area with water. Students will have access to gloves due to the use of the concentrated ammonia solution during the lab period.

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 lab section's waste bottle is full, please alert your teaching assistant.

Prior to Class

Please read the following sections of the Introductory Material:

• Plotting a Calibration Curve
• Centrifuges

Please complete your WebAssign prelab assignment. Check your WebAssign account for due dates. Students who do not complete the WebAssign prelab assignment are required to print and hand in the prelab worksheet.

Lab Procedure

Please print the worksheet for this lab. You will need this sheet to record your data.

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