New Type of Antimony Electrode for pH Measurements T. R. BALL,WEBSTERB. SCHMIDT, AND KARLS. BERGSTRESSER, Washington University, St. Louis, Mo.
A
N INVESTIGATION of the antimony-antimony oxide electrode for measuring pH was begun by Schmidt in this laboratory in the fall of 1931. The work was practically completed when the publication of Parks and Beard (3) appeared. Aside from the fact that these authors used a vacuum-tube potentiometer for their potential measurements, the two investigations were practically identical. Parks and Beard found that the electrode functions between a pH of 2 and 7 , which is in perfect agreement with results in this laboratory. However, in calculating the pH Parks and Beard use the equation E = -0.007 0.05195 pH a t 25", while the present authors found that the 0.05813 pH a t 20" gives the proper equation E = -0.020 pH within about 0.02 unit when E is measured against a saturated calomel electrode with a Leeds and Korthrup portable potentiometer. The difference in the value of the constant cannot be accounted for by the small difference in temperature unless the antimony electrode has a very large temperature coefficient, and probably lies in the apparatus used in making the measurements. Using 48 different electrodes prepared and annealed in various ways, very little difference was found in their performance, so that the discrepancy probably does not lie in the method of preparation or purity of the electrodes. If it is assumed that the antimony-antimony oxide electrode functions by virtue of the changing antimony-ion activity in solutions of varying acidity, other antimony compounds should serve as a substitute for the oxide coating. The sulfide is immediately suggested. Mr. Schmidt prepared electrodes with sulfide coatings by suspending cast electrodes for 1 hour in hot 0.30 N nitric acid and then saturated the solution with hydrogen sulfide. The metal became coated with a thin yellowish film that was not removed with a stream of water from a wash bottle. These electrodes were found to function in the same manner as the ordinary antimony electrode and have the advantage of a greater range on the alkaline side, No claim is made that they are useful for precision measurements, but they will indicate the pH within *0.05 unit in the range from 2 to 10. They are slightly better than the oxide electrodes in the rapidity of establishing equilibrium and the difference between individual electrodes is usually not more than 2 or 3 millivolts. Time did not permit of a thorough investigation of the antimony sulfide electrode by Mr. Schmidt, but the work was carried on by Mr. Bergstresser. The data hereafter presented are taken from the work of both.
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EXPERIMENTAL PREPARATION OF ELECTRODES. Five methods were used in preparing the sulfide coated electrodes. 1. Polished stick electrodes were immersed in 0.30 N nitric acid which was kept on the steam bath for 1 hour. The hot solution was then saturated with hydrogen sulfide while it cooled to room temperature. When not in use the electrodes were kept in this saturated hydrogen sulfide solution. 2. The procedure was identical with that of the first method, except that 0.50 N nitric acid was used. 3. Polished electrodes were placed in a long Pyrex tube which was sealed at one end and connected t o a vacuum pump at the other. A small amount of sulfur was placed in the sealed end of the tube. The tube was evacuated and the portion containing
the electrodes was heated to 400" C. in an electric tube furnace. In a second run a temperature of 490' C. was used. The sulfur was then distilled onto the electrodes by applying heat t o the sealed end. The tube was allowed t o attain room temperature before the vacuum was released. 4. Polished electrodes were heated to 400" C. in a stream of dry hydrogen sulfide. 5. A heavy coating of oxide was formed on the electrodes by heating them t o 500' C. in a stream of air. The cooled electrodes were then treated with hydrogen sulfide while suspended in 0.30 N nitric acid. A complete study of these various electrodes showed that those prepared by the first method gave more nearly reproducible results and all the data in this paper were obtained from electrodes so prepared, RESULTS. Table I contains typical sets of data obtained with buffers which were standardized with quinhydrone up to a pH of 7 and with a hydrogen electrode in the more alkaline range. TABLEI. CALCULATION OF CONSTANT EO OF EQUATION 1 E. M. F.
PH
VS.
BUFFER HC1-KC1 Phthalate-HC1 Phthalate Phthalate Phosphate Phosphate Phosphate Phosphate Phosphate Borate Borate Ph o s p h a t e
PH
SAT.CAL. t o
6.97 8.01 9.20 10.00
0.0540 0.1100 0.2125 0.2136 0.2930 0.3074 0.2370 0.3739 0.3870 0.4480 0.5160 0.5616
Phosphate Phosphate Phosphate Phosphate Phosphate
10.28 10.82 11.41 11.77 12.11
0.5935 0 . 596.6 0.6438 0.6714 0.7046
1.12 2.20 3.92 3.97 5.29 5.64 6.91 6.66
20 20 24 20 20 24 20 24 20 20 20 20
PH
Eo,,. CALCULATED ERROR 0"
( -0.2406)
-0.2339 -0.2358 -0.2345 -0.2377 -0.2353 -0.2355 -0.2358 -0.2337 -0.2347 -0.2331 -0.2324 Av , -0.2348 -0.2437 23 -0.2166 22 -0.2472 23 22 -0.2359 -0.2473 23
1.21 2.18 3.93 3.96 5,33 5.54 5.92 6.67 6.95 8.00 9.17 9.96
+0.09 -0.02 +0.01 -0.01 +0,04 =to.00 +0.01 +0.01 -0.02 -0.01 -0.03 -0.04
10.43 10.51 11.60 11.79 12.40
+o. 19
4-0.15 -0.21
+0.02 $0.19
The e. m. f. values of column 3 are the averages obtained from four or more electrodes measured against the saturated calomel electrode. The Eo values of column 5 are calculated from the equation EO= pH [0.0577 0.0002 (t - IS)] - e. m. f. - Eoai. (1)
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The value of Ecai.a t any given temperature was computed from the equation (2) Esat.oal, = 0.2504 0.00065 (t - 18) (2) The average value of EOis computed over the p H range of 2.2 to 10 only, since in the more acid and alkaline ranges the electrodes behave erratically. Having established the value of the constant EOas -0.2348 volt, the pH of each solution was calculated to the closest 0.01 unit from Equation 3.
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pH
=
+
E. m. f. Ecal.- 0.2348 0.0577 (t - 18)
(3)
This equation would, of course, apply to any type of calomel electrode. The last two columns show the calculated pH and the deviation from the correct value of column two. Over the range of 2.2 to 10 the pH may be relied upon to about 0.05 unit. Individual electrodes rarely differed from others in the same solution by more than 2 or 3 millivolts if the pH was 10 or less. This is somewhat better than the oxide electrodes of Parks and Beard (3). A study of their data shows a variation between electrodes of from 3.2 to 5.2 millivolts, even in the 60
January 15, 1934
INDUSTRIAL AND ENGINEERING CHEMISTRY
most favorable range. In solutions of p H above 10, the sulfide electrodes showed a maximum variation among themselves of 11.5 millivolts. Table I shows that the sulfide electrode is not reliable in this range even when average values of several electrodes are taken. Sugar, starch, and nitrates have no deleterious effect upon the electrodes between pH 2.2 and 10. However, hydroxy acids, such as tartaric, citric, and lactic, render the electrode useless. This has been found to be true of the antimony oxide electrode as well ( 1 ) . The sulfide electrodes were used in the back-titration of alcoholic potassium hydroxide in determining the saponification number of oils. As would be expected, the strongly alkaline solution quickly removes the sulfide film, but the electrode still functions. The end point as determined by finding the maximum value of AE/AV agrees with the phenolphthalein end point. I n this respect the sulfide electrode offers no advantage over the ordinary antimonyantimony oxide electrode. SUMMARY 1. Antimony electrodes coated with antimony sulfide have been prepared by five different methods.
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2. Electrodes prepared by suspension in hot 0.30 N nitric acid for 1 hour, followed by saturation with hydrogen sulfide, may be used to determine the pH of solutions in the range from 2 to 10. 3. Electrodes so prepared agree among themselves within about 3 millivolts if the p H is 10 or less, but may differ by as much as 11.5 millivolts in more alkaline solutions. 4. Starch, sugar, and nitrates have no deleterious effect in the range over which the electrode functions in their absence. 5. The electrode, like the oxide electrode, is useless in the presence of hydroxy acids. 6. The electrode should be useful in determining the saponification of highly colored oils, but offers no advantage over the ordinary antimony electrode. LITERATURE CITED (1) Barnes and Simon, J. Am. Soc. Agron., 24, 156 (1932). (2) Kolthoff and Furman, “Potentiometric Titrations,” 3rd ed., p. 446, Wiley, 1931. (3) Parks and Beard, J. Am. Chem. Soc., 54, 586 (1932). R ~ C E I V ESeptember D 1 2 , 1933.
Colorimetric Determination of Fluorine 0. M. SMITHAND HARRISA. DUTCHER, Oklahoma Agricultural and Mechanical College, Stillwater, Okla.
T
HE.fluoride content of natural water has taken on new significance, since the researches of Smith, Lantz, and Smith (4), Churchill (2), and Kehr (3) have shown that it may be the cause of the tooth defect known as mottled enamel. Of the numerous methods which have been devised for the detection and determination of fluorides, few find satisfactory application in the field of water analysis, where a very sensitive method is required. In the modified Casares-DeBoer method used by Thompson and Taylor ( 5 ) ,the fluorides are determined by the degree of fading of a zirconium-alizarin lake. Willard and Winter (6) suggest the use of quinalizarin (1, 2, 5, 8-tetrahydroxyanthraquinone) as an indicator in their method. Quinalizarin seems to the writers to have advantages over alizarin when used as in the Casares-DeBoer method, in that it is more sensitive to small changes in fluoride content and the change in color is easier to distinguish. For example, a difference in color between samples containing 0.2 and 0.4 part per million is greater in the case of zirconium-quinalizarin than zirconium-alizarin reagent. The best range of the colorimetric standards is from 0 to 2 p. p. m. or 0.0 to 0.1 mg. of fluoride per 50 ml. Above this concentration the fading is too great, and comparisons are not easily made. RECOMMENDED METHOD The zirconium-quinalizarin reagent is prepared by mixing equal parts of a 0.14 per cent solution of quinalizarin (1,2, 5 , 8-tetrahydroxyanthraquinone)and an 0.87 per cent solution of zirconium nitrate, and diluting the mixture 1 to 40. The quinalizarin is dissolved in a 0.30 per cent sodium hydroxide solution, as it is insoluble in water. METHODOF ANALYSIS. Precipitate the sulfates by the addition of 5 ml. of 2 per cent barium chloride solution to 100 ml. of the sample. After settling several hours, draw off a 50-ml. portion for the test. The barium sulfate may be filtered off if desired. Add 3 ml. of 1 t o 1 hydrochloric acid and 5 ml. of the zirconium nitrate-quinalizarin reagent and mix thoroughly. After 20 minutes compare with standards made at the same time and in the same manner. Comparisons are easily made in
American Public Health Association tubes with standard fluoride solution containing from 0 to 2 p. p. m. in steps of 0.2 part. The fading of the color of the zirconium-quinalizarin or zirconium-alizarin lake is a function of time, temperature, and acidity. After 15 minutes a t room temperature the change is very slow and the color remains sufficiently constant for comparisons. Care must be taken that exact amounts of indicator (zirconium-quinalizarin reagent) and acid are added to the sample and to the standards and that the time of fading is the same for unknown and standard. Of the commonly occurring ions, none in the quantities occurring in natural or treated water affect the results except aluminum, iron, sulfates, and phosphates. Less than 20 p. p. m. of sulfates have no effect and may easily be removed by barium chloride, since the Ba++ ion has no effect. Iron above 10 p. p. m. changes the color, making comparisons impossible. Aluminum has no effect up to 0.2 p. p. m. as Al; above this amount and up to 0.6 p. p. m. as A1 the results will be low by 0.1 p. p. m. of fluorine. Aluminum is rarely present in water in amounts greater than 1 p. p. m. expressed as A1203or 0.5 p. p. m. as Al, and may thereby be neglected in most cases. This is further confirmed in the comparative results between those obtained by direct colorimetric and distillation methods. Phosphates affect the color when from 0.3 to 0.4 p. p. m. or more are present. The color is different and is easily recognized by one experienced in the determination. Fortunately these two substances rarely occur in water in amounts greater than 0.5 p. p. m. except in certain localities. When they do occur the distillation method seems to be the reliable procedure. In case phosphates and aluminum are present, distillation with perchloric acid as recommended by Boruff and Abbott (1) will be necessary. Place the sample containing approximately 0.2 ml. of fluoride in a 125-ml. distilling flask, and add a few glass beads and sufficient dilute sodium hydroxide to make it just alkaline t o litmus. Reduce the volume t o 10 to 15 ml. by distilling off the water, and obtain a 50-cc. distillate according t o the procedure outlined by Boruff and Abbott (1).
