I N D UXTRIAL A N D ENGINEERING CHEMISTRY
July, 1924 ~-
665
Effect of Hydrogen-Ion Concentration on t h e Submerged Corrosion of Steel’ By G . W. Whitman, R. P. Russell, and V. J. Altieri MASSACHUSETTS INSTITUTE OF TECHNOLOGY, CAMBRIDGE, MASS.
T h e m a i n factors in the corrosion of steel submerged in natural waters, dilute alkalies, and dilute acids are the protectiveness of f i l m s of corrosion products and the rate of oxygen diffusion. T h e degree of f i l m protection i s largely determined by the p H of the solution direct1.y adjacent to the corroding metal, since this pH a#ects the solubility of ferrous hydroxide and hence the physical condition of the precipitated film. I n the alkaline region increased p H makes the liquid next to the metal mare alkaline and causes the formation of a more protective f i l m , although the initial rate of corrosion i s a s great a s the rate in natural uaters. T h i s f i l m requires considerable time to develop its full protectiveness, and it i s not readily destroyed when the alkalinity of the solution i s reduced. T h e determining factor, therefore, i s increased f i l m protection rather than insuficient hydrogeni o n concentration. T h e results of L y o n s and others, who obtained a m a x i m u m rate of corrosion in beaker tests at a “critical” alkalinity, are invalidated by failure to protect the solution f r o m carbon dioxide in the atmosphere. Over the natural water field the f i l m protection in a given water is constant, because the solution next to the metal is maintained at a constant pH of 9.5 by the solubility of ferrous hydroxide. T h e
corrosion is therefore independent of the hydrogen-ion concentration in the m a i n solution and i s determined by the rate of digusion of dissolved oxygen to the corroding solid. O n the acid side of the natural water region hydrogen gas evolution starts when the alkalinity of the liquid next to the metal is neutralized by the acid in the m a i n solution. T h e chief factors in determining the point at which gas evolution starts are ( I ) the rate of mixing of m a i n solutiqn with surface liquid, ( 2 ) the total acidity i n the m a i n solution, and (3) the rate of production of alkali in the surface liquid by the process of corrosion. Because of the second factor a weak acid will cause hydrogen gas evolution to commence at a lower concentration of H + than i s required with a strong acid. Because of the third factor a decrease in oxygen concentration (which will decrease corrosion and ( O H ) - ion production) will permit readier neutralization, and hydrogen gas will be evolved at lower acidities. T h e rate of oxygen reaction in the acid region is greatly increased by the facts that gas evolution causes turbulence in the liquid adjacent to the metal surface and thereby increases the rate of oxygen d i s u s i o n and that the protective f i l m i s absent. T h e results of this work are i n accordance with a general theory of electrochemical corrosion.
HE mechanism of the corrosion of steel under water has been variously explained, but a t the present time most investigators agree on the electrochemical theory. Wilson2 considered the whole question of submerged corrosion under natural water and formulated his theory on the basis of hydrogen-ion concentration. He divided submerged corrosion into three fields-alkaline, natural water, and acidand advanced his explanations of the effect of variables such as pH, temperature, velocity, oxygen concentration, and overvol1,age on the rate of corrosion. Many of Wilson’s postulat,es were admittedly unsupported by experimental evidence, but his paper must be regarded as a most valuable contribution to the theory of corrosion. The method of attack in the present paper is based on the same division of submerged corrosion into alkaline, natural water, and acid fields. The authors, however, disagree with many of Wilson’s explanations, and advance the following concepts of the processes. THEORY
reductions are possible under very special conditions, such as the deposition of copper, the solution of dissolved chlorine to form chloride ions, and the reduction of nitric acid. The cathode reaction for hydrogen deposition may be expressed as
T
The initial reaction in the submerged corrosion of iron is electrochemical and operates through a corrosion cell. A t the anodic electrode of this cell the corroding metal sends ferrous .ions into solution according to the reaction Fe
+ 2@
= Fe++
(1)
A corresponding reduction occurs a t the cathode area-usually the deposition of hydrogen ions as atomic hydrogen (and subsequent liberation as hydrogen gas), or the reaction of dissolved oxygen to form hydroxide ions. While t.hese two are the most common cathode reactions, a number of other 1 Received April 1, 1924. Presented before t h e Division of Industrial and Engineering Chemistry a t t h e 67th Meeting of t h e American Chemical Society, Washington, D. C., April 21 t o 26, 1924. 2 THIS JOURNAL, 16, 127 (1923).
2 H f = 2H
+ 2@
(2)
and for oxygen solution as
+
‘/io2t H20 = 2(OH)2@ (3) It should be noted that the oxygen effect has usually been expressed as an oxidation (depolarization) of the atomic hydrogen formed by Reaction 2 ; thus Wilson gives the process in the following two reactions :
+
2 H + = 2EI 2@ 2H ‘/to2= HzO
+
If these reactions are written in ful!, showing the hydrogen ion coming from the dissociation of water, they would have the form 2H
++ ++
2Hz0 = 2 H f f 2(OH)2H 2@ 2(OH)2@ 2(OH)‘/202 = € 1 2 0 2@ 2(OH)-
+
+
+
By adding the two equations and canceling 1 mol of water, the reactions reduce to the form HzO
+
‘/‘202
= 2(OH)-
+ 2@
which is identical with Reaction 3, given above. It is thus evident that the net result of Wilson’s series reactions is the same as is obtained when we consider oxygen to react directly a1 the cathode. The authors themselves have used Wilson’s equations in earlier papers but now believe t h a t Reaction 3 expresses the facts in a more direct and convenient form, and also allows the direct computation of the electrode potential a t the cathode. Furthermore, the corrosion of a metal like copper, which is below hydrogen in the electromotive series, can be more readily visualized on the basis of a direct oxygen electrode reaction than by considering that hydrogen must first be deposited.
The extent of corrosion is chiefly determined by two factors: (1) the rate of the cathode reduction (which in the
INDUSTRIAL A N D ENGINEERING CHEMISTRY
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case of oxygen is limited by the rate of oxygen diffusion), and (2) the formation of protective films. Protective films of corrosion products serve to prevent contact between metal and corroding agent, and the degree of their protectiveness varies with the conditions under which the film is formed. In particular, the p H of the solution PH = log(&) against the metal surface affects the character of the film, because the corroding metal first forms ferrous ions which are then precipitated as ferrous hydroxide. Since the character of the film will vary with the solubility and the rapidity with which the hydroxide is precipitated, it will therefore be a function of the pH of the solution at this point. No3F~o7fr~i
Pnt%”(D &%mw ,Cq ABsaefaY.
FIG. A APPARATUS FOR BEAKERTESTS
The process of corrosion results in the formation of ferrous hydroxide against the metal surface. This ferrous hydroxide will tend to maintain a saturated solution in the liquid film a t the metal surface with a normal pH value of 9.5, but dilution and intermixing with the main body of the liquid will tend to change this pH. As a net result, the hydrogen-ion concentration of the liquid film next to the metal (where the actual corrosion is taking place) may be considerably different from the pH of the main solution itself. This fact is of vital importance in explaining the mechanism of underwater corrosion. ALKALINECORROSION (beyond pH = 9.5)-Above a p H of 9.5 increasing the alkalinity of the solution tends to make the liquid film more alkaline. As a result, the solubility of ferrous hydroxide is decreased and the film is precipitated in a more protective form. Corrosion should therefore decrease with increased alkalinity in this range. The progress of corrosion is, of course, dependent upon a supply of dissolved oxygen to react a t the cathode. NATURALWATERCORROSION (pH = approximately 9.5 to 4.5)-1n the natural water region the liquid film tends to hold a pH of 9.5 due to the solubility of ferrous hydroxide. While it is true that intermixture with the main solution would have a tendency to neutralize this alkaline film, the actual acidity of the solution is so very low that neutralization cannot be effective. The process can be pictured as a conflict of two opposing forces-the corrosion reaction constantly forming Fe(OH)s and thus creating alkali to maintain a p H of 9.5, opposed by the process of intermixing which would dilute and neutralize this surface alkalinity. It follows, therefore, that a change in one of these processes which would not affect the other would change the ease with which this liquid film could be neutralized. As the pH of the natural water is decreased, it finally becomes acid enough to neutralize the surface liquor film so that a pH of 9.5 can no longer be maintained against the metal. The hydroxide will then become more soluble, the protective film will be destroyed, and the liquid a t the metal surface will be sufficiently acidic to cause hydrogen gas evolution. This point marks the limit of the so-called “natural water” region and occurs a t a pH on the acid side, which is approximately equivalent to 9.5 on the alkaline side.
Vol. 16, No. 7
From the above it can be seen that the p H of the liquid adjacent to the metal in the natural water region is constant a t 9.5 and is independent of the pH of the solution. It can, therefore, be concluded that the protectiveness of the film will be constant over this region in any given water and that p H value will have no direct effect on the rate of corrosion. The most important variable affecting the corrosion in a given natural water is the rate of oxygen diffusion from the main liquid in to the cathode surface. (Film protectiveness varies with different natural waters, thus varying the rate of oxygen diffusion to an effective cathode area.) Increase in oxygen concentration, in temperature, or in velocity should increase this diffusionand thereby accelerate corrosion. ACIDCORROSION (below a pH of about 4.5)-1n nonoxidizing acids the protective film is absent and hydrogen gas evolution comes into play. The oxygen reaction still accounts for a part of the total corrosion, and its effect is accentuated because the gas evolution disrupts the liquid film and makes diffusion easier. The authors3 have shown that dissolved oxygen may be of major importance in nonoxidizing acids a t high velocities and that the total corrosion is the sum of the corrosion by hydrogen gas evolution and that by oxygen reaction. I n contradistinction to the two preceding fields, the composition of the steel is an important factor in acid corrosion, since overvoltage, which affects gas evolution, is a function of composition.
PREVIOUS WORK Although a large number of investigations have been carried out on the corrosion of iron and steel by varying strengths of alkali, only a few can be considered of importance. Heyn and Bauer4conducted experiments on the action of dilute sodium hydroxide solutions, but their tests were made in open vessels and the absorption of carbon dioxide from the atmosphere converted the hydroxide into carbonate. Similar experiments by Friend6 are open t o the same criticism. Friend carried out other experiments in sealed evacuated vessels, which consequently contained so little dissolved oxygen as to make the corrosion negligible in any case. I,yon,B in a comprehensive investigation, included some work on corrosion by sodium hydroxide solutions, and found that a “critical” concentration of maximum corrosion occurred a t 0.008 N a n d an “upper limit” of no corrosion a t 0.026 N . These experiments were also conducted with free access to the air, and are therefore open to the same objection as the preceding tests. Speller and Texter’ have recently studied the corrosion of iron pipe when exposed to caustic solutions of varying concentration at 150’ F. Their research made a number of important advances over the experiments reported by earlier writers. They found that there was no “critical” concentration of maximum corrosion, but that a t all the alkalinities studied the final corrosion rate was less than in the raw water. The results also indicated that the corrosion rate was dependent upon the type of rust film formed. The data, however, were obtained a t a relatively high temperature and did not include corrosion in the acid ranges.
It was recognized that a quantitative study of the submerged corrosion of steel, which would embrace the acid as well as the alkaline and natural water fields, would be of considerable value in determining the true mechanism of the process, The following experiments, which included a study of the effect of varying temperature, as well as varying hydrogenion concentration, were initiated with this end in view. METHODOF ATTACK The program included determinations of corrosion rate through and slightly beyond the ranges of hydrogen-ion concentration ordinarily encountered in natural waters. It
* THISJ O U R N A L , 4 8 6
7
15, 672 (1923). Mttt. kgl. Matrrialprufungsamt, 26, 2 (1908). “The Corrosion of Iron and Steel,” 1911, p. 123. J . A m . SOC.N a v . Eng., 34, 845 (1912). TKISJOURNAL, 16, 392 (1924).
July, 1924
I N D U S T R I A L A N D ENGINEERING CHEMISTRY
was decided to make runs a t two temperatures (20' and 40" C.) with constant velocity and controlled oxygen concentration, a t pH's ranging from an acidity of about 3 to an alkalinity of about 13.
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wooden tank, and oxygen concentration was maintained by air agitation. I n the prolonged runs more than one tankful of solution was necessary and in the interval during preparation of fresh solution the cell was fed from a barrel containing the same solution. In certain instances it was impossible to keep the cell running continuously over the whole run, and it was left over night, filled with quiet solution. I n a large number of instances, at the completion of a run the cell was left over night with ordinary tap water running through a t constant velocity. I n a series of runs made a t 40" C. the temperature was maintained by blowing in live steam.
BEAKERTESTS
The beaker tests were run in (15-cm.) (6-inch) crystallizing dishes containing 450 cc. of solution. The test pieces (51 x inch) (cold-rolled steel) were 13 X 3.2 mm.) (2 X l / 2 X washed in gasoline to remove oil and grease, slightly polished with emery to give a uniform surface, and immersed in the solutions, resting on U's made of glass rod, thus being kept out of contact with the bottom of the crystallizing dish. The dishes were then placed in a galvanized iron container (Fig. I) which had a water seal to prevent evaporation of solution and infiltration of air around the bottom. Two series of tests a t room temperature were run in parallel, those on one side of the partitioned container being protected against carbon dioxide absorption by soda lime, and the other being un& ~ o o+~k o , ~ o &, -1o ~ o ~ ~ ~ ~ L , I months. protected. The tests ran for 2&
% & L A . ,
'
Two methods of measuring corrosion rate were selected: the first, the decrease in dissolved oxygen method devised by Speller;8 and the second, a modified form of the beaker tests used by previous investigators. The oxygen drop method has the following advantages: ( a ) it is rapid; ( b ) it allows the building of a rust film; (c) it follows the change in corroi:ion rate with time and changing rust film; and ( d ) it allows controlled conditions of oxygen concentration and velocity. Each of these factors is of marked importance, and it was decided to devote the major part of the experimental work to determinations by this method. It was also planned to check the oxygen drop method by actual loss in weight determinations.
METHOD OF
OBTAINING AND D E T E R M I N I N G p H VALUES
Throughout the runs technical sodium hydroxide was used to prepare the alkaline waters, and commercial hydrochloric acid and carbon dioxide from cylinders in work in the acid ranges. The solutions were made up roughly to the desired concentration, and during a run samples for pH determinations were taken about every half hour on both entrance and exit water. When the p H was between 10 and the most acid concentration studied (PH = 2.7), the hydrogen-ion concentration was obtained by colorimetric methods. The method used was a modification of that recommended by Clarke and Lubs,ll the values of the buffer solutions having been checked
APPARATUS AND EXPERIMENTAL METHOD The corrosion cell employed in the oxygen drop method has been described in an earlier paper.g It consists essentially of a wooden box containing steel baffle plates which are corroaed by liquid passing through the box, the rate of corrosion being determined by the flow and by analyses for dissolved oxygen in the entrance and exit solutions. The cell was so constructed as to be equivalent to 150 feet of standard a/8-inch pipe. I n tho case of runs made in the acid ranges, in addition to the oxygen analyseslo made by the Winkler method, the dissolved gases on a water sample of known volume were driven off and collected over mercury, the amount of hydrogen present being determined by the decrease in volume after passing over heated copper oxide in a Burrell gas-analyzing apparatus. This method also gave a check on the Winkler analysis, since, after passing through caustic solution and prior to contact with copper oxide, the oxygen was absorbed in alkaline pyrogallol. I n thc runs with acid or alkaline waters (prepared as described below), 1000 liters of solution were made up in a 8
9 10
p. 59.
Speller and Kendall, T H I S JOURNAL, 15, 134 (1923). T H IJOURNAL, ~ 16, 276 (1924). Am. Pub. Health Assoc., Standard Methods of Water Analysis, 1923,
I
I
by electrometric methods. I n three cases in the acid range the value obtained colorimetrically was checked to less than 0.2 pH electrometrically. I n work with alkaline solutions I1 Clark, "The Determination of Hydrogen Ion," 1920, P. 83. Williams and Wilkins Co.
INDUSTRIAL A N D EN(XNEERING CHEMISTRY
fjK8
of pH mare than 10, it was found that the values as calculated from double indicator titrations were within 0.1 pH of those measured electrometrically. CALCULATION OF CORROSION RATES In a previous publicationg the authors reported corrosion rates obtained in natural waters in the form advocated by Wilson2-i. e., specific corrosion rate as milligrams iron corroded per square centimeter exposed surface per year per cubic centimeter oxygen in 1 liter. However, since the main value of corrosion results lies in estimating the life of metal of given thickness, it has been decided to report all corrosion rates as average specific penetration in inches per year per cubic centimeter oxygen per liter. That is, to calculate the actual average penetration, it is only necessary t o multiply by the average oxygen concentration encountered. K
Vol. 16, No. 7
solution was'as high as that in natural water, but this rate decreased with time. It was also found that after alkaline runs the corrosion rate would not return to the relatively high rate obtained in natural water without prolonged operation with untreated water .(more than 2 weeks). For this reason after one alkaline run and prior to the next (unless a more concentrated solution was to be used) it was necessary to pass a small amount of acid solution through the cell and allow the corrosion to return to its normal value with water as indicated in Fig. 4.
= average penetration in inches per year when the oxygen
concentration is 1 cc. per liter
-
205 (liters solution per minute) X log
entering O2 exit
0 2
sq. cm. exposcd iron surface
The specific penetration reported is in all cases determined from the decrease in dissolved oxygen, though in work with acid solutions where hydrogen was evolved a certain additional amount of corrosion was also progressing by hydrogen evolution. This corrosion was small in the experiments to be reported. RESULTS CORROSION CELL-The results obtained with varying hydrogen-ion concentrations are given in the table. The figures represent data taken after corrosion rate had become constant. These same results are plotted on Fig. 2 , which shows the changes in corrosion rate with varying hydrogenion concentration a t 22" and 40' C. a t the velocity used (about 0.3 foot per second).
pH 13.0 12.4 11.3 10.5 10.3 10.0 10.0 7.4 7.1 6.5 6.5 4.4 4.1 3.6 3.5 3.1 2.7 13.05 12.86 11.05 9.46 8.30 6.74 4.50 3.80 3.50 5.96 5.40 5.84 4.50 5.2 5.0 5
TABULATED RESULTS Rate K" Liters/ Cc. Oz/Liter Inches/Year/ T ( " C . ) Min. In Out Cc. On/Liter R~MARKS Runs at 20' C. usina NaOH and HC1 4.37 0.00019 20.0 2.31 4.44 0.00042 21.0 2.31 6.30 6.09 0,00122 20.5 2.29 5.55 6.15 0.00155 2.31 5.14 5.85 21.0 0.00163 2.26 4.90 20.0 5.62 0.00171 21.0 2.31 5.07 5.84 0.00195 4.72 20.5 2.26 5.56 0.00191 2.22 3.80 20.0 4.47 0.00200 2.29 4.79 21.8 5.66 0.00200 2.27 4.57 22.0 5.41 0,00200 4.79 21.6 5.66 2.29 No hydrogen 0.00191 2.34 4.76 20.0 5.56 Hydrogen detected 20.5 4:iO 0,00214 20.9 2:34 5148 0,00237 20.5 2.35 5.50 4.54 0,00281 21.1 5.36 4.26 2.34 20.0 0.01270 4.15 1.40 2.24 Runs at 40' C. using NaOH and HCl 0,00059 39.2 2.13 3.87 3.67 3.91 3.42 0.00140 40.2 1.99 0,00232 4.45 40.3 2.10 3.60 0,00300 2.26 4.36 3.38 38.0 0.00367 40.0 2.28 4.16 3.06 0.00362 3.06 40.0 2.28 4.14 0.00369 No hydrogen 2.43 41.1 1.46 3.94 0.00414 Hydrogen detected 40.0 2.36 4.04 2.89 2.03 1.74 0.00780 39.1 3.64 Runs a2 P O o C. using COP (high 0 2 concentvation) 5.29 0.00196 No hydrogen 23.9 2.04 6.36 4,44 0,00201 Hydrogen detected 1.71 5.56 23.0 5.10 0,00243 2.05 6.40 12.6 2.37 0.00892 21.0 1.97 4.21 Runs a f ZOO C . with COZ (low Oz concentvation) 1.64 0.46 0.0118 21.6 1.77 3.14 1.20 0.0113 ' 21.7 2.24 K = ( 0 . 0 1 ~ 0 (liters ) per min.) log o2OUt for this cell.
Fig. 3 shows the change in corrosion rate with time encountered in making a run with alkaline solutions. After a run with natural water the initial corrosion rate in an alkaline
Fig. 5 shows the results obtained when using carbon dioxide in comparison with hydrochloric acid to give acid pH values. To check the oxygen drop method of obtaining corrosion, 35 weighed plates were installed in one pass of the cell, and after 31 days' running were removed, cleaned, and reweighed. The average results of the weighings are as follows : Average loss in weight per plate.. . . . . . . . 2 . 7 8 grams
Total loss in weight for 35 plates.. . . . . . . 9 7 . 3 grams Average oxygen content of water.. ...... 5.89 cc. oxygen per liter Temperature of water.. . . . . . . . . . . . . . . . . 19.6'' C. Average specific penetration-calculated from above data. 0.00184 inch per year per cc. oxygen per liter Average specific penetration from oxygen analyses.. . . . . . . . . . . . . . . . . . . . . . . . . . . 0.00194 Percentage variation from oxygen analyses 5.1 per cent
...................
................................
BEAKERTESTS-The results of varying hydrogen-ion concentration on the corrosion of steel in beakers is s h h n in Fig. 6, where corrosion is plotted against initial pH (at p H = 7 the average penetration was 0.00245 inch per year). Since the amount of dissolved oxygen present in the beaker tests was not under control, the authors feel that the results obtained are of only comparative significance. For this reason Fig. 3 is plotted in terms of relative corrosion. The original data can, however, be recalculated from figures given above. DISCUSSION OF RESULTS Fig. 2 shows that a t both temperatures there is a fairly wide range over which variations in hydrogen-ion concentration of the solution have no effect on the rate of corrosion. This range, which Wilson has termed the "natural water field," extends from a pH of about 10 to 4.1 a t 22" C. and from 9 to about 4.3 a t 40" C. with the oxygen concentrations employed in these tests and using sodium hydroxide or hydrochloric acid to vary the pH. On the alkaline side of these limits the corrosion rate progressively decreases with increasing alkalinity, falling to a very low value at a p H of
July, 1924
I N D C S T R I A L A N D ENGINEERING CHEMISTRY
13. Beyond the acid limit of natural water corrosion the rate rises very rapidly and hydrogen gas can be detected in the exit solutions above this point. The decreased corrosion in solutions that are more alkaline than p H = 10 is undoubtedly due to the formation of films of increasing protectiveness. This film formation is demonstrated by Fig. 3, which shows the time required to build up the film a t 40" C. with a pH of 13, and by the fact that a return to natural water flow does not completely destroy the film except over comparatively long intervals of time. These observations on film formation check the data which Speller and Texter obtained in similar experiments, although we find the film to be much more permanent than they do. It is significant that the break into the alkaline region (at pH = 10 a t 22" C.) closely approximates the pH value of a saturated solution of ferrous hydroxide in water (pH = 9.5). This is in entire accord with the proposed theory that the protective character of the hydroxide film is largely governed by the hydroxide-ion concentration of the liquid adjacent to the metal. Increased alkalinity above pH = 9.5 should therefore make the liquid next to the metal more alkaline and increase its precipitating power. Wilson7 and others have attributed the decreased corrosion in alkalies to the very small concentration of hydrogen-ions in the solution, using the concept that hydrogen deposition is the first stage in the oxygen reaction and that the insufficiency of hydrogen ions limits the reaction. This viewpoint appears entirely untenable, on two grounds. In the first place, the potential for either the oxygen or the hydrogen reaction is independent of the alkalinity, provided a saturated ferrous hydroxide solution is maintained against the metal. I n the second place, it has been found, both by Speller and Texter7 and by the authors, that the rate of corrosion in an alkaline solution is initially as high as if the solution were neutral but that this rate progressively decreases with time. These data strongly indicate that the action is one of protective films and that the depletion of hydrogen ions per se cannot be the controlling factor. It will be noted that there is no "critical value" of maximum corrosiveness on the alkaline side. This is in agreement with the results of Speller and Texter, but does not check the findings of Lyons and others from beaker tests. I n the natural water range the constant corrosion rates which are obtained with solutions of varying pH value show that the film protection is constant, and offer confirmation to the theory that the pH of the solution adjacent to the metal does not change in this region. I t is not until the main solution becomes ,somewhat acidic that the alkalinity of saturated ferrous hydroxide solution against the corroding metal is destroyed by neutralization from the main body of liquid. The mihin variable in this region is the rate a t which dissolved oxygen diffuses in to the metal surface. I t is well known that diffusion processes through liquid films speed up with velocity and temperature, the latter effect being due primarily to decreased viscosity. The effect of temperature in this case can be largely ascribed to increased oxygen diffusion caused by decreased viscosity. Fig. 4 shows that the rate of corrosion in natural water on freshly cleaned steel surfaces decreases 50 per cent in one week. This is attributable to the protective rust film which is built up on the steel surface with continued operation. The comparison of hydrochloric acid and carbonic acid in Fig. 5 is especially interesting. Carbonic acid causes hydrogen gas evolution and a break into the acid range at a p H of about 5.4, while hydrochloric requires a p H of 4.1. Carbonic acid is, however, only 6 per cent dissociated a t a pH of 5.4, while hydrochloric acid is practically 100 per cent
669
dissociated a t all these dilutions. As a result, the total acidity of the carbonic acid a t pH = 5.4 is sixteen times as great as the actual hydrogen-ion concentration, and consequently its capacity to neutralize the alkaline film against the metal should be as great as that of a hydrochloric acid solution with sixteen times the hydrogen-ion concentrationi. e., a hydrochloric acid solution with a pH of 4.2. The close' check between this value of 4.2 calculated for hydrochloric acid from the carbon dioxide data and the observed value of 4.1 from the actual runs with hydrochloric acid is a strong argument for the theory which has been here presented. The data-and the theory both point to the conclusion that total acidity is more important than the actual concentration of dissociated hydrogen ions in the solution-i. e., the pH value. Although the data obtained with carbon dioxide in Cambridge water showed that there was no increased corrosion until hydrogen gas was evolved (at pH = 5.4), it is probable that variations in carbon dioxide content may change the film protection in waters that build carbonate scales, thereby affecting the corrosiveness of such waters a t a lower acidity. This means that a weak acid will cause hydrogen gas evolution at lower hydrogen-ion concentration than is required by a strong acid. The data also show that hydrogen gas can be evolved when the H+ concentration next to the metal is as low as a pH of 5.4, since gas was detected when the main solution had this pH with carbon dioxide. In all probability the Hf concentration against the metal was even lower than this, and the pH a t this point may have been 7 , or even slightly on the alkaline side.
A few runs were made with very low oxygen concentration, but the results were somewhat inconsistent and are not included in this paper. On the other hand, these data indicated that a t very low oxygen .concentration hydrogen gas
Vol. 16, No. 7
INDUSTRIAL A N D ENGINEERING CHEMISTRY
670
could be detected in hydrochloric acid solutions at pH valuee of 5 . 0 4 . e., a t one-tenth the concentration required with high oxygen content. Other experimenters have detected gas evolution in water which was practically neutral in the absence of dissolved oxygen. There is no question, therefore, but what gas can be evolved from solutions of lower acidity when the oxygen content is cut down.
these (OH)- ions are neutralized by mixing with the main solution. A decrease in oxygen content will cut down the rate of corrosion and the consequent formation of (OH)ions, and will therefore make it easier to neutralize the surface alkalinity. We should predict, then, that hydrogen gas would be evolved with increasing readiness as the amount of oxygen reaction decreases. In this connection it will be noted from Fig. 2 that the critical H concentration for gas evolution at 40' C. is less than it is a t 20" C. This is due, a t least in part, to the fact that the solutions at 40" C. contained less oxygen than those at 20' C. This is a very important point in the corrosion of hot water systems where the absence of dissolved oxygen permits some hydrogen gas evolution. The corrosion rate in the acid region, as determined by the oxygen drop method, shows a rapid increase over that obtained with natural water. The actual total corrosion is really greater than these data indicate, since that part of the hydrogen which is removed as hydrogen gas is not included in the calculations. The reasons for increased oxygen depolarization in acids are undoubtedly that the evolution of hydrogen gas causes turbulence a t the metal surface and thereby increases the rate of oxygen diffusion, and also that the surface film has been neutralized and dissolved. +
ACKNOWLEDGMENT This fact is readily explainable in the light of our theory. The surface film of liquid is maintained alkaline a t a pH of 9.5 only because the rate of formation of (OH)- ions from the initial corrosion reaction is greater than the rate a t which
Much of the experimental work in this paper was carried out under a fellowship from the National Tube Company, directed by F. N. Speller. The authors wish to acknowledge the valuable assistance and cotiperation afforded by this arrangement.
United States Patents on Bone Black and Decolorizing Carbons' By W.D.Horne 175 PARK AvE., Y O N K E R SN. , Y.
HE very general interest in decolorizing carbons that has
T
developed in recent years makes any fundamental information on the subject a matter of importance. The gradual improvement with regard to their composition, means of preparation, methods of revivification, and technic of operation has made them more and more of value. All these later carbons have naturally been compared with bone black, itself a decolorizing carbon supported upon a framework of calcium phosphate and introduced into use in sugar refining by Derosne in 1812. The later-developed carbons will inevitably have much in common with their century-old predecessor and the history of its evolution must logically be a part of any comprehensive consideration of their development. The literature of bone black is surprisingly spare, while that of the decolorizing carbons is much more highly developed. A very good resum6 of the whole subject may be had, however, by reference to the files of the United States Patent Office on the subject, and it is with this in view t h a t the present compilation has been prepared. The subjoined lists present with comparative fulness the United States patents relating to bone black and decolorizing carbons, arranged under various more specific headings and according to their serial numbers. A full index, probably to be published later, has also been prepared, giving serial number, date, patentee, and title, together with the first claim in full of each patent, so that a quick means is afforded of ascertaining the general character of an invention, while further details can be seen by reference to the patent itself. U. S. PATENTS O N B O N EBLACKA N D DECOLORIZING CARBONS , Black pigment: 148,775 Bone black composition: 122,526 Bone black driers: 329,324; 336,137; 335,586; 341,497; 343,666; 749,723; 769,421; 871,705; 984,931; 1,058,369 Bone black manufacture: 153,741; 155,919; 165,344; 178,315 Charcoal furnace; 12,602 1
Presented before the Division of Sugar Chemistry a t the 67th Meeting
of the American Chemical Society, Washington, D. C., April 21 t o 26, 1924.
Chemically treating bone black: 831,805; 1,177,725 Collecting ammonia from char: 287,570; 297,948 Cooling bone black: 61,851; 68,915; 77,935; 93,208; 101,019; 186,327; 199,118; 278,356 Decarbonizing bone black: 530,632; 585,658; 586,278; 592,547; 1,184,397 Decolorizingcarbons: 739,104; 1,133,049; 1,135,216; 1,151,553; 1,195,720; 1,200,713; 1,219,438; 1,249,041; 1,250,228; 1,251,546; 1,262,770; 1,286,187; 1,287,592; 1,290,002; 1,308,526; 1,314,204; 1,358,162; 1,359,094; 1,362,064; 1,368,957; 1,383,756; 1,385,826; 1,396,773; 1,402,007; 1,413,446; 1,438,113; 1,440,194 Discharging devices: 254,474; 257,114; 268,951; 335,586; 351,929 Filter discharges: 329,305; 335,602; 335,603 Furnaces and kilns: 35,212; 85,700; 88,701; 58,702; 89,492; 173,989; 178,286; 221,725; 236,458; 265,876; 303,375; 303,379; 308,476; 314,866; 316,610; 320,110; 337,411; 345,324; 345,968; 407,912; 407,976; 412,781; 414,608; 523,248; 556,603; 557,498; 584,071; 612,319: 624,510; 638,177; 708,898; 796,303; 940,520; 972,023; 978,625; 1,004,176; 1,160,687 Oxidizing charcoal or char: 429,682; 526,180; 538,025; 617,080 Recarbonizing bone black: 509,460; 530,632 Refining substances: 1,435,972; 1,442,372; 1,447,452; 1,448,846 Retorts: 587,057; 644,507; 796,304; 851,409; 912,644 Revivifying bone black: 22,734; 26,457; 27,462; 32,679; 35,160; 36,230; 39,637; 39,638; 40,371; 47,308; 53,534; 54,771; 60,492; 62,537; 62,927; 65,359; 65,457; 65,597; 68,915; 77,935, 93,668; 96,899; 113,279; 114,780; 134,686; 150,521; 161,253; 165,992; 167,235; 178,256; 179,579; 188,006; 188,029; 190,676; 235,942; 249,004; 260,486; 265,723; 279,335; 293,430; 350,170; 450,209; 1,189,896; 1,184,398; 1,207,178 Revivifying decolorizing carbon: 1,074,337; 1,189,896; 1,269,050; 1,326,159; 1,327,222; 1,423 Steaming bone black ,962; 320,110; 391,335; 447,313; 526,180; 947,503 Sugar and bone black filters and filtration: 13,740; 329,184; 329,185; 329,210; 329,306; 329,329; 329,330; 329,332; 335,622; 335,763; 340,005; 961,180; 1,430,200 Washing bone black: 263,710; 360,581; 529,469
