INDUSTRIAL A N D ENGINEERING CHEMISTRY
370
Vol. 1s. N o . 4
Factors Other than Dissolved Oxygen Influencing the Corrosion of Iron Pipes' By John R. Baylis MONTPBELLO
FILTERS, BALTIMORE, MD.
When a bright iron surface is exposed to the water of any of our public water supplies, dissolved oxygen a t 6rst has the greatest influence on the rate of corrosion, but after a film of iron rust has formed other factors become of major importance. The saturation point of ferrous hydroxide in the absence of dissolved oxygen is very low for natural waters when the pH is above about 8. This does not confirm the work of others who have used pure distilled water. Ferrous carbonate is only very slightly soluble a t the saturation point of calcium carbonate, but the solubility concentration increases very rapidly as the pH is decreased below calcium carbonate equilibrium. Iron will not corrode t o t h e extent of producing hydrogen gas bubbles in pure distilled water a t room temperature or below even after standing for several months. The liberation of hydrogen gas is due to the presence of negative ions other t h a n the (OH)- ions. These negative ions are concentrated a t the metal surface by the difference in electrical potential of the metal and the water. As a result of this concentration iron salts are formed. The final equilibrium of iron and carbonates is a very insoluble ferrous carbonate if the pH of the solution is sufficiently high. Other salts of iron, such as the sul-
.. .. .. I T H I N the past few years so much has been written on the corrosion of iron, with such close agreement as to theories, that another article relating to this phase of the problem would be out of place at the present time. As most of the work done heretofore has been to establish theories, and has been done largely upon freshly exposed iron surfaces, it seems to be now time for water-works chemists to give more consideration to corrosion as it is actually taking place in water pipes. Most of us realize there is yet much unknown when we try to explain all corrosion by the theories now generally accepted. The present investigation, which has been under way for several years, and quite actively for the past two years, has covered several phases of the problem. A great deal of time has been given to studying the products of corrosion, for it is believed that in this way the most can be learned as to how and why corrosion is progressing. Fortunately, an abundance of large-diameter cast-iron pipe has been available for examination. Changes in the distribution system of Baltimore have made it necessary to remove a considerable number of old pipes, some of which have been in service for thirty-five years. This has afforded an excellent opportunity for close inspection and the collection of many samples for test. Additional old pipes were examined, and sarpples collected in other cities where waters of different characteristics are used for public supplies.
W
Electrochemical Theory of Corrosion
fates and chlorides, are very soluble and will exist only a t a pH much less t h a n t h a t formed by ferrous carbonate. When such salts are present they tend to maintain a pH of from 6 to about 7.5 against the iron surface. Some of the iron oxides or hydroxides are very magnetic and are attracted to the metal surfaces with considerable force, resulting in a fiber formation and the building up of a very porous precipitate. Pits with overlying tubercles s t a r t to form when the rust becomes so thick t h a t the diffusion of soluble iron compounds to the water outside of the precipitate is retarded to the extent of. allowing the dissolved oxygen zone to extend to the surface of the precipitate. Precipitation then begins to take place within the existing precipitate and soon forms a n impervious or nearly impervious membrane near the outer surface of the rust. While the membrane is being formed there is considerable concentration of iron salts such as the sulfates and chlorides. Such salts are probably essential for pitting and tuberculation to continue. Active tubercles maintain a pH near 6 on the inside next to the metal surface and dormant ones have a pH somewhat higher. . #
*.
....
R l ~ K a y , Aston,6 ~ Wilson,' Speller,s Whitman, Russell, and Altie~-i,~ and Shipley, McHaffie, and Clare.lo The chief criticism of these very excellent articles, if any may be offered, is that they leave us very much in the dark as to how to lessen corrosion on a broad scale, which is the problem of the water-works chemist. I n fact, the theories probably have retarded progress in some instances, for certain authors have tried to show that within the ranges of our natural waters the pH has no influence on the rate of corrosion. This may be nearly true for a freshly exposed surface, and also for corrosion rates under some conditions, but the p H does a t least have considerable influence in keeping the water satisfactory for domestic uses. The iron exposed to water in our public water supplies probably constitutes the greatest surface of iron exposed ta water where it is impossible to protect with paint or other artificially applied coatings after the original protective coating has corroded or has otherwise been removed. With the hope of throwing a little more light upon the solubility equilibrium of certain iron compounds which are produced when corrosion is taking place, the concentration of certain soluble salts in tubercles, and changes produced in the chemical composition of the water, the present article is submitted. One encouraging fact is that corrosion of the iron pipes in Baltimore has been materially reduced by chemical treatment of the water. This is pot only saving thousands of J . SOC.Chem. I n d . , 48, 320T (1924); THISJOURNAL, 17, 363 (1925). Trans. A m . Electrochem. Soc., 41, 201 (1922); Thompson and McKay, TEISJOURNAL, 16, 1114 (1923). 8 Trans. A m . Electrochem. Soc., 41, 201 (1922). 7 THISJOURNAL, 16, 127 (1923). 8 I b i d . , 17, 339 (1925). 0 I b i d . , 16, 665 (1924). 10 I b i d . , 17, 381 (1925). 4
There is fairly general agreement that corrosion of iron is electrochemical, as suggested by Whitney2 in 1903. Valuable contributions to theory are found in articles by Rancroft,3
* Received J
8
January 4, 1926.
J . A m . Chem. Soc., 36, 394 (1903). T i a m . A m . Electrochem. Soc., 9, 13 (1908).
8
I N D CSTRI,iL A N D ENGINEERING CHEMISTRY
April, 1926
dollars annually for pipe repairs, but is giving the consumer a mater which is satisfactory in every respect. However, the problem of preventing corrosion is reserved for another publica tion. According to the electrochemical theory of corrosion a slight amount of iron will go into solution when exposed to the water of any of our public water supplies. This is apparently what does happen, but it does not lclng remain in solution as the first product of corrosion if certain compounds are present. As to d i a t happens a t this point there is some difference of opinion. The general opinion is that the first
rn
F e r r o ~ s Carbona+e
wih
6 0
7.
4
F i g u r e 1-Relation
I3
B
6
5 ~ ~ 1 ~ s I-R O EN
PPM
Fe
iz
14
Fe
of pH t o S o l u b l e I r o n in t h e Absence of Oxygen
371
row hydroxide of about 3.7 parts per million of Fe a t a pH of 8.0, about 3.4 a t a p H of 9.0, and about 1.8 a t a pH of 10.0. This is a large amount of soluble iron and it would be objectionable in water for domestic use. If the actual solubility of ferrous hydroxide were as given by these authors, it would be essential for some industrial uses to aerate the water to precipitate the soluble iron, rather than deaerate to reduce corrosion as is sometimes recommended. T a b l e I-Soluble Iron i n t h e Absence of Oxygen Freshly c u t lathe turnings or other bright metal surfaces used. T h e samples of iron and steel were placed into Pyrex flasks and the flasks sealed without any air space. T h e solution was filtered through three thicknesses of No. 40 Whatman filter paper, using extra precautions t o not allow the solution to come in contact with the air until after filtration. AFTER STANDING IN COXTACT -APPLIED WATERDisWITH THE IRON Dnys Alka- solved AlkaSoluble standlinity oxygen linity iron COz P . p . m . P . p . m . F e Con P . p . m . P . p . m . pH pH Test ing 1 7 7.2 1.0 6 1.0 8.5 0 10 0.25 8.5 0 10 0.10 2 16 6.8 3.0 6 7.0 3 2 6.3 27.3 50 9.0 7.5(?)18.9 60 6.00 49 1.00 4 6 9.0 7.8 6.0 6.3 29.0 48 5 14 7.4 3.0 50 2.0 9.0 0 39 0.05 6 14 6.6 1.0 2 2.0 8.8 0 10 0.10 7 14 6.6 1.0 2 1.5 8.2 0 0.15 7.0 7.1 3.40 47.0 0 8 9 4.8 9 4.8 1.20 21 47.0 0 7.0 7.3 10 48 4.8 47.0 0 7.0 7.6 1.4 0.05 11 1 6.8 1.0 2 1.0 8.6 0 0.10 12 8 8.4 0 46 10.0 8.8 0 33 0.05
The amount of soluble iron that will exist in water a t various pH values is of considerable importance; consequently much time has been given t o the est'ablishment of approximate solubility curves. Out of one hundred and twenty-five t'ests on various kinds of commercial iron and steel it was not found possible to duplicate the curve of Whitman, Russell, and Davis under any condition of exposure, a t room temperature, to the Baltimore water, or t.o several other
product of corrosion is ferrous hydroxide, a slightly alkaline compound reported to give pure water a pH of about 9.5 in the absence of oxygen. Owing to the alkalinity of ferrous hydroxide it has been assumed by several writers that the pH of the water between about 5.5 and 9.5 has no influence on corrosion rates. It will be shown that while this may be nearly true for freshly exposed surfaces and under the particular conditions of the tests, it is not true when the surface has been exposed for a long time, or under most natural conditions of exposure. I n fact, it is likely that there are no natural conditions which will duplicate experiments conducted by several authors to determine the solubility of ferrous hydroxide, especially when commercial productjs are used. The effect of dissolved oxygen on corrosion rates has been covered very completely by a number of workers, such as Speller, Wilson, Evans, and Whitman. Solubility of Ferrous Hydroxide
The compound of greatest importance in the corrosion of iron is ferrous hydroxide. Whitman, Russell, and Davis" give the following for its solubility: Calculating from the solubility in pure water, we find 6.7 X
0
IO
10-5 mol of ferrous hydroxide (dissociated plus undissociated),
in solution and that it has a pH of 9.6, or a hydroxyl-ion concen(thus checking the values of Krassa,l3 tration of 4 X and Lamb14).
The accuracy of the results obtained by these authors is not questioned, but it appears to be something never realized in practice when commercial products of iron are exposed to natural waters. From the curve given by Whitman, Russell, and Davis (see Figure 1) we should expect iron, in the absence of oxygen, to produce a soluble concentration of ferJ . A m . Chcm. SOC., 47, 70 (1925). 2. Elektochcm., 14, 77 (1908). 18 Ibid., 16, 491 (1909). 14 J . A m . Chem. Soc., 32, 124 (1910). 11 1)
20
30 40 ALKALINITY
50
6 0
70
of Calcium C a r b o n a t e S o l u b i l i t y to t h e S o l u b i l i t y of F e r r o u s C a r b o n a t e 0.00 to 0.10 p. p. m. F e 1.0 to 10.0 p. p. m. F e X 0.10 to 1.0 p. p. m. F e A 10.0 to 20.0 p. p. m. F e
F i g u r e 2-Relation
0
0
natural waters. Various times of exposure from a few hours to over a month were tried. The chief variation of these tests from those by others is that a small amount of precipitated iron oxide or hydroxide was present in most of the tests. This, however, is comparable to conditions in actual service. The points shown in Figure 1are from 1 to 30 days' exposure to iron in sealed containers-some glass and some iron (various lengths of 1/4- and 11/2-inch pipes). The alkalinity of the natural waters was largely calcium carbonate. Fer-
INDUSTRIAL A N D ENGINEERING CHEMISTRY
372
Vol. 18, No. 4
T a b l e 11-Solubility of Ferrous Carbonate Ten grams of FeS01.7Hz0 dissolved in 250 cc. of water a n d 20 grams of NanCOs in 100 CC. of water were rapidly mixed a n d then tightly stoppered t o prevent absorption of oxygen. After standing several hours t h e precipitate was washed several times with oxygen-free distilled water. T h e flask when stoppered holds 350 cc. and the amount of solution withdrawn each time was immediately replaced with nearly oxygen-free distilled water a n d t h e flask tightly stoppered with n o air space a t t h e top.
APPLIEDWATER
Distilled water
10 cc. 0.05 N NaOH a n d distilled water Water with n .. = -H - 6 - .A. -,alk. 34 Wate; with p H 8 . 6 , alk. 50 Water with p H 8 . 6 , alk. 50 Water with p H 8 . 7 , alk. 59 J cc. 0 . 0 5 N NaOH and distilled water 5 cc. 0 . 0 5 N NaOH and distilled water 3 cc. 0 . 0 5 N NacDH and distilled water 2 cc. 0.05 N NacDH and distilled water 2 cc. 0 . 0 5 N N a O H and distilled water 1 . 5 cc. 0 . 0 5 N NaOH a n d distilled water 1 cc. 0.05 N NaOH and distilled water
a
SOLUTION WITHDRAWN FROM FLASK -Alkalinity, P. p. m.Time standing 4 hours 2 hours 17 hours 3 hours 3 hours 18 hours 1 day 2 days 2 days 1 day 1 day 5 hours 19 hours 1 day 3 hours I day 1 day 1 day 1 day 1 day 2 days
Cc. 210 210 200 210 210 210 210 225 230 ~. 250 250 250 250 250 250 250 250 250 250 260 250 ~
pH
9.6 8.0 7.3 7.0 6.6 6.7 6.88.6 7.0+ 7.1 7.1 7.148.1 9.6 9.1 8.6 8.4 7.2 7.3
Phenolphthalein 336 116 20 0 0 0 0 0 0
1 0 0
Co1
-
M.O. 968 388 160 83 55
P. p. m.
‘14 .-
n
0
0 0
0 2.0 13.2 24.6 24.0 33.4
54 44 46 48
60 62 68 60 40 42 40 30 30 36 40
Soluble iron. P. p. m. F e Trace Trace Trace Trace 0.8 6.5 16.0 12 :o
i4.0
0 3.6 4.0 4.8 3.5 Trace Trace 0.10 Trace 0.10 4.0 1.0
12.5 15.0 15.0 10.6 0 0
0 0 0
8.5 5.0
Five grams of CaCOa crystals added to t h e flask withFeCOa.
rous carbonate gives an alkalinity reaction and when it is present it is included in the total alkalinity, but the results are expressed in parts per million of calcium carbonate. I n many of the tests the water was deaerated until practically no dissolved oxygen was present, and in all instances the containers were thoroughly sealed with no air space. When exposed for more than 3 or 4 days to the Baltimore water, and to all water containing a small amount of free or halfbound carbon dioxide, there was a liberation of hydrogen gas if a slight amount of precipitated iron hydroxide was present, indicating that the solution was saturated with this gas. So far as could be observed the amount of hydrogen gas present had no appreciable influence on the solubility equilibrium of ferrous hydroxide. KO compound of iron tried gave as much as 3.4 parts per million of soluble iron a t a pH of 9.0 in the absence of oxygen. From the large number of tests made no conclusion can be formed other than that the saturation point of soluble iron under natural conditions of exposure in water pipes is quite low for all pH values above about 7.5 to 9.0, depending on the alkalinity of the water. I n fact, the point at which only about 0.1 part per million of soluble iron exists appears to be a curve closely following the calcium carbonate solubility curve, shown in Figure 2. It is probably the variable oxidation or precipitation rate that causes alkaline waters to corrode apparently as fast as more acid waters. A certain amount of soluble iron at a p H of 9.5 will oxidize to ferric hydroxide and form a precipitate probably a hundred times as fast as it will a t a pH of 5.5. There are no definite figures on the effect of pH on oxidation rates, but it evidently has great influence on the time of forming a ferric hydroxide precipitate, and it is believed that the slowness is due a t least partially to a slowness of oxidation. If the oxidation of soluble iron compounds is rapid and is not affected by pH, the reduction of the supersaturation of the oxidized compound by the formation of a solid precipitate is affected by pH, and for all practical purposes i t probably is the same or a t least in the same direction. With a low solubility concentration and a fast oxidation the rate of corrosion may under some conditions be even greater than with a high solubility Concentration and slow oxidation. Solubility of Ferrous Carbonate
Ferrous carbonate was produced by adding a solution of sodium carbonate to a solution of ferrous sulfate and washing
practically free from sulfates with oxygen-free distilled water. This may not be the best way to prepare ferrous carbonate and the solid compound obtained may not be the pure carbonate, but i t probably is as pure as ferrous carbonate produced in the process of corrosion. Tests show that ferrous carbonate in the absence of oxygen is practically insoluble at a p H and alkalinity closely following the calcium carbonate solubility curve. This is shown by Figure 1 and also by Tables I1 and 111. A number of tests have been made, but there are not yet sufficient points to locate accurately the lines of variable soluble iron concentrations. Very likely, however, they will not vary greatly from those shown in Figure 2. Most of the recent literature on corrosion seems to have ignored the solubility of ferrous carbonate, or that it is formed as a product of corrosion. It is believed that this is a very serious omission, for even if carbon dioxide does not T a b l e 111-Solubility of Ferrous C a r b o n a t e NazCOa solution was added to excess of FeS04.7HzO solution and washed ten times with distilled water; three-fourths of solution were withdrawn each time.
-SOLUTION WITHDRAWN FROM FLASK-
APPLIEDWATER Distilled water Distilled water Distilled water 2 cc. 0.05 N NaOH and distilled water 10 cc. 0.05 A’ NaOH and distilled water 10 cc. 0.05 N NaOH and distilled water 10 cc. 0.05 N NaOH and distilled
water 10 cc. satd. Ca(0H)z and water 10 cc. satd. Ca(0H)z and water 12 cc. satd. Ca(0H)z and water 15 cc. satd. Ca(0H)z and water 12 cc. satd. C a ( 0 H ) r and water 10 cc. satd. Ca(0H)n and water 10 cc. satd. Ca(0H)z and water Distilled water Distilled water Distilled water Distilled water Distilled water Distilled water
Time standing Days 14 13 11
pH 6.8 6.8
6.8
Soluble Alkairon P. p. m. linity COz P.p.m. P.p.m. Fe 22 25.5 16.0 38 46.6 20.0 35.0 18.0 38
2
6.8
40
36.0
5
7.0
80
8.8
2.4
3
7.8
92
1.7
Trace
86 7.5 6 . 6 4 - 44 6.5 40 6.7 64 6 . 8 + 64
1.3 34.2 44.0 33.4 26.2
0.10 20.0 20.0 11.0 8.8
7.0
42
9.7
4.0
7.3
48
7.9
2.0 0.8
distilled distilled distilled distilled distilled distilled distilled
6 6 5 5 22 20
16.0
7.5
48
5.3
7.4
44
8.8
8.6
40
0
Trace
5.6
42
0
Trace
8.6
41
0
0.10
5.6 8.4
40 39 29 39 36 3s 45
0 0
Trace Trace
8.0
8.0 7.5 7.5 7.0
n i.0
1.5 3.5 16.7
2.5
0.20
Trace 0.30 1.0 8.8
INDUSTRIAL A N D ENGINEERING CHEMISTRY
April, 1926
react directly with metallic iron, ferrous carbonate is produced as a product of corrosion in varying quantities if the acid is present either as free or half-bound carbon dioxide. The rapid decrease in concentration of carbon dioxide when water is exposed to corroding iron is ample proof that it is entering into the combination in some manner. This is well illustrated by Tables IV and V. The decrease takes place either with or without dissolved oxygen in the water, indicating that it is a t least partially independent of oxygen. Approximately two hundred tests made with waters containing either free or half-bound carbon dioxide have in every case shown a decrease of carbon dioxide when the solution wm protected from air containing carbon dioxide. I n a solution with bicarbonate alkalinity less than 150 nearly all the half-bound carbon dioxide will be removed, producing a pH of about 10 in some instances. Definite proof that ferrous carbonate is the first product of the reaction with iron has not been found, and the formation of this compound appears to be fairly slow if this can be measured by the rate of liberation of hydrogen gas; yet the removal of a large portion of the carbon dioxide in the solution is fairly rapid under most conditions. There is considerable evidence of adsorption of the carbonate ions by the oxides or hydroxides of iron, and it is believed that this is largely responsible for the rapid initial decrease of carbon dioxide with the formation of very little hydrogen gas. If the solution is allowed to stand sealed for a long time enough hydrogen gas will eventually be produced to account for a direct chemical combination with the iron. T a b l e I\'-Reduction of Free a n d H a l f - B o u n d COz by C o r r o d i n g I r o n About 30 grams of lathe turnings from a black wrought-iron pipe were placed into a 250-cc. glass tube, and the tube tightly stoppered. Any increase in p H with the alkalinity remaining constant means a reduction of free or half-bound COz. I n many instances there is a reduction of alkalinity caused from the half-bound COz being reduced to the point where it creates a supersaturation of calcium carbonate and some of the calcium carbonate is precipitated, AFTERSTANDING I N CONTACT APPLIEDWATER WITH THE I R O N AlkaAlkaSoluble iron Date linity COn Time linity COz P. p. m. 1924 p H P . p . m. P. p. m. standing p H P. p. m. P. p. m. Fe 5- 3 5.0 23 2 0 . 0 20 hours 9 . 2 23 0 5- 4 30 0 24hours 9 . 2 32 0 8.0 30 0 2dnys 9.0 32 0 5- ti 8.0 5- 9 8.1 30 0 3days 9.0 32 0 5-12 8.1 30 0 3 days 9.0 33 0 5-13 i.9 30 4 lday 8.9 31 0 5-27 32 0 8.2 1 day 9.4 33 0 5-30 X.3 33 0 3days 9.6 33 .O 0.1 6- 1 34 0 8.2 2days 9.5 35 0 6- 4 8.4 34 0 3days 9.6 32 0 6- 9 8.0 34 0 5days 9.6 32 0 0.1 6- 9 38 0 8.7 30minutes9.0 36 0 0.1 6-17 8.6 38 0 8 days 9.6 27 0 6-28 8.8 40 0 11days 9.626 0 7- 6 8.1 38 0 8days 9.6 25 0 7-10 8.3 40 0 4days 9.4 26 0 0.1 7-13 8.6 41 0 3days 9.620 0 0.1 7-14 9.3 43 0 lday 9.625 0 0.1 7-15 8.8 43 0 lday 9.6 22 0 0.1 7-16 9.0 44 0 lday 9.4 30 0 0.1 7-17 8.0 40 0 1 day 9.2 28 0 0.1 7-13 8.2 40 0 lday 8.6 29 0 0.1 7-10 8,2 38 0 1 day 9.0 29 0 0.1
.
T a b l e V-Reduction of Free a n d H a l f - B o u n d COz b y C o r r o d i n g Iron Twenty grams of iron lathe turnings from a black wrought-iron pipe were placed into a 250-cc. tightly stoppered glass tube. T h e applied water was filtered water before the application of lime and was approximately saturated with dissolved air a t a temperature of b2' C. The quantity of hydrogen gas bubbles liberated is only approximate. -APPLIED WATER-SOLUTION WITHDRAWN FROM TUBEAlkaTime AlkaSoluble linity COz standing linity COz iron Hn gas p H P . p . m . P . p . m . Days p H P , p . m . P . p . m . P , p . m . F e Cc. 6.6 30 9.7 1 7.4 34 3.5 0 6.6 30 9.7 2 7.5 32 2.6 0.60 0 0 0.30 0.5 7.0 38 4.6 4 8.0 36 6.6 31 9.7 5 2.6 37 1.0 0.05 0.5 6.6 31 9.7 7 1 . 3 34 2.0 0.20 1.0 6.5 31 11.4 7 7.5 36 5.7 2.80 1.0 6.6 33 11.9 6 7.5 35 4.8 2.70 1.0 6.6 36 10.1 11 7.4 34 6.0 1.60 3.0 6.6 35 11.4 14 7.9 38 1.0 0.40 5.0 6.9 40 6.0 14 8.5 44 0 0.05 5.0 6.8 37 7.5 30 10.0 5Za 0 0.05 10.0 6.8 39 5.1 7 9.4 44 0.05 1.0 0 6.8 39 6.2 11 9.5 43 0 0.05 2.0 a Most of the alkalinity was NazCOa caused from corrosion of the glass tube, for the soap hardness was only 15;6 as compared with 40.0 for the applied water.
373
The amount of carbon dioxide adsorbed for a certain amount of the iron hydroxides depends largely on t h e hydrogen-ion concentration of the solution. When a certain amount of carbon dioxide is introduced into a solution containing metallic iron and the precipitated compounds of iron, part of it is rapidly adsorbed and some enters into chemical combination, probably with soluble ferrous hydroxide, forming ferrous carbonate or the bicarbonate. An equilibrium with the hydroxides and the concentration of carbon dioxide in the solution is soon reached in which no additional carbonates are adsorbed. The pH of the solution at this equilibrium is at a point where carbon dioxide will continue to combine slowly with the iron to form ferrous carbonate. This reduces the hydrogen-ion concentration of the solution, upsetting the equilibrium of the adsorbed carbon dioxide, and causing it to be released. I n this manner practically all carbon dioxide introduced into a solution exposed to metallic iron will eventually unite to form ferrous carbonate if certain negative ions, such as the sulfates and chlorides, are absent or present in only very small quantities. The adsorption of the sulfate and chloride ions will be treated more fully later. Ferrous carbonate may not be formed in the presence of dissolved oxygen, but dissolved oxygen probably never reaches the surface of corroding iron after a thin film of iron hydroxide has formed. If ferrous carbonate is formed in the absence of dissolved oxygen beneath a coating of iron rust, which it most assuredly does, and then diffuses to the surface where dissolved oxygen is present, the carbon dioxide released, a n d , this carbon dioxide, probably together with more that is naturally in the water, is' driven back towards the iron surface by the difference in electrical potential of the iron and water, the solubility of ferrous carbonate becomes of considerable importance in determining corrosion rates. Whitman and his co-workers have assumed that the pH within the film of iron rust is above that a t which ferrous carbonate is only very slightly soluble. The writer finds no proof that this is true with most natural waters, except possibly a t the start of film formation. If the surrounding water has a pH of 6.0, the solution within the coating of iron rust against the iron surface may also have a pH near this point in many instances when sulfate or chlorine ions are present in the water, but if carbonate ions are the only negative ions present other than the (OH)- ions the pH will be somewhat higher. It is usually the attraction to the metal surface of the negative ions other than the (OH)- ions and their combination with the ferrous ions that causes the increase in pH, and not the concentration of soluble ferrous hydroxide as reported by Whitman, Russell, and Altieri.9 It is not to be inferred that the pH of ferrous hydroxide is not about 9.4 to 9.6, as given by these authors and by Shipley, McHaffie, and Clare,lo but that the increase to such a high point by corroding iron is not ordinarily caused by soluble iron compounds. Tables IV, VI, and VI1 show this very clearly, for there is the increase in pH without the presence of an appreciable amount of iron. The importance of such a phenomenon in the role that corroding iron itself plays in forming protective coatings becomes apparent, especially when we note the tremendous reduction in alkalinity in many instances. I n such cases there is a precipitation of calcium carbonate because the corroding iron has removed all the free carbon dioxide and enough of the half-bound carbon dioxide to increase the pH above the equilibrium point of calcium carbonate. The problem in treatment of water to prevent corrosion is to produce conditions nearest ideal for the precipitation of all the products of corrosion a t the point where corrosion is taking place-that is, where every tendency to corrode will quickly form insoluble compounds. Much depends upon the characteristics of the precipitate formed, so that con-
374
INDUSTRIAL A N D ENGINEERING CHEMISTRY
Vol. 18, No. 4
Table VI-Reduction of Alkalinity b y Corroding Iron w h e n t h e Water I s Alkaline About 50 grams of lathe turnings from wrought-iron pipe were placed into a 500-cc. Pyrex beaker and tightly stoppered. An air space of about 1 inch was left. The water was frequently agitated, except when it stood overnight. APPLIEDWATER Minutesiron -WATER AFTERSTANDING IN CONTACT WITH IRON Alkalinity, P. p. m. exposed to Alkalinity, P. p. m. Sqluble Date PhenolSoap the air Time PhenolSoap 1ron 1924 phthalein M.O. pH hardness while wet standing phthalein M.O. pH hardness P . p , m . Fe 12-22 0 39 7.2 38 5 80 minutes 1 30 8.2 29 0.2 12-22 0 39 7.2 38 5 30 minutes 2 25 9.0 22 0.1 12-23 0 39 7.1 38 5 18 hours 7 24 9.4 22 0.1 12-23 12 51 9.4 52 5 20 minutes I9 42 9.6444 0.1 12-23 12 51 9.4 62 5 90 minutes 11 40 12-23 12 51 9.4 62 60 10 minutes 6 36 9.2 33 0.2 20 minutes 7 31 9.4 27 0.2 12-23 12 51 9.4 62 6 60 minutes 15 32 9.6+ 33 0.1 12-23 11 50 9.4 50 6 60 minutes 11 32 9.6 29 0.1 12-24 11 50 9.4 60 5 18 hours 11 23 9.6+ 23 0.1 12-24 11 50 9.4 50 30 20 minutes 10 30 9.6 26 0.1 12-25 11 50 9.4 50 5 22 hours 11 25 9.6422 0.2 12-26 11 50 9.4 48 5 22 hours 5 26 9.2 25 0.1 12-26 10 50 9.4 48 45 3 hours 12 24 9.6+ 23 0.1 12-27 11 51 9.4 48 90 20 hours 5 26 8.8 22 0.2
.
centration cells will not be formed and all corrosion will not be stopped. Ferrous carbonate apparently aids in protecting iron surfaces where the p H and alkalinity are above calcium carbonate equilibrium, especially if it is overlaid by some kind of a coating that keeps out the dissolved oxygen whereas a t a lower pH it is much more soluble. At a pH and alkalinity where ferrous carbonate is nearly insoluble water will not stain white bathroom fixtures or clothing being laundered, and therefore under such conditions the most objectionable result of corrosion is removed. Table VII-Chemical Changes Produced in the Baltimore Water w h e n Exposed to Wrought-Iron Pipe Filtered water at the Montebello Filters both before and after the addition of lime was allowed t o pass slowly through 50 feet of l/r-inch black wrought-iron pipe for 15 months before starting the tests. Approximately 300 cc. of water were withdrawn each time for test, con quently according t o the theoretical capacity of the pipe the actual time of %anding was longer than the time between tests. All water was filtered through three thicknesses of thoroughly washed No. 40 Whatman paper in such a manner as t o allow practically none of the water to come in contact with the air. -BEFORE PASSINGPIPE- -AFTER STANDING I N PIPEDis- Time Alka- solved standAlka- Soluble Temp. Cog linity oxygen ing COz linity iron OC. pH P . p . m . P . p . m . P . p . m . D a y s p H P . p . m . P . p . m . P . p . m . 22 6.6 10.6 40 4.9 7.2 3.1 27 0.20 20 6.3 30 6.4 7.2 3.1 0.30 17.0 26 19 6.3 31 3.2 7.4 2.2 0.10 26 18 37 6.5 7.4 2.2 21 13.2 0.20 6.3 16 6.5 34 6.5 7.6 1.3 23 15.4 0.30 16 6.4 35 8.1 8.0 0 20 14.5 0.15 14 8.8 56 7.2 17 0.05 0 8.6 0 38 9.3 17 0.05 10 6.6 11.0 8.5 0 36 9.8 15 10 6.5 9.0 8.4 0 0.05 38 10.4 20 16 6.6 8.4 6 8.4 0 0.05 39 Trace 14 6.8 6.0 10.5 16 6 8.6 0 4s 9.4 0 0.05 19 7.6 3.3 3 20 47 6.2 3 20 9.3 0 0.10 21 7.5 47 21 4.8 3 9.2 0 0.05 7.5 20 50 18 7.5 4.9 5 21 9.2 0 Trace 50 17 7.5 6.4 5 20 9.2 0 Trace 36 17 6.6 3.2 20 9.2 0 0.10 54 7.8 6.5 18 9.4 0 15 0.25 35 6.6 6.5 16 0.20 9.2 0 17 58 7.8 8.1 15 0.35 14 8.4 0 8.8 7.2 13 0.40 13 8.5 0 56 51 8.1 9.3 0.40 13 8.5 0 10 7 0.15 52 9.8 13 10 8.6 8.4 0 6 50 10.4 20 17 8.2 8.5 0 0.15 49 10.5 16 12 8.2 0.25 6 8.4 0
i
Compounds Formed in the Corrosion of Iron
Recent literature on corrosion gives little consideration to the formation of compounds other than ferrous and ferric hydroxides. This may lead some to believe that they are the only compounds formed, whereas it is probably the intention of the authors to show that corrosion can progress indefinitely with the formation of only these two compounds. That other compounds are formed, and frequently in large proportions, is well known to many students of corrosion. Examination of the products of corrosion reveals a number of iron compounds. Some of the iron rust is gelatinous in appearance and much of it is crystalline. I n using the terms “gelatinous” and “crystalline” it is not intended to convey the impression that the gelatinous compounds are not composed of a meshwork of very minute crystals. Nearly all the rust forming
in the absence of oxygen is black and quite magnetic, and is probably an oxide of iron rather than the hydroxide. (Figures 3, 8, and 12) The soluble products extracted from tubercles overlying pits show the presence of sulfates, carbonates, and chlorides of iron. Sometimes the amount of sulfates and chlorides extracted from tubercles is over 1 per cent of the weight of the entire wet tubercle. A little calcium and magnesium are frequently present on the inside of tubercles if the exposure has been to water that is fairly alkaline, but the amounts of these constituents account for only a small percentage of the negative ions. A large percentage of these ions are combined with iron but part may be adsorbed by the solid constituent of the tubercles. There is good evidence that practically any compound of iron may be produced by corrosion if the necessary negative ions are present and the hydrogen-ion concentration is favorable. This does not mean that the first product of corrosion is other than ferrous hydroxide, for, although there is some evidence to the contrary, the writer hasnotfoundpositive , proof. The concentration in tubercles of such large quantities of iron salts such as sulfates and chlorides is very significant and will be treated more fully in a later parag r a p h . Heretofore many chemists have adhered strongly to the idea that only d e f i n i t e compounds are formed in a crystalline p r e ci p i t a t e . Evidence is accumu- Figure 3-Particles of Black Iron R u s t Attracted b y Metallic Iron lating to show that The attractive force is quite noticeable along c r y s t a l l i n e the entire surface, but appears to be somewhat many greater at the ends. When near metallic iron compounds may have the suspended particles are usually oriented as shown by the drawing. various amounts of other ions present. That this is in the iron oxide compounds seems possible, especially when we try to account for definite compounds on the basis of chemical analyses. Does Ferrous Hydroxide Precipitate during Corrosion?
Whitman, Russell, and Altierig state that “the corroding metal first forms ferrous ions which are precipitated as ferrous hydroxide.” This implies that we might expect a supersaturation of soluble iron sufficient to produce a precipitate in the absence of oxygen before the hydrogen overvoltage, or the concentration of H (nascent) or Hz gas, is sufficient to stop solution of the metal. There are indications that
I N D USTRI-4L ,4-VD ENGINEERING CHEMISTRY
April, 1926
something of this nature happens. However, the saturation point of ferrous hydroxide does not seem to be nearly so great under most natural conditions of corrosion as is reported by several authors. Spellers gives the same impression by assuming the formation of nascent hydrogen in solution and that there is a tendency for it to form hydrogen gas. If hydrogen gas is actually produced to the extent of being liberated as bubbles, which is the J < case when iron is exposed to most natural waters containing no dissolved oxygen, then it seems that it might come close to the saturation point in acid-free distilled water. This apparently is not the case, for experiments conducted under almost complete vacuum showed no appreciable increase in corroFigure &Device Used for Measuring the sion rates over what Gas Liberation of Corroding Iron m i" e h t b e exDected a t atmospheric pressure. If an appreciable amount of hydrogen gas were produced it certainly would have been released under the vacuum. From this we might assume that practically no hydrogen gas is formed in the absence of oxygen when iron is exposed to water free from any negative ions other than the (OH)ions. Can the reaction 2 H f = 2H @ take place to any material extent without the formation of Hz gas? This may be questioning the electrochemical theory of corrosion as now explained, which is not the purpose of this article, but the theory apparently does not fully account for some of the known facts and needs additional study. The writer has been unable to produce appreciable amounts of what is believed to be ferrous hydroxide by corroding iron, unless it is a crystalline and not a gelatinous precipitate. To show the influence of dissolved oxygen when fresh metal surfaces are exposed, more precipitated oxides of iron can be produced in 30 minutes by agitating iron lathe turnings in a flask partially filled with water than will be produced in a week under nearly complete vacuum. It seems possible that there are two routes by which corrosion of iron takes place. One is strictly electrochemical as now generally explained, in which oxygen is necessary to convert the hydrogen to water and the ferrous hydroxide to some compound of iron of greater oxygen content and less soluble than ferrous hydroxide. The other route is either by first combining with some negative ion such as (COS)-- or (SO,)--, and later being converted to ferrous or ferric hydroxide when the soluble salt diffuses out to water containing dissolved oxygen, or the combination with negative ions other than the (OH) - ions may merely be a secondary reaction in which ferrous hydroxide is still the first compound formed. With a high concentration of certain negative ions it is possible that the combination may be direct in some instances, though theoretically it probably should not be. Whatever be the process by which acids combine with iron, this process accounts for most of the corrosion taking place in water pipes that have been in service for a number of years, The miter considers this to be of fundamental importance, and contrary to the views of most recent students of corrosion. It is not essential for initial corrosion, but unless it takes place in this manner iron surfaces soon become so passive that corrosion is almost completely stopped. Y
+
375
From the large number of experiments conducted it is concluded that ferrous hydroxide is not ordinarily precipitated in the natural process of the corrosion of iron unless certain negative ions are present. If this were true i t would be possible for corrosion to progress in the absence of oxygen merely by the liberation of hydrogen gas, which is not the case in the absence of negative ions other than the (OH)ions a t a temperature below 22" C. It is possible for considerable ferrous hydroxide, or the ferrous ions, to be adsorbed by ferric hydroxide, and this no doubt accounts for the throwing out of solution of appreciable quantities of soluble ferrous iron when considerable ferric hydroxide is present. There is also the possibility of a n electrical double layer in which ferrous ions are first adsorbed, giving the surface a positive charge, then a layer of negative ions, if any are present. Influence of the Negative Ions
When iron is submerged a t room temperature in distilled water, practically free from dissolved oxygen and negative ions other than the (OH)- ions, no hydrogen gas is liberated as gas bubbles even after several months' storage. If to this same water a small amount of some salt such as calcium sulfate or sodium chloride is added, there will usually be a liberation of hydrogen gas to the point of producing gas bubbles. With some form of iron which has a large surface area, such as lathe turnings, the bubbles will usually begin to form within 3 or 4 days. The addition of such salts causes the p H value of the water to rise and equilibrium will be reached a t a pH between 8 and 10, usually near 9. At this point very little soluble iron will be found in the M M. solution, usually less than 0.1 part per mil3 0 MINUTES lion. This is also a t a point where no ferrous sulfate or chloride is supposed to exist; yet hydrogen 4 HOU R S gas may continue to be liberated in small quantities for some time, if not indefinitely. The addition of bicarbonates 8 HOURS such as sodium or calcium bicarbonate will c a u s e considerable liberation of hydrogen gas and a n in2 4 HOURS crease in pH to over 9.6 if a s u f f i c i e n t a m o u n t is added. The reaction in this Figure 5-Characteristic Formation of case can be explained Iron Hydroxide Precipitates o n Submerged Iron as the formation of an equilibrium between the bicarbonates and ferrous carbonate. The continued liberation of hydrogen gas under some conditions is difficult to explain on the basis of the assumption that ferrous hydroxide is the first product of corrosion unless there is some reaction between the Fe++ and certain negative ions, for pure water will not liberate hydrogen gas to the extent of producing bubbles. If it could be assumed that the hydrogen gas equilibrium with pure water were nearly at the point of gas liberation, then it would not be difficult to assume gas liberation by a slight change in the concentration of soluble salts. However, tests under almost complete
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INDUSTRIAL A N D ENGINEERING CHEMISTRY
vacuum indicate that water containing no negative ions other than the (OH)- ions does not produce hydrogen gas in more t h a n very minute quantities, if at all. The apparatus shown in Figure 4 was used to measure the r a t e of gas liberation a t atmospheric pressure. Lathe turnings or other specimens of iron and steel were placed into d 5 0 100 Pyrex flasks, the solition to de tested added, and the flask sealed with no air space. It is believed that this gives very good protection from the outside g a s e s , especially after gas bubbles begin to be liberated. To reach t h e inside solution the gas must be dissolved by the water in the beaker, diffused through a catillarv * " Figure 6-Fibrous Formation of Iron Rust tube "several inches a t Surface of Tubercle l o n g , t h e n through t h e measuring tube which is partially filled with hydrogen &s i f any is being liberated. There is yet a mercury seal and another gas gap to pass. With these precautions it is possible for the hydrogen gas liberated to pass from the flask without the inside solution becoming contaminated from the outside. A number of natural waters were tried and all produced hydrogen gas bubbles when some precipitated iron hydroxide was present with the metallic iron. A few tests, however, with natural waters low in soluble salts, deaerated, and new lathe turnings failed to produce the gas bubbles. These natural waters varied from a very low concentration of dissolved compounds to very hard waters containing nearly 1000 parts per million. Waters with high concentrations u s u a l l y produced hydrogen gas faster than those with the lower c o n c e n t r a tions, b u t t h e amount of free and half-bound carbon d i o x i d e had the greatest influence. Gas formed by the presence of negative ions other than the (OH)- ions beneath a coating of iron rust is believed to have considerable effect on corrosion rates .when. the surPipe submerged in water having pH of 8.0 face has been exposed for some time, for it is largely independent of the dissolved oxygen in the surrounding water. The measurement of corrosion rates by the loss of dissolved oxygen, as proposed by Speller and Kendall,15does not account for corrosion where there is hydrogen gas liberation, and it is believed that this constitutes a large amount of the corrosion taking place in old pipes. We have the outstanding fact that iron submerged in pure water soon becomes very passive, even though there is plenty of dissolved oxygen in the solution, whereas i f negative ions such as the carbonates, sulfates, and chlorides ~~
are present it does not become nearly so passive. Fujihara's finds this to be true for (C03)--, and Shipley, McHaffie, and Clare'O apparently observe the same phenomenon with soil containing considerable calcium sulfate. If the explanation is offered that the negative ions merely change the potential difference, as is supposed to be the case with acids, then we are confronted with the disappearance of a small amount of these ions from the solution in the case of sulfates and a large amount in the case of carbonates. If the (COS)-- only ions disappeared the cause might be attributed to the formation of very insoluble ferrous carbonate, but when the sulfate and chlorine ions are also reduced and there is hydrogen gas liberation at a pH between 9 and 10, then it is hard to explain a direct combination with ferrous hydroxide, for neither ferrous chloride nor ferrous sulfate exists a t such a high pH. Possibly, there are places somewhat of the nature of pits where the pH is still below the point where considerable soluble iron exists and the hydrogen gas is liberated a t these points. The power to concentrate the negative ions against the metal surface beneath a coating of iron rust, making the solution a t this point slightly acid rather than alkaline as has been assumed by several authors, offers a more plausible explanation. The influence of negative ions, especially certain polyvalent ions such as sulfates, upon precipitation of the iron hydroxides has been studied by Miller" and others. Miller finds that the ratio between the amount of ferrous and ferric hydroxide precipitated and the amount of sulfates precipitated with the hydroxides varies with the pH of the solution. Considerably more sulfates are precipitated with ferrous hydroxide a t a certain pH than with ferric hydroxide. At a pH of 7.5, the mol ratio of SO4 to Fe is about 0.100 for ferrous hydroxide and about 0.010 for ferric hydroxide. When these hydroxides are being produced by corroding iron, any negative ions present will tend to combine with or be adsorbed by them; that is, if the solution contains dissociated Fe++ and dissociated (SO4)--, as will be the case if neutral sulfate salts are present, it is possible under certain conditions of concentration and pH for these ions to unite. Such a combination, however, is not generally regarded as taking place. If we assume the dissociation of ferrous sulfate into Fe++ and (SO*)-- ions, and also assume that Fe is being thrown off from the metal and immediately takes a positive charge from the hydrogen, then must we assume that this same Fe + + first combines with (OH)- as it takes up the positive charge only to be immediately replaced with sulfate ions? The writer has no objection to such an interpretation, as it probably makes little difference on corrosion rates which way the reaction takes place. If the ferrous ions unite only with (OH)- ions to form ferrous hydroxide, and this compound is precipitated as is assumed by some writers, it then possesses a positive charge and negative ions will be adsorbed. When the ferrous hydroxide has adsorbed negative ions to its equilibrium with the pH and the concentration of dissolved salts in the water, and it is then oxidized, the equilibrium will probably be different and negative ions will be given off. Regardless of how the change from ferrous compounds to ferric hydroxide takes place, negative ions are released if any neutral or newly neutral salts such as sulfates or chlorides are present in the solution. These negative ions tend to be drawn toward the metal surface by the difference in electrical potential, and when reaching the zone of no dissolved oxygen they again encounter ferrous ions and the process is repeated. Negative ions, such as the sulfates and chlorides, which form highly soluble compounds tend t o concentrate iron salts near the metal surface and thus bring about corrosion by the liberation of hydrogen gas even though dissolved oxygen may be present in the nearby solution. The 16
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THISJOURNAL, 16, 134 (1923).
Vol. 18, No. 4
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Chcm. Met. Eng., 39, 810 (1925). Public Health RPts., 40, 1413 (1925).
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