Spectrophotometric microdetermination of sulfate

Spectrophotometric Microdetermination of Sulfate Joe B. Davis1 and Frederick Lindstrom Department of Chemistry and Geology, Clemson University, Clemson, S.C. 29631

A spectrophotometric method has been developed for the selective determination of sulfate in the 0-10 and 0-100 microgram (ppm in water) ranges. The aqueous samples are treated with a mixture of hydriodic acid, acetic anhydride, and sodium hypophosphite and heated in a modified countercurrent reaction apparatus to evolve hydrogen sulfide. This gas is swept by nitrogen into a buffered solution of ferric ion and 1,lO-phenanthroline, where it reduces ferric ion to ferrous ion. The bright orange tris(1,lO-phenanthroline)iron(ll) complex is formed and is measured spectrophotometrically at 510 nm. Of over 20 common ions, only nitrite and those ions capable of yielding hydrogen sulfide under the same conditions interfered. Sulfonated surfactants did not interfere. SULFATE REDUCTION to hydrogen sulfide is a very attractive basis for a simple, selective, and sensitive method for the determination of this extremely important anion. Use of this one step eliminates the shortcomings of barium sulfate and related precipitation methods. Hydrogen sulfide is a gas, so its quantitative transfer presents no problems. It is highly reactive, so its measurement is also easy. Determination of small amounts of sulfate thus need be no more difficult than the determination of small amounts of nitrogen by the Kjeldah1 method. Sulfate reduction, however, is a problem because the reaction is highly irreversible. Thermodynamically, sulfate should be reduced in a Jones reductor. In reality, few reagents are known which can satisfactorily accomplish the task, and reduction has best been carried out by the use of hydriodic acid under drastic conditions. It has been found in this work that a mixture of hydriodic acid and acetic anhydride containing a small amount of sodium hypophosphite will reduce sulfate to hydrogen sulfide rapidly, conveniently, and quantitatively in a two-stage reaction system. A method developed earlier for sulfur dioxide has been modified for microgram measurement of the evolved hydrogen sulfide. The combination has proved very successful. St. Lorant was the first to employ hydriodic acid as a reductant in sulfate analysis ( I ) . The amount of sulfide produced was determined spectrophotometrically by forming methylene blue. Luke determined macro amounts of sulfate by hydriodic acid reduction followed by iodometric titration of the evolved products (2, 3). He determined micro amounts by forming a colloidal suspension of lead sulfide and measuring the absorbance of the suspension (4). The latter method required very precise control of conditions to obtain reproducible results. Gustaffson made an extensive study of sulfate reduction by hydriodic acid ( 5 ) and the methylene blue reaction (6). She chose to reduce dry samples in a medium of hydriodic acid, acetic acid, and sodium hypophosphite.

Present address, Department of Chemistry and Physics, Winthrop College, Rock Hill, S.C. 29730 I. St. Lorant, 2.Phys. Chem., 185, 245 (1929). C. L. Luke, IND.ENG.CHEM., ANAL.ED.,15, 602 (1943). Ibid., 17,298 (1945). C. L. Luke, ANAL.CHEM.,21, 1369 (1949). Lilly Gustaffson, Tulunra, 4, 236 (1960). (6) Ibid., p 227.

(1) (2) (3) (4) (5)

524

~~~~

ANALYTICAL CHEMISTRY, VOL. 44, NO. 3, MARCH 1972

~ ~

~

~~~

Although the methods of these authors are applicable in the microgram range, they do possess disadvantages. One of these is intolerance for water, which requires time-consuming preliminary sample drying with the dangers of contamination and loss. Another is that the methylene blue reaction requires that hydrogen sulfide first be trapped by an absorbant, then reacted with other reagents in strongly acid solution. Color development is slow and there is the possibility of loss of unreacted hydrogen sulfide. A simpler procedure applicable to aqueous samples without pretreatment was desired. The hydrogen sulfide detection step should produce a color that may be measured without delay. The procedure developed by Stephens and Lindstrom for sulfur dioxide produced a color immediately and also responded to hydrogen sulfide (7). Most methods of sulfate analysis at the microgram level have been based on precipitation reactions (8). A sulfate selective electrode with an advertised range of 1-10,000 ppm sulfate and containing a membrane of barium sulfate suspended in a silicone rubber matrix has been marketed (9). This type of sulfate electrode, however, has been reported to be unsatisfactory because of low selectivity (10). An instrument based on the principle of chemiluminescence has been reported to be applicable to sulfate analysis down to 2 ppm (11). THEORY

Sulfate Reduction. The following half reactions and potentials are of interest (12): SO4*-

+ 4H+ + 2e = H 2 S 0 3+ H 2 0 E"

=

+0.17 V (1)

+ 4H+ + 4e = S + 3H20 E" = +0.45 V E" = f0.141 V S + 2H+ + 2e = HzS

H2S03

Iz + 2e HSO,

=

21-

E"

=

+0.536V

(2)

(3)

(4)

+ 2H+ + 2e = H3P02 + HzO E"

=

-0.50 V (5)

Combination of Equations 1, 2 , and 3 gives the half reaction for the reduction of sulfate to hydrogen sulfide: S042-

+ 10H+ + 8e = HzS + 4H20 E"

=

+0.303 V (6)

Where the reduction is carried out by hydriodic acid, the reaction is, combining Equations 4 and 6: (7) B. G. Stephens and F. Lindstrom, ANAL.CHEM.,36, 1308 (1964). (8) M. J. Fishman and B. P. Robinson, ibid., 41, 323R (1969). (9) National Instrument Laboratories, Inc., 12300 Parklawn Drive, Rockville, Md. 20852. (10) G. A. Rechnitz, Chem. Eng. News, 45, (25), 146 (1967). (11) W. L. Crider et al., Abstracts, 20th Annual Pittsburgh Con-

ference on Analytical Chemistry and Applied Spectroscopy, March 1969, No. 69. (12) W. M. Latimer, "Oxidation Potentials," 2nd ed., PrenticeHall, Englewood Cliffs, N.J., 1952.

Sod2-

+ 10H+ + 81-

= HzS

+ 412 + 4Hz0 E"

= -0.233 V

(7)

The negative standard potential and the free energy of reaction of f42.99 kcal indicate that the reduction is thermodynamically unfavorable. The free energy of reaction, calculated by use of the Gibbs-Helmholtz equation and the enthalpies of formation of reactants and products in their standard states at 25 "C (13), decreases to +28.4 kcal at 127 "C, the temperature of refluxing acetic acid and hydriodic acid used in this work. Several factors are capable of adjustment to favor the reduction: rapid removal of hydrogen sulfide, highly acidic conditions, a high ratio of iodide to iodine, a low water activity, and high temperature. These conditions are mutually attainable in hot concentrated hydriodic acid. One danger, however, is increased by such conditions: possible loss of partially reduced sulfate in the form of sulfur dioxide. The reduction of sulfate to sulfurous acid : S04'-

+ 4H+ + 21-

= HzS03

Figure 1. Digestion unit for preparation and storage of reducing solutions

+ Iz + HzO Eo = -0.366V

(8)

is thermodynamically favored (AGO = f16.88 kcal) over reduction to hydrogen sulfide. Thus either or both products could result from the reaction. The free energy of reaction for reduction to sulfurous acid also decreases with temperature increase, becoming 1-11.43 kcal at 127 "C. As will be shown later, the gas-phase reduction of sulfur dioxide to hydrogen sulfide with hydrogen iodide: SOz

+ 6HI = HzS + 312 + 2H20

(9)

is of interest. The free energy of this reaction is -33.35 kcal at 25 "C and -30.40 kcal at 127 "C. The use of other reducing agents for sulfate is thermodynamically possible. Chromous ion, with a reduction potential of -0.408 V (12), should also reduce sulfate to hydrogen sulfide. Stone and Forstner found that such reduction occurred but at an extremely low rate except in the presence of iodide (14). Combination of Equations 5 and 6 produces an equation for reduction of sulfate by hypophosphorous acid: Sod2-

+ 4HaP02 + 2H+ = HzS + H8POi Eo

=

+0.80 V (10)

The standard potential of this reaction indicates a very favorable equilibrium constant but hypophosphorous acid alone does not prove to be a satisfactory reductant. Moreover, it decomposes at high temperatures. Hydriodic acid containing a small amount of hypophosphorous acid or sodium hypophosphite has been found to reduce sulfate rapidly. Gustaffson (5) found the presence of water detrimental to the reduction but an acetic acid medium proved satisfactory. Because of the desirability of using aqueous samples in most work, use of acetic anhydride was considered. Reaction with the anhydride would effectively remove the water and produce a solution of composition similar to that successfully used by Gustaffson. Hydrogen Sulfide Detection. Stephens and Lindstrom used a solution of ferric ammonium sulfate and 1,lO-phenanthroline to determine sulfur dioxide in air and found that (13) F. D. Rossini et al., Nut. Bur. Stand. (US.)Circ. C500, Washington, D.C., 1952. (14) H. W. Stone and J. L. Forstner, J . Amer. Chem. Soc., 79,1840 (1957).

the same solution also responded to hydrogen sulfide (7). Ferric ion is reduced to ferrous ion by hydrogen sulfide or other reducing agents. The resulting ferrous ion forms a stable complex with 1,lO-phenanthroline (15, 16). K = 2 X loz1 (11) Fez+ 3(phen) = Fe(phen)2f

+

The high stability of the tris(1 ,lo-phenanthroline)iron(II) complex has the effect of raising the reduction potential of ferric ion to a more positive value, facilitating its reaction with reducing reagents, Spectrophotometric measurement of the highly colored complex, E = 11,100, affords a very sensitive method for determination of the amount of hydrogen sulfide produced in the reduction step and thus the amount of sulfate. EXPERIMENTAL

Apparatus. Absorbance measurements were made at 510 nm os. water. A Perkin-Elmer Model 4000-A spectrophotometer was used with 10.00-cm cylindrical silica cells for the 0-10 microgram calibration curve and a Beckman Model B spectrophotometer was used with 1.000-cm square silica cells for all other measurements. A Leeds & Northrup Model 7401-A1 pH Indicator and electrodes was used for pH measurements. Digestion units, Figures 1 and 2, were fabricated from borosilicate glass. Side arms were made from threaded culture (15) H. Diehl and G. F. Smith, "The Iron Reagents," The G. F.

Smith Chemical Co., Columbus, Ohio, 1960. (16) L. G. SillCn and A. E. Martell, "Stability Constants of MetalIon Complexes," The Chemical Society, London, 1964. ANALYTICAL CHEMISTRY, VOL. 44, NO. 3, MARCH 1972

525

0.I A

0

2

4 6 0.1 M CH,CICOOH, rnl.

8

IO

Figure 3. Effect of chloroacetic acid in hydrogen sulfide detector

N2 I

I

b

a

Figure 2. Sample digestion train First digestor b. Second digestor a.

c.

Gas washing trap

d. Hydrogen sulfide detector

tubes (Kimble 45066-A, 16 X 100 mm) and were closed by Teflon-lined screw caps. Gas inlet tubes extended nearly to the bottom of each unit. Two units, as in Figure 1, were used for preparation and storage of reducing solutions and had capacities of 75 and 150 ml. The units in Figure 2 were used in the tandem arrangement shown for sample digestions. Detector solutions were placed in unit d, Figure 2, which consisted of a gas dispersion tube (Corning 39533, coarse porosity) fitted through a 2-hole No. 4 rubber stopper into a 25 X 200mm test tube. Units were interconnected as shown with Tygon tubing for glass-to-glass joints. A stream of nitrogen was passed through all units during use and the outlets were vented outside the laboratory. Condenser portions were cooled with tap water during digestions. The bottoms of digestion units were passed through holes drilled in asbestos boards to permit heating by Bunsen burners while shielding the remainder of the units from direct heat. Dectector units, Figure 2 4 were immersed in a constant temperature water bath fabricated from an expanded polystyrene chest of 12-liter capacity. Holes were cut in the lid to accommodate the 25-mm test tubes, a thermometer, and a 75-watt aquarium heater (Aquastock Model 4203). Temperature stability proved better than b0.2 OC per week. Glass-tip syringes were fitted with Chaney-type adaptors. A 1-ml syringe with a 22-gauge, 1.5-inch needle was adjusted to deliver 1.000-ml samples. Glass tubing delivery tips bent to reach the bottoms of digestion units were attached to two 5 ,

526


ml syringes with short lengths of poly(viny1 chloride) tubing. One syringe was adjusted to deliver 1.0 ml and the other to deliver 2.0 ml of reducing solutions. Reagents. ACS Reagent Grade chemicals were used except as noted. The hydriodic acid was Baker & Adamson Reagent for Methoxyl Determinations, assay: 57 %. This product does not contain the usual preservative, hypophosphorous acid. J. T. Baker hypophosphorous acid, purified 50% solution, was used in preliminary experiments. Sodium hypophosphite was B & A Reagent monohydrate, 99% pure, sulfate assay: 0.010% maximum. Ammonium bifluoride was B&A technical grade. Chloroacetic acid was Baker Analyzed Reagent. Welding Gas Products Co., Greenville, S.C., Extra High Purity nitrogen was used without further purification as a carrier gas. Known sample solutions were prepared by dilution of standardized sulfuric acid solutions with distilled deionized water. The sources of the sulfur-containing surfactants used in the interference study are as follows: Texize Linear Alkyl Benzene Sulfonate, 97 % active, Texize Chemical Co., Greenville, S.C.; Santomerse ME-B Slurry, Lot No. 9109, solids 62.10%, active 59.76 %, Monsanto Co., Inorganic Chemicals Division, St. Louis, Mo.; Aerosol OT, loo%, American Cyanamide, and Sodium Lauryl Sulfate U.S.P., Lot NO. 772441. Both of the latter were supplied by Fisher Scientific CO. U.S.P. XVI allows a total of 8 % of sodium sulfate and sodium chloride in U.S.P. Sodium Lauryl Sulfate. Special Solutions. Several special solutions were prepared according to the following directions:

A. 0.012M 1,lO-phenanthroline: Dissolve 0.595 gram of ClzHaNz.HzOin 150 ml of hot water, cool, and dilute to 250 ml. B. 0.1M ferric ammonium sulfate-0.1M chloroacetic acid: Dissolve 12.057 grams of FeNH4(S0J2.12H20 and 2.362 grams of CHzCICOOH in 150 ml of hot water, cool, and dilute to 250 ml. C. 0.1M sodium chloroacetate: Dissolve 4.725 grams of CH2C1COOHin 250 ml of water, add several drops of 0.5 phenolphthalein indicator, neutralize with 6 M NaOH until faintly pink, and dilute to 500 ml. (The reagent grade salt was not commercially available.) D. Buffer Solution: 0.015M chloroacetic acid-0.015M sodium chloroacetate: Mix 150 ml of Solution C and 150 ml of 0.1Mchloroacetic acid and dilute to 1 liter. E. Reducing Solution I : Mix, in order, 10 ml of 57% hydriodic acid, 10 ml of acetic anhydride (with cooling), and 5.0 grams of sodium hypophosphite in the digestion unit in Figure 1. Boil for 30 minutes in a stream of nitrogen and allow to cool without interruption of nitrogen flow. Store the solution in the digestion unit under nitrogen. F. Reducing Solution 11: Mix, in order, 20 ml of 57% hydriodic acid, 20 ml of acetic anhydride (with cooling), and 3.25 grams of sodium hypophosphite in the digestion unit in Figure 1. Boil and store as for Reducing Solution I.

A

0.3

-

0.2

-

0.1

-

, A

2.4

2.5

2.6

2.7

2.8

0. O'/I

2.9

-o.o

PH

Figure 4. Effect of pH of hydrogen sulfide detector

G . SuUide Detector Solution: Pipet into the detector unit, Figure 2d, 5 ml of 0.012M 1,lO-phenanthroline (Solution A), 1 ml of 0.1M ferric ammonium sulfate-0.1M chloroacetic acid (Solution B), and 20 ml of buffer solution (Solution D). Outline of Proposed Procedure. An aqueous sulfate sample was to be refluxed with a mixture of 57% hydriodic acid, acetic anhydride, and a small amount of hypophosphite in the digestion apparatus a shown in Figure 2. The hydrogen sulfide and any sulfur dioxide evolved would be carried by nitrogen through a second similarly charged digestor b and finally into collector d containing a mixture of ferric ion, 1,lo-phenanthroline, and buffering agents to form the bright orange ferrous-phenanthroline complex. The remaining ferric ion would be bleached with fluoride ion. Variations were to be made in the outlined procedure in the direction of obtaining the maximum amount of color for a given amount of sulfate consistent with high reproducibility and low interference. Each variable was to be optimized by testing the variable, while holding all others constant, to find the value of the variable which would produce maximum absorbance in the level region of its own curve of variable us. absorbance of ferrous phenanthroline. Optimum values are used in the Proposed Procedure in the next section. Analysis of Reagents for Sulfate. Using an early version of the Proposed Procedure, a digestion mixture of 5 ml of 57 % hydriodic acid, 5 ml of acetic anhydride, and 1 ml of 5 0 z hypophosphorous acid was prepared and digested to remove sulfate present in the reagents. The effluent gas when passed through a detector solution indicated that the reagents contained a total of 186 pg of sulfate. When milliliter portions of water were added, the blank was then 0.017 absorbance unit in a 1-cm cell. One milliliter of each of the reagents was tested in turn. Each milliliter of hydriodic acid yielded 1.7 pg of sulfate; acetic anhydride, 0.2 pg; and hypophosphorous acid, 179 pg. The chief source of sulfate in the reagents was obviously the hypophosphorous acid. Sodium hypophosphite was adopted as a source of hypophosphite because of its lower sulfate level. Effects of Variables in Sulfide Detector. The effects of varying the composition of the hydrogen sulfide detector solution were determined by following the outlined procedure, using conditions approximating those in the Proposed Procedure, and digesting samples containing 38.6 pg of sulfate. The effects of changing the concentration of chloroacetic acid, Figure 3, the pH of the solution, Figure 4, the concentration of ferric ion, Figure 5 , and the concentration of 1,lOphenanthroline, Figure 6, were determined. It was necessary to prepare the ferric ammonium sulfate in an acidic solution so some of the chloroacetic acid was incorporated in this solution. Figure 4 indicated that a rather low pH was desirable and chloroacetic acid (pK, 2.87) and sodium chloroacetate was chosen as a buffer. The amount of buffer solution added, together with that contained in the ferric ammonium sulfate solution, as specified in the Proposed Procedure, is sufficient to produce a pH in the optimum range,

I

2

3

4

5

0.1 M Ferric Ion, ml.

Figure 5. Effect of ferric ion in hydrogen sulfide detector

041

A

I

I

I

I

I

I

I

I

1

0.2.

t

O'I 0.01 0

4 6 8 0.012 M I,lO-Phenanthrollne, ml, 2

I IO

Figure 6. Effect of 1,lO-phenanthroline in hydrogen sulfide detector

and at the same time to produce a chloroacetic acid concentration within the optimum range. The concentration of ferric ion, Figure 5 , was not at all critical. The amount chosen as optimum is near the minimum amount required to reduce the blank due to ferrous ion impurity. The concentration of 1,lo-phenanthroline, Figure 6, for optimum response was not critical either. Approximately the minimum was used for economy of reagents. Concentrations of solutions were adjusted so that convenient quantities could be pipetted. The temperature coefficient of the detector s o h tion specified in the Proposed Procedure was determined over the range of 10 to 60 "C. It was linear and amounted to +1.0 % per degree, Although higher temperatures produced higher sensitivity, the blank also increased with temperature. A temperature somewhat above room temperature was chosen to simplify water bath requirements. No cooling supply was needed at this temperature (30 "C) and only a small heater was required. Effects of Variables in Reducing Solutions. The efl'ects of varying the composition of the solution in each digestor were studied. Preliminary experiments indicated that a single digestion unit presented problems in reproducibility. Incomplete reduction of sulfate to hydrogen sulfide was SUSpected. It was determined, by direct injection of the gases, that the gas washing trap, Figure 2c, did not affect hydrogen sulfide but completely absorbed sulfur dioxide when filled with 0.05M potassium biphthalate. Digestion of samples, using a trap filled with 0.05M potassium biphthalate and a single digestion unit, indicated that the optimum ratio of 57 hydriodic acid to acetic anhydride was 1:l (v/v). Variation of the concentration of sodium hypophosphite established an optimum concentration of this reagent of about 0.13 g/ml, Figure 7. Data for 50% hypophosphorous acid, also shown for comparison, indicated that it was not as effective, probably because of the larger amount of water introduced by using the

z


527

Table I. Interferences PB sod’-

0.01 0

I

0.04

I

I

I

0.08 0.12 0.16 Hypophorphite, g./ml.

I

0.20

I

0.24

clodc103-

Figure 7. Effect of hypophosphite in first digestor Sodium hypophosphite monohydrate 0 50 hypophosphorous acid 0

Iona None c1Na+ K+

z

Io 3-

Pbz+ Sn2+ HzOz H20z NH4+

caodZCNMn04Cr042NOaMOOS Pod3AsOzBa2+ FSCNNOa-

0.3A

-

0.2

0.1

-

sos2S 2-

Source

Taken

Found

...

Re1 error,

z

0.0

48.3 48.3 48.4 HCl 48.3 48.2 NaCl 48.3 48.3 KC1 48.3 48.3 HClOa 48.3 KC103 48.3 48.1 48.3 KIO, 48.3 48.3 Pb(0Ac)z 48.3 SnClz.2Hz0 48.3 48.3 48.5 HzOz 48.3 41.1 HzOz(30,ooO ppm) 48.3 NH~OAC 38.6 38.8 38.4 (NHa)zC204.HzO 38.6 38.4 KCN 38.6 KMn04 38.6 38.2 38.1 KzCr04 38.6 31.8 NaN03 38.6 (NH~)~MO?O~~ 38.6 . ~ H ~ O39.5 39.5 NaHZPO4 38.6 31.5 NaAsOz 38.6 37.0 BaClz.2Hz0 38.6 36.8 NaF 38.6 43.0 NH4SCN(50 ppm) 38.6 30.3 KNOz 38.6 60.1 Na2SO3(25 pprn) 38.6 102.9 (NH&S (23.5 ppm) 38.6

-4.1 -4.1 $11.4 -21.5 +55.1 $166.6

50.0

+29.5

+0.2 -0.2 0.0 0.0 -0.4 0.0 0.0 0.0 +0.4 -1.2

+O.S -0.5 -0.5

-1.0 -1.3 -2.1 $2.3 +2.3

-2.8

Surfactants6 0.0 0

0.02

0.04 0.06 0.08 Hypophoaphlle, p h i .

0.10

0.12

Figure 8. Effect of hypophosphite in second digestor 0

Sodium hypophosphite monohydrate hypophosphorous acid

0 50

z

R-sOs-

1

38.6

44.1 +14.2 1 (100 ppm) 38.6 43.8 I1 38.6 $13.5 0.0 38.6 11 (100 ppm) 38.6 41.8 $8.3 R’-SO8I11 38.6 39.0 I11 (100 ppm) 38.6 +1 .o 72.3 +87.3 R”-oSOaI V (100 ppm) 38.6 loo0 pprn unless noted 6 Surfactant kev I Texize, liiear alkyl benzene sulfonate, sodium salt. I1 Monsanto, Santomerse ME-B Slurry, linear alkyl aryl sodium sulfonate. 111 American Cyanamide, Aerosol OT, di(2-ethylhexyl) ester of sodium sulfosuccinate. I V Sodium lauryl sulfate. 0

solution. Comparison with Figure 8 indicates that incomplete reduction was indeed the case, only about 50% of the sulfate being reduced to hydrogen sulfide in a single step. The second digestor was inserted into the stream as shown in Figure 2 and its effect as a function of the composition of the reducing solution was studied, using the same amount of sulfate and the optimum reducing solution in the first digestor. The same ratio of 57% hydriodic acid to acetic anhydride as was used in the first digestor was again found best but results for hypophosphite from sodium hypophosphite and 50 % hypophosphorous acid were somewhat different, Figure 8. A lower concentration of hypophosphite was required for optimum results and the better source was again the sodium salt. The solutions specified in the Proposed Procedure result in approximately optimum compositions in each digestor, compared with the above results, but are the result of slight modifications to minimize the effects of variations in nitrogen flow rate and digestion time. The solutions used in the Proposed Procedure are also designed to take advantage of the economies of a countercurrent flow of reagents, in which part of the solution used in the second digestor is again used in the first digestor for the succeeding sample. Samples containing less than 40 Mg of sulfate produced essentially identical results when digested with the specified solutions for times of 10-20 minutes and at nitrogen flow rates of 39-97 ml per minute. Samples containing much more sulfate gave slightly low results at nitrogen flow rates in excess of 50 ml per minute and digestion times shorter than 12 minutes. Interferences. Solutions were prepared to contain 2000 parts per million by weight of most ions tested. For each test, 0.5 ml of the solution being tested and 0.5 ml of a known sulfate solution were used as a sample. The effective con528

ANALYTICAL CHEMISTRY, VOL. 44,

NO. 3, MARCH 1972

centration of the ion being tested was then 1000 ppm in the sample, with a few indicated exceptions, Table I. The higher concentration of hydrogen peroxide corresponds to a sample which is 3 % in hydrogen peroxide. The error in the case of sulfite amounts to 72% of the calculated amount for quantitative interference and that for sulfide amounts to 103% of the calculated amount. RESULTS AND DISCUSSION

Proposed Procedure. Assemble the digestion train as shown in Figure 2. Place a flow meter and a check valve (Pioneer Plastics P575-0332) in the nitrogen line, close the stopcock, and add a few small silicon carbide boiling stones to each digestor. Place about. 15 ml of 0.05M potassium biphthalate and 3 or 4 drops of 0.05 Thymol Blue indicator in the trap, Figure 2c. Place the detector unit, Figure 2d, containing Sulfide Detector Solution G in a constant temperature water bath at 30 i 0.3 “C. Add 10 ml of Reducing Solution I1 to the second digestor. Add 2 ml of Reducing Solution 11, 1 ml of Reducing Solution I, and 1 ml of distilled deionized water to the first digestor. Allow nitrogen to flow at a rate of 45 ml per minute and boil both digestors for 15 minutes. Continue the nitrogen flow for an additional 5 minutes.

Figure 9. Calibration curve, 1-cm absorption cells

Sulfatr, pg.

I

0

.

0

0

I

I

I

I

I

I

I

I

I

1

2

6

4

8

IO

Sulfatr, pg.

Remove the gas dispersion tube from the detector and rinse adhering solution into the detector with 3 or 4 small portions of distilled water. Add 2 ml of 2M ammonium bifluoride to the detector solution, dilute to 50 ml, and measure the absorbance at 510 nm US. water. For subsequent samples, the following procedure should be followed. Drain the previous solution from the first digestor. This may be easily accomplished by blocking the digestion train immediately after the trap for a short period of time to allow the nitrogen flow to build up a slight pressure in the apparatus, then opening the stopcock and allowing the solution to drain through the gas inlet tube. An alternate procedure is withdrawal through the side-arm with a syringe fitted with a curved glass tip, in which case the stopcock is unnecessary. Then transfer 2 ml of Reducing Solution I1 from the second digestor to the first and replenish the second digestor with 2 ml of fresh Reducing Solution 11. Add 1 ml of Reducing Solution I and 1 ml of sample solution to the first digestor, digest the sample, and process the detector as outlined above. Samples should consist of distilled deionized water until the blank is approximately 0.010 absorbance unit in a 1-cm cell. New or long-unused apparatus may require that several blank samples be run to remove contamination. The amount of sulfate is preferably read from a calibration curve, similar to Figure 9 or 10, prepared from samples of known sulfate

content. The gas washing trap should be emptied and refilled when the color changes to red-orange. Accuracy and Precision. A Ringbom plot, Figure 11, was made to evaluate the range and accuracy of this method (17). The ideal range, indicated by the straight portion of thecurve, is about 25 to 55 pg of sulfate when 1-cm absorption cells are used. The slope of this portion indicates a relative error of about 2.8% per 1% photometric error. Over a range of 10 to 100 pg, relative error is 5.6% for the same photometric error. Absolute error is dependent upon instrumental error, of course. The calibration curve of Figure 10, made with a 10-cm absorption cell, is linear up to almost 10 pg of sulfate and demonstrates the utility of the method for trace levels. Caution was necessary in this range to avoid Contamination and a high blank. The results of the interference studies of Table I were used to determine the precision of the method in lieu of a separate series of samples. Standard deviation of the first 10 samples of this series is h0.105 pg of sulfate and the relative standard deviation is k 0 . 2 2 z . The relative standard deviation of a larger group, composed of the first 20 samples of Table I is

h1.23z. (17) A. Ringbom, Z.AnaI. Chem., 115, 332 (1939); see also G. H. Ayers, “Quantitative Chemical Analysis,” 2nd ed., Harper & Row, New York, 1968, pp 460 ff. ANALYTICAL CHEMISTRY, VOL. 44,NO. 3, MARCH 1972

529

Figure 11. Ringborn plot, 1-crn absorption cells

I

3

10

30

100

300

1000

Sulfate, pg.

Hydrogen Sulfide Detector. The slope of Figure 9 corresponds to the production of 4.70 ions of the tris(1,lO-phenanthroline)iron(II) complex per sulfate ion or a 4.70-electron oxidation of hydrogen sulfide by the detector solution. That of Figure 10 corresponds to a 4.94-electron oxidation of hydrogen sulfide. A barely discernible deviation from Beer's law in the form of a slightly steeper slope appears in Figure 9 at low levels of sulfate. Some other less sensitive detector compositions also displayed the same type of deviation. In each case, the initial slope corresponded to a 5-electron oxidation of hydrogen sulfide. A similar deviation was also observed during the development of the method for sulfur dioxide but was attributed to experimental error at the time (18). A possible explanation is more vigorous reaction between ferric ion and the reducing gas at very low levels. The nearly linear temperature coefficient of the detector indicates that more than one reaction is involved. The sensitivity of the sulfide detector was unchanged by relatively wide variations in chloroacetic acid, Figure 3, and ferric ion, Figure 5 . Chloroacetic acid was added to the ferric ion solution to prevent hydrolysis and the quantity and concentration of buffer solution were adjusted to obtain the optimum pH of 2.63 in the detector solution, Figure 4. The low pH would seem to favor loss of hydrogen sulfide from the detector solution. However, a second detector in series demonstrated that no measurable amount of hydrogen sulfide escaped. At this low pH, the tris(1 ,lo-phenanthroline)iron(II) complex does not form during digestion of the sample. Addition of ammonium bifluoride solution serves a twofold purpose. First, the color of unreacted ferric ion is removed by formation of the colorless hexafluoferrate ion and, second, the pH is increased to approximately 4, where the iron(I1) complex forms immediately. The ferric ammonium sulfate-chloroacetic acid solution and the mixed sulfide detector solutions were found to deteriorate after preparation. This was probably due to photochemical reduction (19). Freshly prepared ferric ion solutions gave an absorbance of 0.002 to 0.004 when treated with the specified amounts of 1,lo-phenanthroline, buffer, and ammonium bifluoride solutions. Absorbance gradually (18) B. G. Stephens, Wofford College, Spartanburg, S.C., personal communication, 19(i4. (19) J. Novak and H. Arend, Tahnta, 11, 898 (1964). 530


increased to about 0.009 eight days after preparation. Mixed sulfide detector solutions increased in absorbance much more rapidly. Blanks were not appreciably affected for periods of several hours, however. Weekly preparation of ferric ion solutions and twice daily preparation of mixed detector solutions proved satisfactory. Other solutions used in the detector solution were unaffected by standing. After treatment with ammonium bifluoride solution, no absorbance change occurred in the detector solutions within 15 minutes. Thereafter, absorbance slowly increased about 3 in 12 hours. Reducing Solutions. As is evident from Figures 7 and 8, the optimum levels of hypophosphite were not the same in the first and second digestors. The functions of the two digestors are actually quite different. The equilibrium is principally between dissolved species in the first digestor. In this hot, highly acidic solution, any sulfur dioxide produced would be expected to escape rapidly. Sulfur dioxide is reduced in the second digestor to hydrogen sulfide in a very favorable reaction (Equation 9) which can take place in the gaseous phase. The temperature in this digestion unit was about 115 "C (388 OK) and the free energy of the gaseous phase reaction near this temperature (400 OK) is -30.40 kcal. In addition to producing higher reduction efficiencies, the use of sodium hypophosphite in preference to hypophosphorous acid also resulted in lower blanks without excessive preliminary boiling of reducing solutions. This was due to the higher level of sulfate in commercial hypophosphorous acid, which is usually prepared as the barium salt and the barium is then precipitated with sulfuric acid. Prolonged boiling resulted in lowered activity of reducing solutions because of decomposition of hypophosphorous acid at high temperatures. The procedure of countercurrent transfer and replenishment of reducing solutions was developed for two reasons. Different solutions proved best in each digestor, which necessitated two different solutions. Boiling reduced the efficiency of digestion solutions because of decomposition of hypophosphorous acid. In order to maintain a constant efficiency in the second digestor, replenishment was necessary. No increase in time or effort was necessary to simply transfer the solution to the first digestor, which operated best with a solution more concentrated in hypophosphorous acid, than to add a small amount of another solution to the first digestor. This transfer of solution also significantly lowered reagent costs. A single reducing solution could be used for

both digestors with but little sacrifice in sensitivity or precision. Reproducibility was excellent using other reducing solutions and detector solutions which varied widely from optimum compositions. However, the use of the solutions specified resulted in highest sensitivity and least dependence on those variables least easily controlled in a manual process: quantities of reducing solutions, carrier gas flow rate, and digestion time. In an automated version of this method, where such variables could be more precisely controlled, comparable results could be expected with a single reducing solution for both digestors. Gas Washing Trap. The sulfide detector is sensitive to iodide as well as a variety of reducing gases. Although this sensitivity is the basis for a method for the analysis for iodide (20), it did necessitate insertion of a trap in the gas stream to prevent carry-over of hydrogen iodide to the detector. Deionized water, 0.1M hydrochloric acid, and 0.05M potassium biphthalate all absorbed hydrogen iodide satisfactorily but all became ineffective after a number of samples were run. The buffer solution proved to have the highest capacity and the end of its effectiveness could be readily detected by a decrease in pH. Addition of Thymol Blue indicator, which changes color from yellow to red-orange near pH 1.7, was a satisfactory method of determining when the solution approached exhaustion. Phosphine, one of the decomposition products of hypophosphorous acid (21), was not absorbed in the trap and did not react with the detector solution. It did serve, however, to reverse any slight air oxidation of iodide in the trap and prevent its reaction with hydrogen sulfide. Interferences. Few serious interferences were found in this procedure, Table I. Although it interfered severely in other similar methods (3,nitrate presented no serious difficulty. This was probably due to the second reduction step. Barium caused low results but in this solution the solubility product of barium sulfate was exceeded by more than ten thousandfold. Such a high level of barium could not occur in a sulfate sample solution, so its interference is more academic than actual. Low results in the presence of fluoride were probably due to volatilization of hydrogen fluoride and its subsequent reaction with ferric ion in the sulfide detector solution. The only serious non-sulfur-containing interference was nitrite. Rapid volatilization of nitrous acid from the boiling reducing solution would be expected. Unreduced nitrous acid would then oxidize hydrogen sulfide prior to its reaction with the sulfide detector. Sulfite interference was somewhat less than quantitative. As in the case of nitrite, much of the sulfite would be rapidly evolved as sulfur dioxide before complete reduction, and passage through the second digestor would be so rapid that reduction would still be incomplete. Quantitative results could be obtained by prior oxidation to sulfate or by direct evolution of sulfur dioxide from hot acidic solution into a detector previously calibrated with standard sulfite solutions. Interference by both sulfite and nitrite could be eliminated by acidification and boiling of the sample to volatilize sulfur dioxide and nitrous acid before reduction. Sulfide interfered quantitatively. Gustaffson (6) reported calibration of methylene blue detectors directly against sulfide solutions but several attempts at similar calibrations were not

(20) J. B. Davis, A. S. Wilson, and N.G. Golden, presented at the 42nd Annual Meeting of the South Carolina Academy of Science, Clemson University, Clemson, S.C., April 26, 1969. (21) E. S. Gould, “Inorganic Reactions and Structure,” Henry Holt and Co., New York, N.Y., 1955, p 251.

successful because of rapid decomposition of standard solutions. Precautions were taken to exclude oxygen and bacteria from the refrigerated diluted solutions, but low results were obtained unless samples were digested with reducing solution. Sulfide interference in sulfate determinations could be eliminated by heating acidified samples before reduction to volatilize hydrogen sulfide. To test the interference of sulfonated surfactants, samples of biodegradable linear alkyl aryl sulfonates were obtained from the Texize and Monsanto Companies. These samples were of industrial-grade surfactants used by compounders to build consumer detergent products familiar in the retail market. These compounds did not interfere except in amounts which could be attributed to the amount of sulfate contained in the surfactant as a result of sulfonation with an excess of oleum and neutralization with base. Apparently, the carbonsulfur bond in these sulfonates is resistant to the attack of the reagents used in this method. Aerosol OT, the di(2-ethylhexyl) ester of sodium sulfosuccinate, is used as a wetting agent in industry. It also does not interfere except to a level which could be attributed to impurity levels of sulfate. This compound is an alkyl sulfonate and the carbon-sulfur bond in this compound also seems resistant to attack in this method. Sodium lauryl sulfate, a sulfate ester and not a sulfonate, interferes severely. The sulfur is bonded to oxygen in this case and this ester linkage is obviously much weaker than the carbon-sulfur bond of a sulfonic acid. Fortunately for the analyst, this surfactant cannot compete with the linear alkyl aryl sulfonates economically as an ingredient in detergents for general household and industrial purposes. The sulfate ester surfactants are used in drug and cosmetic products. The reader is reminded that if tests of this method for sulfates is made on commercial detergent preparations, the results will be positive for sulfate for these products are often built with up to 40 sodium sulfate. Safety Precautions. Phosphorus acid decomposes when heated with evolution of phosphine, a poisonous gas with a garlic-like odor (21). No danger exists with cold solutions, but the digestion units should always be vented outside the laboratory during heating. Prepared reducing solutions present the usual hazards associated with strong, highly reactive acids. During preparation, reagents should be added in the order listed to prevent a vigorous undesired reaction of acetic anhydride with hypophosphite (22). Caution should be exercised in addition of acetic anhydride to concentrated hydriodic acid. The digestion units should be cooled in a beaker of ice water to prevent excessive temperature rise and the solution should be agitated by a slow stream of nitrogen during the slow addition of acetic anhydride. Distillation of hydriodic acid should only be carried out in apparatus that is being swept by nitrogen or a violent reaction may occur with the oxygen of the air to produce water and iodine. The addition of sizeable quantities of such oxidizing agents as nitrate and perchlorate to refluxing reducing solutions is not at all advisable. CONCLUSIONS

The recommended procedure should serve well for the determination of sulfate in samples from a variety of sources. The application to the determination of sulfate in ground and (22) A. Michaelis and M. Pitsch, Ann. Chern.,310,59 (1900). ANALYTICAL CHEMISTRY, VOL. 44, NO. 3, MARCH 1972

531

surface waters should be particularly attractive, especially since sulfonated surfactants do not interfere. Determination of sulfate in biological fluids should also be quite practical. The low level Of sulfate this method permits One to reach should offer an opportunity to study the role of this common ion in many biological processes. While the method should be able to accept most such Samples without any treatment, it might be well to precipitate the protein material and run the analyses on the filtrate to avoid fouling the digestion vessel with organic material. The possible interferences due

to oxidizing agents and the ions listed in Table I should be kept in mind when choosing a protein precipitating agent.

for review May 19, 1971. Accepted October 29, 1971, Presented in part at the Combined Southeast-Southwest Regional Meeting, A.C.S., Memphis, Term,, December 1965. Taken in part from the dissertation,@ 1965, submitted by Joe B. Davis to the Graduate School of Ciemson University in partial fulfillment of the requirements for the degree of doctor of philosophy, July 1965. Work supported in part by National Defense Education Act Fellowship.

Spectrophotometric Analysis of Phenols and of Sulfonates by Formation of an Azo Dye L. Ronald Whitlock,I Sidney Siggia, and Janice E. Smo1a2 Department of Chemistry, University of Massachusetts, Amherst, Mass. 01002 Phenols and sulfonic acids are analyzed by a method based on the coupling reaction between a diazotized amine and the phenol. The concentration of the azo dye formed was measured by visible spectrophotometry. Sulfonic acids and their salts were converted to phenols via alkali fusion at 360 O C using potassium hydroxide, prior to coupling. Diazotized sulfanilic acid and p-phenylazoaniline, a previously unreported reagent, were particularly useful reagents for quantitative phenol analysis. With the latter, the optimum pH and reaction time for the phenols were remarkably similar and the molar absorptivities of the dyes formed were significantly higher. The coupling reaction conditions for a wide variety of phenols and the c and Xmsx values for the azo dyes are presented. The limit of detection was about 1 pmole of phenol per liter of solution and the minimum amount detectable was 0.3 pg. This method for sulfonic acid analysis will extend the useful range of application of the alkali fusion method to include polysulfonates and higher molecular weight sulfonates. Alkali fusion of halogenated sulfonates resulted in substitution of both the halogen and sulfonate groups. To help determine which sulfonates can be analyzed by the fusion procedure, the thermal stabilities of a large number of sulfonates and several phenolates are given.

SULFONIC ACIDS have represented a difficult analysis because of their limited reactivity and lack of volatility. A method for their analysis has been described by Siggia, Whitlock, and Tao (1) based on alkali fusion giving quantitative conversion to sulfite and the corresponding phenol. By utilizing direct gas chromatographic analysis for the phenols formed, many sulfonic acids and salts including their mixture could be analyzed with speed, and good precision and accuracy. However, with the alkali fusion method some sulfonates give phenols of low volatility that cannot be measured by gas chromatography. These include the polysulfonated compounds, those sulfonates containing other very polar functional groups, and certain higher molecular weight sulfonates. To extend the alkali fusion method to include these sulfonates, l Present address, Research Laboratories, Eastman Kodak Company, Rochester, N.Y. 14650 Present address, W. R. Grace & Co., Cambridge, Mass. 02140

(1) S. Siggia, L. R. Whitlock, and J. C. Tao, ANAL.CHEM.,41, 1387 (1969).

532


a method for phenol measurement based on the phenolic group reactivity has been developed. It is commonly known that substitution of a hydroxyl group on an aromatic ring, as occurs during alkali fusion of sulfonates, facilitates rapid electrophilic substitution. The position of attack is either para or ortho to the hydroxyl group. This effect has resulted in several methods for phenol measurement (2-6). The method described here is based on the coupling reaction between the phenol and a diazonium salt, producing a highly colored azo dye. The resulting dye concentration is measured quantitatively by visible spectrophotometry. This method of phenol analysis has been known for many years (7-11). The work concerning the diazo coupling reaction reported here was carried out, in part, so that a complete and workable procedure for use with the alkali fusion method for sulfonate analysis could be developed in detail. This includes a determination of the proper coupling reaction conditions for a wide variety of phenols and the E and Amax values for the azo dyes formed. In addition, a critical survey of several diazo compounds, including some not previously used, was made to determine which give the best results for phenol measurement. The problems commonly found in the analytical utility of the coupling reaction are nearly all derived from the coupling reaction conditions that must be controlled. A more extensive use of this method for phenol measurement has been hindered for two reasons. First, there is the general lack of information defining the proper reaction conditions needed for the many possible phenol-diazo combinations. Second, (2) B. Smith, Acta Chern. Scand., 10, 1589 (1956). (3) L. Tykken, R. S. Treseder, and V. Zahn, IND.ENG.CHEM., ANAL.ED., 18, 103 (1946). (4) H. H. Willard and A. L. Wooten, ANAL.CHEM., 22, 423, 670 (1950). (5) C. U. Houghton and R. G. Petty, Analyst, 62, 117 (1937). (6) C. J. B. Smit, M. A. Joslyn, and A. Lukton, ANAL.CHEM.,27, 1159 (1955). (7) R. H. DeMeio, Science, 108, 391 (1948). (8) G. A. Lugg, ANAL. CHEM., 35, 899 (1963). (9) W. Lee and J. H. Tumbull, Talarzfa,3, 318 (1960). (10) T. Takeuchi, M. Furusawa, and Y.Takayama, Jap. Anal., 4, 568 (1955). (11) G. B. Crump, ANAL. CHEM., 36,2447 (1964).

Spectrophotometric microdetermination of sulfate

Recommend Documents