ISDCSTRI.4L A S D ESGIXEERIXG CHEJfISTRY
January 15, 1930
55
Quantitative Determination of Small Amounts of Hydrogen Peroxide and of Ozone' Nelson Allen CENTRECOLLEGE, DANVILLE, KY.
HE literature shows that a large amount of work has been done on the perfection of very delicate methods for the detection of hydrogen peroxide and of ozone but that very few processes have been applied t o the quantitative determinations of these substances. The purpose of this research was t o try out these methods and determine if they are applicable to micro-quantitative work.
T
Determination of Hydrogen Peroxide
I n the case of hydrogen peroxide several procedures were found to give excellent results, and naturally these methods were based on color comparisons since the amounts of hydrogen peroxide involved were very minute. The peroxide used in the whole work was prepared from c. P. barium peroxide and was standardized with acid potassium permanganate. Stock solutions of the hydrogen peroxide were prepared from the standard by dilution with distilled water, and these solutions were re-prepared daily from the original solution which was re-standardized frequently. The method of DenigPs ( 2 ) using an acid solution of ammonium molybdate was found to give a perceptible yellow color, due t o the formation of permolybdic acid, with as small a quantity of hydrogen peroxide as 1 part in 200,000 parts of solution. The reaction of hydrogen peroxide with a solution of titanium sulfate is more sensitive than this, being accurate to 1 part in a million. The orange-yellow color of the pertitanic acid that is formed can be easily matched with standards containing known amounts of hydrogen peroxide. The comparison can be carried out either in Kessler tubes or in a colorimeter. Several investigators have suggested the oxidation of ferrous to ferric ion by hydrogen peroxide and the reaction of the ferric ion with thiocyanate ion t o produce the wine-red ferric thiocyanate as a quantitative method. Horst ( 6 ) describes such a process and Quartaroli (11) gives one that involves the catalytic action of hydrogen peroxide on the oxidation of ferrous ion by nitric acid. Keither method lends itself t o colorimetric determinations, but both are very sensitive qualitatively. The reaction of hydrogen peroxide with iodide ion has been used by many for determining the peroxide. The reaction is represented thus : 21-
+ Hz02 + + 2 0 H 12
and since a base is one of the products of the reaction a further reaction takes place: 312
+ 60H- +
1 0 3 -
+ 51-
f 3Hz0
The iodine that is liberated can be titrated with sodium thiosulfate or treated with starch solution and the blue color compared with standard solutions of potassium iodide treated with hydrogen peroxide. I n either case the solutions should be acidified first in order to reverse the second reaction above and liberate all of the available iodine. Since the reaction of hydrogen peroxide with acid potassium iodide solution proceeds rather slowly, many catalytic agents, notably ferrous salts and lead acetate, have been proposed, but better quantitative results were secured without the aid of such reagents. The minimum amount of the peroxide that could be detected 1 Received September 9, 1929 Abstracted from a thesis submitted in partial fulfilment of the requirements for the degree of master of science at the University of Chicago, August, 1929.
by color comparisons was about 1 part in 10 million. Iodide ion is very easily oxidized t o free iodine, so oxidizing agents other than hydrogen peroxide must be absent in order to secure accurate results. Keiser and N c l l a s t e r ( 7 ) and Engler and Wild ( 4 ) give reports on the use of various organic compounds as qualitative reagents for hydrogen peroxide, ozone, and the oxides of nitrogen. Csher and Rao (Is),in their work on the determination of ozone in the atmosphere, made use of the extremely delicate Griess-Ilosvay reaction with nitrites. They found that both ozone and hydrogen peroxide oxidize nitrite ion t o nitrate ion, the first in alkaline solution and the second in acid. The reactions are represented thus: 03 f NOzH20z NOa-
+
+Oz --f
+ XOj+ NO,-
HzO
On the basis of these equations they worked out an indirect colorimetric method for both substances. This procedure was found to give accurate results for hydrogen peroxide down to a concentration of 1 part in 10 million but, when applied to ozone, was more difficult to carry out and proved less sensitive than other methods for ozone. It was found that the best method for hydrogen peroxide was one involving the decolorization of a very dilute acid solution of potassium permanganate Such a solution containing a small amount of magnesium sulfate (9) was treated with minute quantities of hydrogen peroxide and compared with a blank of the permanganate. It was shown in this way that 1 part of the peroxide in 10 million parts of solution could be readily detected. For quantitative work the unknown is compared with standards containing known amounts of hydrogen peroxide. The magnesium sulfate acts as a catalyst for the reaction and enables the colors to be more readily and accurately compared. This method is rather a simple one in that only one solution is required and its concentration does not have to be known; it is merely adjusted by dilution t o suit the amount of hydrogen peroxide that is being determined. Determination of Ozone
The purpose of the work done on ozone was t o find a method for determining the minute amounts of the substance t h a t are produced when ultra-violet .light is thrown on a stream of pure oxygen. A large number of methods hare been proposed for the detection of ozone in the atmosphere and the majority of these make use of solutions of iodide ion. The reaction between ozone and iodide ion is usually given as 0 3 HzO 21- + Iz 2 0 H - 3. Oz A base is formed in the reaction, so a part of the iodine reacts with it to produce iodide ion and iodate ion as in the reactions with hydrogen peroxide. Lechner (8)points out that if a current of ozonized oxygen is passed through wash bottles containing neutral potassium iodide solution there is a loss of iodine due to volatilization, and he recommends the use of a n alkaline solution to avoid this loss. Mellor (10) quotes Lechner to the effect that ozone reacts with iodide ion in alkaline solution to produce iodate ion, according to the equation, I0 3 ----f 1 0 8 -
+
+
+
+
56
dSd L Y T I C AL EDI TIO-\-
If this equation were correct, one molecule of ozone would liberate three molecules of iodine, for on acidifying 103-
+
51-
+
6H’
--f
312
+
3H20
I n a neutral solution, however, one molecule of ozone liberates only one molecule of iodine as shown in the equations above. Since Lechner does not give the equation reported by hlellor the relative oxidizing poJyer of ozone in neutral and in basic solution was investigated. A Siemens ozonizer was set u p and arranged so that the stream of ozonized oxygen could be received in glass-stoppered flasks. The current was kept constant and the flow of oxygen t o the ozonizer maintained a t a constant rate. By allowing the ozonized oxygen to flow for 15 seconds into the flasks successively, it was found that practically equal amounts of ozone could be collected. To show that equal amounts were being secured, 25 cc. of alkaline potassium iodide Tvere added from a pipet to each flask, the flasks shaken thoroughly, and the solutions rinsed out and titrated with 0.0113 N sodium thiosulfate after acidifying with sulfuric acid. Almost exactly equal volumes of thiosulfate were required for each sample, showing that equal amounts of ozone were present. By treating one sample with 25 cc. of neutral potassium iodide and one with 25 cc. of alkaline, it was shown that the same amount of iodine was liberated in each case. Oxidizing Power of O z o n e i n N e u t r a l a n d Basic Solution VOLUMEOF THIOSULFATE TAKEN TRIAL Neutral K I Basic K I cc. CC. 1 3.75 4.00 2 3 75 3.78 3 3.47 3.73
T a b l e I-Relative
This proves that the ozone reacts with iodide ion in alkaline as in neutral solution and the equation reported by Rlellor is incorrect. One molecule of ozone liberates one inolecule of iodine in each case and the equivalent weight of ozone as a n oxidizing agent is 24. When a n acid solution of potassium iodide was used, it was found that more than one molecule of iodine was produced by 1 molecule of ozone. I n a series of experiments run as above with an acid solution of iodide ion, 1.23 molecules of iodine were liberated by 1 molecule of ozone. T a b l e 11-Relative
TRIAL
3,71
+
Oxidizing Power of Ozone i n Acid a n d Basic Solution VOLUME O F THIOSULFATE T.4KEN Acid K I Basic KI
cc
Average 3 . 7 1 3.00 = 1.25 approximately.
CC.
2.62 2.30 2.40 4.15 3.50 3.00
This result accords with the observations of Riesenfeld and Bencker (12) and with a n equation reported by Llellor, 405
+ lOHI + Ha02 + 51, + 4Hz0 + 302
For the use of potassium iodide in determining very small amounts of ozone a colorimetric niethod was used. The stream of ozonized oxygen from a photochemical cell was run slowly through Fisher wash bottles containing dilute alkaline potassium iodide. An aliquot part of the solution was taken, acidified, and starch solution added, then the colors were compared with those produced by standard solutions of iodine in potassium iodide. The amount of iodine liberated was estimated in this way and from this value the quantity of ozone was calculated. If the current of ozonized oxygen is run through too long or if the potassium iodide solution is too concentrated, some iodine is liberated by the oxygen and errors are introduced. However, this method gave the most satisfactory results of all those tried for minute amounts of ozone. About 0.000001 gram can be detected in this way. Benoist ( 1 ) recommends a slightly alkaline solution of
T’ol. 2, K O . 1
fluorescein as a n extremely dklicate reagent for ozone. The green fluorescein color is dissipated by the ozone and he worked out a quantitative ratio of 2 molecules of ozone to 1 of fluorescein in the reaction. Benoist’s work was repeated and his procedures were followed exactly, but his ratio could not be checked. Ozone mas collected in flasks. One sample was treated with alkaline potassium iodide, acidified, and titrated with sodium thiosulfate, while the other was titrated directly ryith a standard solution of fluorescein until the green color ceased to disappear. The assumption that the two samples contained the same amount of ozone was proved to be correct on treating several samples with alkaline potassium iodide and titrating later with thiosulfate. The ratio secured from a series of experiments was 11 molecules of ozone t o 1 of fluorescein. Some typical data are given in Table 111. T a b l e 111-Comparison
TRIAL
of S o d i u m T h i o s u l f a t e a n d Fluorescein as R e a g e n t s for Ozone
THIOSULFATE SOLUTION CC.
FLUORESCCIU SOLUTION CC
0.00395 liter X 0.0113 N X 24 (equivalent weight of ozone) = 0.00107 gram of ozone in each flask Fluorescein solution contained 0.00001822 gram per cc.: 0.00001822 gram X 36.52 cc. = 0.0006654 gram fluorescein T h e molecular ratio would be 0.0006654 : 0,00107 :: 332 : X Whence X equals 534, and 534 i 48 = approximately 11 molecules
Benoist makes no statements as t o the product of the reaction between ozone and fluorescein, but probably assumes the formation of a di-ozonide in the quinoid ring. Ozone may react with organic compounds in two ways (5)-it may act as a st’raight oxidizing agent or it may combine directly to form ozonides. In this reaction between ozone and fluorescein several such reactions probably take place due to the breaking down of the complex molecule and consequently no exact quantitative ratio can be established. For this reason and also because the two substances react with each other only on prolonged shaking, fluorescein is not recommended as a quantitative reagent. A solution of crystal violet or methyl violet is almost as delicate as fluorescein for the qualitative detection of ozone. An amount of ozone that will not give a blue color to starch-iodide solution will give a percept’ible diminution in the color of fluorescein or crystal violet solutions. A very recent method reported for ozone is that of Egorov (3)in which fluorescin, the non-fluorescent, leuco compound of fluorescein, is oxidized back to that substance by ozone. Egorov reports a rat’io of 1 part of fluorescein by weight produced by 0.96 part of ozone, which is a molecular ratio of about 1 to 7. This method cannot be very accurate as a quantitative process, since fluorescein itself is decolorized by ozone and the only reason that a green color appears is that the first reaction proceeds faster than the second. For the direct oxidation of fluorescin to fluorescein only 1 molecule of ozone should be required. Literature Cited (1) Benoist, Comfit. r e n d . , 168, 612 (1919). (2) DenigCs, Bull. sot. chim., [31 7, 4 (1892). (3) Egorov, Untersuch. Lebensm., 66, 355 (1928). (4) Englrr and Wild, Ber., 29, 1940 (1896). (5) Harries, “ D a s Ozon,” Springer, 1916. (6) Horst, Chem.-Ztg., 46, 572 (1921). (7) Keiser and McMaster, Am. Chem. J., 39, 96 (1908). (8) Lechner, 2. Elekfrochem., 17, 412 (1911). (9) Macri, Boll. chim.farm.. 66, 417 (1917). (10) Mellor, “Comprehensive Treatise on Inorganic and Theoretical Chemistry,” Vol. I, p. 905, Longmans, 1922. (11) Quartaroli, Gam. chim. ilal., [ll 48, 102 (1918). (12) Riesenfeld and Bencker, Z . anot‘g. Chem., 98, 167 (1916). (13) Usher and Rao, J. Chem. Soc.. 111, 789 (1917).
