CORROSION OF ELECTRODEPOSITED NICKEL Resistance to Water Containing Dissolved Carbon Dioxide and Air LEONARD C. FLOWERS AND JAMES B. KELLEY Westinghouse Electric Corporation, Springjield, Mass.
-
A n experimental investigation showed that nickel electrodeposits from plating baths containing sal ts of naphthalene disulfonic acids as brightening agents are dissolved rapidly and continuously at room temperature by distilled water through which carbon dioxide and air are bubbled. Corrosion rates range from 25 to 96 mg. per square dm. per day. Electrodeposits from nickel plating baths containing no brightening agent, as exemplified by Watts’ nickel plate, may be pitted under these conditions but do not corrode rapidly; neither does rolled sheet nickel.
The presence of air, or probably oxygen, in the bubbling gas is shown to be of paramount importance. In terms of electrochemical theory, dissolved oxygen acts mainly as cathode depolarizer in the case of bright nickel plate, and when oxygen is freely available the corrosion reaction is fast. In the case of sheet nickel or Watts’ nickel plate, it is believed that a passive condition is produced when dissolved oxygen exceeds an undetermined limit; thus the corrosion process is stifled. No explanation can be given at this time for the lack of passivity of bright nickel plate.
M
nickel plates deposited from baths containing the sodium salts of 1,5-naphthalene disulfonic acid and of 2,7-naphthalene disulfonic acid; these are within the category claimed by Schlotter. Other tests were run with proprietary bright nickel baths and it is the writers’ understanding that many of the proprietary baths now in general use employ such substances as brightening ingredients (9). When a specimen plated with bright nickel is immersed a t room temperature in a loosely-covered vessel containing distilled water through which carbon dioxide gas is bubbled at atmospheric pressure, a corrosive attack is soon observed, usually within 2 or 3 hours. The first indication is a dulling of the bright surface; jet black spots appear soon afterwards. These enlarge and extend gradually over the surface of the metal but never cover it completely. As the attack progresses the blackness seems to become less intense as the metal takes on a mottled and nonuniform darkgray appearance, often with brownish or reddish streaks that may be iridescent. A green color develops in the water and eventually nickel carbonate separates from solution as an apple-green precipitate. A yellow-brown or yellow-orange precipitate may sometimes appear earlier; qualitative tests have shown such precipitates to contain iron, probably from codeposition with nickel as a result of iron salt impurities in the plating bath. The attack may then diminish, but if the specimen is placed in a fresh portion of water through which carbon dioxide is bubbled, more nickel dissolves and this process can be repeated until the bright nickel coating is almost completely gone, exposing the basis metal over most of its area. If the exposed basis metal is copper or brass, it seems to remain free from any substantial attack, judging by its appearance. However, no quantitative experiments on the basis metal have been made and the known properties of copper and copper alloys in such an environment make it seem probable that corrosion should be occurring t o some degree. No experiments have been carried out with iron or steel as basis metal. An entirely different result is obtained by exposing a specimen on which nickel has been plated from a bath containing no brightening agent. Two different types of behavior have been observed: Frequently, there is no measurable attack on the nickel coating from the beginning to the end of the exposure period, in some cases aa long as 200 days. I n other instances the nonbright plate does discolor at the start and nickel begins to dissolve, but unlike bright nickel plate, it stops dissolving within 2 days to a week.
ETALLIC nickel is generally considered to be resistant to water containing free carbon dioxide, and it is also amatter of common belief that electrodeposited nickel coatings will possess this good resistance if they are heavy enough to be free from porosity (12). Yet there are a number of reported instances in the literature ( I , 4-6,8, 10, I S ) which indicate that nickel metal itself may be noticeably attacked by water solutions of carbon dioxide, especially when the water also contains dissolved air. McKay (4)has reported that the metal is extremely variable in its behavior, sometimes being strongly attacked and sometimes showing a complete absence of attack. The recent work of Wesiey and Copson (IS) shows that mixtures of carbon dioxide and air dissolved in hot water under a gage pressure of 35 pounds per square inch can cause appreciable corrosion of nickel in a critical range of composition-namely, from about 55 t o 90% carbon dioxide by volume. Nowhere, however, does there seem to be any reference t o a possibility that pronounced differences may be encountered in the corrosion susceptibility of electrodeposited coatings produced by different electroplating processes when immersed in cold water containing dissolved carbon dioxide and air. T h e present work shows that such differences do exist, and therefore consideration must be given t o this possibility whenever electrodeposited nickel coatings, no matter how thick, are contemplated for underwater exposure. QUALITATIVE DESCRIPTION OF CORROSION PHENOMENA
At the present time nickel electroplates may he classified into two groups, those which are deposited from solutions containing organic brightening agents and those plated from the somewhat older baths in which no such addition agents are used. The “bright nickel” processes, as the former are called, have been generally adopted where there is need for a lustrous and brilliant appearance, as for decorative purposes. The older type, best exemplified by the well known Watts’ nickel plate, have been widely used for eorrosion protection but must be mechanically buffed and polished when a bright finish is desired. A number of the bright nickel processes seem to stem from a patent issued to Schlotter in 1934 (11). The brightening agents cited as examples are nickel benzene disulfonate and nickel naphthalene trisulfonate; a general claim is made to an momatsic sulfonate containing more than one sulfonic acid group. A number of the experiments to be described here were carried out on 219
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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
Vol. 42 No. 4
ence of numerous holes during the latter part of the exposure period. The average rate for the second spwimen is 53 mg. per square dm. per day. Measurements on 25 other specimens of bright nickel plate gave rates ranging from 25 to 59 mg. per square dm. per day, with a composite average of about 40. These specimens were plated from eight different bright nickel baths and no significant differences were seen among any of the specimens. Rates of this order are much too high to be safe when the service life of an electroplated coating is to be predicted. FACTORS INFLUENCING CORROSION RATE-LEVELING-OFF EFFECTS
NICKEL PLATE
WATTS' NICKEL PLATE (A,TABLEII)
TIME IN DAYS
Figure 1. Corrosion of Electrodeposited Nickel by Bubbling Carbon Dioxide through Distilled Water in Beakers
Thereafter it remains unattacked during many tveeks of exposure. At other times, tiny pits are formed soon after immersion, and then pcnet'rate quickly to tho basis metal. I n no case, however, is the nonbright coating corroded sufficiently to expose any significant area of underlying metal. Numerous experiments have been performed with specimens plated froni various bright nickel baths and from nonbright or Watts' baths, as well as with specimens cut from rolled nickel sheet. All plated specimens were scrubbed with soap or Bon Ami and water just before they were to be placed on corrosion test; the rolled sheet nickel specimens were pickled t o a smooth, uniform, matte surface in a hot solution of sodium nit,rate, sodium chloride, and sulfuric acid. Jn some experiments electroplates from Watts' baths were mechanically polished and buffed before the tests were made; these are not reported in detail since no difference was ever observed in the behavior of buffed and unbuffed deposit's.
The dissolving process is influenced by several factors, one of which apparently is saturation of the water by dissolved nickel bicarbonate. This is shown in Figure 1, which represents the cumulative corrosion of several specimens of bright nickel plate, Watts' nickel plate, and rolled nickel sheet. I n these experiments, 5 X 10 em. panels were immersed in 450 ml. of distilled water in separate loosely-covered beakers. Carbon dioxide gas from a commercial cylinder was bubbled through the water. Dissolvcd nickel was measured a t intervals shown by each point on t h e curves. The paths of these curves are discontinuous-for example, the graph for bright nickel plate shows three distinct steps in the corrosion process. Nickel dissolved a t a rapid rate from the start up to point A on the curve, a t which timc the water was green in color and an apple-green salt had begun to precipitate. Analysis of the water, including the precipitate, gave a nickel concentration of 1.47 grams per liter. Continued exposure in this saturated water resulted in some further salt precipitation, but the corrosion rate between A and B flattened out to a low value. -\t point B , the specimen was removed, rinsed with water, and immersed in a fresh portion of water through which carbon dioxide was bubbled. The specimen then corroded along the path, BC; the rate was just as high as it had been before. The new portion of water became saturated at point C where the nickel concentration was found to be 1.27 grams per liter and the specimen stopped corroding until the test water was again replaced a t point D. Subsequent tests for dissolved nickel gave the path DEF, although it is likely that additional experimental points between E and Ii' would have followed the broken line. A t point F the concentration of nickel was 1.1 grams per liter, whereupon the panel was again immersed in fresh water, restoring the conditions for rapid attack. This water-changing process was continued beyond the scope of Figure 1 until the bright nickel plate was almost entirely dissolved away.
TABLEI. CORROSION OF SPECIR.IENS OF ELECTRODEPOSITED BRIGHThTICKEL REMOTED FROM BASISMETAL
EXI'ERIRIEXT WITH BRIGHT NICKEL PLATE HEMOVEU FHOllI BASIS METAL
Lest i t be thought that the pronounced attack on bright nickel occurred because the plate was too thin or by galvanic action between the coating and the basis metal through pores left in the plate by the plating process, it should be mentioned that many of the coatings were 0.002 inch or more in thickness. Furthermore, bright nickel plate corroded just as rapidly when plated over rolled sheet nickel as when the basis metal was copper or brass. The most convincing proof that the underlying metal had nothing to do with the corrosive attack was finally obtained when specimens of bright nickel plate were removed from their basis metal and then were exposed to t'he carbonated water conditions. For this purpose, removable bright nickel plate was prepared by plating on fusible alloy (melting point, 7OoC.), which was subsequently melted out in hot water. Corrosion data for these specimens are given in Table I. The average corrosion rate for specimen 1, which had the longer exposure and lost more weight, is 33 mg. per square dm. per day, but this value, calculated on the basis of a constant specimen area, may be too low since the actual area was diminished by the pres-
Specimen thickness, inch Surface area (both sides), sq. dm. Oriminal weight, grams FinEl weight, grams Loss in weight, % Exposure time days Av. corrosion ;ate, m d s q . dm./day
Spec. 1 (Bubbling COz through Distilled Water in Beaker) 0.0037 1.02
4.36 0.78 82
108 33
Spec. 2 (Bubbling
35% CO2,and 6 5 7 , Air through Distilled Water in Test Cell)
0,0037 1.05
4.497 3,336
26 21
53
~~
The solubility of nickel bicarbonate in water under carbon dioxide pressure has been investigated by Muller and Luber (8). They found the saturated solution under a pressure of 50 atmospheres to contain 4.45 grams of dissolved nickel per liter, but on reducing the pressure t o atmospheric and allon ing the solution to stand for 2 days, only 1.39 grams per liter remained in solution. This latter figure is roughly in accord with the quantities of dissolved nickel found in the present work at points A and C where the corrosion rate falls off so abruptly.
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INDUSTRIAL AND ENGINEERING CHEMISTRY
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The nickel salts used in the plating bath were commercial electroplating grade, the boric acid was reagent grade, and the sulfonic acid salts were the technical grade supplied b y the Eastman Kodak Company. There was no attempt to purify the individual ingredients. Solutions were made up in 3- to 4-liter quantities, warmed to 120’ F., and filtered to clear and sparkling condition through filter paper on a Buchner funnel. Before being used t o plate panels for test, the solutions were “plated out’’ on dummy cathodes for several hours to remove soluble metallic impurities, though none were definitely suspected. Watts’-type baths, B and C, were treated with hydrogen peroxide to prevent pitting. In baths A , D, E, F , and G, no antipitting agents were used. Panels plated in baths A , D and E were electrocleaned in alkaline solution; those plated i n baths B , C, F , and G were bright dipped beore plating. Bath composition F is comparable t o those specified in the original Schlotter patent. Compositions D,E,and G differ from Watts’-type baths only by addition of naphthalene disulfonates. The behavior of a specimen of commercial quality rolled nickel sheet, as shown in Figure 1, differed very little from t h a t of Watts’ nickel plate. I n both cases the flattening of the corrosion curve took place when about 200 mg. of nickel had been dissolved. Changing the test water, as indicated a t points H and K , had only a slight influence on the corrosion rate. No differences were observed between specimens taken from two lots of nickel sheet, one of which had been obtained 15 years before the other. When Muller and Luber (8) dissolved pure nickel powder in water under high carbon dioxide pressure, they also found a leveling-off effect after less than 200 mg. per liter had dissolved, whereas the solubility determined by dissolving the salt (nickel carbonate) was much higher. The similarity between their figure and the leveling-off data found in the present work is striking, but there is insufficient information to tell whether it is significant.
TIME IN DAYS
Figure 2.
Corrosion of Nickel Deposits Plated with and without Brightener
Other curves on Figure 1 are drawn t o show the corrosion behavior of Watt’s nickel plate and rolled nickel sheet under similar exposure conditions. The two graphs for Watts’ plate illustrate the two types of behavior mentioned earlier; the specimen plated from the bath designated as A in Table I1 began t o corrode initially a t about the same rate as was found for bright nickel, but the rate fell practically t o zero within a week; the other Watts’ plate ( B in Table 11) apparently suffered no initial attack and remained unchanged throughout the exposure. The flattening of the curve for Watts’ pIate, A , occurred when the test water contained only about 200 mg. of dissolved nickel; therefore, it cannot be attributed to the effect of a saturated solution. Microscopic inspection of the surface a t the end of the exposure period revealed a scattered attack over fairly large areas, where seemingly a “skin” of metal had dissolved away, and also two pin-point pits which had penetrated t o the basis metal near the edge of a punched hole in the specimen. Replacing t h e test water with a fresh portion a t point G resulted in no increase in the corrosion rate. It has not been possible t o explain the difference between the two Watts’ plates b y taking into account the minor differences in composition between baths A and B in Table 11.
EFFECT OF ORGANIC BRIGHTENERS IN PLATING BATH
The supposition that i t is the use of organic brightener in the plating bath that leads to the rapid and continuous corrosion of bright nickel is amply confirmed by the data presented in Figure 2, Brass panels plated as described in Table I1 were completely immersed in loosely-covered beakers of water through which car-
TABLE 11. PLATING BATHCONDITIONS Test Panel Designation
A B C D E F U Waits Waits Waits Briiht Briiht Briiht Brikht Bath composition, grams/liter Ni s04.7HzO 240 Ni Cl,. 6HzO 30 Boric Acid 30 Sodium naphthalene. 1-5 disulfonate. Sodium naphthalene, 2-7 disulfonate. p H (adjusted by HzSOa) Current density, amp./sq. dm. 3.8 and
. . ..
Plating temp.,
O
F.
5.4
..
240 56 37
311 57 33
240 30 30
.. ..
..
..
..
3.75
240 30 30
0.375
160
..30
300 60 41
30
3.75
..
4.5
2.0
..
..
3.2
2.5
2.7
4.4
3.8
3.8
2.3
4 4
140
and 140
5.4
..
and
and
5.4
5.6 130
..
TIME IN DAYS
130
Figure 3.
Effect of Frequent Changing of Water in Beakers
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INDUSTRIAL AND ENGINEERING CHEMISTRY
Figure 4. Apparatus for Corrosion Study under Controlled Atmosphere in Flowing Water
bon dioxide was bubbled a t about 15 ml. of free gas per minute. The significant difference between the two groups of curves in Figure 2 is simply that those specimens plated from baths cont.aining naphthalene disulfonic acid salts corrode continuously at, rapid rates clustering around 40 mg. per square dm. per day, whereas specimens plated from baths of practically t,he same coniposition, except for the absence of such brighteners, fail to corrode in this manner. Variations of current density and pII as recorded in Table I1 apparent,ly have no effect. I t is possible to produce conditions of exposure in water containing carbon dioxide under which the spectacular difference between bright, nickel and Watts' nickel or sheet nickel largely disappears. This was done in one set of experiments illustrated by the curves in Figure 3. In these tests bhe wat,er in the beakers was removed and replaced by fresh water a t very frequent intervalsthat is, every day for 11 days and every 2 or 3 days thereafter. Carbon dioxide had been bubbled through the replacement water for a t least 24 hours before use. Bubbling in the test beakers was maintained a t about 15 ml. per minute. Under these conditions, Watts' niclrel and sheet nickel corroded continuously, there being little or no diminution of t.he rate, which in these tests was about half that shown by bright nickel. Replenishing the water and rinsing the panels a t frequent intervals had exposed the specimens to room air and had also disturbed the layer of solution immediately adjacent t o the metal surface. I n some wag this treatment. may have prevented the formation of a corrosion-resisting ability on the part of Watts' nickel and sheet nickel, The significance of these data, will be discussed later.
Vol. 42, No. 4
apparatus used is shoyn in Figure 4. In this a p p a r a h s Tyater is replenished every 2 days without removing the test panels or disturbing their immediate environment, and the gaseous atmosphere can be changed a t will. Rat,ches of previously saturated distilled water are poured a t convenient' intervals into the 4-gallon reservoir through the water filling inlet. A constant head tube, or standpipe, within the reservoir is kept full by the gas lift pump which operates on the same gas mixture used in the tesB cells, thus maintaining saturated condit'ions in the water flowing out of the reservoir. The constant head tube leads to a manifold through the tubulature in the bottom of the reservoir. Water flows from the manifold into four test cells, one of which is shown in Figure 4, connected to the manifold by pure gum rubber tubing. The inlets to the test cells are made of capillary glass tubing of a size to allow a water flow rate of approximately 500 ml. per day, The test cells are made from 1000-ml. wide-mouthed Erlenmeyer flasks. The height of the sealed-on overflow tube is chosen t o make the volume of water retained in the test cell approximately 1000 nil. A slow stream of gas is always flowing to room air from the Y-vent tube in the equalizer line between the test cell and t,he reservoir. The gas mixture employed in the experiments is produced by combining a measured &earn of commercial grade carbon dioxide with a measui,ed stream of compressed air. Both streams buhble through water and mix in a 2-lit,er flask, not, shown in the diagram. From this flask the mixed gases pass to the gas lift pump and to the test cells; bubbling is maint,ained in the cells a t about 4 t o 5 ml. per minute. From time to time, the concent,ration of cnrbon dioxide is measured by absorption in potassium hydroxide solution. Water from the test cells overflows into the receivers a t about 500 mi. per day; therefore, the 1000-ml. volume may be said t o be renewed every 2 days. The receivers are changed a t 1-, 2-, or 3day intervals, depending on the current state of the corrosion process. The nickel content and volume of thc collcctcd water
I/
CELL N0.2
30
EXPERIMENTAL METHOD FOR DETERMINING CORROSIOS RATES IN FLOWING WATER
1 TIME IN DAYS
When it became apparent that simple exposure tests in beakers were so greatly influenced by such a factor as removing the specimens a t daily intervals, a scheme was devised for conducting experiments in a more controlled environment. A sketch of the
.
Figure 5. Corrosion of Sheet Yiclcel Specimens in Test Cells Water s i t u r n t e d w i t h carbon d i o x i d e i n r e s e r v o i r . r i o t l i u b l i l i n g through c e l l s
April 1950
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INDUSTRIAL AND ENGINEERING CHEMISTRY
are measured a t each change. Experience has shown that when corrosion is proceeding a t a constant rate, the concentration of nickel in the receivers is practically identical with that in the test cells. Nickel determinations were made by a modification of the colorimetric method described by Mitchell and Mellon (7). It was found that reagent grade acetone could be substituted for the ethanol prescribed in the original method and thus reduce the cost of the solvent, a not inconsiderable item when many determinations are t o be made. It was also found that the colorimetric readings would be too low if the bromine water employed as oxidant was allowed to become unsaturated, even though there appeared to be an excess in the sample after oxidizing; for this reason, an ample quantity of free bromine was always maintained in the stock bottle. Colorimetric readings were taken with a Fisher electrophotometer, using the 425B filter, after the prepared unknowns had been allowed to stand from 28 to 32 minutes to develop maximum color. Experiments were first conducted with four pickled rolled sheet nickel specimens to determine whether the four test cells would all perform in the same way. The data are shown on a greatly enlarged scale in Figure 5. The gas flow consisted only of carbon dioxide, bubbling through the water in the reservoir but not through the test cells. A small but continuous dissolving of nickel was observed; the agreement among the specimens was as consistent as might be expected for different samples of the same metal. These four curves would show little deviation from each other if plotted to a scale such as was used for Figures 1,2, or 3. EXPERIMENTS WITH BRIGHT NICKEL PLATE, WATTS’ NICKEL PLATE, AND SHEET NICKEL
Three specimens of bright nickel, heavily plated over sheet nickel, and one specimen of pickled rolled sheet nickel were then placed in the test cells. The data are represented in Figure 6, which shows that the composition and available supply of the dissolved gases exert a determining influence on the nature and progress of the corrosion. An over-all glance at the figure indicates that the progress of the corrosion can be separated into two regions: I n the first, designated “quiet conditions,” in which the gas stream was not bubbled through the test cells but only through the gas lift pump in the reservoir, all specimens were very much alike in their behavior. I n the second region, under the heading of “bubbling conditions,” the gas stream was caused to bubble through the test cells as well as through the reservoir and there is a notable distinction in corrosion rate between bright nickel plate and rolled nickel shcet. Quiet conditions wePe maintained from the start of this experiment until 43 days had elapsed. Presumably the water flowing into the test cells was saturated with carbon dioxide and air in equilibrium with the various gaseous compositions. At first the atmosphere consisted wholly of carbon dioxide; after 14 days, 20% air was introduced and after 35 days, 70% air. The corrosion rates seemed to rise somewhat with increasing air concentration but, in general, all specimens were corroded to the same extent as long as the quiet conditions were maintained. At the end of the quiet period, shown by point A on Figure 6, nickel was dissolving from all specimens a t a rate of about 10 mg. per square dm. per day. The gas mixture, without changing its composition from 30% carbon dioxide and 70% air, was then allowed to bubble through the test cells. Every bright nickel plated specimen began immediately to dissolve much faster, as shown between points A and B. The rate became constant a t about 76 mg. per square dm. per day, somewhat higher than had usually been found in beaker tests but of the same order of magnitude. On the other hand the corrosion rate of the rolled nickel sheet specimen suddenly diminished and after a few days was reduced almost t o zero. Seven days later the carbon dioxide was shut off and a sodalime absorber was placed in the air line, thus delivering air free from carbon dioxide to the apparatus. This art of the experiment is shown on the curves from B to C, a n t i t is obvious that
7’23
after the water had been freed of carbon dioxide, which required about, 3 days, the bright nickel almost stopped corroding. Rolled nickel sheet, having stopped corroding during the previous period, was unaffected by the change in atmosphere. The next step is shown a t point C, where the composition of the gas mixture was shifted to the other extreme limit-that is, back to 100% carbon dioxide as used a t the start. No increase in any of the corrosion rates was observed and none of the specimens dissolved to any extent, as shown from C to D. Apparently both air and carbon dioxide are necessary in the corrosion process. To confirm this belief, air was again added to the gas stream. The composition was set at 30% carbon dioxide and 70% air, as used before. The result can be seen by following the progress of the curves beyond the dividing line Do’ on Figure 6. The bright nickel specimens again began to dissolve rapidly, reaching a constant rate of 96 mg. per square dm. per day, and were apparently undergoing the same type of attack as had been witnessed previously. The slight jump shown a t D’for rolled nickel sheet was insignificant and was soon followed by a return t o zero rate. In one of the cells containing a bright nickel specimen, the bubbling gas was shut off a t point E, thus restoring quiet conditions in this cell only. The corrosion rate for this specimen then diminished along the path EF. The final rate was about 10 mg. per square dm. per day, almost identical with that observed for similar exposure conditions just prior to point A . On repeating the experiment with a specimen of Watts’ nickel plate in one of the test cells and with bright nickel and rolled sheet nickel in the others, it was found that the corrosion behavior was reproducible and that the dissolving of the Watts’ plate closely followed the path of the rolled sheet. All specimens corroded alike as long as the gas mixture (35% carbon dioxide and 65% air) did not bubble through the cells, but when bubbling conditions were initiated the bright nickel corrosion rate rose to 86 mg. per square dm. per day whereas the rates for sheet nickel and Watts’ plate fell off practically to zero. However, the nature of the slight corrosion that did occur on the Watts’ plate would be important
/
1800
CONDITIONS
B TIME IN DAYS
Figure 6. Effect of Changes in Atmosphere on Corrosion of Bright Nickel Plate and Sheet Nickel
INDUSTRIAL AND ENGINEERING CHEMISTRY
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Vol. 42, No. 4
in their test vessel, but the red product found by the authors could not be identified as this substance, since iron has not been found in the M7atts' plate. Further evidence is furnished by the data represented in Figure 7. Here the air content of the gas stream was kept a t 65% froin the start to the end of the experiment. Under these conditions bright nickel plate began to dissolve immediately, although in other similar experiments, a few days' delay has been observed. Watts' plate and sheet nickel have never dissolved under such conditions to yield a measurable amount of nickel in the test water. However, microscopic examination has shown the presence of a few isolated tiny pits on the Watts' plate, similar to those previously described. Corrosion a t such pits never seemed to spread nor did the number increase, even when the evposure was continued for 71 days.
Nm ROLLED SHEET NICKEL G= BRIGHTENED WATTS' PLATE TABLE P ) S * PROPRIETARY BRIGHT NICKEL PLATE
CHEMIC.4L COMPOSITION O F WATER IN TEST C E L L S
TIME IN DAYS
Corrosion of V a r i o u s TJ-pes of Kickel in Flowing W a t e r Kept S a t u r a t e d by Mixture of 3570 COa and 65% Air Figure 7 .
in evaluating this plate for service in water containing dissolved carbon dioxide. Severe pit,ting was observed to be taking place in the early stages of the experiment when the at,mosphere consisted wholly of carbon dioxide, not bubbled through the cell, Examination of the plate at the close of the experiment showed that these pits had penetrated to the basis metal. Each pit contained a small amount of insoluble corrosion product, which, being brickred in color, was first thought, t'o be an iron compound derived from impurities in the plat,e or a copper compound resulting from corrosion of the basis metal, but repeated spot test>sfailed t o show the presence of any met'al other than nickel. However, the quantity of material available for test had to be handled under the microscope and one cannot discount the possibility that relatively minor amounts of iron or copper salts incorporated in a larger amount of pale green nickel salt may dominate the coloration, yet he undet'ected by spot test separations. Wesley and Copson (13) found a deposit of bright red or rust,y colored substance on their corroding nickel specimens when iron samples were included
The foregoing experiments indicate that air plays an important part in the corrosion process. A further elucidation of this point can be had by considering the chemical composition of the water in the test cells. Just before t'he end of the experiment represented by Figure 6 the test water was analyzed for nickel, carbon dioxide, and pH. A4tthis time, as noted in Table 111, one of the bright, nickel specimens was dissolving a t the characteristic rapid rate of about 96 mg. per square dm. per day, but the dissolving of another bright nickel specimen had been slowed down to 10 mg. per square dm. per day by stopping the bubbling proce,-s. The sheet nickel specimen was apparently not dissolving at all. In measuring dissolved carbon dioxide care was taken to avoid loss of gas to room air. The method employed was t,o draw off a quantity of cell water into a closed flask containing a measured volume of 0.02147 sodium hydroxide solution and 10 ml. of 394 barium chloride, then titrate with 0.02.V hydrochloric acid to a phenolphthalein end point. Measurements of pH were made with a Beckman p H meter on samples exposed only a few seconds to room air. Results are shown in Table 111. The bicarbonate ion concentration was calculated by assuming equilibrium conditions for the dissolved salt, as Si(HCO3)2 in the cells containing dissolved nickel. For the cell containing no dissolved nickel, bicarbonate concentration was calculated from the titrated quantity of carbon dioxide, using the data of Harned and Davis ( 2 ) for the ionization constant in pure water. Lndissociated carbon dioxide was calculated for each cell by subtracting the bicarbonate equivalent from the titration data. Act,ivities vc-ere a.lways assumed to be equal to the molal concentrations. Accepted data for the solubilities of oxygen, nitrogen, and carbon dioxide and for the ionization constants of carbonic acid may also be used to compute the equilibrium composition of the water entering the test cells. JTThenthe gaseous atmosphere is composed of 307, carbon dioxide and 70y0air, the water should contain:
OF RATER IN TESTCELLS TABLE 111. COMPOSITION
Cell atmosphere Dissolving rate a t time of sampling, mg./sq. dm./dav Dissolved nickel, mg./ liter CO, titrated, me./liter
HCOP ion, calcd., mg./ liter Undissociated COz, calcd., mg./Iiter PH, determined PH, calculated from carbonic acid equilibrium
Bright Bright Kickel Piickel Sheet N i c lie1 30% COz. 70% 30% C 0 1 , 7 0 % 30% Cog, 7070 air bubbling air n o t bubair hnbhling through cell bling through cell
96
10
172 25.6 28.2
29 18 9 17 8
Too low t o measure 0.0 20. 20.2
357
GO
440
38Z
6.2
s.4
4.4
G.1
5 3
4 .2
4
1-13
Dissol\.ed oxygen Dissolved nitrogcn Undissociated COS H C O I - ion p H value
Mg./Litei 5 8 9.7 437 0 4 1
4.2
The reasonable agreement, betryeen t,he value for undissociated carbon dioxide in this list and the calculated values in t,he second and fourth columns of Tabld I11 leaves little or no doubt concerning the adequacy of the assumptions. Further support is given by t,he close correspondrnce of the calculated and determined pH values, I t is therefore reasonable to conclude that nickel dissolves under these conditions as bicarbonate. This conclusion was also reached by LLIiiller and Luber (8) for t,he di-o'ring of nickel in solutiol~sof carbon dioxide under pwssuro.
April 1950
INDUSTRIAL AND ENGINEERING CHEMISTRY
Since nickel bicarbonate is so soluble, it would not be expected t o stifle the corrosion by forming a protective coating. Instead, the process should continue until the metal had all dissolved. Therefore, it would seem that the normal and expected behavior of nickel metal exposed t o carbonated water would be t o dissolve away, as does happen in the case of bright nickel plate. Since Watts’ plate and rolled nickel sheet do not follow this pattern, it must be assumed that some factor is causing the metal surface to be protected, or to be made “passive.” ROLE OF OXYGEN IN CORROSION PROCESS *
Is
The function of air, or probably oxygen, appears t o be of the utmost importance in the corrosion process. On referring again to point A on Figure 6, it was possible t o produce a spectacular rise in the corrosion rate of bright nickel and a simultaneous decrease, practically to zero, of the rate for nickel sheet, simply by changing from quiet, or nonbubbling conditions, to bubbling conditions. These effects cannot be attributed to differences in carbon dioxide concentration as (Table 111) the cell water through which the gas mixture does not bubble should contain something like 382 mg. of undissociated carbon dioxide per liter. This is only about 13% less than the 440 mg. per liter calculated for the bubbling condition. Other experiments have shown the rapid attack to occur at much lower concentrations. There would be, however, a significant difference in the supply of oxygen available to the corrosion process. The rapid increase in the corrosion rate of bright nickel plate can be explained by considering the differences in oxygen concentration. Although there was plenty of carbon dioxide in solution prior to point A , no such statement can be made for the oxygen concentration, which would be limited practically to the quantity entering with the incoming water, since gases can diffuse only slowly through the quiet water surface from the low volume above the liquid. Approximately 500 ml. of water entered each cell per day; this would be expected to bring in about 2.9 mg. of oxygen. Subsequent to point A the bubbling gas would bring in many times that amount of oxygen. Among the possibilities for electrochemical reaction in the corrosion process are the following:
+ 2e O2 + 2H30+ + 2e -+
Anode: N i ” --+N i + + Cathode:
3Hz0
When a weight balance is computed for these reactions, not more than about 11 mg. of nickel can dissolve each day when only 2.9 mg. of oxygen are consumed a t the cathode. The measured increments of dissolved nickel just prior to point A and also at point F on Figure 6 were found to vary from 9 to 12 mg. per day. Such agreement suggests that the electrochemical reactions are correct as chosen and that the function of oxygen is that of cathode depolarization. If this theory is correct, one would expect to find no corrosion when oxygen is excluded from the test cell. The data show the attack to be very slight when the specimens were exposed to a supposedly 100% carbon dioxide atmosphere and even this slight attack may be due to traces of oxygen in the gas stream since no special attempt was made to ensure an oxygen-free condition.
If the principal function of oxygen is t o act as cathode depolarizer, a rapid increase in corrosion rate should be observed when the oxygen supply is abruptly increased. Such abrupt increase occurs at point A because the bubbling gas stream brings into each cell a quantity of oxygen far in excess of the saturation value and also provides a rapid means for getting the oxygen into solution. As a result of the electrochemical reaction, nickel should then dissolve continuously as fast as oxygen can be replenished a t the metal surface. It is thus possible to account for the rapid and almost constant dissolving rate for bright nickel plate between A and B and also beyond point D. One can also believe that cathode depolarization may account for the corrosion of sheet nickel and Watts’ nickel plate prior
725
to point A , since in this region there is no material difference in corrosion behavior between these specimens and bright nickel plate, A similar assumption may also explain the moderately high, yet not greatly differing, corrosion rates found for all types of nickel in the experiments, represented in Figure 3, in which the specimens were exposed briefly to room air each time the test water was changed, every 2 or 3 days. However, some other explanation must be sought t o account for the sudden flattening of the corrosion curves for sheet nickel and Watts’ nickel plate a t point A on Figure 6, when the oxygen supply was increased by introducing the bubbling condition. Apparently, the great increase in available oxygen affects these metal surfaces in such a way that cathode depolarization is no longer the controlling factor. ROLE O F OXYGEN IN PRODUCING PASSIVITY
When a metal surface does not dissolve a t a faster rate, even though the environmental conditions might be expected to favor an increase, i t is commonly stated that a passive condition has been produced. It is difficult t o explain this further, but it seems almost certain that the stifling of the corrosion process is a result of a plentiful oxygen supply, and the authors believe that the stifling is associated with the formation of a protective oxygencontaining skin on the metal surface. A similar view has been expressed by Wesley and Copson in discussing the passivity of nickel in steam condensate (19). However, other theories can be used to explain the formation of a passive surface, and the evidence here is too limited to establish the validity of any one view. The pitting observed on several of the Watts’ plates gives further support to the belief that the absence of general corrosion is due to a passive or protected surface. In such instances the metal surface seems to be thoroughly protected except for two or three tiny spots which happen to be anodic. This type of corrosion is well known in other instances when metal is believed to be passive-that is, in the pitting of stainless steel and the occasional failure of chromate when used to inhibit the corrosion of iron and steel. Whether this possibility of pitting is a practical hazard or not depends on the service application and local conditions of exposure. It would seem unwise to generalize at the present time except to say that any type of nickel may be susceptible to some form of attack when immersed in water containing a substantial amount of carbon dioxide and a limited quantity of air. It is entirely possible, of course, that the hazard to Watts’ plate will be confined t o artificially high concentrations of carbon dioxide and that exposure to natural water will rarely or perhaps never lead to corrosion failure for this reason. The same cannot be said for bright nickel plate, however, since rapid failure has been observed in one instance of exposure to natural water having the composition given in Table IV.
TABLE IV. COMPOSITION OF NATURAL WATERIN WHICHBRIGHT NICKELPLA~TE CORRODED Free ammonia Bicarbonate HCOsCalcium, Ca Carbonate, COa’ Chloride, CI Hydroxide, O H Iron, Fe lMaynesium, A!lg Nitrate, NosNitrite, NOeSilica, Si02 Sulfate,, 804Alkalinity Hardness Total solids Loss on ignition Fixed solids pH water as received pH: water after boiling 15 min.
Parts per Million 0.012 67.0 19.2 0.0 5.5 0.0
0.07 4.7 0.40 0.004 8.7 31.3 53 69 135 53 82 7.1 9.6
INDUSTRIAL AND ENGINEERING CHEMISTRY
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MINIMUM CONCENTRATION O F CARBON DIOXIDE FOR RAPID ATTACK ON BRIGHT NICKEL
The normal amount of carbon dioxide in the atmosphere (0.03%) has no appreciable effect on bright nickel plate during long time immersion in an open vessel of distilled water. A certain minimum quantity of dissolved carbon dioxide seems to be necessary before the attack can become progressive. In an attempt to determine the approximate location of this minimum,
Vol. 42, No. 4
bath makes the electrodeposited metal more susceptible to attack in carbonated water. Apparently, something in the plate prevents oxygen from exerting its customary stifling effect. When the stifling effect is absent, oxygen acts principally as a cathode depolarizer and allows nickel to dissolve as soluble bicarbonate a t a rapid rate. The “something” in the plate may be codeposited organic matter, may be a different crystal form or size, may be associated with the striated layer structure described in the literature (3,I C ) , or may be none of these things. The detrimental effect of traces of codeposited iron has been suspected in some of the specimens plated from commercial baths, but suspicions of this type do not help t o account for the marked differences observed between specimens plated from laboratory Watts’ baths, with and without brighteners. Even though the mechanism remains unknown, it is hoped that the experiments described in this paper may arouse further interest in a more thorough study of the factors influencing the passivity of bright nickel, since many natural waters contain appreciable quantities of carbon dioxide. Bright nickel plate seems admirable in so many respects as a protective coating that it deserves much attention and effort in this adverse situation.
350m 300
i
PROPRIETARY BRIGHT NICKEL PLATE
W D
50
-
TIME IN DAYS
Figure 8. Minimum Content of Carbon Dioxide Producing Rapid Corrosion of Bright Nickel Plate
the proportion of carbon dioxide was increased stepwise a t weekly intervals, starting with a composition of 1 to 3%. Data for various specimens of bright nickel plate, Watts’ nickel plate, and rolled sheet are shown in Figure 8. No appreciable attack occurred on any specimen during the first 2 weeks’ exposure, when the carbon dioxide content did not exceed 5 t o 6%. However, when the carbon dioxide content was stepped up to 11 to 12%, a specimen plated from a proprietary bright nickel bath immediately began to dissolve a t the characteristic rapid rate. Calculation shows that the water then contained about 160 mg. of dissolved carbon dioxide per liter. Another specimen had been plated from a Watts’ bath containing brightener (bath G, Table 11)’but the plate did not dissolve rapidly at this time. Instead, pits began to form; some 20 to 25 were estimated to be present. The experiment was continued beyond the 30-day period shown in Figure 8, and eventually when the carbon dioxide content had been increased to 35% and was maintained at that level for 3 weeks, the brightened \X7atts’ specimen suddenly began to corrode rapidly in the pattern common to all bright nickel specimens. Evidently some degree of passivity is developed by bright nickel, but it possesses no ability to withstand for long the corroding influence of the carbon dioxideair-water mixtures employed in these experiments. The authors cannot give a satisfactory explanation as to why addition of naphthalene disulfonic acid salts to a nickel plating
s~-nnt.mY
1. Bright nickel plates deposited from eight different baths mere rapidly and continuously dissolved a t room temperature in distilled s a t e r through which mixtures of carbon dioxide and air were bubbled a t atmospheric pressure. The rapid attack occurs on specimens plated from baths in which naphthalene disulfonic acid salts are used as brightening agents. The corrosion rates ranged from 25 to 96 mg. per square dm. per day, Rapid corrosion was not observed when the gas mixture contained less than about 10% carbon dioxide. 2. Under similar conditions, rolled sheet nickel and nickel plates deposited from l17atts’ baths did not corrode rapidly o r continuously. I t is believed that the difference in behavior is a result of a differing reaction to dissolved oxygen in the water. When the oxygen supply is limited, rolled shcet nickel and Watts’ plate are dissolved to about the same extent as bright nickel; rates of 10 mg. per square dm. per day have been observed and can be interpreted on a basis of the amount of oxvgen available for the cathode depolarization reaction. When the oxygen supply is plentiful, rolled sheet nickel and Watts’ plate become passive; bright nickel plates dissolve more rapidly. The lack of‘ passivity on the part of bright nickel plate apparently results from the use of brightening agents in the plating bath but the mechanism cannot be explained. 3. Electrodeposits from Watts’ baths have been found to be pitted through to the basis metal a t tiny, isolated and widely-separated spots in many of the experiments. In such instances, the remainder of the surface appears to be unattacked and dissolved nickel cannot be measured in the test water. It is believed t h a t the passive state has been broken down a t these localized spots. 4. The corrosion of bright nickel diminishes to a low rate when the test water becomes saturated with dissolved nickel. Under approximately 1 atmosphere of carbon dioxide pressure, saturation is reached when the water contains about 1200 mg. of dissolved nickel per liter. 5. -4test method is described in which specimens can be subjected to a slow constant flow of corrodent under controlled atmospheric conditions and by which corrosion rates are determined by measuring the amount of metal dissolved by the flowing liquid. ACKNOWLEDGMENT
Credit for preparing and plating the numerous specimens required for the experiments is owed t o Myron Ceresa, S. Barnartt, N. E. Pingel, H. E. Ricks, Colin Hastie, and P. W. Prouty, all of the Westinghouse Electric Corporation.
INDUSTRIAL AND ENGINEERING CHEMISTRY
April 1950
LITERATURE CITED (1) Guertler, W., and Liepus, T., 2. Metallkunde, 17, 310-5 (1925). 65, 2030-7 (2) Harned, H. S., and Davis, R., Jr., J. Am. Chem. SOC., (1943). (3) Hothersall, A. W., and Gardam, G. E . , J . Electrodepositow’ Tech. SOC.,15, 127-40 (1939). (4) McKay, R. J., IND.ENG.CHEM.,21, 1283-7 (1929). (5) McKay, R. J., and Worthington, R., “Corrosion Resistance of Metals and Alloys,” p. 369, New York, Reinhold Publishing Corp., 1936. (6) Mellor, J. W., “A Comprehensive Treatise on Inorganic and Theoretical Chemistry,” Vol. XV, p. 156, London, Longmans, Green & Co., 1936. (7) Mitchell, A. M., with Mellon, M. G., IND. ENG.CHEEK., ANAL. E D . , 17, 380-2 (1945).
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(8) Miiller, E., and Luber, A., 2. anorg. allgem. Chem., 187, 209-30 (1930). (9) Pinner, W. L., Soderberg, G., and Baker, E. M., “Modern Electroplating,” p. 242, New York, Electrochemical Society, Inc., 1942. (10) Robl, R., 2.angew. C h a . , 37, 938-9 (1924). (11) Sohlatter, Max, U. 5. Patent 1,972,693 (Sept. 4, 1934). (12) Uhlig, H. H., ed., “Corrosion Handbook,” p. 254, New York, John Wiley & Sons, Inc., 1948. (13) Wesley, W. A., and Copson, H. R., J . Electrochem. Sac., 95, 22641 (1949). (14) Young, C. B. F., Proc. Am. Electroplaters’ Soc., 28, 124-35 (1940). RECEIVE August D 22, 1949.
Separation of Acetic Acid and Water by Distillation __
J
EFFECT OF CALCIUM CHLORIDE ADDITION LEO GARWIN AND KENTON E. HUTCHISONl Oklahoma Agricultural and Mechanical College, Stillwater, Okla.
B
acetic acid solution a t Experimental data are presented on the vapor-liquid ECAUSE acetic acid about 6.5 weight % lithium equilibrium of the system acetic acid-watercalcium and water are not too chloride, 10 weight % calchloride at 1 atmosphere. These data were obtained with readily separated by ordicium chloride, and 12 weight a view to ascertaining the possibility of separating acetic nary distillation, methods acid and water under conditions of reversed relative vola% sodium chloride. i n v o l v i n g auxiliary techIn order for this relative tility by extractive distillation with calcium chloride. niques have been used for volatility reversal t o take The results show a considerable effect of calcium chloride some time. These methods place throughout the distiladdition, with a reversal taking place at approximately 8 include (10) a z e o t r o p i c lation column, it is necescalcium chloride in the liquid phase. weight distillation with a watersary that the extractive disi m m i s c i b l e organic comtillation agent be present pound such as butyl acetate in the liquid in the proper concentration on all of the trays of (Othmer process),liquid-liquid extraction with ethyl ether or ethyl the column. That is to say, i t must be soluble in glacial acetic acetate, followed by the removal of the solvent from the extract by acid as well as in water. Semiquantitative solubility studies by fractional distillation, and simple extractive distillation (without Davidson (1) show sodium chloride and potassium chloride t o reflux) using a wood oil (Suida process). In the last-named be rather insoluble in glacial acetic acid. On this basis, it might method, the water is removed overhead and the acetic acid-wood be expected that lithium chloride, an alkali chloride, would also be oil bottoms mixture is separated by a second distillation under insoluble. Calcium chloride, however, is quite soluble in acetic vacuum. acid and data for its solubility as a function of temperature (6) The aqueous acetic acid solution t o be separated is very freare given in Figure 1. It was selected, therefore, as the salt for quently a dilute one, and it was thought worth while t o investigate further investigation. further the separation of the components of such a mixture by an extractive distillation process in which the acetic acid would be EXPERIMENTAL taken overhead and the bulk of the mixture (water) would be removed as bottoms. I n order t o do this-i.e., reverse the normal A11 chemicals used in this work were analytical reagent grade. relative volatility of acetic acid and water-it would be necessary The glass, electrically-heated equilibrium still employed was esto use, as the extractive distillation agent, a substance which sentially the one described by Jones, Schoenborn, and Colburn would tend to form a loose combination with the water. Inor(C), but modified in the following respects: ganic salts seemed t o offer good prospects for this purpose. The condensate chamber was filled with glass beads to reduce McBain and Kam (6)reported some work on the distillation of its volume relative to t h a t of the residue chamber to the greatest dilute solutions of acetic acid in water in the presence of lithium possible extent. During operation the condensate-residue volume chloride, sodium chloride, potassium chloride, potassium thiocyaratio was about 1 to 4. A wick of glass wool was substituted for the wire helix in the nate: sodium sulfate, potassium nitrate, and sodium acetate. flash boiler. This permitted better distribution of the distillate Quartaroli (9) did a somewhat similar study with sodium chloover the boiler heating surface, avoiding local overheating, and ride, lithium chloride, calcium chloride, and sodium bromide. minimizing the danger of the glass cracking. Calculations based on the data of these investigators showed that, The pressure on the still was maintained a t 760 * 0.5 mm. by of the salts posseseing commercial possibilities, lithium chloride, means of a Model No. 5 industrial Cartesian manostat (The Emil calcium chloride, and sodium chloride were the most effective, Greiner Company), actuated by compressed nitrogen gas from a with expected relative volatility reversals taking place in dilute cylinder. 1 Present address, Kerr-McGee Oil Industries, Ino., Oklahoma City, Okla.
