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Lab 27 - Determination of K sp for a Sparingly Soluble Salt

Lab 27 - Determination of K sp for a Sparingly Soluble Salt . Introduction

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Lab 27 - Determination of K sp for a Sparingly Soluble Salt

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  1. Lab 27 - Determination of Ksp for a Sparingly Soluble Salt Introduction Inorganic substances are broadly classified as acids, bases, or salts. With few exceptions, all common salts are strong electrolytes. However, the solubilities of salts span a broad spectrum, ranging from slightly or sparingly soluble to very soluble.

  2. This experiment is concerned with heterogeneous equilibria of slightly soluble salts. For true equilibrium to exist between a solid and solution, the solution must be saturated. Thus the solubility equilibrium constant can be expressed solely on the concentrations of the ions.

  3. The rule for writing the solubility product expression states that Ksp is equal to the product of the concentration of the ions in solution, each raised to the power of the coefficient in the equilibrium equation. For the salt MA, the dissociation proceeds as: MA M+ + A-Ksp = [M+][A-] For the salt MA2, the dissociation proceeds as: MA2 M+ + 2A-Ksp = [M+][A-]2

  4. To determine the Ksp for a sparingly soluble salt • we need only determine the concentration of one of the ions, because the concentration of the other ion is related to the first ions concentration by a simple stoichiometric relationship. In this experiment, we will determine the Ksp for silver chromate, Ag2CrO4, by measuring the concentration of the yellow chromate ion, CrO42-, using a spectrophotometer.

  5. Ag2CrO4(s) 2Ag+(aq)+ CrO42-(aq) • for which the Ksp = [Ag+]2 [CrO42-].

  6. A. Preparation of calibration curve • Initial [CrO4-2] 0.0024M

  7. Molar absorption • Absorption / molarity 1. 0.041 / 2.4e-5 = 1708 plot the absorbance versus concentration to construct the calibration curve

  8. B. Determining Ksp of Ag2CrO4 • 1. Using pipettes, accurately prepare 3 separate solutions, • in a 150 mm test tube 5 mL of 0.004 M AgNO3 to 5 mL of 0.0024 M K2CrO4. Stopper each test tube. Shake the solution thoroughly at periodic intervals over a 10 min period. This will establish equilibrium between the solid phase and ions in solution.

  9. B. Determining Ksp of Ag2CrO4 • 2. Transfer approximately 3 mL of each solution and most of the insoluble Ag2CrO4 to 75 mm test tubes and spin in a centrifuge for two minutes. Discard the supernatant and retain the precipitate.

  10. B. Determining Ksp of Ag2CrO4 • 3. To the test tube add 2 mL of 0.25 M NaNO3. Shake as before, then centrifuge.

  11. B. Determining Ksp of Ag2CrO4 • 4. Transfer the clear, pale yellow supernatant liquid from the test tube to a clean, dry cuvette. Measure and record the absorbance of the solution. Using the calibration curve, calculate the molar concentration of CrO42- in the solution

  12. Analysis • Using the calibration curve, calculate the molar concentration of CrO42- in the solution.

  13. Lab 28 – Heat of neutralization Introduction Most chemical reactions are accompanied by a chance in energy, usually in the form of heat. The energy change of a reaction is called the heat of reaction or enthalpy. The symbol, DH, is used to denote the enthalpy change. A chemical change that has heat given off during the reaction is said to be exothermic andDH is negative (<0); if heat is absorbed, the reaction is endothermic and DH is positive (>0).

  14. In this experiment you will measure the heat of neutralization when an acid and a base react to form water. The quantity of heat given off during the exothermic reaction will be estimated using a thermally insulated vessel called a calorimeter. The temperature change of the calorimeter and its contents is measured before and after the reaction. The change in enthalpy, DH, is equal to the negative product of the temperature change, DT, times the heat capacity (C) of the calorimeter and its contents

  15. . Heat Capacity of the Calorimeter • Procedures • 1. Your instructor (me) has determined the heat capacity (C) of the styrofoam calorimeter in J/K. • For each degree rise in temperature, the calorimeter gains 10.2 Joules of energy.

  16. Heat of Neutralization of HCl – NaOH • 1. Add exactly 50 mL of 1 M NaOH to the dry calorimeter. Measure out exactly 50 mL of 1 M HCl and place it near the calorimeter. Wait 3-4 minutes.

  17. Heat of Neutralization of HCl – NaOH • 2. Measure the temperature of the acid; rinse the thermometer with tap water and wipe dry. • Insert the thermometer into the calorimeter and measure the temperature of the NaOH solution. The temperatures of the two solutions should not differ by no more than 0.5oC.

  18. Heat of Neutralization of HCl – NaOH • Lift the lid and carefully add the 1.0 M HCl all at once. • Stir the solution carefully with the thermometer and record the highest temperature. • Calculate the heat of neutralization per mole of water formed.

  19. calculations Heat gained by soln • T x mass x specific heat (4.18 J/k-g) Heat gained by calorimeter • T x heat capacity of calorimeter (10.2J / K)

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