mal ratio and a 4.8:l theoretical Hg:Cl ratio) did not prevent the oxidation of some chloride, thereby yielding erroneously high COD results. These results, however, can be corrected by collecting, and accounting for, the liberated chlorine. Table I data would indicate the feasibility of determining dichromate reflux COD in salt water without chloride sequestration by mercuric sulfate. While this possibility is intriguing, the care and extra trouble required in the manipulations would render this variation unsuitable for routine COD determination. Table I1 illustrates the serious error to which the standard method is prone. COD was determined in samples of sea water obtained from Marineland of the Pacific, Palos Verdes Peninsula, Calif. Having established the repeatability of the recovery method on synthetic COD and chloride mixtures as well as on sea water, we used the method to determine COD on waste water from a tuna-canning operation. Table I11 shows some apparent and corrected COD values covering chloride concentrations between approxi-
mately 12,000 and 19,000 mg/l. Table III also shows that a 10:1 ratio of HgS04:C1 will not entirely prevent some chloride from being oxidized. The fraction of chloride oxidized does not appear to be dependent on chloride concentration, sample us. dichromate and acid volumes, nor the amount of mercuric salt added; rather, it appears to vary with the substrate. True values for oxygen consumed from dichromate can be determined in highly saline wastes and sea water by the use of the chlorine-recovery method outlined herein without the need for prior chloride determination or dependence upon a large excess of mercury which, as demonstrated, will not achieve complete sequestration of chloride.
ACKNOWLEDGMENT For inspiration and guidance in the conduct of this study, the author is indebted t o Richard D. Pomeroy. Received for review January 7, 1974. Accepted March 25, 1974.
Use of a Potentiometric Sulfur Dioxide Electrode for Lamp Sulfur Determinations John A. Krueger Orion Research lncorporated, I 1 Blackstone Street, Cambridge, Mass. 02139
Sulfur in petroleum is commonly determined by the Lamp Combustion Method ( I ) , in which the sample is burned under controlled conditions and the gaseous sulfur combustion products are collected and analyzed. In the past, the instability of sulfur dioxide solutions and the unavailability of interference-free methods for sulfur dioxide measurement have made it more convenient to use a hydrogen peroxide absorbing solution, and to measure the sulfate formed by oxidation of sulfur oxides by acidimetric titration or gravimetric determination of precipitated barium sulfate. The recent introduction of a gas-sensing electrode which is selective for sulfur dioxide makes possible the measurement of sulfur combustion products directly as sulfur dioxide in the absorbent. The theory of gas sensing electrodes has been discussed in another publication (2). Briefly, the sulfur dioxide electrode is a gas detecting device which senses the level of dissolved sulfur dioxide in aqueous solutions. A gas permeable membrane separates the sample solution from an internal filling solution in contact with a pH-sensing electrode. As the partial pressure on each side of the membrane equilibrates, the pH of the filling solution varies. The internal filling solution contains a high level of bisulfite, so that the variation in pH is reversible and reproducible. The electrode has a slope of 58 mV per decade over a concentration range of 10-6 to 2 x 10-2M. (1)
"Sulfur in Petroleum Products by Lamp Combustion,'' in "Standard Methods of Chemical Analysis," F. J. Welcher. Ed , (Abstracted from former ASTM Methods D90-55T) p p 2031-33, Vol. 2. part B.
1963 ( 2 ) J W . Ross, J r . , J H Riseman, and J. A . Krueger. "Potentiometric Gas Sensing Electrodes," J Pure Appi. Chem., 36, 473-387 (1973) (paper presented at the IUPAC Conference in Wales. 1973).
1338
This paper describes a method of collection and analysis of oxidized sulfur products which is considerably faster than earlier methods. The sulfur is measured directly as sulfur dioxide in the absorbing solution-particulate and ionic by-products of the lamp combustion do not interfere.
EXPERIMENTAL Apparatus. A VWR Scientific Lamp Sulfur Apparatus, Cat. No. 51934-000, was used for sample combustion, with a single absorber or, in later experiments. two absorbers in series. Air was drawn through the system a t about 300 ml/minute using a vacuum pump; the air flow was controlled by a 5-turn stainless steel valve. Sulfur dioxide measurements were made using a n Orion Model 95-64 sulfur dioxide electrode and Orion Model 801 digital pH/mV meter, reading to 0.1 mV. A liquid membrane chloride electrode, Orion Model 92-17, p H electrode, Orion Model 90-01, and double junction reference electrode, Orion Model 91-02, with 10% K N 0 3 in the outer chamber, were also used. Reagents. Sulfur dioxide standards were made by appropriate dilution of Orion sulfur dioxide molarity standard, Cat. No. 9564-06. Osmolality and p H of samples were adjusted using Orion "SO2 buffer," made by adding 190 grams of anhydrous N a ~ S 0 4to about 800 ml of water in a 1-liter volumetric flask, mixing thoroughly, then adding 53 ml of concentrated (96-9770) and diluting to the mark with distilled water. The p H of the buffer was about p H 1.2 after 1 + 10 dilution. All other chemicals were Fisher Certified Reagents. The sodium hydroxide used was tested and found to be free of sulfur dioxide. Procedure. Electrode Operation and Storage. Correct operation of the electrode was confirmed by preparing a calibration curve (3).Between measurements, the electrode was kept in the recomniended storage solution to reduce changes in osmolality inside the electrode. The storage solution was prepared by adding about (3) Sulfur Dioxide Instruction Manual, Orion Research Incorporated, 1973
ANALYTICAL CHEMISTRY, VOL. 46, NO. 9, AUGUST 1974
T a b l e I. D a t a O b t a i n e d with SO2 a n d C1- Electrodes, to Approximate KTCX IHgC141
2
x
10-2
[Cl-] before adding SO?
PH
2.0
1.6
x
[SO21 added
10-2
RESULTS AND DISCUSSION Choice of Absorption Solution. The absorption solution must be capable of absorbing all the sulfur dioxide produced during the combustion step and of storing it without loss until measurement. When 1M sodium hydroxide was used as an absorbing solution, a subsequent titration with standard HCl indicated that the solution had absorbed 0.5M of carbon dioxide. During the analysis, the solution rapidly became more acidic, and less absorbant for S O z . Acidification of this solution to release surfur dioxide was further complicated by effervescence of carbon dioxide. The use of tetrachloromercurate (TCM) as an absorbing solution for sulfur dioxide has been studied thoroughly by West and Gaeke (J), and EDTA has been reported as a stabilizer for SO2 ( - 5 ) .T o eliminate carbon dioxide absorption, a 5 x 10-2M tetrachloromercurate absorbing solution containing 0.06 gram NazHz EDTA per liter was used, a t a neutral pH. Freeing Absorbed SO2 for Measurement. Before the collection efficiency of this solution could be measured, a means was required for freeing combined SO2 for measurement. To investigate the strength of the disulfitomercurate complex, the formation constant for the disulfomercurate equilibrium
+ ?SO2 + L”,O
===
Hg(SO,)?’-
+ 4C1-
2
10-2
4.4 grams of NaZS03 to 100 ml of water in a 250-ml Erlenmeyer flask, adding 1.5 ml concentrated HCI, and filling with water to the neck. The electrode was stored in the flask. Sulfur Dioxide Meczsurement. Samples were burned in the apparatus typically fur 5-10 minutes. At the end of each run, the absorption solution (100 ml) was transferred to a 200-ml volumetric flask. and 0.5 gram of sulfamic acid added to destroy absorbed nitrogen oxides which would react with sulfur dioxide (20 ml of “SO2 buffer” were added to adjust p H below the first ionization constant for HzS03 (pK1 = 2.0) and bring the sample to the same osmolality as the electrode internal filling solution). The sample was diluted to 200 ml. Sulfur dioxide was measured by a known addition technique. The electrode was placed in 100 ml of the solution, and the initial potential recorded. A known volume of sulfur dioxide standard solution was added, and the initial concentration read from a published table relating the change in electrode potential to the initial concentration under the conditions of the experiment (3). The known addition method has three major advantages for sulfur dioxide determinations. No electrode standardization is required. The determination is quick and independent of earlier data, so that the effects of electrode drift are negligible. Finally, the technique gives accurate results even in the presence of complexing agents for the species being determined; it is necessary only that the same proportion of the species is complexed after the addition of standard solution is made. This facilitates the use of strong complexing agents for sulfur dioxide, such as tetrachloromercurate, as collection media.
(HgC1,)’-
Free SO?
i4H+
(1) was determined under the conditions found in the absorbing solution from data obtained with the chloride, sulfur dioxide and pH electrodes. Substitution of the data in Table I in Equation 1 shows that C1- is displaced as SO2 is added. ( 4 ) P West andG Gaeke A n d / Chem 28, 1816 (1956) (5) F P Scaringelli L Elfers D Norrts and S Hochhetser Anal Chem 42. 1818 (1970)
x
[Cl-I after adding SO?
1.6 X
10-3
lo-*
10
09
08
07
06
05
04
03
02
01
00
pn
2.0
1.9
1.8
17
1.6
1.6
1.4
13
12
11
1.0
0.9
OB
07
Figure 1. Ratio of free sulfur dioxide in TCM to free sulfur dioxide in a T C M - f r e e solution, based on 10-4M SO2 pH range 2.0 to 0 . 3 for
K =
and 10-2M TCM solutions
[Hg(SO,),‘-][Cl-]‘[H+]‘ [HgCl,’-][SO,]’
(-log K ) is approximated by measuring the chloride, sulfur dioxide, and hydrogen ion concentrations. Since the chloride electrode showed a slight transient response to free Hg2+, an excess of chloride was added. As the total tetrachloromercurate and disulfitomercurate can be determined from the data in Table I, -log K has been approximated a t 11. Therefore, a t low pH’s, the effect of varying pH on the sulfur dioxide complexation is much greater than the effect of varying TCM concentration. Figure 1 shows the ratio of free sulfur dioxide in TCM to free sulfur dioxide in a TCM-free solution, over a pH range of p H 2.0 to pH 0.3 for 10- 3 and 10- 2M TCM solutions. These data were obtained by measuring the change in potential when known levels of TCM were added to the sulfur dioxide solution, and using a known subtraction table to calculate the loss of free SOz. Figure 1 shows that the ratio of free SO2 in TCM to free SO2 without TCM is not greatly affected by changes in TCM concentration, and that a plateau exists where SO2 measurements can be made. When the pH is adjusted to pH 1.2, as by addition of “SO2 buffer,” about 20% of the SO2 is still complexed. The total SO2 concentration can be measured by known addition. Absorption Solution Efficiency. The efficiency of a range of absorbing solutions was tested by adding a known amount of sulfur dioxide to the solution and bubbling the products from a 30-minute combustion of sulfur-free petroleum ether through the solution. The results are shown in Table 11. Where two traps were used, 50 ml of 2 x 10-5M SO2 and TCM were in the first trap and 50 ml of TCM in the second trap. When TCM/EDTA solution was used, the pH fell from 7.2 to 4.8 in 20 minutes. As collection efficiency depends on pH, the solution was buffered by making it 0.025M in
ANALYTICAL CHEMISTRY, VOL. 46, NO. 9. AUGUST 1974
1339
Table 11. Ability of Different Collection Solutions t o Hold SO2 During a Simulated L a m p Combustion Recovery, Absorbing solution
4 X lo-* T C M 0.06 EDTA
4 x 10-1 T C M
Recovery,
R
Initial SO.
10 - 4
70
10 - 5 10 - 4
75 86
10 --j 10 - 4
93 82
10 --j 5 x 10-6
88 94
Absorbing solution
4 X TCM 0.025 NaH2P04 0 . 0 2 5 Na?HPO4 0.06 E D T A 1. 5% glycerol 4 X TCM 0.06 EDTA 0.025 NaH2POa 0 . 0 2 5 Na,HPOI 2 . 0 % glycerol 2 flasks
0.06 EDTA 4 X lo-* T C M 0.06 EDTA 0,025 NaH2P04 0.025 N a 2 H P 0 4
x
10-1 T C M 0 , 0 2 5 Na,HPO, 0.025 NaH,POi 0.00 E D T A 3 . 0 % glycerol 4 x 10-2 T C M 0.06 EDTA 0.025 N a H p P 0 4 0 , 0 2 5 Na2HPOl 2 . 0 % glycerol
4
4
10 --j
c /o
x lo-:
94
5
x lo-' T C M
0.025 NaH2POI 0,025 NapHP04 0.06 EDTA 2 flasks 4 X lo-' T C M 0.025 NaH2P04 0.025 Na2HP04 0.06 EDTA
93
10 - 4
Initial SO?
10 - 4
99
10 -5 5 x 10-6
99 93
6
x
10-6
96
96
Table 111. SO, 'Total SO, for Two Simulated Fuels Theoretical Simulated fuel
1 g dimethylsulfoxide 99 g methanol
6%
s
Amount of fuel burned
0.31
0.97 g
Exp i o S
27 3 m V
0 623 X 1 0 - -
0 413
o
61
0 405
98 8
o
441
0 161
96 9
0.165
98.0
on adding 10 ml of
0.41
1.00g
10 - 2 so, 22 1 m V
on adding 10 ml of
1 g toluene
Concentration of so M 1.
Known addition
1.75 g
0.168
sulfonyl chloride 99 g methanol
10-2 so, 2 3mV
x x
10-3
10-3
SO SO,
100
+
on adding 10 ml of
0.168
1.75 g
10 - 3 so2 2 8 . 0 mV
0 . 4 5 x 10-3
on adding 10 ml of 10 - 2
Table IV. Comparison of Electrode Method t o ASTM M e t h o d JP4 and JP5 Fuels
% S by
% S by SOi
ASTM D90-55T
Electrode, Analysis 1
0.34
A B
0.016 0.19
c
0.360 0,0180 0,201
Difference between S by SO2 Difference electrode Electrode, between two analyses, Analysis 2 methods, % %
0.367 0,175 0.200
5.8 11.0 5.0
1.9 2.3 0.5
Na2HP04 and NaH2P04. Using the phosphate buffer, the pH fell from 6.87 to 6.85 in a 20-minute burn. The addition of glycerol is believed to inhibit the chain reaction mechanism by which free sulfur dioxide is oxidized (6). While the addition of 2% glycerol improved the yield, increasing the proportion of glycerol had no incremental effect. A tenfold increase in the TCM concentration led to a 12% increase in collection efficiency. How(6) P. Haller, J. SOC.Chem. Ind., 38, 52-62 (1969). 1340
*
so,
ever, TCM levels above 5 x 10-2M were considered dangerous for everyday use. An optimum absorbing solution was chosen for further experimental work, which contained 10.86 grams HgC12, 5.9 grams KC1, 0.06 gram NazEDTA, 3.42 grams NaH2P04.H20, 3.52 grams Na2HP04, and 20 ml glycerol per liter. Comparison of Measured and Theoretical Sulfur in Simulated Fuels. Two simulated fuels containing known amounts of sulfur were burned using 50 ml of an optimum absorbing solution in each of two collection flasks. The composition of the fuels and theoretical and measured percentages of sulfur are shown in Table 111. These results indicate that sulfur combustion products were almost completely in the form of sulfur dioxide. At higher temperatures, smaller yields of sulfur trioxide are obtained:
+
O1
-
+
SO, 45.210 calories (3) A chamber or contact process is required to obtain high yields of SO3 from S02. Comparison of Electrode and Gravimetric Methods. Finally, three samples of JP4 and JP5 jet fuels were ana-
ANALYTICAL CHEMISTRY, VOL. 46, NO. 9, AUGUST 1974
?SO2
lyzed by the electrode method, using the optimum absorbing solution described above. The analysis was compared to results obtained using the barium sulfate gravimetric method. Analyses using the ASTM procedure were conducted by an independent laboratory. The analyses by the independent laboratory were obtained with the information that a t %S levels below 0.270, the range of accuracy for any result may exceed 10%. These results are shown in Table W . While these results show some scatter, the discrepancies between the two methods were within the range of accuracy obtained in gravimetric determination. A second analysis of each sample using the electrode methods showed less than 2.5% range of accuracy. This has been expressed as a % difference [(A - B/A) X 1001. The electrode methods can be used to determine 70s a t levels below 0.005%. Storage of Sample Solutions. Our work confirms the low rate of air oxidation of disulfitomercurate reported in the original work of West and Gaeke. Storage in plastic bottles is adequate for up to a month. For longer periods,
higher levels of TCM should be used, or the sample should be stored in Barex 210 or glass bottles with little or no air space above the solution. Further Applications of Electrode Analysis. This procedure could be adapted for measurement of SO2 in air, where the few interferences and method simplicity would make it even more useful. The air would be drawn through a midget impinger containing TCM-absorbing solution and metered. The sulfur dioxide in the solution would be measured by electrode. Similarly, recent work by Reiszner and West (7) has suggested collection methods for SO2 in the atmosphere using TCM-absorbing solutions. The use of a sulfur dioxide electrode method would allow analysis of SO2 in the ambient atmosphere with much shorter collection times. Received for review November 26, 1973. Accepted March 1, 1974. (7) L. Reiszner. and P. West, Environ. Sci. Techno!..7, 526 (1973)
I CORRESPONDENCE Determination of Organic Carbon in Water Sir: I would like to point out an error in the “Water Analysis” Review by Fishman and Erdmann ( I ) . The error occurs on p 386R a t the top of the second column. In describing the method of Croll [reference I O Q in (I)] for the determination of organic carbon in water, they give the relative standard deviations of measurements a t 0.5, 5.0, and 50 mg C/1. as f38, *7, and *1.7%, respectively. This is incorrect. In fact, the paper ( 2 ) described experiments where organic carbon was added to water already containing 0.15 mg/l. organic carbon and states that “levels of added organic carbon of 0, 0.5, 5.0, and 50 mg/l. were analyzed with relative standard deviations of *38, f 7 , f l . 5 , and *1.7%, respectively.” Thus, the absolute figures would be that levels of organic carbon of 0.15, 0.65, 5.15, and 50.15 mg/l. were determined with relative standard deviations of *38, f 7 , f l . 5 , and f1.7%, respectively. Since the publication of the article in question, these figures have been improved as shown in Table I. The limit of detection of the apparatus at the 95% confidence level has been found to be 0.02 mg C/1. This ap-
(1) M . J . Fishrnan and D. E. Erdrnan, Anal. Chem.. 45 (5). 361R (1973). (2) B. T. Croll, Chem. Ind., 386 (1972).
Table I. Analysis of Replicate Samples Using the Total Organic Carbon Analyzer Organic carbon, mul.
50.0 5.0 0.5
s,w / l . 1-0.7
Re1 std dev,
=0.08
-1.4 11.6
0.1
= O . 006
r4.3 &6
0.05 0.025
50,004 10.003 10.004
Blank
=to.025
70
19 ill
paratus is now being manufactured by Phase Separations Ltd., Deeside Industrial Estate, Queensferry, Flintshire, United Kingdom. As you will see, the figures quoted in the review indicate considerably poorer performance than the instrument actually gives. B. T. Croll Resources Group The Water Research Association Medmenham, Marlow, Buckinghamshire SL7 2HD, England Received for review, July 20, 1973. Accepted January 24, 1974.
A N A L Y T I C A L CHEMISlHY, VOL. 46. NO. 9, AUGUST 1974
1341
