Effect of Dissolved Oxygen

7.2.1 Effect of Dissolved Oxygen

7.2.1.1 Air - Saturated Water. In neutral or near - neutral water at ambient

AQUEOUS ENVIRONMENTS

air - saturated water, the initial corrosion rate may reach a value of about 10 gmd. This rate diminishes over a period of days as the iron oxide (rust) fi lm is formed and acts as a barrier to oxygen diffusion. The steady - state corrosion rate may be

1.0 – 2.5 gmd, tending to be higher the greater the relative motion of water with respect to iron. Since the diffusion rate at steady state is proportional to oxygen concentration, it follows from (7.3) that the corrosion rate of iron is also propor- tional to oxygen concentration. The concentration of dissolved oxygen in air -

saturated, natural fresh waters at ordinary temperatures is 8 – 10 ppm. In the absence of dissolved oxygen, the corrosion rate at room temperature is negligible both for pure iron and for steel.

In the absence of diffusion - barrier fi lms of corrosion products on the surface, the theoretical corrosion current density, i (A/cm 2 ), of steel in stagnant, air - saturated fresh water is [cf. Eq. (5.5) , Section 5.4 ]

DnF

− = 3 c × 10 (7.6)

where D is the diffusion coeffi cient for dissolved oxygen in water (cm 2 /s), δ is the thickness of the stagnant layer of electrolyte next to the electrode surface, c is the concentration of oxygen, and the other terms have their usual signifi cance. Using D = 2

× 10 −5 2 cm /s, n = 4 eq/mole, F = 96,500 C/eq, δ = 0.05 cm, and c = 8/32 × 10 −3 mole/liter at 25 ° C, we obtain i = 38.6

× 10 −6 A/cm 2 , corresponding to a cor- rosion rate of 0.45 mm/y [1] .

7.2.1.2 Higher Partial Pressures of Oxygen. Although increase in oxygen concentration at fi rst accelerates corrosion of iron, it is found that, beyond a critical concentration, the corrosion drops again to a low value [2] . In distilled water, the critical concentration of oxygen above which corrosion decreases

again, is about 12 mL O 2 /liter (Fig. 7.1 ). This value increases with dissolved salts and with temperature, and it decreases with an increase in velocity and pH. At pH of about 10, the critical oxygen concentration reaches the value for air -

saturated water (6 mL O 2 /liter) and is still less for more alkaline solutions. The decrease of corrosion rate is caused by passivation of iron by oxygen, as shown by potentials of iron in air - saturated H 2 O of − 0.4 to − 0.5 V versus S.H.E. and 0.1 to 0.4 V in oxygen - saturated H 2 O (28 mL O 2 /liter). Apparently, at higher partial pressures, more oxygen reaches the metal surface than can be reduced by the corrosion reaction — the excess, therefore, is available to form the passive fi lm; increasing the cathodic reaction rate by higher oxygen concentration increases polarization of anodic areas until the critical current density for passiv- ity is reached (see Fig. 6.1 , Section 6.2 ) [3] .

Because passivity accompanies higher oxygen pressures, passive – active cells are established in the event that passivity breaks down locally (e.g., at crevices). Such breakdown is accompanied by severe pitting, particularly at higher tem- peratures, in the presence of halide ions, or at a critical pressure of oxygen where passivity is on the verge of either forming or breaking down. This behavior

118 IRON AND STEEL

Figure 7.1. Effect of oxygen concentration on corrosion of mild steel in slowly moving dis- tilled water, 48 - h test, 25 ° C [2] . ( Reproduced with permission. Copyright 1955, The Electro- chemical Society .)

corrosion of steel. In appreciable concentration of chlorides, as in seawater, pas- sivity of iron is not established at all, and in such media increased oxygen pressure results in an increased corrosion rate.

7.2.1.3 Microbiologically Infl uenced Corrosion. Microbiologically infl u- enced corrosion (MIC) is corrosion that is caused by the presence and activities of microorganisms — that is, organisms that cannot be seen individually with the unaided human eye, including microalgae, bacteria, and fungi [4] . Microbiologi- cally infl uenced corrosion can cause various forms of localized corrosion, includ- ing pitting, dealloying, enhanced erosion corrosion, enhanced galvanic corrosion, stress corrosion cracking, and hydrogen embrittlement [4] . As a result of MIC, corrosion can occur at locations where it would not be predicted, and it can occur at very high rates. All engineering alloys, with the exception of titanium and high chromium – nickel alloys, have been reported to undergo MIC. Furthermore, MIC has been documented to take place in seawater, fresh water, distilled/demineral- ized water, hydrocarbon fuels, process chemicals, foods, soils, human plasma, saliva, and sewage [4] .

Although sulfate - reducing bacteria (SRBs), active only in anaerobic (oxygen - free) environments, are a very common cause of corrosion and have been exten- sively studied, MIC can also be caused by other types of microorganisms — for example, Thiobacilli [sulfur - oxidizing bacteria (SOB), which oxidize sulfur com- pounds to sulfuric acid] and other acid - producing microorganisms, including both bacteria and fungi.

AQUEOUS ENVIRONMENTS

Microbes can adhere to metal surfaces forming a biofi lm, consisting of a community of microorganisms, leading to corrosion [4, 5] . When the acidic prod- ucts of bacterial action are trapped at the biofi lm – metal interface, their impact on corrosion is intensifi ed [6] .

Although iron does not corrode appreciably in deaerated water, the corro- sion rate in some natural deaerated environments is found to be abnormally high. These high rates have been traced to the presence of sulfate - reducing bacteria (e.g., Desulfovibrio desulfuricans ). Their relation to an observed accelerated cor-

rosion rate in soils low in dissolved oxygen was fi rst observed by von Wolzogen K ü hr in Holland [7] . The bacteria are curved, measuring about 1

× 4 μ m, and are found in many waters and soils. They thrive only under anaerobic conditions in the pH range of about 5.5 – 8.5. Certain varieties multiply in fresh waters and in soils containing sulfates, others fl ourish in brackish waters and seawater, and still others are stated to exist in deep soils at temperatures as high as 60 – 80 ° C (140 – 175 ° F).

Sulfate - reducing bacteria easily reduce inorganic sulfates to sulfi des in the presence of hydrogen or organic matter, and they are aided in this process by the presence of an iron surface. The aid that iron provides in this reduction is prob- ably to supply hydrogen, which is normally adsorbed on the metal surface and

which the bacteria use in reducing SO 4 . For each equivalent of hydrogen atoms they consume, one equivalent of Fe 2+ enters solution to form rust and FeS. The bacteria, therefore, probably act essentially as depolarizers.

A possible reaction sequence can be outlined as follows: Anode :4 Fe → 4 Fe 2 + + 8 e − (7.7) Cathade :8 HO + 8 e 2 − → 8 H ads on Fe + 8 OH − (7.8)

8 H ads bacteria + Na SO 2 4 ⎯ ⎯⎯⎯ → 4 2 H O Na S + 2 (7.9)

Na S 2H CO 2 + 2 3 → 2 NaHCO 3 + HS 2 (7.10) Summary : 4 Fe + 2 2 H O Na SO + 2 4 + 2 2 H CO 3 → 3 Fe OH ( ) 2 + FeS + 2 NaHCO 3

(7.11) Ferrous hydroxide and ferrous sulfi de are formed in the proportion of 3 moles

to 1 mole. Analysis of a rust in which sulfate - reducing bacteria were active shows this approximate ratio of oxide to sulfi de. Qualitatively, the action of sulfate -

reducing bacteria as the cause of corrosion in a water initially free of sulfi des can

be detected by adding a few drops of hydrochloric acid to the rust and noting the smell of hydrogen sulfi de.

120 IRON AND STEEL

In addition to the cathodic reaction listed in (7.8) , other cathodic reactions

could be considered [8, 9] , such as the reduction of H 2 S:

2 (7.12) Severe damage by sulfate - reducing bacteria has occurred particularly in oil - well

2 HS 2 + 2 e −

→ − 2 HS + H

casing, buried pipelines, water - cooled rolling mills, and pipe from deep water wells. Within 2 years, well water in the U.S. Midwest caused failure of a galvanized

water pipe 50 mm (2 in.) in diameter by action of sulfate - reducing bacteria, whereas municipal water using similar wells, but which was chlorinated before- hand, was much less corrosive.

A combination of low temperature and low humidity is one approach to controlling the growth of bacteria, but fungi may be capable of growing under such conditions [10] . Regular cleaning is a good practice to prevent biofi lm for- mation and subsequent corrosion. Chlorination is used to eliminate bacteria that cause corrosion, but this treatment can produce byproducts that are environmen- tally unacceptable [10] . Aeration of water reduces activity of anaerobic bacteria

since they are unable to thrive in the presence of dissolved O 2 . Addition of certain biocides can be benefi cial, but microorganisms are capable of becoming resistant to specifi c chemicals after long - term use. Eradication of microbial popu- lations may be achieved by combining several chemicals or by increasing the concentration of a biocide [10] .