Effect of Dissolved Salts

7.2.6 Effect of Dissolved Salts

The effect of sodium chloride concentration on corrosion of iron in air - saturated water at room temperature is shown in Fig. 7.11 . The corrosion rate fi rst increases with salt concentration and then decreases, with the value falling below that for distilled water when saturation is reached (26% NaCl).

Since oxygen depolarization controls the rate throughout the sodium chlo- ride concentration range, it is important to understand why the rate fi rst increases, reaching a maximum at about 3% NaCl (seawater concentration), and then decreases. Oxygen solubility in water decreases continuously with sodium chlo- ride concentration, explaining the lower corrosion rates at the higher sodium chloride concentrations. The initial rise appears to be related to a change in the protective nature of the diffusion - barrier rust fi lm that forms on corroding iron. In distilled water having low conductivity, anodes and cathodes must be located

132 IRON AND STEEL

Figure 7.11. Effect of sodium chloride concentration on corrosion of iron in aerated solu- tions, room temperature (composite data of several investigations).

relatively near each other. Consequently, OH − ions forming at cathodes in accor- dance with [cf. Eqs. (7.3) and (7.4) , Sec. 7.1 ]

O + − HO → 2 OH − 2 e 2 − 2 (7.15)

are always in the proximity of Fe 2+ ions forming at nearby anodes, resulting in a fi lm of Fe(OH) 2 adjacent to and adherent to the metal surface. This fi lm provides an effective diffusion - barrier fi lm. In sodium chloride solutions, on the other hand, the conductivity is greater; hence, additional anodes and cathodes can operate much further removed one from the other. At such cathodes, NaOH does not react immediately with FeCl 2 formed at anodes; instead, these substances diffuse into the solution and react to form Fe(OH) 2 away from the metal surface. Any Fe(OH) 2 so formed does not provide a protective barrier layer on the metal surface. Hence, iron corrodes more rapidly in dilute sodium chloride solution because more dissolved oxygen can reach cathodic areas. Above 3% NaCl, the continuing decreased solubility of oxygen becomes more important than any change in the diffusion - barrier layer; hence, the corrosion rate decreases.

Alkali - metal salts (e.g., KCl, LiCl, Na 2 SO 4 , KI, NaBr, etc.) affect the corro- sion rate of iron and steel in approximately the same manner as sodium chloride. Chlorides appear to be slightly more corrosive in the order Li, Na, and K [29] .

Alkaline - earth salts (e.g., CaCl 2 , SrCl 2 ) are slightly less corrosive than alkali - metal salts. Nitrates appear to be slightly less corrosive than chlorides or sulfates at low concentrations (0.2 – 0.25 N ), but not necessarily at higher concentrations [30] . The small differences for all these solutions may arise, for example, from their specifi c effect on the Fe(OH) 2 diffusion - barrier layer, or perhaps from the different adsorptive properties of the ions at a metal surface resulting in differing anode – cathode area ratios or differing overvoltage characteristics for oxygen reduction.

Acid salts , which are salts that hydrolyze to form acid solutions, cause cor- rosion with combined hydrogen evolution and oxygen depolarization at a rate

AQUEOUS ENVIRONMENTS

paralleling that of the corresponding acids at the same pH value. Examples of such salts are AlCl 3 , NiSO 4 , MnCl 2 , and FeCl 2 . Ammonium salts (e.g., NH 4 Cl) are also acid, but produce a higher corrosion rate than corresponds to their pH (at NH 4 Cl concentrations > 0.05 N ) [30] . Increased corrosivity is accounted for by the ability of NH 4+ to complex iron ions, thereby reducing activity of Fe 2+ and increasing the tendency of iron to corrode. Ammonium nitrate in high concentra- tion is more corrosive (as much as eight times more) than the chloride or sulfate,

in part because of the depolarizing ability of NO − 3 .

In the presence of excess NH 3 , common to some synthetic fertilizer solutions, the corrosion rate in ammonium nitrate at room temperature may reach the very high value of 50 mm/y (2 ipy). The complex formed in this case has the structure

[Fe(NH 3 ) 6 ](NO 3 ) 2 [31] . Since coupling mild steel to an equal area of platinum has no effect on the rate, the reaction is apparently anodically controlled. Metal- lurgical structure affects the rate, a cold - worked mild steel reacting much more rapidly than one quenched from elevated temperature. This indicates that the reaction is not diffusion - controlled, but depends instead on the rate of metal - ion formation at the anode and perhaps also to some extent on the rate of depolar- ization at the cathode.

Alkaline salts , which hydrolyze to form solutions of pH > 10, act as corrosion inhibitors. They passivate iron in the presence of dissolved oxygen in the same manner as NaOH (Fig. 7.3 , Section 7.2.3 ). Examples of such salts are trisodium

phosphate (Na 3 PO 4 ), sodium tetraborate (Na 2 B 2 O 7 ), sodium silicate (Na 2 SiO 3 ), and sodium carbonate (Na 2 CO 3 ). In addition to favoring passivation of iron by dissolved oxygen, they may form corrosion - product layers of ferrous or ferric phosphates in the case of Na 3 PO 4 , or analogous compounds in the case of Na 2 SiO 3 , with such layers acting as more effi cient diffusion barriers than hydrous FeO. They may, on this account, also inhibit corrosion below pH 10 and may provide

better inhibition above pH 10 than NaOH or Na 2 CO 3 .

Oxidizing salts are either (a) good depolarizers, and therefore corrosive, or (b) passivators and effi cient inhibitors. Examples of the fi rst class are FeCl 3 , CuCl 2 , HgCl 2 , and sodium hypochlorite. They represent the most diffi cult class of chemicals to handle in metal equipment. Examples of the second class are Na 2 CrO 4 , NaNO 2 , KMnO 4 , and K 2 FeO 4 . The differences in properties accounting for an oxidizing salt being either a depolarizer or a passivator are discussed in Chapter 17 .

7.2.6.1 Natural - Water Salts. Natural fresh waters contain dissolved calcium and magnesium salts in varying concentrations, depending on the source and location of the water. If the concentration of such salts is high, the water is called hard; otherwise, it is called soft. For many years before the causes were clearly understood, it was recognized that a soft water was more corrosive than

a hard water. For example, a galvanized - iron hot - water tank was observed to last

10 – 20 years before failing by pitting in Chicago Great Lakes water (34 ppm Ca 2+ , 157 ppm dissolved solids), whereas in Boston water (5 ppm Ca 2+ , 43 ppm dissolved solids), a similar tank lasted only one to two years.

134 IRON AND STEEL

The two parameters that control corrosivity of soft waters are the pH and the dissolved oxygen concentration. In hard waters, however, the natural deposi- tion on the metal surface of a thin diffusion - barrier fi lm composed largely of

calcium carbonate (CaCO 3 ) protects the underlying metal. This fi lm retards dif- fusion of dissolved oxygen to cathodic areas, supplementing the natural corrosion barrier of Fe(OH) 2 mentioned earlier (Section 7.2.3 ). In soft water, no such pro- tective fi lm of CaCO 3 can form. But hardness alone is not the only factor that determines whether a protective fi lm is possible. Ability of CaCO 3 to precipitate on the metal surface also depends on total acidity or alkalinity, pH, and concen- tration of dissolved solids in the water. For given values of hardness, alkalinity, and total dissolved salt concentration, a value of pH, given the symbol pH s , exists

at which the water is in equilibrium with solid CaCO 3 . When pH > pH s , the

deposition of CaCO 3 is thermodynamically possible.

Langelier divided natural fresh waters into two groups: those oversaturated with CaCO 3 and those undersaturated [32] . Since only near - or oversaturated waters tended to form a protective fi lm of CaCO 3 on iron, an estimate of the corrosivity of a water was established through analytical criteria for under - or oversaturation. Using certain simplifi cations, Langelier showed that the value of

pH s — at which a water is in equilibrium with solid CaCO 3 (neither tends to dis- solve nor precipitate) — can be calculated from the relation *

pH s = ( p K

2 ′− p K

s ′+ ) p Ca + p alk (7.16) where K ′ 2 is the ionization constant ( H + )( CO 2 3 − )( / HCO − 3 ) , K ′ s is the solubility

product of calcium carbonate [( Ca 2 + )( CO 2 3 − )] , the concentration of calcium ions (Ca 2+ ) is in moles/1000 g H 2 O, and alk (alkalinity) represents the equivalents per liter of titratable base to the methyl orange end point (often reported as ppm CaCO 3 ) according to the relation

( alk )( + H + ) → 2 CO 2 − + HCO − + OH 3 − 3 The letter p refers to the negative logarithm of all these quantities. The saturation

index, also known as the Langelier index, is defi ned as the difference between the measured pH of a water and the equilibrium pH s for CaCO 3 , or

Saturation index = pH measured − pH s

A positive value of the saturation index indicates a tendency for the protective CaCO 3 fi lm to form, whereas a negative value indicates that it will not form. Charts (see Appendix, Section 29.3 ) and nomographs have been prepared to obtain values of pH s for waters varying widely in composition and at various temperatures [33] .

* For a derivation of this and a more exact equation, see Section 29.3 in the Appendix.

AQUEOUS ENVIRONMENTS

Tabulated data, originally organized by Nordell, for calculating the Langelier saturation index are presented in Table 7.3 [34, 35] . After studying hundreds of concentrated cooling waters, Puckorius and Brooke [36] proposed a new index, the practical saturation index (PSI),

PSI = 2 pH s − pH e

pH e = . 1 485 × log total alkalinity) ( + . 4 54

CaCO 3 precipitates at PSI < 6.0, according to Puckorius and Brooke. An estimate of over - or undersaturation can also be obtained in the labora- tory by measuring the pH of a water before and after exposure to pure CaCO 3 powder for a time adequate to achieve equilibrium. An increase in pH corre-

T A B L E 7.3. Data for Calculating the Langelier Saturation Index [34, 35] Total Dissolved

C M. O. Alkalinity D Solids (ppm)

A Ca Hardness

(ppm as CaCO 3 ) 50 – 300

(ppm as CaCO 3 )

0.9 18 – 22 1.3 Temperature ( ° C)

2.6 890 – 1000 3.0 1. Obtain values of A , B , C , and D from this table.

2. pH = (9.3 + A + B ) − ( C + D ). 3. Langelier saturation index = pH – pH s .

136 IRON AND STEEL

sponds to undersaturation. Because the composition of the water is altered by any CaCO 3 taken up or precipitated during the test, the fi nal pH is not necessarily the same as the calculated pH s . Alternatively, the water after saturation with CaCO 3 can be titrated with acid, with the increase in required milliliters of acid compared to that required by the original water being a measure of undersatura- tion (marble test) [37] .

Any fresh water can be placed in one of the following categories:

Saturation Index

Characteristic of Water Positive

Saturation Level

Protective fi lm of respect to CaCO 3 CaCO 3 forms Zero

Oversaturated with

In equilibrium

Negative

Undersaturated with

Corrosive

respect to CaCO 3

Chicago water, for example, has an index of 0.2, whereas the value for Boston water is − 3.0.

A soft water with negative saturation index can usually be treated with lime, sodium carbonate, or sodium hydroxide, in order to raise the saturation index to

a positive value, thereby making the water less corrosive. For very soft waters, the required pH may be too high for uses such as tap water unless Ca 2+ ions are added simultaneously. A saturation index of +0.1 to about +0.5 is considered to

be satisfactory to provide corrosion protection. At higher values of the saturation index, excessive deposition of CaCO 3 (scaling) can occur, particularly at elevated temperatures. Waters with a slightly negative Langelier index may deposit CaCO 3 because of pH fl uctuations. In general, for waters of negative saturation index, the more negative the index, the more corrosive the water. The relationship between the Langelier index and the corrosion rates of water pipes from a fi eld test at a city water works is presented in Fig. 7.12 [38, 39] .

At above - room temperatures, the saturation index may become more posi- tive, and, in any case, the rate of CaCO 3 deposition is higher, should there be any tendency toward deposition. Hence, corrosion protection at all temperatures without excessive scaling requires adjustment of water composition to a satura- tion index that is at least constant over the entire operating temperature range. Powell et al. [33, 40] showed that, for any specifi c alkalinity, there is a correspond- ing pH value at which the decrease of measured pH with temperature is almost

exactly compensated by a decrease in the factor ( p K 2 ′− p K s ′ ) . Under these condi- tions, the saturation index is nearly constant with temperature, and scaling tends to be the same in hot or cold water (Table 7.4 ). The amount of possible scaling, of course, depends on the value of the saturation index. The alkalinity and the pH of waters can be adjusted to bring them into the proper composition range

using Ca(OH) 2 , Na 2 CO 3 , NaOH, H 2 SO 4 , or CO 2 .

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Figure 7.12. Relationship between the Langelier index and the corrosion rates of water pipes. (Reproduced from Fujii [39] with permission of Japan Association of Corrosion Control.)

T A B L E 7.4. Alkalinity – pH Limits for Uniform Scale Deposition at Various Temperatures [40]

Alkalinity (ppm as CaCO 3 )

pH Measured at Room Temperature

Limitations of Saturation Index. If a natural water contains colloidal silica or organic matter (e.g., algae), CaCO 3 may precipitate on the colloidal or organic particles instead of on the metal surface. If such is the case, the corrosion rate will remain high even though the saturation index is positive. For waters high in

dissolved salts, such as NaCl, or for waters at elevated temperatures, the CaCO 3 fi lm may lose its protective character at local areas, resulting in corrosion pitting. Furthermore, if complexing ions are added to a chemically treated water, such

as polyphosphates, which retard precipitation of CaCO 3 , the saturation index may no longer apply as an index of corrosivity. Other than these exceptions, the saturation index is a useful qualitative guide to the relative corrosivity of a fresh water in contact with metals for which the corrosion rate depends on diffusion of dissolved oxygen to the surface, such as iron, copper, brass, and lead. The index does not apply to corrosivity of a water

IRON AND STEEL

in contact with passive metals that corrode less the higher the surface concentra- tion of oxygen, such as aluminum and the stainless steels.