June, 1919
T H E J O V R N A L OF I N D U S T R I A L A N D E N G I N E E Ri X G C H E M I S T R Y
for practically all of t h e material of t h e woods. I n t h e case of t h e hardwoods, t h e method for lignin is not applicable, due t o t h e entirely different and highly complex nature of t h e non-cellulose constituents. SUMMARY
I--Methods for t h e summative analysis of woods are described and analyses of five California woods by these methods given. 2--Sawdust is found t o be t h e most satisfactory mechanical condition of wood for analytical purposes. 3--Cellulose in wood is determined by a modificat i o n of t h e Cross and Bevan method involving chlorination. in vacuo. T h e cellulose residues are tested for t h e presence of lignin and corrections applied when necessary. 4-Lignin is determined by Konig’s method with 7 2 per cent sulfuric acid. The probable relation of lignin so obtained t o true lignin is discussed. ;-In t h e analysis of coniferous woods by these methods 96 t o 9 7 per cent of t h e wood constituents are accounted for. I n t h e case of t h e hardwoods examined, t h e lignin determinations fail and only 83 t o 9 1 per cent of t h e wood constituents are obtained. 6-Cutin is not a constituent of importance in wood tissue. REFISRENCES 1-Cross and Bevan, “Cellulose,” 2nd ed., p. 244. JOURNAL, 9 (1917), 557. 2-Schorger, THIS 3-Dean and Tower, J . A m . Chem. Soc., 29 (1907). 1125. 4-Sieber and Walter, Chem. A b s . , 8 (1914), 1202. 5-Johnsen and Hovey, Paper, 21 (1918), 40. 6-Schorger, THISJOURNAL, 9 (1917), 564. 7-Renker, J . Sac. Chem. I n d . , 28, 1269. 8-Konig and Rump, Z . N a h r . Genussm., 28 (1914), 177-222. 9-Johnsen and Hovey, P a p e r , 21 (1918), 42. 10-Cross and Bevan, “Cellulose,” 2nd ed., p. 95. 9 (1917), 564. 11-Schorger, THISJOURNAL, 12-Schorger, I b i d . , 9 (1917), 562-3. 13-Klason, Cross and Bevan’s “Researches on Cellulose,” 3, p. 105. 14-Cross and Bevan, “Cellulose,” 2nd ed., p. 102. 15-Cross and Bevan, I b i d . , p . 96. 16-Konig, Chem.-Ztg., 36, 1101. 17-Ost and Wilkening, I b i d . , 34, 461. 18-von Fellenberg, Chem. A b s . , 11 (1917), 2122. 19-Johnsen and Hovey, Paper, 21 (1918), 50. 20-Konig and Rump, Z . Nahv. Genussm., 28 (1914), 186. JOURNAL, 9 (1917), 560. 21-Schorger, THIS 22-Johnsen and Hovey, P a p e v , 21 (1918), 48. DIVISIONOF AGRICULTURAL CHEMISTRY UNIVERSITY O F CALIFORNIA AGRICULTURAL EXPERIMEXT STATION BERKELEY, CALIFORNIA
T H E DETERMINATION O F IODIDE IN MINERAL WATERS AND BRINES’ By W P. BAUGIIMANAND
w. w.
SKINXER
IKTRODUCTIOX
A large number of methods have been proposed for t h e determination of iodine i n t h e presence of bromide and chloride and this i n itself would seem t o indicate t h a t t h e problem has not been satisfactorily solved. A colorimetric method, depending on t h e liberation of iodine by nitrous acid and determining 1 Read before the Division of Water, Sewage, and Sanitation, 56th Meeting of the American Chemical Society, Cleveland, September 10 to 13, 1918.
563
t h e amount by t h e intensity of t h e color imparted t o a few cubic centimeters of carbon bisulfide, developed by Struve,l modified by Haywood,2 has been adopted as a tentative method for t h e determination of iodide in mineral waters by t h e Association of Official Agricultural chemist^.^ This method has been studied by t h e Association of Official Agricultural Chemists and t h e results reported by one of US.^ This report shows t h a t t h e method gives fairly satisfactory results when applied t o mineral waters and brines containing only one or two milligrams of iodine per liter, b u t t h a t i t is not satisfactory when larger quantities are t o be determined. The palladious iodide and thallous iodide methods are reported t o give satisfactory results, b u t t h e reagents required, palladious nitrate and thallous sulfate, are expensive and not always available. The indirect methods are not suitable where one halogen greatly predominates. Most of t h e other proposed methods are based on t h e behavior of solutions of t h e haloids towards oxidizing agents. Very few, if any, oxidizing agents have been overlooked in t h e search for a solution of t h e problem, b u t some have been only superficially investigated and their limitations not established. An endeavor has been made in this investigation t o determine t h e limitations and suitability of a few of t h e methods which seemed most promising and t o modify them as necessary so t h a t results obtained by their use might be accurate and reliable. CHEMICALS USED
CHLORIDE-The chemically pure reagent was dissolved i n water and reprecipitated with hydrochloric acid gas, removed from t h e liquor b y filtration, and finally heated in a porcelain dish until all of t h e hydrochloric acid had been driven off. It was further purified by several recrystallizations from distilled water. SODIUM BROMIDE-This was a stock reagent and was purified by recrystallization several times from distilled water. S T A N D A R D POTASSIUM IODIDE SOLUTIOX-Potassium iodide was synthesized according t o t h e method employed by Gooch and B r o ~ n i n g . ~Pure iodine was prepared by twice resubliming C. P . iodine from a small quantity of potassium iodide. Three-fourths of t h e iodine was added t o a n excess of electrolytic iron and covered with distilled water. When t h e iodine had disappeared, t h e solution was decanted from t h e excess of iron, t h e remainder of t h e iodine added, t h e solution filtered into a large volume of boiling water t o which a quantity of potassium bicarbonate, exactly equivalent t o t h e iodine, had been added and t h e precipitated iron oxide removed by filtration. It was then made up t o a convenient volume and standardized by precipitating and weighing t h e iodine as silver iodide from aliquot parts by volume. SODICM
Z . anal. Ckem., 8 (1869), 230. Bureau of Chemistry, Bulletin 91. s “LMethods of Analysis,” 1916, p. 47. 4 “Report on Water,” W. W. Skinner, Referee A . 0. A . C., Bureau of Chemistry, Bullelin 162. 5 A m . J . Scz., 39 (1890), 196. 1
2
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T H E J O U R N A L OF I N D U S T R I A L A N D ENGINEERING CHEMISTRY
P E R M A N G A N A T E OXIDATION O F I O D I D E T O I O D A T E
Iodide is oxidized t o iodate according t o the following equation: K I 4- z K M n 0 4 HzO = KIOs zKOH zMnOz Pean de Saint Gilles,l who originated the method, added a n excess of potassium permanganate t o the neutral or slightly alkaline solution of t h e haloids, heated a few minutes, added an excess of ferrous sulfate containing sulfuric acid, and titrated the excess of ferrous iron remaining in the solution, or in other words, he determined the amount of potassium permanganate required t o oxidize t h e iodine t o iodate. Barneby2 calls attention t o the important fact t h a t when ferrous iron is titrated with potassium permanganate in the presence of hydrochloric or hydrobromic acids, free halogen is liberated. He recommends, therefore, adding manganese sulfate and phosphoric acid t o prevent the liberation of chlorine or bromine. The results reported by him are very satisfactory, iodine being accurately determined in the presence of as much as 2 0 g. potassium bromide or I O g. sodium chloride. Klemp3 determined the amount of potassium permanganate necessary t o effect the oxidation by observing the point when t h e permanganate color appears in the solution after adding zinc chloride t o facilitate the precipitation of the manganese oxide. This end-point, however, is not very sharp. Reinige4 determined the excess of potassium permanganate by titrating with thiosulfate. Sonstadtj precipitated the iodic acid as barium iodate and estimated the iodine in the precipitate, while others recommend destroying t h e excess of permanganate with a1cohol16or hydrogen peroxide,’ removing the manganese oxide by filtration, and titrating the iodic acid in the filtrate iodimetrically. Grdger’s method appears t o be the simplest and least complicated and the results in Table I were obtained by the use of this principle which depends upon t h e relative affinity of t h e halogens for oxygen, which is different from their affinities for hydrogen. Thus, iodide is most readily oxidized t o iodate, a n d bromide much less readily t o bromate, while chloride occupies an intermediate position.8 This method is, therefore, especially suitable for determining small amounts of iodine in the presence of large amounts of bromine. I n the experiments, aqueous solutions of potassium iodide or potassium iodide mixed with the other halogen salts, were adjusted t o a volume of about I O O cc. in an Erlenmeyer flask, I cc. sodium hydroxide solution (4 g. per I O O cc.) added, heated t o boiling, excess of potassium permanganate added, and again heated until t h e precipitated manganese oxide coagulated. It was then allowed t o cool, sufficient alcohol added t o cause t h e permanganate color t o disappear, and after allowing the precipitate t o settle on the steam
+
+
+
Compt. rend., 46 (1858). 624. J. Am. Chem. Soc., 57 (1915). 1496. * 2.anal. Chem., 20 (1881), 248. 4 Ibid., 9 (1870), 39. 8 Cham. News, 26 (1872), 173. eGrBger, Z. aagcw. Cham., 7 (1894), 52; Rnudsen, Chem. E m . , 17 1 2
(1913). 119; 7 Auger, Bull. SOC. chim., 11 (1912). 615. 8 Roscoe and Schorlemmer, 1, p. 230; Ostwald, “Principlw of Inorganic Chemistry,” pp. 277-8.
Vol.
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No. 6
bath, it was filtered through a Gooch crucible with suction and washed with -hot water. The filtrate, t o which I or z g. of potassium iodide had been added, was acidified with hydrochloric acid and the iodin,e titrated with sodium thiosulfate. The reactions which take place are expressed by the following equations: KI03 SKI 6HC1 = 6KC1 3Hz0 31z 312 6Na2SzOs = 6 N a I gNa&Os Since only one-sixth of the iodine represents t h a t which was originally present in the sample, the number of cubic centimeters of thiosulfate solution required must be divided by six. The method is very rapid and the results very accurate. Unfortunately, however, since no separation of t h e iodine is effected, the solution is not left in a suitable condition for the determination of bromine.
+
+
+
+ +
+
METHOD TABLEI-PERXANOANATE )8
NaCl NaBr I a s K I Grams Grams Gram Taken Taken Taken 0.0201 0.0602 5 0.0402 5 0.0201 5 5 0.0201 10 0.0201 10 0.0602 5 5 0.0012 10 0.0012 10 0.0000 N / 2 0 NazSzOa, 1 cc. =: 0.006320 g. iodine. N/10 NazSzOa, 1 CC. =: 0.012676 g. iodine.
....
.. .. ....
I Found Gram 0.0200 0.0609 0.0400 0.0201 0.0202 0.0202 0.0606 0.0011 0.0011 0.0000
....
.. ..
Rrror Gram -0.0001 +0.0007
4.0002
0.0000
+o .OOOl
+O,OOOl +0.0004 -0.0001 -0.0001 0 * 0000
O X I D A T I O N O F I O D I D E T O I O D I N E B Y M E A N S O F POTASSIUM B I C H R O M A T E
Dechanl proposed a method for the separation of iodides from bromides and chlorides based on t h e liberation of iodine from neutral solutions of iodides by potassium bichromate according t o the equation: 5KzCrz0, 6 K I = 8KzCr04 CrzOs 312 T o t h e sample dissolved in I O O cc. water, he added 40 g. potassium bichromate and distilled with steam. He reports only three experiments and used only small quantities of salts t o establish t h e method, a n d t h e results are given below.
+
+
C1 Taken Gram 0.0123 0.0560 0.1940
Br Taken Gram 0.0126 0,0252 0.0504
I Taken
Gram 0.01443 0.02880 0.05760
+
I Found Gram 0.01441 0.02833 0.05628
It will be noted t h a t the results for iodine are low, the variation from theory being greater when t h e quantity is increased. Dechan attributed this t o t h e use of cork instead of glass stoppers in his apparatus, a n d subsequently described a glass-stoppered apparatus without mentioning whether or not the results were improved by it. Freidheim and Meyer2 repeated the experiments of Dechan, using larger quantities of iodide and found i t impossible t o set free all of the iodine. I n view of t h e fact t h a t the statement is made in at least one well-known textbook* t h a t potassium bichromate will liberate t h e iodine completely from J . Chem. SOC.,49 (1886), 682; 61 (1887), 690. 12. anofg. Chcm., 1 (1892), 407. 1
:Prescott and Johnson, “Qualitative Analysis,” 7th ed., pp. 347 and
367.
T H E J O U R N A L OF I N D U S T R I A L A N D E N G I N E E R I N G C H E M I S T R Y
J u n e , 1919
iodide, t h e experiments reported in the table were made. The distillation flask was made from a 2 50 cc. ground glass-stoppered wash bottle and t h e iodine distilled into a 2 5 0 cc. salt bottle containing I O O cc. of a potassium io&ide solution ( I O g. potassium iodide per I O O cc.), which was kept well cooled. A second bottle was connected with this in series, b u t no iodine went over into the second bottle except a trace in t h e first three experiments where aeration was employed instead of steam distillation. For steam distillation, the steam was generated from distilled water in a liter Florence flask. I n this apparatus, t h e iodine had no opportunity t o come in contact with cork or rubber stoppers. Distillation was continued until about I O O cc. oi water had been collected, whereupon i t was interrupted and the receiver replaced by another containing a fresh solution of potassium iodide. This was continued until no more iodine was evolved. An inspection of Table I1 shows t h a t t h e results are all low a n d are not consistent. They confirm Freidheim’s a n d Meyer’s statement t h a t iodine cannot be completely set free from iodide by potassium bichromate. TABLE 11-DISTILLATION
15
75
.. 0.0802
11/i
I1V 2 15
75
.. 0.0802 211/x 1‘/a 1
10
75
20
75
OF
IODINs
WITH POTASSIUM
3.82 0.0701 -0.0101 1.54 0.21 Tota15.57 4.94 0.0781 -0.0021 0.98 0.26 0.02
RICHROMATE
Boiled and aerated
Boiled and aerated
-
Total 6.20 . , 0.0802 11/i 4.60 0.0740 4 . 0 0 6 2 Boiled and aerated 1 0.61 1 0.45 1 0.21
..
Total 5.87 5.90 0.0772 4 . 0 0 3 0 Distilled in steam 0.18 l1l2 0.05 1
0.0802 3
-
40 100
, . 0.0602 1 1
Total 6.13 4.21 0.0568 -0.0034 0.30
Distilled in steam
Distilled in steam
-
40 100
5 0.0602 1
Total 4.5 1 4.30 0.0580 -0.0022 0.30
40 100
5 0.0802
.................
1
Total 4.60 4.0010.0719 -0.0083 1.35 0.36
a/r 1
Total 5.71 4.61 0.0584 4 . 0 0 1 8 0.02
.................
1
40 100
10 0.0602
40 100
10 0.0802
a/4
1
*/4
1
-
Total 4.63
6.05k0.0789 -0.0013 0.22
-
.................
Total 6 27 NazSzOa solution, 1 cc. = 0.0126Yg. iodine. OXIDATION O F I O D I D E TO I O D I N E BY MEANS OF F E R R I C SULFATE
The use of ferric sulfate or ferric chloride for the liberation of iodine from iodides according t o the equation
+
+
565
+
Fe2(S04)3 2KI = 2FeS04 KzSO4 IZ has been advocated by many.1 This method is recommended for the determination of iodine in t h e presence of bromide and chloride also in practically all modern textbooks on quantitative analysisSa With the exception of Gooch and Mar and Gooch and Ensign, the authors mentioned have not given the method a very thorough investigation. Gavaxzi states t h a t when bromides are present, the use of ferric chloride is not permissible, since it will 2berate bromine, b u t t h a t this difficulty is overcome by the use of ferric sulfate, iodine alone being evolved. Gooch and Mar, working on the problem of a direct determination of chlorine in a mixture of chlorides and iodides, desired a method by which the iodine could be set free and removed from t h e solution by boiling. They found t h a t from a volume of 300 cc. containing I O ‘cc. sulfuric acid ( I : I ) , 5 g. ferric sulfate, and 0.0059 g. potassium iodide, every trace of iodine had disappeared after 5 minutes’ boiling. However, with I g. of potassium iodide, all of the iodine was not driven off after boiling for one hour, which they a t tributed t o the establishment of an equilibrium between t h e ferric sulfate and the ferrous, and they remedied this by adding a small amount of nitric acid t o change t h e ferrous iron t o the ferric state. Gooch and Ensign found the mixture of ferric sulfate, sulfuric acid and nitric acid, recommended above for the separation of iodides and chlorides, unsuitable for the separation of iodides and bromides, because bromine was also liberated. It was therefore decided t o investigate the behavior of ferric sulfate alone. I n the following experiments t h e apparatus employed was the same as in the experiments with potassium bichromate. The ferric sulfate was a C. P. reagent. This forms a clear solution in hot water, but hydrolyzes slightly on boiling. The hydrolysis is considerably increased when salts, sodium chloride for instance, are present. WITH FERRIC SULFATE Iodine Standard Fez(SO4)s NaCl Taken NazSzOa Iodine as K I ExPT. Taken Taken Soln. Found Error No. Grams Grams Gram cc. Gram Gram 1 1.5 0.0803 12.581 0.0798 4.0005 0.1004 15.64 0.0992 4.0012 0.1004 15.72 0.0997 -0.0007 0.0602 9.40 0.0596 -0.0006 0.0596 0.0602 9.40 -0.0006 0.1002 0.1004 15.80 -0.0002 0.0803 12.60 -0.0004 0.0799 0.1004 15.76 0.0999 4.0005 0.0012 0.0020 0.961 -0.0008 0.0020 1.10 0.0014 -0.0006 11 i.5 .. 0.0016 0.0020 1.28 4.0004 12 1.5 5 0.0008 0.20 0.0003 4.0005 13 2.0 0.0020 1.22 0.0015 --0.0005 14 2.0 10 0.0020 0.94 0.0012 -4.0008 1 Experiments 1 to 8, 1 cc. = 0.00634 g. iodine. 0.001268 g. iodinLe. 3 Experiments 9 t o 14 , 1 cc. Volume of solution in distillation flask, 75 cc.
TABLE 111-LIBERATION OF IODINE
..
..
-
Experiments I t o 14 in Table I11 show t h a t in a solution of 7 5 cc., iodine as potassium iodide in quantities as great as one-tenth of a gram alone, or mixed with as much as 5 or I O g. of sodium chloride, can be 1 Carnegie, Chem. News, 1891, pp. 60-87; Fallieres, Z. anal. Chem., a5 (1886), 554; von Topf, Ibid., a6 (1887), 299; Weiss, Re9erl. anal. ckim., 6 (1885). 238; Gavazzi, Gaaetta, 13, 327, 454; Gooch and Mar, Am. J. Sci., 89 (1890), 293; Gooch and Ensign, Ibid., 40 (1890), 145. 2 Olsen, 4th ed., p. 98; Sutton, “Volumetric Analysis,” 10th ed., p. 224; Scott, “Standard Methods of Chemical Analysis,” 2nd ed., p. 205; Treadwell, “Quantitative Analysis,” 3rd ed., p. 656.
T H E J O U R N A L OF I N D U S T R I A L A N D E N G I N E E R I N G C H E M I S T R Y
566
determined quite accurately. The iodine is completely removed after 20 or 30 min. distillation. Experiments 15 t o 23, Table IV, show t h a t in a volume of 7 5 cc., 2 g. of ferric sulfate will liberate bromine from 0.5 g. or more of sodium bromide, b u t not from 0.4 g. The liberation of bromine proceeds slowly and is not continuous, b u t ceases as soon as a n equilibrium between ferric sulfate and ferrous sulfate is established. TABLR IV-ACTION
EXPT
OF
Fen(SO4)a NaBr G r a m s . Grams 10.0 5 .o 2.0 1.0
FERRIC SULFATE O N RROMIDE Iodine as KI Gram
Standard hTazS?Oa Solution cc. 2.90 1.41 0.40 0.10 0.37 0.01
N 0. 15 16 17 18 1 .O 19 20 0.5 21 0.5 0.4 22 0.4 23 NazSzOs solution, 1 rc. = 0.00634 g. iodine. Volume of solution in distillation flnsk, 75 cc.
0.01 0.00 0.00
Iodine Found Cram 0.0184 0.0089 0.0025 0.0006 0.0023 0.00006 0,00006 0,0000 0.0000
LTpon distilling solutions of mixtures of iodide and bromide with ferric sulfate, one would expect t h a t t h e ferrous sulfate formed by t h e oxidation of t h e iodide would prevent or a t least hinder t h e oxidation of t h e bromide and i t is apparent upon comparing t h e results obtained in Expts. 24 t o 30 (Table V) with t h e results tabulated in Table IV t h a t t h e supposition is true. Thus, 2 g. ferric sulfate liberated from I O g. sodium bromide in t h e absence of iodide, bromine equivalent t o 0.0184 g. iodine, and in t h e presence of 0.0803 g. iodine as potassium iodide, bromine equivalent t o only 0.00g.i g. iodine (Expts. 1 5 and 24); also from 5 g. sodium bromide in t h e absence of iodide, bromine equivalent t o 0.0089 g. of iodine, and in t h e presence of 0.0803 g. of iodine as potassium iodide, bromine equivalent t o only 0.0042 g. iodine (Expts. 16 and 25). This can also be demonstrated by distilling a solution containing in 7 5 cc. I O g. sodium bromide with a mixt u r e of 2 g. ferric sulfate and 0.2 g. ferrous sulfate. TABLEV-EFFECT
BROMIDEON THE LIBERATIONOF IODINE WITH FERRIC SULRATE Iodine Standard NaCl NaBr as K I NazSzOs Iodine Error EXPT Fe.(SOi)a Present Present Present Solutlon Pound Gram Cc. Grams Grams Gram Gram Grams No. 14.211 10.0 0.0900 4-0.0097 24 2.0 5.0 13.33 0,0845 + O . 0042 25 2.0 12.76 3.0 0.0809 4-0.0006 26 2.0 12.68 0.0804 +0.0001 1.0 27 2.0 12.51 0.0793 -0.0010 .. 1.0 28 1.5 12.52 0.0794 -0.0009 1 .o .. 29 1.5 6.39 0.0405 +0.0003 1 .o 2.0 30 3.30 0.0209 +0.0008 1 .o .. 31 2.0 3.18 0.0202 +O.OOOl 0.5 .. 32 2.0 9.612 0.012’2 +0.0002 0.4 1.5 .. 33 6.55 0.0083 +0.0003 0.4 1.5 .. 34 1.58 0.0020 0.0000 0.4 1.5 35 1.68 0.0021 +O.OOOl 0.4 .. 2.0 36 1 . 2 0 0.0015 -0.0005 0.1 .. 1.5 37 1.78 0.0022 4-0.0002 0.4 10 2.0 38 0.55 0.0007 -0.0001 0.4 2.0 10 39 0.00 0.0000 .... 40. 2.0 10 0.4 1 Experiments 24 t o 32 1 cc. = 0.00634 g . iodine. 2 Experiments 25 to 40,’ 1 cc. = 0.001268 g. iodine. Volume of solution in distillation flask, 7 5 cc. OF
..
..
~~
No bromine will be liberated. On t h e other hand, when iodide is present in much smaller quantities t h a n i n t h e above experiments, a much smaller quantity of ferrous sulfate is formed and its influence is overcome and superseded by another influence, and t h e stability of t h e bromide toward t h e oxidizing influence of ferric sulfate is decreased. More bromine is liberated when a small amount of iodide is present t h a n when iodide
Vol.
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No. 6
is entirely absent. I n t h e experiments tabulated in Table 111, solutions of pure potassium iodide were distilled with ferric sulfate and i n all cases t h e results obtained for iodine were slightly less t h a n theory. Now i t has been demonstrated in Expts. 2 2 and 23, Table IV, t h a t 0.4 g. of sodium bromide is not decomposed by as much as 2.0 g. of ferric sulfate. However, when a mixture of 0.4 g. of sodium bromide and 0.0120g. of iodine as potassium iodide was distilled with 1.5 g. ferric sulfate, there was a positive error in t h e iodine result of 0.0002 g., and a ‘mixture of 0.4 g. of sodium bromide and o.oo80 g. iodine as potassium iodide gave a positive error of 0.0003 g. These errors are small and negligible from a n analytical point of view, t o be sure, b u t they show t h a t t h e stability of t h e sodium bromide has been decreased. I n Expts. 11 and 13, only 0.0016 g. of iodine was obtained from 0.0020g. of iodine as potassium iodide, while from mixtures of 0.4 g. of sodium bromide and 0.0020 g. of iodine as potassium iodide, o.oo20 g. and 0.0021 g. were obtained (Expts. 3 5 and 36, Table V). Apparently t h e accuracy of these last two results leaves nothing t o be desired, b u t their apparent accuracy is probably due t o a slight decomposition of t h e sodium bromide. I n order t o investigate this further, t h e experiments tabulated in Table VI were made. The specified quantity of sodium bromide was introduced into a distillation flask, 7 5 cc. of water poured in, and then t h e ferric sulfate added. Pure iodine (Col. 3) was weighed in a weighing tube, t h e t u b e introduced into t h e neck of t h e distillation flask where t h e t u b e was opened and t h e iodine allowed t o fall t o t h e bottom of t h e flask. T h e flask was then quickly stoppered and distilled with steam as before. The results show t h a t t h e presence of iodine reduces t h e stability of t h e bromide. We have seen before (Expts. 2 2 and 23, Table IV) t h a t no bromine will be liberated from 0.4 g. of sodium bromide distilled with 2.0 g. of ferric sulfate, yet in t h e presence of iodine, 1.5 g. of ferric sulfate decomposed slightly 0.3 g. of sodium bromide (Expts. 4j and 46), also considerably more bromine is liberated from 5 g. of bromide when iodine is present t h a n when i t is absent. It seems possible t h a t this is due t o t h e formation of a small amount of bromiodide. TABLEVI-INFLUENCE
OF
IODlNE
ON THE
OXIDATION
FERRIC SULFATE
NaBr Iodine NazSnOa Fez(SO4)a Taken Taken Taken Solution Gram cc. Grams Grams 0.0486 7.60 0.0 0.0 0,0439 6.72 0.0 1.5 0.0948 8.48 3.0 0.0 0.0390 6 10 0.1 1.5 0.0418 6.78 0.3 1.5 0.0512 8.20 0.3 1.5 6.96 0.0431 1.5 0.5 0.0586 11.45 1.5 5.0 8.04 0.0385 1.5 5.0 49 NaZSzOa solution, 1 cc. = 0 00632 g. iodine. Volume of solution in distillation flask, 75 cc.
EXPT. No. 41 42 43 44 45 46 47 48
OF
BROMIDEWITH
Iodine Found Gram 0,0480 0,0425 0 .05 36 0.0386 0.0428 0.0518 0.0440 0.0724 0,0508
Error Gram -0.OOOh -0.0014 -0.0012 -0.0004 +O.OOlO 4-0.0006 4-0.0009 +0.0138 t0.0123
Bergland,’ working on a method for estimating bromide in t h e presence of chloride by liberating t h e bromine with potassium bisulfate and potassium permanganate and removing i t by aspiration, also found 12:
anal. Chem., 1886, p 184.
'
J u n e , 1919
T H E J O U R N A L OF I N D V S T R I A L A N D ENGINEERING C H E M I S T R Y
t h a t chloride was less stable toward t h e oxidizing agent i n t h e presence of a bromide. T h e reaction between ferric sulfate and a solution containing iodide and bromide is, therefore, somewhat complicated. I t is a reversible reaction. Using quantities of ferric sulfate and potassium iodide given in Table 111, we see t h a t t h e reaction proceeds completely t o t h e right, b u t by adding ferrous sulfate it can be prevented from proceeding entirely t o t h e right and not all of t h e potassium iodide will be oxidized. Substituting sodium bromide for potassium iodide, we see in Expts. 15 t o 21, Table IV, t h a t only a small amount of bromide is decomposed and t h a t a n equilibrium is established as soon as a small amount of ferrous sulfate is formed. With mixtures of iodides and bromides, if there is present sufficient iodide t o form a n appreciable amount of ferrous sulfate, this ferrous sulfate will prevent, or a t least hinder, t h e decomposition of t h e bromide. Now we have also seen t h a t upon distilling a mixture of bromide, iodine, and ferric sulfate, there is certain evidence which points t o t h e formation of a brom-iodide. Ferrous sulfate will hinder t h e formation of this brom-iodide, so if iodide is present i n amounts equivalent t o from 0.040 t o 0.100 g. of iodine, t h e iodine may be separated, by distillation with ferric sulfate, from as much as I or z g. of sodium bromide without any appreciable error due t o t h e liberation of bromine or t h e formation of bromiodide. Sufficient ferrous sulfate is formed by t h e oxidation of t h e potassium iodide t o prevent t h e liberation of bromine or t h e formation of brom-iodide. If t h e amount of iodine as iodide present is less t h a n 0.040 g., then bromine will be liberated from t h e above amount of bromide along with t h e iodine and probably also some brom-iodide will be formed because not sufficient ferrous sulfate is formed t o prevent it. And now, finally, if we have a very small amount of iodine, equivalent t o only 0 . 0 0 2 or 0.003 g. of iodine i n solution with 0.3 or 0.4 g. of sodium bromide, i t is not likely t h a t any free bromine is liberated b u t a small amount of brom-iodide is formed and distils over with t h e iodine. I t is apparent, therefore, t h a t in distilling mixtures of bromide and iodide with ferric sulfate t h e stability of t h e bromide depends upon t h e absolute concentration of bromide and also upon t h e ratio of t h e concentration of iodide t o t h e concentration of bromide, and t h a t this method is not as reliable as t h e permanganate method. If, however, one is only interested in removing iodine so t h a t i t will not interfere with t h e subsequent bromine determination no difficulty will be encountered. The sample of brine or mineral water usee for t h e bromine determination will not as a rule contain more t h a n 0.1 g. of bromine and usually less on account of t h e small content of bromide compared with t h e content of other salts and t h e disadvantage of working with more t h a n I O g. or so of salts. Experience has shown t h a t iodine can be accurately separated from such a quantity of bromide. This will be more fully brought out in a subsequent paper. Iodide in t h e presence of a large amount of bromide may be determined by making a double distillation
567
with ferric sulfate. The first distillate consisting of iodine and a small amount of bromine is collected in I O O cc. of a sodium hydroxide solution (50 g. p e r liter) containing 5 t o I O cc. of 3 per cent hydrogen peroxide. The distillate is then boiled down t o 40 t o 50 cc. and made slightly acid with sulfuric acid. If any iodine is liberated during acidification, i t is reduced with a few drops of a solution of sodium sulfite or sodium thiosulfate. It is then washed into t h e distillation flask, made up t o a volume of 7 5 cc., 1.5 - 2 . 0 g. ferric sulfate added, and distilled with steam into a potassium iodide solution ( I O g. potassium iodide per I O O cc.) which is then titrated with sodium thiosulfate. T h e results tabulated in Table V I I were obtained by this method. TABLEVII-DOUBLE DISTILLATION WITH FERRIC SULFATE (1.5 0 . F e r r i c Sulfate Used for Each Distillation) NaCl lTaBr I as Taken Taken K I Taken I Error Grams Grams Gram Found Gram
.. .. ..
... ... ...
5.0
.5
5.0
..
5.0 5.0 10.0
.. ..
0,0803 0,0402 0.0201 0.0803 0.0803 0.0402 0.0602 0.0201
0.0790 0.0389 0.0196 0.0784
-0.0013 -0.0013 -0.0005
0.0784
-0.00 19 -0,001 1 -0.0024
0.0391 0.0578 0.0196
-n.nnig
-0,0005
These results are fairly satisfactory b u t not as good as those obtained by t h e permanganate method, while t h e procedure is more time-consuming and difficult t o carry out.
c 0N c L us I0 N s The permanganate method is recommended for t h e determination of iodine in mineral waters and brines and accurate results can be obtained by observing t h e following directions: Take such a quantity of the brine or water as will contain not more than 0.1g. iodine a s iodide or more than I O g. total salts, adjust the volume t o IOO t o 150 cc. and boil i t with a sufficient amount of sodium hydroxide and sodium carbonate t o precipitate the calcium and magnesium. Filter off t h e precipitate and wash with hot water, introduce t h e filtrate into a n Erlenmeyer flask, adjust the volume t o about IOO cc., neutralize with dilute sulfuric acid, and add I cc. of a solution of sodium hydroxide (4 g. per 100 cc.). Heat t o boiling, add an excess of potassium permanganate, continue the heating until the precipitate begins t o coagulate, and then allow t o cool. Add sufficient alcohol to cause the permanganate color to disappear and allow the precipitate to settle on t h e steam bath. Filter and wash with hot water. After cooling, add one or two grams of potassium iodide, acidify with hydrochloric acid, and titrate with standard thiosulfate The number of cubic centimeters required, divided by six, represents the number of cubic centimeters required by the iodine originally present in the sample.
The iodine may be removed from t h e sample in which bromine is t o be subsequently determined according t o t h e following method: The neutral or slightly acid sample which should contain not more than o I g. of bromine or I O g . total salts is introduced into the distillation flask and adjusted t o a volume of approximately 75 cc., 1.5-2.0 g. of ferric sulfate added, the liberated iodine distilled with steam into IOO cc of a potassium iodide solution ( I O g. potassium iodide per IOO cc.). The potassium iodide solution may be titrated with sodium thiosulfate solution and the result used t o check the figure obtained by the permanganate method. The bromine may be determined in the liquid remaining in the distillation flask according t o a method which will be described in a later paper.
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T H E JOURNAL OF INDUSTRIAL A N D ENGINEERING CHEMISTRY
Vol.
11,
No. 6
SUMMARY
The permanganate method for the determination of iodine in the presence of bromides and chlorides gives very satisfactory results. Distillation with potassium bichromate will not completely liberate iodine from iodides. Iodide in the presence of as much as I O g. sodium chloride can be determined with satisfactory results by distilling with ferric sulfate. Iodide can be determined quite accurately in t h e presence of bromide by a single distillation with ferric sulfate if the quantity of bromide present in t h e sample does not greatly exceed 0 . 4 g. This method is especially recommended for removing iodine from the sample which is to be used for the bromine determination. Iodine may be determined in t h e presence of large amounts of bromide by making a double distillation with ferric sulfate. However, this method is more tedious and time-consuming than the permanganate method and the results obtained not quite so accurate.
U
cu
BUREAUOF CHEMISTRY OF AGRICULTURE DEPARTMENT WASHINCYI’ON, D. C.
STUDIES ON CANNING’ AN APPARATUS FOR MEASURING THE RATE OF HEAT PENETRATION By W. T. BOVIE AND J. BRONBENBRBNNER Received December 17, 1918
The changes in t h e rate of destruction of bacteria a t various temperatures is no exception t o the general rule t h a t t h e speed of reactions greatly increases with temperature. For example, according t o Miss H. Chick, the temperature coefficient,2 p, for the killing ~ for the of Bacillus typhoszcs in hot water is 9 2 , 0 0 0 , and destruction of t h e Bacillus paratyphoszrs by means of phenol E.C = 48,000. For comparison, t h e value p for the hydrolysis of cane sugar by acid is given by Arrhenius as 25,600. Since the rate of sterilization is inversely proportional t o the time necessary for killing, we have, in any problem of sterilization by heat, a t least two variables-time and temperature. In the case of sterilization of canned foods by heat a third variable, namely, the distance of any particular organism from t h e surface of the can, must be considered. It is impossible from the d a t a now a t hand t o formulate a law connecting these three variables-time, space, and temperature-for, due to the heterogeneous nature of most canned goods and the possible changes in t h e physical properties as sterilization proceeds, there is a n ever varying amount of heat transferred by con1 This work is a part of the investigation of food poisoning, conducted under the direction of Dr. M. J. Rosenau, Professor of Preventive Medicine and Hygiene, Harvard Medical School. The investigations were done under the auspices of the Advisory Committee of the National Research Council on the Toxicity of Preserved Foods, and under a grant t o Harvard University from the National Canners’ Association. The relation between the rates of destruction Ka and KIfor two differe n t temperatures To and TI, respectively, is given by the formula
*
in which a is the base of the natural system of logarithms, and p i s a constant which is greater the more rapid the rate of destruction. Calculated from the data given by Miss H. Chick, “The Process of Disinfection by Chemical Agencies and Hot Water,” J . Hyg., 10 (1910), 237.
*
FIG.
1
vection from the outside of the can towards t h e center. A mathematical treatment of the problem such a s might be given for the transfer of heat in a solid conductor, as a disc of iron, is quite out of the question.’ Jn t h e present paper a method is described for determining experimentally the rate of heat transfer in canned foods, during the process of sterilization. The apparatus t o be described below has passed through
a
FIG.2 1 R . E. Buchanan gives a formula which contains the diffusivity of the canned food as a constant term. Experiments which will be reported later show that, especially with starchy foods, the diffusfvity changes during the process of cooking. The change is undoubtedly caused by the absorption of water by the various components of the food. Since the amount of this absorption is influenced by the ripeness of the food, the acidity, etc., the possibility of a practical application of Buchanan’s formula seems very doubtful. (R. E. Buchanan, et al., “Notes on Conditions which Influence Thermal Death Points,” Abslr. of Bad., 2 (1918), 5.)
