INDUSTRIAL A N D E N G I N E E R I N G C H E M I S T R Y
748
For the same reasons, if the final product were assumed to be a mere mixture of carbon and filter aid rather than the composite substance we believe it to be, the flow rate and clarity produced by a unit dose of final product A should be substantially identical with the flow rate and clarity produced by the same dose of an equivalent mechanical mixture, D, These predictions were fully realized for the mixtures but not for product A: -1% TreatmentFlow rate Clarity Brilliant Filter aid 5 1.00 Fair Carbon U 0.04 Poor Carbon H 0.10 Poor Carbon I 0.08 Poor Charcoal C 0.10 Brilliant Final product A 0.75
TreatmentFlow rate Clarity 1.20 Brilliant 0.42 Fair 0.28 Brilliant Brilliant 0.82 Brilliant 0.42 0.90 Brilliant
-2%
Filter aid B Mixture J Mixture K Mixture L Mixture D Final product A
On a 2 per cent basis the flow rate of the final product is more thandouble that of the correspondingmixture and substantially the same as the flow rate of the 1 per cent filter aid addition, while the clarity of the filtrate from the product is substantially identical with that of the filtrate from the filter aid alone. The advantage in flow rate of final product A on a 1 per cent dosage is even more apparent if this material is compared to equal percentages of carbons C, G, H , and I . These carbons, however, are not used by themselves as a general rule, and the better comparison is furnished on the mixture basis. Figure 2 illustrates the results of filtration with filter aid B and with final product A . It appears evident that this new product is in no sense a mixture of decolorizing carbon with diatomaceous filter aid but is, on the contrary, a substance having properties entirely different from and more valuable than those of the corresponding mixture.
M-23 in the Sugar Factory Raw cane sugar liquors contain certain percentages of suspended, insoluble, finely divided solids that are ordinarily removed by filtration using diatomaceous silica filter aids.
Vol. 34, No. 6
I n addition to the suspended solids there is a small percentage of dissolved impurities and the sirup is fairly dark. The color as well as some of the dissolved impurities are removed in char filters, When M-23 is used in the sugarhouse, it is proposed that this product be substituted for the diatomaceous filter aid ordinarily employed. Since M-23 product has equal filter aid efficiency, it will effectively remove the suspended solids by the normal filtration steps and in addition will show a good decolorizing effect on the plant liquors. This treatment will give results comparable to bone char and will permit simplified handling. A considerable proportion of the load on the char houses is thus removed and a considerable saving in the refining process results. In plants not employing a char house for decolorization of their washed sugar liquors, this step can be accomplished by the use of M-23. I n these plants effective and efficient treatment of the washed sugar liquor ii3 obtained by a threestage countercurrent decolorizing step. The sugar liquor is decolorized in large mechanically agitated tanks, and the M23 is removed by filtration in a pressure flter. M-23 does not require additional filter aid to obtain satisfactory filtration of the bleached sugar liquor. The results are considerable saving in the decreased amount of solids to be handled, less press labor, less wash water, and fewer sugar losses. Small-scale tests indicate that as little as 6.0 pounds of M-23 per ton of sugar solids yields a decolorized washed sugar liquor of excellent quality, and a filtration rate of 10 to 20 gallons per square foot per hour is obtained. Another advantage obtained on the hI-23 material is the comparative ease of wetting. The product does not require the long stirring period usually associated with the use of vegetable carbons but instead wets itself readily with a minimum of labor and dusting. PREBENTID in a group of papers on Decolorizing Carbons and Analysis before the Division of Sugar Chemistry and Technology a t the 102nd Meeting of the AMERICAN CHEXICAL SOCIETY~ Atlantic Cily, N. J.
Corrosion of Steel by issolved Carbon Dioxide and Qxygen G . T. SIWERDAS' AND H. H. UHLIG2 Department of Chemical Engineering, Massachusetts Institute of Technology, Cambridge, Mass. NEXPECTED corrosion difficulties have been found in parts of the condensate return lines of central steam heating systems. The condensates carried by these return lines are essentially dilute solutions of carbon dioxide and oxygen. As a result, the problem of steel corrosion in solutions of carbon dioxide and oxygen and the relative importance of each gas has received renewed interest and is the subject of the present study. The major factors to be considered in this problem are temperature, dissolved carbon dioxide and oxygen concentrations, pH, circulation of corroding medium, metal composition, and duration of attack. I
Present address, The M. W. Kellogg Company, New York. N. Y. Present address, General Electria Comprtny, Sahenectady, N. Y.
Previous Investigations
h number of incidental experiments bearing on the effect of carbon dioxide and oxygen on the corrosion of iron and steel have been reported. Rhodes and Clark (8)reported that up to 200 pounds per square inch, corrosion a t 22" C. increases rapidly with carbon dioxide pressure but reaches a constant value a t about 300 pound:; per square inch, further increase in pressure having no eflect. Whitman, Russell, and Altieri (14) found that dissolved carbon dioxide increased corrosion as the pH was lowered, and that a t a given pH more corrosion was caused by carbon dioxide than by hydrochloric acid. Girard (4) confirmed these results, as did Groesbeck and Waldron (6) who investigated a wide range of
INDUSTRIAL AND ENGINEERING CHEMISTRY
June, 1942 ~
Controlled laboratory experiments indicate that both dissolved carbon dioxide and oxygen in steam condensates can be effective in causing failure of return pipes and that oxygen is relatively five to ten times more effective than equal concentrations of carbon dioxide. The two gases simultaneously may be considerably more corrosive than would be expected for the two gases acting individually. Condensates at 90" C. are two to two and a half times as corrosive as those at 60' C. When corrosion is caused largely by dissolved carbon dioxide, the structure and composition of the steel are factors in determining its life.
Procedure The corrosion tests were conducted in three steps: (a) Gas previously analyzed for oxygen and carbon dioxide was passed into the cell until the distilled water was saturated. By analysis this was found to require from 1 to 2 hours. Saturation by carbon dioxide mixtures could conveniently be followed by noting when the p H of the solution reached a steady value. (b) Three auxiliary specimens identical with the test specimens in every respect were then inserted to produce steady-state conditions. The p H in carbon dioxide solutions increased as much as 1.0 unit upon intro~
FLOWMETER
D f
In this investigation corrosion under controlled conditions was carried out in a glass reservoir 14.5 cm. in diameter and 30 cm. high, immersed in a thermostat maintained at i.0.1' C. This reservoir with auxiliary apparatus is illustrated inFigure 1. A known mixture ofcarbon dioxide, oxygen, and nitrogen to supply the desired gas concentrations was allowed to flow into
distilled water contained in the reservoir in which the test specimens of steel were totally immersed. Each gas was obtained from a commercial cylinder and was allowed t o flow in the desired amount under pressure t o a common cylinder. The gases were then thoroughly mixed by tilting the cylinder, into which had previously been placed a number of iron bolts. The flow of mixed gases was regulated a t 125 cc. per minute with the aid of a reregulator, and flowmeter. The passage ducing valve, of gases throug the solution permitted both stirring and maintenance of gas concentrations in the water. The advantage of this over previous arrangements was that, no matter how rapid the consumption of gases by corrosion processes, the aqueous concentrations were maintained constant with time. A sintered glass plug through which the gases entered the water ensured rapid absorption, and a chimney located in the center of the jar directed the flow of gas bubbles. To effect closer correspondence of temperature within the reservoir with that of the thermostat, the gas mixture waa humidified by first passing through separate containers of distilled water located in the thermostat. This was to avoid cooling by evaporation in the liquid producing corrosion. A further reason was that the dissolved gas concentration depended on the water vapor content of the gas, and best control could be obtained by using gas saturated with water a t the test temperature. The temperature was measured by a thermometer immersed in the reservoir. Three steel specimens, 6.3 X 2.5 X 0.318 cm., with a hole of 0.318-cm. diameter at one end, were SUR ended from the coyer plate by means of small glass hooks. T i e specimens were immersed 5.1 om. from the axis of the reservoir. The average weight loss of the three specimens served as a measure of corrosion for a specific gas concentration. Separate gas mixtures were prepared for each test to cover the desired range of concentrations. In preparing specimens for the tests, the adherent mill scale on the steel was removed by abrasion using No. 60 followed by No. 120 emery cloth. The specimens were finally degreased with benzene and boiling acetone. After corrosion the specimens were removed from the test solution and were freed of corrosion products by immersion in boiling caustic soda solution containing zinc powder. This treatment was found t o have no effect on the weight of uncorroded or clean s ecimens. The specimens were then again cleaned in organic sogents before weighing. The steel of the test specimens was the same as that commonly used for steam return ipes conforming to specification S2 of the A. S. M. E. Power $)oiler Code. The analysis given by the Carnegie-Illinois Steel Corporation, which kindly supplied the steel, was as follows: carbon, 0.15; manganese, 0.49; phosphorus, 0.013; and sulfur, 0.031 per cent.
?
oxygen concentration and pH. In these investigations the main emphasis was on pH, and the carbon dioxide concentration was not measured. The results indicate, however, that total acidity rather than p H of the corroding medium is better correlated with corrosion losses of iron. A number of field investigations have been made in which the National District Heating Association (N. D. H. A.) tester was used. This tester, consisting of three helical steel coils, was mounted axially in some part of the condensate system, which was always maintained full of condensate. The weight loss of these specimens was determined after a definite period of exposure ( I , IS). Walker showed that, although under actual operating conditions oxygen in the condensate (0 to 0.30 cc. per liter) is a factor in heating system corrosion, it has a minor effect compared with amounb of carbon dioxide present (1 to 14 cc. per liter). He also emphasized t h a t velocity of flow was an important variable. Seeber and Holley (9) inserted testers in lines a t different temperatures and found a marked increase in corrosion with increase in temperature. The above data have mainly qualitative significance because of the impossibility in systems of this kind of controlling the several factors responsible for corrosion. Finnegan, Corey, and Jacobus (3) conducted a series of experiments covering the effect of dissolved carbon dioxide and oxygen on the corrosion of mild steel a t 60' C. They corroded helical coils of steel wire identical with the N. D. H. A. tester, mounted in series in a glass tube through which the corroding solution was circulated at 2.5 cm. per minute by thermal convection. They reported that corrosion increased with both carbon dioxide concentration and oxygen concentration and that oxygen was approximately fourteen times as corrosive as carbon dioxide a t the same molar concentration.
Apparatus
749
GAS SUPPCY CYLINDER
VALVE CONSTANT PRESSURE OUTLET
SATUPATORS
1
THERMOSTAT
FIG. 1.
APPARATUS
,THERMOMETER
CHIMNEY
SPECIMEN
750
INDUSTRIAL AND ENGINEERING CHEMISTRY
TABLEI. EFFECT OF GASRATE,DISSOLVED GASES,AND TEMPERATURE" Run
N0.b
Hours of Test
C
XO
43 44 50 63 22 23 24 25 26 65 27 28 29 30 45 46 47 48
5.00 5.03 4.98 4.62 5.06 5.35 5.26 5.02 5.22 4.83 5.00 5.12 4.90 4.93 5.00 5.00 4.95 4.97
470 125 70 30 125 125 125 125 125 125 125 125 125 125 125 125 125 125
0.66
53 55 56 54 57 58 59
4.98 4.98 5.00 4.98 5.00 4.98 5.00 5.00 4.92
125 125 125 125 125 125 125 lZ5 125
1.09 1.03 1.08 0.69
0.71 0.69
..
1.40 1.26 1.50 0.00 0.00 0.00
0.55 0.61 0.67 0.79 4.10 3.82 3.91 3.80
0.61
0.62 0.00 O.O0 0.00
PH Previous t o Test
PH during Test
P
a. d.
A. D.
2"
Temperature 60' C. 4.9 5.1s 5.62 5.52 5.6 5.17 5.51 4.9 5.13 5.35 4.6 5.20 7.2 5.09 5.51 15.7 4.94 5.51 0.5 6.02 6.16 28.0 4.87 5.84 14.2 5.14 6.04 7.2 0.0 18.8 4.92 5.48 5.70 8.1 5.17 23.4 4.84 5.42 0.4 6.78 7.6 0.3 6.9 6.9 26.1 4.75 5.18 8.2 5.00 5.38 5.36 15.7 4.86
17.7 13.7 13.3 10.6 21.1 28.8 11.9 17.0 10.3 0.9 21.3 15.7 23.4 4.7 31.4 48.2 36.4 40.2
1.7 0.7 0.7 0.4 1.4 0.8 1.7 0.5 0.4 0.1 0.5 0.1 1.0 0.4 1.0 1.6 2.9 0.4
9.5 5.4 5.2 3.8 6.8 2.8 14.4 3.1 4.3 14.0 2.3 0.8 4.1 7.8 3.1 3.3 7.9 1.0
18.1 12.9 13.4 11.0 20.6 29.5 10.8 17.0 10.3 0.9 22.8 16.3 23.4 4.0 30.0 48.7 36.4 40.9
Temperature 90" C. 0.1 6.5 6.6 11.3 4.83 5.08 5.20 5.7 5.06 0.2 6.2 6.2 12.2 4.88 5.06 5.18 5.7 5.07 14.8 4.81 5.66 70 . 9 0 4.98 5:!6
19.9 46.4 37.8 8.5 45.3 33.0 28.4 ::!l
1.0 2.5 3.8 1.0 1.9 1.7 0.7
5.1 5.3 7.5 12.0 4.2 5.2 2.6
19.7 47.1 37.7 7.8 45.8 33.4 28.4
XC
..
Vol. 34, No. 6
salts, however, present niainly as the bicarbonate, can be detected qualitatively in the solution after corrosion has begun. They alter the pH of the carbon dioxide solutions to more alkaline values and increase the apparent solubility of carbon dioxide. When the rate of oxidation of ferrous ions in the solution by dissolved oxygen equals their rate of formation by corrosion, a steady state prevails, and p H and gas concentrations attain constant values. In addition, in order to transfer oxygen from the gas to the solution a t a rate equal to the rate of consumption by corrosion, an appreciable driving force is needed. Hence the oxygen concentration of the liquid drops, when corrosion starts, from the saturation value to a lower value which is determined by the laws of diffusion. The oxygen concentration then stays constant. It is for these reasons tha,t auxiliary specimens are necessary before the corrosion test period is started.
Analytical Methods
Dissolved carbon dioxide was determined by precipitationwithanexcessof barium hydroxide and titration of the excess with hydrochloric a T h e nomenclature is as follows: acid, using phenolphthalizin as indicator. Only C = gas rate, cc /minute (at 20' C 760 mm H g ) dissolved 0; concentration, cc.7liter (at Oo. C.,*76() mm. Hg). Xo ferrous ion in sufficient' concentration total dissolved Con concentration, including ionic and undissociated forms, cc,/ Xo liter (at 0' C., 760 mm. H@. possibly interfere, but calculation showed that a t = observed rate of corrosion, inches/year X 1000. P the p H of the titrated solution, the interference a. d. = average deviation from mean, inches/year X 1000. A. D. = average deviation from mean, per cent. The fact that the end point was was = r a t e of corrosion, adjusted from run concentrations t o average concentration for PI group of Funs, inches/year X 1000. permanent supports this oonclusion. The anab Values of PI adjusted in the following manner: runs 43, 44, 50, 63 adjusted t o 0.69 cc. 1YtiCal method was checked against water of o2and 5.0 c ~ co2 . per liter: 22-24 t o 1.34 CC. 02:27-30 and 44 t o 0.67 00. ol; 45-4s to 3.91 cc 0%:53, 55, 56 to 1.07 00. 0 2 ; 54. 57, 58 to 0 64 cc. Or per liter. known carbon dioxide concentration and was found to give results within 0.4 cc. carbon dioxide per liter. Dissolved oxygen was determined by the wellduction of the specimens, for reasons noted later, and reached known Winkler method, using soluble starch as indicator. Fera constant value within 1to 4 hours. ( c ) When the pH was ricioninterfereswith this analysis byoxidizingpotassiumiodide constant, the auxiliary specimens were replaced with three to iodine and thus giving high results for oxygen. Since the weighed test specimens. Corrosion then proceeded a t consolution is strongly acid a t the end point of the determination, stant p H and constant oxygen and carbon dioxide concenferric ion is soluble, but fortunately the interfering reaction trations for approximately 5 hours. From the subsequent is a slow one, and it was possible to correct for it by a blank average weight loss and dimensions of the specimens, a coras was done by Urbain and Miller ( I d ) . The method was rosion rate was calculated. Samples of solution were withchecked by analyzing air-saturated water containing various drawn a t intervals during each of the three parts of the run concentrations of ferric ammonium sulfate. The end point and cooled to 25" C. for determination of the pH. Imfaded slowly, but with the blank correction, the analysis was found to be high only to the extent of 0.03 and 0.09 cc. of mediately before the end of the test period, two samples were withdrawn for determination of the dissolved oxygen and oxygen per liter a t 10 and 30 p. p. m. of iron. Since the iron content of the test solutions was less than 10 p. p. m., this carbon dioxide concentrations, care being taken a t all times to avoid exposure of the sample to the atmosphere. difference is not significant. The rate of consumption of oxygen by dissolved ferrous salts between sampling and The increase in pH upon introducing the steel specimens is analysis was sufficiently slow so as to be negligible. explained as follows: The corrosion reaction is essentially the reaction between carbonic acid and iron to form ferrous The pH of the solutions was measured in a MacInnes and bicarbonate and ferrous hydroxide. Calculations making Belcher type glass electrode (7), which prevents exposure of the sample to the atmosphere. The glass electrode was use of observed pH values and the known dissociation constants for dissolved carbon dioxide show that only bicarbonchecked against buffers of 1.99, 3.98, and 6.99 pH, and was a t e concentration is important and carbonate concentration found to read correctly within 0.05 pH unit. negligible. Oxygen has the effect of increasing the rate of corrosion by its depolarizing action a t the metal surface but, Effect of Gas Flow Rate in addition. of continuallv oxidizing the small concentration Four runs were made (Figure 2) in which the gas rate was of ferrous hydroxide to flrric hydrokde in the main body of varied from 30 t o 470 cc. per minute, all other factors being the corroding liquid. The latter is a relatively slow reaction. kept constant. Corrosion increased slightly with gas flow, Calculations show that the resulting orange colored precipithe increase amounting to 0.001 inch penetration per year tate of FezOs.nHpO suspended in the water solution is so when the gas flow increased from 100 t o 200 cc. per minute. insoluble that its saturated solution has no effect on the gas The observed trend is probably causedl by increased stirring concentrations or pH of the aqueous attacking medium and action a t the higher velocities. In all the corrosion tests the hence has no effect on the corrosion rate. Soluble ferrous 60
64
-
E
..
A:; li:E
751
INDUSTRIAL AND ENGINEERING CHEMISTRY
June, 1942
gas rate was held at 125 * 5 cc. per minute. At this rate the ,velocity of the solution past the specimens, estimated from the motion of suspended fibers in the liquid, was about 45 cm. per minute, or 0.025 foot per second.
0
100
300
200
400
GAS FLOW CC, /MIW.
Effect of Duration of Test There is always a question as to whether a corrosion test period is sufficiently prolonged so that calculated rates can apply to longer periods. If there is a transient initial rate of corrosion sufficiently different from a final rate, or if the rate decreases or increases with time, any extrapolation of weight loss is subject to error. Evidence is presented in Figure 3 that up to 8 hours, corrosion is linear with time. Although a
I
CORROSION OF MILD STEEL
II
I
004 0.03
The average deviation from the mean of the three specimens for each run averaged 0.0015 inch per year or about 6 per cent of the average corrosion rates. Corrections introduced because of slight variations in gas analysis from one run to another were less than this mean deviation. Corrosion is tabulated in terms of inches penetration per year X 1000 for convenience in applying the data as well as in comparing them with other corrosion data. To express the results (inches per year X 1000) in milligrams per square centimeter per day, the values should be multiplied by the factor 0.055. Both dissolved oxygen and carbon dioxide acting individually or simultaneously increase corrosion of steel in water. The relative part each gas plays can be conveniently calculated by making use of a term called the "relative corrosiveness", which can be calculated in the following way: The value of oxygen corrosion divided by the oxygen concentration is computed. This is essentially equal to the slope of curves A in Figure 5 at 60" C. (left) or a t 90" C. (right) and is constant over the oxygen concentrations recorded a t any single temperature. If this value is divided by the ratio of carbon dioxide corrosion to carbon dioxide concentration, an estimation of the relative corrosiveness ia obtained. Since the zero oxygen curves of Figure 4 are not linear, the relative corrosiveness would be expected to be a function of the carbon dioxide concentration, as is the case. Values are listed in Table 11. A concentration of 2.5 cc. oxygen per liter a t 80" C., for example, produces six times the corrosion caused by an equal concentration of carbon dioxide. With increasing carbon dioxide concentration, carbon dioxide is relatively less effective in corroding steel. The relative corrosiveness is unaffected by oxygen concentration and decreases with increase in temperature.
TABLE11. RELATIVE CORROSIVENESS OF OXYQENCOMPARED TO CARBON DIOXIDE AT EQUAL CONCENTRATIONS
0.0 2
'
Concn. of COa or 01,oc./l. Relative corrosiveness 60'
C.
900 c.
2.6
6 6
4 7
6
10
20
8 7
10
..
01)I
0 0
2
6 WRS DURATION 4
8
porous layer of iron corrosion products forms on some of the specimens, particularly those corroded in solutions high in oxygen and low in carbon dioxide concentration, this layer is not protective in the usual sense of the term, and whatever resistance it offers to diffusional processes appears to remain constant with time. Only a film whose partial protection varies with time can alter the linear relation shown in Figure 3, and the presence of such a film was not indicated. It is probable, therefore, that the linear relation does not break down beyond the 8-hour period.
Corrosion Data The corrosion of steel in solutions of fixed oxygen concentration with varying carbon dioxide concentrations are summarized a t 80" and 90" C. in Figure 4. Each point is the average of three check specimens run simultaneously. Figure 5, making use of intercepts in Figure 4,shows the effects on corrosion of varying oxygen concentrations for constant carbon dioxide concentration. Detailed data are listed in Table I. Corrosion was uniform, and no pitting of any kind was observed.
It is also of interest to examine the additivity of the corrosion produced by each gas. The relations between the corrosion caused by carbon dioxide and oxygen separately and the two gases acting together are given in Table 111. The fifth column shows, for instance, that, in a solution containing 10 cc. carbon dioxide and 0.67 cc. oxygen per liter, corrosion amounts to 0.0169 of an inch per year. I n a. solution with 10 cc. carbon dioxide alone per liter, the corrosion rate is 0,0081 inch; and when the solution contains 0.67 cc. oxygen TABLE111. ADDITIVITY OF CORROSION^ CARBON DIOXIDE Concn., Co./L. COI 01 P C O , P O a
BY
P ( c o ~ + o ~Pcoi+Por At 60" C. 4.0 10.9 12.1 10 0.67 8.1 10 1.34 8.1 10.0 23.0 18.1 10 3.91 3S.O 8.1 27.6 36.0 20 0.07 13.3 4.0 22.8 17.3 20 1.34 13.3 10.0 31.1 23.3 44.2 20 3.91 13.3 27.6 40.8 At 90° C. 6 0.64 13.8 7.9 30.0 21.7 6 1.07 13.8 18.9 '36.0 32.7 10 0.64 21.0 7.9 43.0 29.6 10 1.07 21.0 18.9 46.8 40.6 0 P aorrosion in inahea per year X 1000.
-
OXYQEN AND q c o , +Od Pco*+Po~ 1.40 1.30 1.07 1.32 1.33 1.08 1.38 1.10 1.46 1.13
INDUSTRIAL AND ENGINEERING CHEMISTRY
752
This diffusion rate is proportional t o the oxygen concentration in the attacking medium, which results in the corrosion being proportional to the oxygen concentration. In corrosion by dissolved carbon dioxide, the equivalent of reaction 2 is the following reaction, 4, which applies to the case (pertinent to this investigation) in which carbon dioxide reacts with water t o form bicarbonate but no carbonate ion: 2C02 2H20 2~ +-2HCOaHP (4) It is seen that one mole of carbon dioxide takes up only one quarter the number of electrons compamd with one mole of oxygen, so that the corrosion accounted for by equal molar concentrations of the IO I5 two gases may be expeoted t o be in the ratio of 1to 4. The relative corrosiveness of oxygen compared with carbon dioxide as defined previously would correspondingly be 4. This assumes that for the same concentration gradient: (a) oxygen and carbon dioxide diffuse through the surface liquid film a t the same rate and ( b ) , having reached the metal surface, react with iron according to reactions 3 and 4 a t such rapid rates that they do not influence the rate of corrosion, and the concentrations at the metal surface are zero. The first assumption has been verified experimentally (6),but the second can be shown to fall short of the actual situation. I n corrosion by hydrogen evolution, an appreciable acidity must be maintained at the metal surface before hydrogen is evolved a t any appreciable rate. This was suggested by Whitman, Russell, and Altieri (14) and would mean, in this case, that a finite carbon dioxide concentration must exist a t the metal surface if hydrogen evollution is to take place. This finite concentration may be looked upon as the concentration necessary to overcome hydrogen overvoltage. Hence, a given concentration of carbon dioxide in the main body of the solution will cause less diffusion through the liquid film than will the same molecular concentration of oxygen, because only a fraction of the carbon dlioxide concentration is available as a driving force, whereas the whole of the oxygen concentration is available for this purpose. It follows, then, that in equal molar concentrations oxygen should be more than four times as active as carbon dioxide. This agrees
+
0
5
15
IO
20
25
30 0
5
C02 CONCENTRATION (C CJLITER)
per liter and no carbon dioxide, the corrosion rate is 0.004 inch. The sum of these two is 0.0121 and is given in the sixth column. The seventh column is merely the fifth divided by the sixth and indicates that the presence of one gas in each case increases the corrosiveness of the other. The effect is most pronounced a t low oxygen concentrations where the increase in corrosion a t 0.67 cc. oxygen and 10 cc. carbon dioxide per liter is as much as 40 per cent. The temperature coefficient of corrosion is relatively small. Expressed as the factor by which corrosion is multiplied for a temperature rise of 30" from 60" to 90" C., the results are given in Table IV.
TABLEIV. TEMPERATURE COEFFICIENT OF CORROSION, FOR 30" C. INCREASE (60-90' C.) Cc. O,/Liter
Cc. COdLiter
Nil Nil
1.00 0.50 Nil Nil 0.65 0.65
15.0 10.0
10.0 5.0
Temp. Coefficient 2.1 2.2 2.6 2.6 2.5 2.5
The temperature coefficient is larger for carbon dioxide corrosion than for oxygen corrosion, and is intermediate in value when the two gases are present. Seeber and Holley (9) in agreement found a temperature coefficient of 2.4 for a 30" rise for the corrosion of mild steel in carbon dioxide-oxygen solutions. Discussion of Results It is known that the corrosion of iron is accounted for by the anodic reaction, Fe + Fe++ + 2e (1) accompanied by the cathodic reaction, 2H+ + 2e + Hz or if oxygen is present, depolarization takes place according to the reaction, 1/*02 Ha0 2~ -+ 20H(3)
+
+
The latter two reactions control the rate a t which iron corrodes. Reaction 3 is so rapid that the oxygen concentration a t the metal surface approaches zero, and the rate a t which depolarization proceeds depends on the rate of diffusion of oxygen through the stagnant film a t the surface of the metal.
Vol. 34, No. 6
+
+
June, 1942
INDUSTRIAL AND ENGINEERING CHEMISTRY
with the facts. The relative corrosiveness of oxygen compared with carbon dioxide, as Table I1 shows, is greater than 4 and increases with carbon dioxide concentration. The increase with carbon dioxide concentration indicates that, as the concentration of this gas increases at the metal surface through concentration increase in the body of solution, the corrosion is not linearly proportional but falls off. The decrease in relative corrosiveness with increase in temperature indicates that a given surface concentration of carbon dioxide is more corrosive, the higher the temperature. This follows from the fact that diffusion of carbon dioxide and oxygen across a given barrier is essentially the same at any one temperature. Hence, any change in relative corrosiveness with temperature must be caused by change in reactivity of one of the gases. Metal surface oxygen Concentration is practically zero; therefore the surface concentration of carbon dioxide must become more reactive, as would be reasonably expected. Another factor enters into the expected value for the relative corrosiveness which arises from the fact that specimens corroded by oxygen solutions, unlike those corroded by carbon dioxide solutions, are covered with a black corrosion product layer consisting largely of hydrated Fea04. A resistance is present, therefore, in oxygen corrosion which is not present in carbon dioxide corrosion. This acts to make the relative corrosiveness of oxygen less than it otherwise would be, opposite to the effect of finite carbon dioxide concentration a t the metal surface discussed above. The presence of this layer, depending on the oxygen concentration of the attacking medium, readily explains the trend of additivity of the carbon dioxide and oxygen corrosion (Table 111). Thus, when oxygen concentration is high, the adherent film which is formed is essentially uninfluenced by the carbon dioxide concentrations. The two gases acting together produce an effect only slightly greater than the sum of corrosion by each gas individually. When oxygen concentration is low, the carbon dioxide concentration prevents formation of the adherent film built up in oxygen solutions, and oxygen can then more readily diffuse to the metal surface so that simultaneous corrosion exceeds the additive effects of corrosion by the two gases acting singly. Cox and Roethelli (2) found that, up to approximately 6 cc. oxygen per liter, corrosion at room temperature is proportional to oxygen concentration, whereas above this value corrosion drops off from the extrapolated linear value. They conclude that the black, granular, corrosion-product film formed below 6 cc. oxygen per liter is essentially nonprotective compared to the reddish gelatinous film formed at higher oxygen concentrations. This conclusion is consistent with the barrier postulated above for the black oxide film because this nonprotective. film nevertheless has a resistance to diffusion which is constant and independent of the extent to which corrosion has proceeded, and is also independent of the oxygen concentration within a deiinite range. Evidence for this is given by the proportionality between oxygen concentration and corrosion. The temperature coefficient of corrosion as summarized in Table IV bears out the mechanism of corrosion postulated above. With increase of temperature, the rates of diffusion of carbon dioxide and oxygen through the interface film increase. I n addition, however, the acidity or finite carbon dioxide concentration a t the metal surface needed for hydrogen evolution is known to decrease, making carbon dioxide more effective. Since this last effect is absent for oxygen, the concentration being effectively zero a t the metal surface, carbon dioxide becomes relatively more active than does oxygen for a given temperature increase. Sherwood and Holloway (IO) found that the temperature coefficient for transfer of carbon dioxide and oxygen through the liquid
753
film in gas absorption in a packed tower is 2.1 for a 30" C. temperature interval. The average temperature coefficient for diffusion of sugars in aqueous solution for the same temperature interval is given as 2.0 to 2.1 (11). Both these figures are in excellent agreement with the temperature coefficients reported in Table IV. There is good evidence, therefore, that the corrosion of iron or steel by water containing carbpn dioxide and oxygen is largely controlled by the diffusion of reactants to the metal surface, modified in the case of carbon dioxide by the retarding effect of the finite surface concentration necessary for hydrogen evolution.
Comparison of Data with Those of Finnegan e t al. Finnegan, Corey, and Jacobus (a), using the N. D. H. A. corrosion tester ,coils, obtained data at one temperature, 60" C., which can be compared with the above results. Their data for corrosion in mixtures of dissolved carbon dioxide and oxygen are at first glance appreciably lower than the data presented in the left-hand graphs of Figures 4 and 5. It is apparent, however, that the discrepancies are most pronounced a t high ratios of carbon dioxide to oxygen concentration, whereas for low ratios the agreement between the two sets of data is excellent. This is illustrated in Table V in which a few of the results are compared according to the ratio of carbon dioxide to oxygen concentration in the corroding medium. TABLEV. DEPENDENCE OF CARBON DIOXIDE/~XYQEN RATIO ON AQREEMENT WITH DATAOF FINNEQAN, COREY, AND JACOBIJS Conon., Co./L. 01 COe
Ratio,COa/Oa
Corrosion, In./Year X 1000 (60°C.) Finnegan et al. Preaent work
At high ratios of carbon dioxide to oxygen, corrosion as in the case of nonoxidizing acids is accompanied by hydrogen evolution, and the rate is governed by the acid concentration of the attacking medium in addition to the various factom related to the composition and physical structure of the steel. At low ratios of carbon dioxide to oxygen almost no hydrogen is evolved, and the rate of depolariaation by oxygen controls the rate of corrosion. For the latter case, composition and structural differences in the steel bear no relation to the corrosion rate, and all low-alloy compositions corrode equally at rates proportional to the oxygen transport to the metal surface. The lack of correspondence in Table V at high ratios of carbon dioxide to oxygen is probably not explained on the basis of experimental .variations such as differences in the rate of flow of solution past the steel test specimens. If difference in velocity of flow were a significant factor, one would not expect the excellent correspondence for oxygen corrosion. I n addition, Whitman, Russell, Welling, and Cochrane (16),in experiments on the effect of velocity on the corrosion of steel in dilute sulfuric acid, showed that velocity differences of the order applying to the above experiments (0.025 foot per second) cause much less change in corrosion than Table V would indicate, and that at higher acid concentrations the effect of low velocity reverses sign. We are led to the explanation that the steel coils used by Finnegan, Corey, and Jacobus were not so corrodible in carbon dioxide or dilute acid solutions as the mild steel plate used in the present experiments. This, in turn, is explained by differences in their composition and structure. The effect of
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INDUSTRIAL AND ENGINEERING CHEMISTRY
composition and structure also indicates a reason for the fact that carbon dioxide is relatively more corrosive with respect to an equal molar solution of oxygen in our measurements (Table 11) than the 1 to 14 ratio reported by Finnegan et al. We conclude, therefore, that steel equipment exposed to water relatively low in oxygen but relatively high in carbon dioxide will corrode a t a rate which, to some extent, is a function of the impuriti&, heat treatment, and mechanical history of the steel.
Summary Corrosion of a typical steam condensate line steel containing 0.15 per cent carbon by dissolved carbon dioxide up to 25 cc. per liter and by dissolved oxygen up to 4 cc. per liter was investigated a t 60" and 90" C. Corrosion was measured by loss in weight. The corrosion tests were conducted in water a t a specified temperature through which a gas mixture of nitrogen, carbon dioxide, and oxygen of controlled composition was bubbled throughout the duration of the test. The gas stream maintained constant the concentration of dissolved gases, determined by analysis of the solution, and in addition served to circulate the solution past the totally immersed metal specimens. Both dissolved carbon dioxide and oxygen increase the rate of corrosion. Solutions a t 60" C. containing oxygen cause six to ten times as much corrosion as carbon dioxide in the same molar concentrations, the comparative effect of carbon dioxide being greater a t low gas concentrations. At 90' C. the increase in corrosiveness of carbon dioxide is relatively greater than the corresponding increase for oxygen by a factor of approximately 20 per cent. If carbon dioxide as well as oxygen concentration were zero a t the corroding steel surface, it is expected that oxygen would be four times as corrosive as carbon dioxide. The higher observed values lead to the conclusion in agreement with previous observations that when corrosion is by hydrogen evolution, a finite acid concentration (in this case carbon dioxide concentration) must be maintained a t the metal surface. The subsequent reduced concentration gradient accounts for less transport of carbon dioxide to the metal surface than is the case for an equal concentration of oxygen. Corrosion by a solution containing both carbon dioxide and oxygen is 10 to 40 per cent higher than the sum of the corrosion by the dissolved gases acting individually. The increase in corrosion observed is greater a t low ratios of oxygen to carbon dioxide concentration. This behavior appears to be related to the nature of the corrosion products. At high ratios of oxygen t o carbon dioxide concentration, the surface layer of hydrated iron oxide observed in oxygen but not in carbon dioxide corrosion is also observed for the gas mixture. This layer acts as a barrier to diffusion of both gases so that its absence a t low ratios increases the relative amount of oxygen diffusing to the surface, and the corresponding total cor-
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rosion. The barrier remains constant with time, which explains the linearity of corrosion with time. Raising the temperature from 60" to 90" C. multiplies corrosion by dissolved oxygen by a factor of 2.1 and corrosion by carbon dioxide by a factor of 2.6. Mixtures of the two gases have an intermediate temperature coefficient. The larger coefficient for carbon dioxide is expected from the relatively greater activity of the surface carbon dioxide concentration a t high temperatures. Independent measurements of the temperature coefficient for diffusion of carbon dioxide, oxygen, and sugars give values of 2.1 for 30" C. rise, in good agreement with the corrosion coefficients reported above. This is additional evidence that the corrosion rate is largely controlled by diffusion of the gases to the metal surface. Comparison of the data with those of Finnegan, Corey, and Jacobus (3) shows good agreement when corrosion is caused largely by oxygen, but discrepancies amounting to a ratio greater than 2.5 occur for corrosion largely by carbon dioxide. The conclusion is that the steel used by Finnegan, Corey, and Jacobus corroded less by dissolved carbon dioxide than the mild steel used in obtaining the present results. Unlike the case for dissolved oxygen, it appears that, when corrosion is largely by dissolved carbon dioxide, the structure and composition of the steel are factors in determining its life.
Acknowledgment The authors take pleasure in expressing their gratitude to W. G. Whitman for the benefit of his valuable criticism and many suggestions.
Literature Cited Collins, L., and Henderson, E., Heating, Piping A i r Conditioning, 11, 539 (1939). Cox, G., and Roethelli, B., IND. ENG.CIIEM.,23, 1012 (1931). Finnegan, T., Corey, R., and Jaoobus, D., Ibid., 27, 774 (1935). Girard, R., Compt. rend., 181, 215 (1925). Groesbeck, E., and Waldron, J., Proc. Am. SOC.Testing M a terials. 31. 279 (1931). Hutchinson,' M., and Sherwood, T., IND. ENQ.CHEM.,29, 836 (1937).
MacInnes, D., and Belcher, D., IRD.EINQ.CHEM.,ANAL.ED., 5, 199 (1933). Rhodes, F.. and Clark, J., IND. ENQ.CFImr., 28, 1078 (1936). Seeber, R., and Holley, M., Heating, Piping A i r Conditioning, 9, 387 (1937).
Sherwood, T., and Holloway, F., Trans. Am. Inst. Chem. Engrs., 36, 39, 181 (1940). Taylor, H. S., "Treatise on Physical Chemistry", Vol. 11, p. 1025, New York, D. Van Nostrand Co., 1931. Urbain, O.,and Miller, J., J . Am. Water Works Assoc., 22, 1261 (1930). Walker, J., Heating & Ventilating, 30, 28 (1933); Proc. Am. SOC.Testing Materials, Preprint 104 (1940). Whitman, W., Russell, R., and Altieri. K., IND. ESQ. CHEM.,16, 665 (1924). Whitman, W., Russell, R.. Welling, C., and Cochrane, J.. I b i d . , 15, 672 (1923).
