Solubility of CaSO4

Solubility of CaSO4 Experiment 8 Major Concepts and Learning Goals ∙ ∙ ∙ ∙ ∙ ∙ Application of the solubility product constant (Ksp) Saturated solutions Le Chatlier’s Principle/Common ion effect Activities and activity coefficients Ion selective electrodes Calibration curves Laboratory Task ∙ Produce a calibration curve using standard solutions of CaNO3 ∙ Measure the [Ca2+] of three different solutions 1) a saturated solution of CaSO4 in H2O 2) a saturated solution of CaSO4 in 0.10 M KNO3 3) a saturated solution of CaSO4 in 0.10 M Na2SO4 ∙ Observe the effect of ionic strength on the [Ca2+] by comparing the results of solution 1 and 2. ∙ Observe the effect of the Na2SO4 on the [Ca2+] (common ion effect). Introduction The solubility of CaSO4 at 25 ºC is described by the following reaction and equilibrium CaSO4(s) ↔ Ca2+ + SO42Ksp(CaSO4) = [Ca2+][SO42-] = 2.4∙10-5 eq. 1 In words, this equilibrium expression implies that the product of the calcium ion concentration and the sulfate ion concentration can be no larger than 2.4∙10-5 in any aqueous solution. Saturated solutions Any aqueous solution in which the product of the calcium ion concentration and the sulfate ion concentration is about 2.4∙10-5 is said to be a saturated CaSO4 solution. If a little more Ca2+ or SO42- is added to a saturated CaSO4 solution the equilibrium will shift to the left to form solid CaSO4 and the value of the product of the calcium ion concentration and the sulfate ion concentration would be restored to about 2.4∙10-5. This statement is the basis of Le Chatlier’s Principle. When an equilibrium position of a reaction is disturbed, a new equilibrium position will be established by shifting the reaction in a direction that alleviates the stress caused by the disturbance Saturated solutions can be prepared by a variety of methods. In this experiment the first saturated solution has been prepared by adding solid CaSO4 to purified water (the water comes from a purification system that includes a carbon filter, an ion-exchange resin and a UV lamp). The solution was mixed for several days and allowed to settle and reach equilibrium for several weeks. Then the solution was filtered to remove any suspended CaSO4 particles that did not dissolve or settle to the bottom. Because the only source of the calcium ions and sulfate ions come from the dissolution of CaSO4, [Ca2+] = [SO42-]. In fact, the same statement can also be made for the second saturated solution, since KNO3 is not a source of Ca2+ or SO42-. When we ignore the impact of ionic strength, the [Ca2+] and [SO42-] concentrations of these first two saturated solutions are expected to be about 5.0∙10-3 M, based on the literature Ksp value for CaSO4 {(2.4∙10-5)1/2}. Le Chatlier’s Principle and the Common Ion Effect One general case in which Le Chatlier’s principle can be applied is when the solution contains a soluble salt of an ion that is in common with the insoluble salt in question. This is the case in the third saturated solution; CaSO4 in dissolved 0.10 M Na2SO4. In this solution there are two sources of the sulfate ion; the Na2SO4 and the CaSO4. The SO42- coming from the Na2SO4 is 0.10 M, because Na2SO4 is a completely soluble salt. The sulfate ion coming from the CaSO4 is equal to the calcium ion concentration. Thus, Ksp = [Ca2+] ([SO42-]CaSO4+ + [SO42-]Na2SO4) = 2.4∙10-5 Letting [Ca2+] = x, we arrive at Ksp = x (x + 0.10) = 2.4∙10-5 If we assume x <<< 0.10 M, then Ksp = x 0.10 = 2.4∙10-5 and [Ca2+] = 2.4∙10-4 M, This is considerably lower than the first saturated solution. It is also worth pointing out that our solution to the problem verifies that our stated assumption was valid. This lowering of the [Ca2+] is due to the common ion effect. Activities and Activity Coefficients In reality, equilibria are affected by the concentration of ions, any ions, in solution. The ionic strength ( ) is used as a measure of the total ion concentration of a solution. It is calculated by incorporating each of the “i” ionic species in solution into the following equation, where C is the concentration in mol/L and Z is the charge on the ion. = ½ ∑(CiZi2) So, why does the ionic strength matter? Well let’s look at the CaSO4 equilibrium as an example. The ionic strengths of the saturated solution 2 and 3 are considerably greater than that of saturated solution 1. Initial estimates of the ionic strength of these three saturated solutions (these are initial estimates because the concentrations of [Ca2+] and [SO42-] were determined ignoring the impact of ionic strength) Saturated solution 1 = ½ ([Ca2+](2)2 + [SO42-](-2)2)) = ½ (5∙10-3(2)2 + 5.0∙10-3(-2)2)) = 0.02 M Saturated solution 2 = ½ ([K+](1)2 + [NO3-](-1)2 + [Ca2+](2)2 + [SO42-](-2)2)) = ½ (0.1(1)2 + 0.1(-1)2 + 5∙10-3(2)2 + 5.0∙10-3(2)2)) = 0.12 M Saturated solution 3 = ½ ([Na+](1)2 + [SO42-](-2)2) + [Ca2+](2)2) = ½ (0.2(1)2 + 0.1(-2)2 + 2.4∙10-4(2)2 ) = 0.20 M Electrostatic interactions between the negative ions and the Ca2+ ions in solution cause the formation of an ion cloud around the Ca2+ ions. The larger the ionic strength of the solution the greater the radius of this ion cloud. A similar ion cloud forms around the SO42- ions from the positive ions in solution. The size of the ion clouds about Ca2+ and SO42- defines the energetics associated with these ions finding each other in solution and forming CaSO4(s). Thus, the [Ca2+] and the [SO42-] in the second saturated solution would be expected to be significantly greater than 5.0∙10-3 M, and in the third saturated solution the [Ca2+] would be expected to be greater than 2.4∙10-4 M. Mathematically, the impact of ionic strength is accounted for by introducing the concept of ion activity, A. The activity of an ion can be thought of as its effective concentration and is given by product of its concentration and activity coefficient, . The activity coefficient depends upon the size of the ion, its charge, and the ionic strength of the solution. For example: ACa2+ = [Ca2+] Ca+ The activity coefficient can be calculated using the Debye-Huckle equation. The most important concept to appreciate in terms of activities is that the definition of the equilibrium expression that we first learned and used (eq. 1) is only a model. It works very well at low ionic strengths, where the activity coefficients are close to unity, but breaks down at higher ionic strengths. The more universal model expresses the equilibrium constants in terms of activities. Equation 1 becomes Ksp(CaSO4) = ACa2+ASO42- = [Ca2+] Ca2+[SO4 2- ] SO42- = 2.4∙10-5 eq. 2 and SO42- for solution 2 are less than for solution 1 because the ionic strength of solution 2 is greater than the ionic strength of solution 1, and as a result [Ca2+] and [SO42-] must be larger in solution 2 for both solutions to satisfy the Ksp expression. However, it is important to appreciate that the ionic strength of solution 1 is large enough to have an impact on the solubility. So, even for this solution, it is expected that the solubility will be significantly larger than the 5.0 mM predicted above. Ca2+ Ion- Selective Electrodes Ion-selective electrodes are electrochemical cells that have been carefully crafted into a probe that is useful for measuring the activity of a specific ion in solution. An ionselective electrode most often consists of two reference electrodes, which give a constant potential. The two reference electrodes are electrically linked via a voltmeter and a salt bridge through the solution being measured. The salt bridge consists of a membrane between the solution being measured and the inner reference electrode solution. This membrane is made of a unique material that preferentially allows the ion of interest to partially penetrate the membrane. This partial penetration leads to the development of a junction potential, the magnitude of which depends entirely on the activity of the ion in the solution being measured. The Ca2+ selective electrode used in this application contains a membrane that preferentially chelates Ca2+. The potential measured by the voltmeter depends upon the potential at each of the reference electrodes and the junction potential. The potential of the reference electrodes is constant and junction potential is only influenced by the activity of the Ca2+ in the solution being measure. Ecell = constant + 59.16/2 log ACa2+ = constant + 59.16/2 log [Ca2+] Ca2+ Typically, one would measure the voltage of several standard Ca2+ solutions in the range of 0.1-10-6 M Ca2+ at a constant ionic strength. To measure the concentration of [Ca2+], the usual procedure is to measure the potentials of the standard solutions with a Ca2+ selective electrode and use the data to produce a calibration plot. To make the analysis of the data straightforward and simple all of the standards and samples are prepared using an electrolyte of high ionic strength. As a result, all of the solutions have the same ionic strength and, as a result, the activity coefficient becomes a constant. Applying the rule of logs log AB = log A + log B Ecell = constant + 59.16/2 log [Ca2+] Ca2+ Becomes Ecell = constant + 59.16/2 log [Ca2+] + 59.16/2 log Ca2+ Ecell = constant1 + 59.16/2 log [Ca2+] + constant2 Combining constants Ecell = constant + 59.16/2 log [Ca2+] Then, a calibration plot is produced by plotting the cell potential as a function of the log [Ca2+]. Then, your samples are prepared by diluting it quantitatively with the same electrolyte. This ensures that the ionic strength of your sample solutions match that of your standards. Experimental Note: In this experiment once you dilute the saturated solutions, they will no longer be saturated. Since all of the left over CaSO4 will be removed by filtration, the [Ca2+] in each of the samples is reduced by a known dilution factor, which will be taken into account during your calculations. In the Lab Work in pairs. Several of these steps can be performed simultaneously, which will allow us to complete this lab within the allotted time. I have assigned the tasks to student A and B. 1. Use a syringe filter to filter approximately 4 or 5 mL of the each of the three saturated solutions in to three clean, dry and labeled 20 mL screw top bottles. (Student A) 2. Obtain about 600 mL of 0.5 M KNO3 in a clean beaker. Rinse your 25 mL pipet at least four times with the DI water and once with 0.5 M KNO3. Add 25 mL of 0.5 M KNO3 to each of twelve clean, dry Erlenmeyer flasks. Label them std 1-6 and test 1-6. (Student B) 3. Obtain about 40 mL of the 0.100 M Ca(NO3)2 master standard in a clean, dry beaker and label it. Pipet 25 mL of the 0.100 M Ca(NO3)2 standard into a clean, dry 100 mL volumetric flask and dilute to the mark with 0.50 M KNO3. Place the cap on and mix thoroughly by inverting the flask back and forth about 20 times. Pour the contents into a clean and dry beaker. Label this solution 0.0250 M Ca(NO3)2. Immediately, rinse your 25 mL pipet at least four times with the deionized water. (Student A) Preparation of the six standards 4. Use the micropipeter to pipet, 1.000 mL of the 0.100 M Ca((NO3)2 master standard into the flask labeled standard 1 and 0.500 mL of the 0.100 M Ca((NO3)2 master standard into the flask labeled standard 2. Swirl each flask to mix thoroughly. (STUDENT B) 5. Switch tips. Use the micropipeter to pipet 1.000 mL of the 0.0250 M Ca((NO3)2 standard into the flask labeled standard 3 and 0.500 mL of the 0.0250 M Ca((NO3)2 standard into the flask labeled standard 4. Swirl each flask to mix thoroughly. (STUDENT B) 6. Switch tips. Use the micropipeter to pipet 2.000 mL of the standard 1 into the flask labeled standard 5 (use the 1 mL micropipeter twice) and 1.000 mL of standard 1 into the flask labeled standard 6. Swirl each flask to mix thoroughly. (STUDENT A) Preparation of the six test samples 7. Switch tips. Use the micropipeter to pipet 2.000 mL of the filtered solution of saturated CaSO4 in DI water into the flask labeled “TEST 1” (use the 1 mL micropipeter twice) and 1.000 mL of the filtered solution of saturated CaSO4 in DI water into the flask labeled “TEST 2”. (STUDENT A) 8. Switch tips. Use the micropipeter to pipet 2.000 mL of the filtered solution of saturated CaSO4 in 0.10 M KNO3 into the flask labeled “TEST 3” (use the 1 mL micropipeter twice) and 1.000 mL of the filtered solution of saturated CaSO4 in 0.10 M KNO3 into the flask labeled “TEST 4”. (STUDENT B) 9. Switch tips. Use the micropipeter to pipet 2.000 mL of the filtered solution of saturated CaSO4 in 0.10 M Na2SO4 into the flask labeled “TEST 5” (use the 1 mL micropipeter twice) and 1.000 mL of the filtered solution of saturated CaSO4 in 0.10 M Na2SO4 into the flask labeled “TEST 6”. (STUDENT B) Summary Tables of Solution Preparation Standards (Calculate the [Ca2+] for these standards and show work to instructor prior to leaving lab) solution 0.0250 M std 1 std 2 std 3 std 4 std 5 std 6 [Ca2+] 0.25 M KNO3 0.100 M Ca2+ 0.0250 M Ca2+ std 1 75.00 25.00 0.000 0.000 25.00 1.000 0.000 0.000 25.00 0.500 0.000 0.000 25.00 0.000 1.000 0.000 25.00 0.000 0.500 0.000 25.00 0.000 0.000 2.000 25.00 0.000 0.000 1.000 Test solutions of diluted saturated solutions solution 0.25 M KNO3 (CaSO4)sat. in H2O (CaSO4)sat. in KNO3 (CaSO4)sat. in Na2SO4 test 1 25.00 2.000 0.000 0.000 test 2 25.00 1.000 0.000 0.000 test 3 25.00 0.000 2.000 0.000 test 4 25.00 0.000 1.000 0.000 test 5 25.00 0.000 0.000 2.000 test 6 25.00 0.000 0.000 1.000 10. Get help from one of your instructors to set up the apparatus. Measure the cell voltages of each of the six standards and the six test solutions. To make the measurement you will have to transfer 15-20 mL of each solution into the 20 mL screw cap bottle. This is necessary to ensure that the salt bridge of the Ca-ion selective electrode is submersed when making the measurements. When switching between samples, thoroughly rinse the 20 mL screw cap bottle with several portion of DI water and then with a couple of milliliters of the test solution. Record all data on the Lab Pro Software and copy and paste the data into an Excel spread sheet. Give each solution at least 3 minutes to equilibrate with the ion-selective membrane prior to taking a reading. Lab Report Data Use the template to do the following calculations: 1. Calculate the [Ca2+] of your six standards. (5 pts) 2. Plot the log[Ca2+] vs. the voltage readings from the Ca-ion selective electrode. (10 pts) 3. Use the Analysis Tools function to obtain the regression parameters and statistics. (5 pts) 4. Using the regression parameters to determine the [Ca2+] of the six diluted test samples that were measured. Determine the uncertainty in these measurements using the regression statistical analysis outline in the “Introduction to Statistics”. This particular analysis is complicated by the logarithm. The Linear Regression statistical analysis (the ugly Sy/m equation) provides the uncertainty in the log[Ca2+]. You need to apply the rule for propagating uncertainty for logarithmic functions (sCa2+ = Slog[Ca2+][Ca2+]) (10 pts) 5. Use the dilution factors to determine the [Ca2+] in the three saturated solutions. You will have two independent measurements for each. Propagate all uncertainties. Report the average of the two results. (10 pts) 6. Based on the [Ca2+] you found for the saturated solution in DI water, estimate the product of the activity coefficients of Ca2+ and SO42- in this solution ( Ca2+ SO42-). (5 pts) 7. Using the [Ca2+] you found for the saturated solution in DI water, calculate the ionic strength of the solution and use the extended Debye-Huckle equation Harris in the Activities chapter) to calculate the Ca2+ and SO42-. (5 pts) log = -0.51z2 1/2/[1+( 1/2/305)] Points to address in the context of your discussion: 1. Were the two independent measurements for each saturated solution in agreement? Calculate the differences between two independent measurements for each of the three saturated solutions. (5 pts) 2. Are the solubilities of CaSO4 in these three different solutions significantly different? Compare and rationalize the relative magnitudes of the [Ca2+] in the three saturated CaSO4 solutions. (5 pts) 3. Does the ionic strength of the solution of saturated CaSO4 in de-ionized water have a significant impact on its solubility? In other words how does the solubility of CaSO4 in de-ionized water compare to the prediction in the writeup above from Ksp? How well does the product of the activity coefficients calculated in #6 above match the activity coefficients estimated from the extended Debye-Huckle Equation in #7? (5 pts) 4. The solubility of CaSO4 is actually more complicated than what has been considered in this experiment because there are some complex ion issues (see section 8-5 in your Harris textbook). Discuss how this complicates the interpretation of the results in this experiment. (5 pts) 5. Is the common ion effect evident in the data for the [Ca2+] of the saturated CaSO4 solution in 0.1 M Na2SO4? Explain your answer. (5 pts)