In the Laboratory
The Purification of Water by Freeze–Thaw or Zone Melting James Oughton and Silas Xu St. Andrews College, Christchurch, New Zealand Rubin Battino* Department of Chemistry, Wright State University, Dayton, OH 45435;
[email protected] Background The purification process known as zone refining or zone melting began with Pfann’s 1952 paper (1), although the basis of the method of separation known as fractional crystallization had been known for centuries. In fact, there is evidence that more than 2000 years ago Scandinavians obtained salt by partially freezing sea water in the winter, discarding the ice, and boiling water away from the remaining more concentrated brine solution. There are three books on the subject (2–4), all old, which go into detail on the physical chemistry of the separations, experimental methods, and illustrations. A number of articles on the subject have been published in this Journal (5–11), including a tested demonstration on the zone refining of a mixture of naphthalene and methylene blue (9). Industrially, the zone melting technique has been of major importance in the production of sufficiently pure germanium and silicon for transistors, and other semiconductors. The traditional zone melting–refining method utilizes a traveling molten zone along a horizontal or vertical column of solid. As the molten zone advances the crystallizing material at the rear of the zone is usually purer solidified solvent than the melt. (In some systems impurities concentrate in the forming crystal rather than the melt.) For NaCl/H2O brine solutions, the ice that forms is purer water than the solution. As described in Schoen’s edited volume (11), zone refining has been used to convert sea water to fresh water and for the separation of heavy water. This paper investigates quantitatively the purification of NaCl/H2O solutions via the process of partial freezing. (This process is related to recrystallization as a purification method.) NaCl concentrations are determined by electrical conductivity. Since two of us (JO and SX) were high school students, we aimed for simplicity, low expense, and duplication in any high school laboratory. Except for the conductivity meter, all of the equipment is widely available. (Some inexpensive ways of measuring conductivity were tried, but were unreliable.) Experimental Method To obtain credible results, the water used must be consistent in its purity, and conductance measurements must be made at the same temperature (conductivity increases about 2%/°C). Christchurch tap water was used for all experiments, since it showed no variation in conductivity when measured over several days at 13 °C—the temperature for all of the measurements. The conductivity of the tap water used at 13 °C was 0.097 mS/cm, which contrasted with a conductivity of 0.0022 mS/cm for distilled water. The exceptional purity of this tapwater made it unnecessary to use distilled water, which may be required in other locations. The NaCl was purchased
from Analar and had a minimum purity of 99%, although supermarket NaCl may also be used. The test solutions were made by weighing the NaCl in a 100-mL beaker to ±1 mg or about 3 significant figures, using an electronic balance (Model ETD series 3). Small amounts of water were added to completely dissolve the salt. This solution was quantitatively transferred to a 100-mL volumetric flask, with appropriate rinsing. The solution was then diluted to the mark. The salt solutions were poured into plastic soft drink cups (standard clear polystyrene, 150 mL) and labeled. This was repeated for the six NaCl solutions tested: 0.5 to 5.5 g/100 mL in 1-g increments. (Note that about 6 g of NaCl will saturate 100 mL of water at 13 °C.) The cups of solution were placed in the freezer compartment of a refrigerator and left there until the solutions were approximately 25, 50, and 75% frozen. As the concentration of the salt increases, the melting point of the solution decreases. This means that it takes longer for solutions of higher concentration to attain the same percentage frozen. Location within the freezer also affects the rate of freezing. The freezer is colder around the perimeter than in the center. The cups must be arranged so that all are around the perimeter or all are in the center. For our experiments, the cups were placed around the perimeter. All solutions experienced the same temperature. Before any readings were taken a trial was undertaken to establish how much water was absorbed on the cloth mesh filter and whether conductivity readings would be affected. A 100-mL solution was made at a concentration of 5.5 g/ 100 mL. This solution was poured through the cloth filter into a funnel for collection in another measuring cylinder. This procedure was repeated 3 times, and each time more than 99 mL was collected. Conductivity was also measured. The readings were within 1 mS/cm on either side of the reference conductivity of 83.1 mS/cm. This showed that the cloth mesh filter does not affect the results of the experiment. For the experiment, the partially frozen solutions were removed from the freezer and the liquid portion was rapidly filtered through a cloth mesh filter (dish towel) in a funnel into a 100-mL graduated cylinder. The amount of liquid accumulated and the percentage frozen were recorded. The conductance of the liquid portion was measured using a conductivity meter (ETD Series 3). Any conductivity meter of 1% accuracy or better (ours was ±0.3% of the reading) and full-scale ranges of 2 to 200 mS/cm will work. The A. Daigger & Co. Catalog lists a variety of such meters, many portable, in the $400–500 range. After the ice melted in the cup, its conductance was measured. Very little ice was caught in the cloth filter; a spatula was used to scrape out what was caught. Some water was absorbed, but it was a small amount and did not affect the conductivity readings.
JChemEd.chem.wisc.edu • Vol. 78 No. 10 October 2001 • Journal of Chemical Education
1373
In the Laboratory
Table 1. Conductivity of Aqueous NaCl Solutions at 13 ⴗC after Partial Freezing and Thawing NaCl/(g/ 100 mL)
Conductivity /(mS/cm) Initial
25% Frozena Liquid
d
9.98
Ice
e
6.88
50% Frozenb
75% Frozenc
Liquid
Ice
Liquid
12.2
7.10
15.4
Ice
0.5
9.4
1.5
25.8
27.6
22.1
30.9
18.8
36.7
22.5
2.5
40.1
40.0
36.3
47.7
30.3
52.4
34.2
3.5
54.8
57.7
38.4
63.5
44.6
69.2
44.7
4.5
66.6
69.8
51.2
81.6
56.8
84.5
62.9
5.5
83.1
89.2
73.1
88.9
69.1
97.7
74.3
aAverage
24.5%, SD = 4.0%. 49.3%, SD = 3.6%. cAverage 75.6%, SD = 2.1%. dPortion of the solution that remained unfrozen. eAfter melting. bAverage
8.44
A calibration curve was prepared of conductivity versus NaCl concentration to determine the concentrations of the salt solutions. Since it was difficult to control the extent of freezing (something that was determined after separating ice and liquid for each sample), the results for runs with percentages of freezing closest to 25, 50, and 75% were averaged. Results The results for the measurements are given in Table 1. It is readily observable that the conductivity of the residual solution is significantly higher than that of the melted ice for each initial concentration and each (average) percentage of frozen solution. This is shown graphically in Figure 1 for the unfrozen water and in Figure 2 for the melted ice. Discussion The results show that zone melting, as carried out in this work as a freeze–thaw approach, is a viable method for the purification of water. An examination of the data and the figures shows that the most effective results were obtained when the salt solution was frozen to 50%. This is probably due to incorporation of relatively more salt into the crystalline ice phase as the remaining liquid solution becomes more concentrated. That the initial concentration of the salt solutions had little effect on the separation process is shown by the relatively straight lines on in graphs. Obvious variations of this experiment would use other aqueous salts. For example, CaCl2 yields three ions in aqueous solution compared to two for NaCl and also has a different conductivity. Tests with other salts, such as CaCl2, would demonstrate the generality of the zone refining process. High school students should find this approach to zone refining interesting and fun. Acknowledgments
Figure 1. Conductivity of unfrozen water at 13 °C.
RB thanks the University of Canterbury for an Erskine Fellowship. SX and JO thank St. Andrews College Science Department for the use of facilities and equipment. Literature Cited
Figure 2. Conductivity of melted ice at 13 °C.
1374
1. Pfann, W. G. Trans. AIME 1952, 194, 747. 2. Pfann, W. G. Zone Melting; Wiley: New York, 1958. 3. Herington, E. F. G. Zone Melting of Organic Compounds; Wiley: New York, 1967. 4. Schildknecht, H. Zone Melting; Academic: New York, 1966. 5. Zief, M.; Ruch, H.; Schramm, C. H. J. Chem. Educ. 1963, 40, 351. 6. Knypl, E. T., Zielinski, K. J. Chem. Educ. 1963, 40, 352. 7. Hinton, J. F.; McIntyre, J. M.; Amis, E. S. J. Chem. Educ. 1968, 45, 116. 8. Lippert, I. S.; Ritter, J. E. Jr. J. Chem. Educ. 1969, 46, 650. 9. Hawkins, M. D. J. Chem. Educ. 1976, 53, A48. 10. Needham, G. F.; Boehme, G.; Willett, R. D.; Swank, D. D. J. Chem. Educ. 1982, 59, 63. 11. Christian, J. D. J. Chem. Educ. 1956, 33, 32. 12. Schoen, H. M. New Chemical Engineering Separation Techniques; Interscience: New York, 1962; p 218.
Journal of Chemical Education • Vol. 78 No. 10 October 2001 • JChemEd.chem.wisc.edu
