October 1952
INDUSTRIAL AND ENGINEERING CHEMISTRY LITERATURE CITED
(1) Austin, J. B., and Day, M. J., IND.ENG.CEIEM., 33, 23 (1941). (2) Beoker, K.‘, 2. Metallkunde, 20, 437 (1928). (3) Gregg, J. L.,“Alloys of Iron and Tungsten,” Chapt. IV, New York, McGraw-Hill Book Co., Inc., 1934. (4) Hurd, D.T., U. S. Patent 2,554,194(June 1950). ( 5 ) Lander, J. J., and Germer, L. H., Metals Technot., 14 (September 1947).
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(6) Norton, F. J . , private communication to D. T. Hurd. (7) Schenk, R., Kursen, F., and Wesselkoch, H., 2. anorg. u. allgevn. Chem., 203, 159 (1931). RECEIVED for review April 14, 1952. ACCEPTEDM a y 20, 1952 Presented at the X I I t h International Congress of Pure a n d Applied Chemistry, New York, September 1951.
Reaction between Ferrous Iron
and Dissolved Oxygen in Brine D. C. BOND AND G. G. BERNARD The Pure Oil Co., Research and Development Laboratories, Crystal Lake, Ill.
I
APPARATUS
N OIL field technology the reaction between ferrous iron
Figure 1 is a diagram of the apparatus used. All rertotions and oxygen in brine is of considerable importance. In an were carried out in a 1-liter Florence flask. A magnet, sealed in oil- or gas-producing well, water from two sands, one containing glass, was placed in the flask and rotatbedby means of a rotating oxygen and the other containing ferrous iron, can mix and react. In a water aeration tower ferrous iron is oxidized to form insoluble magnet under the flask. Preliminary tests showed that if the flask, partially filled with ferric hydroxide, which can be settled and filtered out of the water. brine, was stoppered and allowed to stand, no measurable change In a “closed” system for disposal of oil-well brine, the accidental in oxygen content of the brine in the flask would occur in 3 hours. entrance of air into the system can lead to the oxidation of ferrous In all experiments lasting an hour or less, samples were merely iron in the brine. Brine or fresh water containing oxygen can withdrawn from the flask by means of a pipet, and the flask was be injected into a sand, where it may mix withwater that contains then stoppered. ferrous iron: or in areas where air has been injected into oil In experiments lasting more than an hour, samples of brine sands for secondary recovery of oil, oxygen can be dissolved, were displaced through the capillary tube in the rubber stopper reacting with ferrous iron in the connate water in the sand. in the neck of the flask. This was done by inflating a rubber In all of these cases it ie important to know something about balloon inside the flask with water from a separatory funnel, the rate of reaction between ferrous iron and dissolved oxygen. Tests showed that if the balloon was filled with air-saturated This information is needed in order to predict the retention time water, no measurable amount of oxygen diffused through the required after aeration of water, to tell whether plugging of balloon into the brine in the flask in 3 days. Thus, it was possible filters or injection sands by ferric hydroxide can occur, to predict to stir the reaction mixture and withdraw samples from time to what will happen when various fluids mix in underground rocks time without having air come in contact with the brine in the or in well bores, or to predict whether a given mixture of brines flask. The reaction flask was immersed in a water bath mainwill be corrosive. The reaction between ferrous iron and dissolved oxygen has tained at constant temperature (=k0.Zo F.). Dissolved oxygen determinations were made with a polarobeen studied by many investigators. In most cases a procedure has been used which involves the air blowing of solutions ( 4 , 6, graphic dissolved oxygen meter ( 1 7 ) 7, 1.2, 14, 99). In these cases determination of the rate of PROCEDURE Water reaction between ferrous iron The reaction between ferand oxygen has been obscured rous iron and dissolved oxygen by possible changes in rate of Produces hydrogen ion. Since solution of oxygen and changes it was desired that the experiin solubility of oxygen. With ments be carried out a t constant pII, the brines used were few exceptions quite concenRemoving Samples Slirrer buffered by the addition of trated (6, 13) or strongly boric acid, potassium acid Florence acidified solutions have been phthalate, and potassium hy“--h Nater Healer droxide. Table I shows the used (10,13, 18). J composition of the solutions It appears that no work has used in the experiments. been done on systems comBrine of the desired concenmonly encountered in the oil tration was pre ared with disfield, that is, mixtures containtilled water a n f reagent-grade sodium chloride. The soluing about 1 to 10 p.p.m. distion was buffered a t the desolved oxygen and 1 to 100 sired pH and then air blown p.p.m. ferrous iron with high Regulator for several hours to produce concentrations of salts. The a brine that was nearly satuMagnetic rated with air. The pH valuee present work was undertaken were determined with the glass in order to obtain infsrmation electrode, corrected for sodium about the behavior of such content of solutions tested. Figure 1. Apparatus for Studying Reaction between systems. Dissolved Oxygen and Ferrous Ion The buffered brine solution,
Y
!I Illya
Yh
INDUSTRIAL AND ENGINEERING CHEMISTRY
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TIME
Figure 2.
Vol. 44, No. 10
- MINUTES
Effect of pH on Reaction between Ferrous Ion and Oxygen
Ratio of equivalents of ferrous ion to equivalents of oxygen = 1.07 Sodium chloride concentration, 10 Temperature, 80’ F.
I
Figure 3.
Effect of Temperature on Reaction between Ferrous Ion and Oxygen
Ratio of equivalents of ferrous ion t o equivalents of oxygen = 1.07 Buffered at pH 7 Sodium chloride concentration, 10 70
for 6 minutes. At various time intervals samples of brine were withdrawn or displaced from the flask and analyzed for oxygen by means of the dissolved oxygen meter, In one set of experiments the p H was varied, while temperature, sodium chloride concentration, and the initial ferrous ion-oxygen ratio were held constant. Although the sodium chloride concentration was constant, the total salt concentration varied somewhat from solution to solution, since different buffer salt concentrations mere required to obtain the desired pH values. Table I gives the composition of the buffered salt solutions used. The results of these experiments are given graphically in Figure 2. In another set of tests the temperature was varied, while pH, sodium chloride concentration, and initial ferrous ion-oxygen ratio were held constant (Figure 3). Figure 4 shows the results
containing dissolved oxygen, was then poured into the reaction flask and allowed to stand several hours. A sample of the brine was then displaced from the reaction flask into the test cell of the oxygen meter and the oxygen content of the sample was determined immediately. Table I gives this initial oxygen Concentration for the solutions used. From the concentration of oxygen in the brine and the capacity of the reaction flask the total amount of dissolved oxygen in the flask was calculated. The amount of Mohr’s salt equivalent to the oxygen in the flask was calculated by means of the equation 4Fe++
+ 02 + 4H.20 --+2Fez03 + 8H+
This amount of finely powdered i\lohr’s salt was weighed out on an analytical balance and placed in the reaction flask. The contents of the flask was stirred by means of the magnetic stirrer
Figure 4.
Effect of Initial Ferrous Ion-Oxygen Ratio on Reaction between Ferrous Ion and Oxygen Buffered at pH 7 Temperature, 80‘ F. Sodium chloride concentration, 10 Yo
INDUSTRIAL AND ENGINEERING CHEMISTRY
October 1952
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SO% completed was 170,3.5, and 1.3 minutes for ratios of equivaOF BRINES USEDIN STUDYING REACTION lents of ferrous ion to equivalents of oxygen of 1.07, 2.14, and TABLE I. COMPOSITION BETWEEN DISSOLVED OXYGENAND FERROUS IRON 4.28, respectively. This applies t o 10% sodium chloride solution Reaotion CSHI buffered at p H 7 at 80" F. (COOH) Dissolved Conditions The reaction rate is not NaC1, KOH, HaBOs, COOK Oxygen T$rn$., EFFECT OF SALTCONCENTRATION. %
%
1 10 10 10 10 26 10 10 10 10 10
10
%
3.24 2.08 2.08 2.08 2.08 1.21 0 0 0 0 0.326 0.095
%
M~L.' 4.40
0 0
too
FJH
80 32 80 100 120 80 32 80 80 80 80 80
7.0 7.0 7.0 7.0 7.0
7.0 5.0 5.0 6.0 6.5 8.0 9.0
changed measurably by a change in salt concentration. EFFECT OF HYDROUB FERRIC OXIDE. The apparent reaction rate is decreased considerably by the presence of hydrous ferric oxide, INDUCTION PERIOD.At low p H values, low temperatures, or low ratios of ferrous ion-oxygen the reaction does not begin immediately after the ferrous salt is added to the oxygen solution. The induction period varies from a few minutes to several hours,
ceed to completion under the conditions tested. With ferrous ion-oxygen ratios of 1, 2, and 4, about 18, 6,
obtained in still another set of experiments in which the initial ferrous ion-oxygen ratio was varied, while pH, temperature, and salt concentration were held constant. Figure 5 gives the results obtained in buffered 1, 10, and 26% sodium chloride solutions at constant pH, temperature, and ferrous ionoxygen ratio. Here again it should be noted that the reaction mixture contained buffer salt in addition t o the sodium chloride. I n another set of experiments the effect of hydrous ferric oxide on the rate of the ferrous ion-oxygen reaction was studied (Figure 6). RESULTS
EFFECT OF pH.
io00
10000
The time required for the Time-Minutes Figure 6. Effect of Hydrous Ferric Oxide on Reaction between Ferrous reaction between ferrous iron and oxygen to be Ion and Oxygen 80% completed was 2600,1400,650, 170,10, and Initial concentration of ferrous ion, 0.00043 equivalent per liter '/z minutes at p H values of 5, 6, 6.5, 7, 8, and Ratio of equivalents of ferrous ion to equivalents of oxygen = 1.07 9, respectively. This applies t o buffered 10% Buffered at pH 7 Temperature 80' F. sodium chloride solution with approximately Sodium ohloiide concentration, 10 Q equivalent quantities of ferrous iron and oxygen at 80" F. EFFECT OF TEMPERATURE. The time required for the reaction period is observed. Ten minutes to several hours may be rebetween ferrous iron and oxygen to be 60% completed was 660, quired for the reaction to be initiated, depending uponreaction 12, 4,and l/t minutes a t temperatures of 32", 80°,loo", and 120" conditions. F., respectively. This applies to 10% sodium chloride solution As the temperature is lowered from 80" to 32' F., the rate of buffered at p H 7 with approximately equivalent quantities, reaction is decreased by a factor of about 50 or lOO(Figure3). initially, of ferrous iron and oxygen. I n the interval from 80" to 120" F. the effect of temperature is EFFECT OF RATIOOF FERROUS IRON TO OXYGEN. The time more nearly what might be expected-that is, the rate doubles required for the reaction between ferrous iron and oxygen to be roughly with each temperature rise of 10' to 20" F.
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I N D U S T R I A L A N D E N G I N E E R I N G CHEMISTRY
The reaction rate is greatly influenced by a change in the p H of the reaction mixture (Figure 2). Thus, a change from p H 6 to p H 8 can increase the reaction rate by a factor of 500. One unexpected result is that t h e reaction between ferrous iron and oxygen does not proceed t o completion. With equivalent amounts of ferrous iron and oxygen an equilibrium is reached when about 80 to 82% of the oxygen has reacted (Figure 4). Even with two equivalents of ferrous iron per equivalent of oxygen, about 7% of the oxygen remains, while with four equivalents of ferrous iron per equivalent of oxygen about 4% of the oxygen remains unreacted at equilibrium. These statements pertain to solutions a t 80" F buffered at p H 7 . The same equilibrium is reached at p H 7 , regardless of the reaction temperature (Figure 3). However, it appears that the reaction more nearly approaches completion a t both higher and lower p H values (e.g., p H 9 or p H 5 ) than it does at p H 7 (Figure 2). Under the conditions tested, salt concentration has no measurable effect on the reaction rate (Figure 5 ) . This is in agreement with the results of Lamb and Elder (14j and shows that thegoverning step in the reaction involves a t least one nonionic reactant (16). Calculations of first-, second-, and third-order velocity constants, using the data given here, show that the reaction is complex. The data do not satisfy equations for any simple reaction. COMPOSITION OF EQUILIBRIUM MIXTURE
B meamrable amount of oxygen was found in all of the reaction mixtures, even after the reaction had apparently ceased (Figures 2, 3, and 4). There are several possible explanations of this effect, but none of them fits all of the facts. It might be suspected that the reading obtained with the oxygen meter was erroneous, that is, perhaps the meter indicated that oxygen was present cven though the oxygen concentration was zero. However, it was observed that if the equilibrium mixture was b1oP-n with nitrogen, the oxygen content of the mixture as measured by the ovygen meter gradually dropped to a value near zero. This was taken as an indication that the meter was operating satisfactorily. The results cannot be explained on the assumption that the Mohr's salt used was impure. Even if the salt had contained as little as SO'% of the calculated amount of iron, there should have been a great excess of ferrous iron in the tests in Rhich the calculated ferrous ion-oxygen ratio was twoorfour(Figure4.) Inthese tests oxygen was found in the equilibrium mixture. There is a possibility that ferrous iron was adsorbed by the hydrous ferric oxide formed in the reaction. Results given in Figure 6 show that in the presence of large amounts of ferric oxide the rate of removal of oxygen was diminished greatly. This would be expected if the ferrous ion were adsorbed by the ferric oxide; the reaction of oxygen with ferrous ion on a surface should be much slower than the reaction with ferrous ion in solution. This probably explains partly why oxygen was found in the equilibrium mixtures. However, it can hardly explain the results obtained with ferrous ion-oxygen ratios of two and four. In these cases the ferric oxide produced by the reaction would have needed to adsorb 1 and 3 equivalents, respectively,
Vol. 44, No. 16
of ferrous ion per equivalent of ferric oxide, in order to remove all of the ferrous ion from solution. This is not likely. Further, tests with colored indicators shoir-ed qualitatively that ferrous ion was present in the equilibrium mixtures, at least in casea where the initial ferrous ion-oxygen ratio was two or greater. Adsorption of ferrous ion by the ferric oxide explains some, but not all, of the observed effects. CONCLUSIONS
Changes in temperature and in p H affect the rate of reaction between ferrous iron and oxygen in the manner expected-Le., the rate of reaction decreases as the temperature or the p H of the solution decreases In the range from 80" to 32' F. the effect of temperature is abnormally great. At low temperatures or at low p H values, the reaction exhibits an induction period of several minutes to an hour or more depending upon conditions. The rate of reaction is not influenced by salt concentration within the limits tested. The reaction between ferrous iron and oxygen does not proceed to completion even in the presence of excess ferrous iron. With equivalent amounts of ferrous iron and oxygen at plI 7 the amount of unreacted oxygen a t equilibrium (about 20%) is greater than a t higher or lover p H values. The rate of reaction between ferrous iron and oxygen is greatly decreased by the addition of colloidal ferric hydroxide to the reaction mixture. REFEREKCES
(1) bgde, G., and Sohinimel, F., Z . a m ~ i g u. . allgem. Chem. 225, 29-
32 (1935). (2) Banerji, P. K., J . PTOC. Asiatic SOC.B ~ i i y a ,18, 1255 (1922). (3) Banerji, P. IC, 2. anorg. u . allgem. Citenz., 128, 343-9 (1923). (4) Baskerville, C., and Stevenson, R., J . Am. Chem. SOC.,33, 1104 (1911). (5) Chretien, A., and Rohmer, R., ,4nn. d i m . , 18, 267-85 (1943). (6) Cornog, Jacob, and Hershberger, dlbert, Proc. I o w a Acad. Sci.. 36,264-5 (1929). (7) Ennos, F. R., Proc. Cambi-idgePhil.Soc., 17, 182 (1911). (8) Friend, J. A. X., and Pritohett, E. G . K., J . Chem. Soc., 1928, 3227-32. (9) Halvorson, H. O., and Starkey, R. L., J . Phgs. C'hem., 31,626-31 (1927).
(10) (11) (12) (13)
Jilek, A,, Chem. Lisly, 15,105-9, 138-40 (1921) Just, G., 2.p h y s i k . Chem., 63,385-420 (1908). Karpova, I. F., J . Gen. Chenz. (U.S.S.R.), 7,2613-19 (1937). Kobe, K. A , , and Dickey, IT,,IND.ENG.CHEW, 37, 429--31 (1945). (14) Lamb, A. B., and Eider, L. IT'., Jr., J . Am. Chem. SOC..53, Lli7 (1931). (15) La Mer, V. K., C h e m Revs., 10, 179 (1932). (16) MaoArthur. G. G.. J . Phus. Chem.. 20. 545 (1916) (17) Marsh, G. .4.,A n a l . Chem., 23, 1427 (i95l). (18) Mikhelson, E. M., J . Gen. Chein. (U.S.S.R.), 1, 905-9 (1931) (19) Posnjak, E., Am. Inst. Maiiing M e t . Engis., Contrzbs., No. 1615D (1926). (20) Poulld, J. R,., J . Pilus. Chenz., 43, 955-67, 969--80 (1939). (21) Pounds, J. R., J . Soc. Chenz. Ind. ( L o n d o n ) , 55, 327~-30'1' (1936). (22) Reedy, J. H., and Machin, J . S., IND.ENG.CHEM.,15, 1271-2 (1923). (23) Thomas, R., and Killiains, E., J . Chem. Soc., 119, 749-58, (1921) - 4 C C E P T E D July 16, ! \ ) 5 2 . RECEIVED for review Sepetmber 20, 1951 \ - - ,
