Chapter 5 - Water Corrosivity.ppt

Chapter - 5

Water is commonly described either in terms of its nature, usage, or origin. The implications in these descriptions range from being highly specific to so general as to be non-definitive. Ground waters originate in subterranean locations such as wells, while surface waters comprise the lakes, rivers, and seas.

Fresh Water

Fresh water may come from either a surface or ground source, and typically contains less than 1% sodium chloride. It may be either "hard" or "soft," i.e., either rich in calcium and magnesium salts and thus possibly forming insoluble curds with ordinary soap. Actually, there are gradations of hardness, which can be estimated from the Langelier or Ryznar indexes or accurately determined by titration with standardized chelating agent solutions such as versenates.

Types of Water

• Brackish water contains between 1 and 2.5% sodium chloride, either from natural sources around otherwise fresh water or by dilution of seawater. Brackish water differs from open seawater in certain other respects. The biological activity, for example, can be significantly modified by higher concentrations of nutrients. Fouling is also likely to be more severe as a consequence of the greater availability of nutrients.

• Within harbors, bays, and other estuaries, marked differences can exist in the amount and type of fouling agents present in the water. The main environmental factors responsible, singly or in combination, for these differences are the salinity, the degree of pollution, and the prevalence of silt. Moreover, the influence of these factors can be very specific to the type of organism involved. Apart from differences that can develop between different parts of the same estuary, there can also be differences between fouling in enclosed waters and on the open coast. In this respect the extent of offshore coastal fouling is strongly determined by the accessibility to a natural source of infection. Local currents, average temperature, seasonal effects, depth, and penetration of light are operative factors. The presence of pollutants can also be quite important and highly variable in coastal areas.

Brackish Water

Drinking Water Brakish WaterSea Water

&Brine

TDS : Less than 1000 mg/L TDS : 1000 – 30,000 mg/L TDS : More than 30,000 mg/L

Classification of Oilfield Water

Produced Water Injection Water Cooling Water

Boiler Water Potable Water

• A major use of water in industry is the transfer of heat and the production of steam. There is extensive use of cooling water in almost every manufacturing process, in commercial air conditioning, and even a substantial percentage in domestic air conditioning. Water is used in the passive sense for potable and for fire control purposes. Fossil and nuclear fuel steam plants are encountered in the heating and power generating fields.

• Water can be corrosive to most metals. Pure water, without dissolved gases (e.g., oxygen, carbon dioxide, and sulfur dioxide) does not cause undue corrosion attack on most metals and alloys at temperatures up to at least the boiling point of water. Even at temperatures of about 450 oC, almost all of the common structural metals, except magnesium and aluminum, possess adequate corrosion resistance to high-purity water and steam. In summary, the factors influencing the corrosion of materials in water systems are:

1. Physical configuration of the system;2. Chemistry of the water (hardness salts, chlorides, and dissolved gases being

the most important);3. Presence of solids in the water;4. Flow rate;5. Temperature of the water;6. Presence of bacteria.

• The concentrations of various substances in water in dissolved, colloidal or suspended form are typically low but vary considerably. A hardness value of up to 400 ppm of calcium carbonate, for example, is sometimes tolerated in public supplies, whereas 1 ppm of dissolved iron would be unacceptable. In treated water for high-pressure boilers or where radiation effects are important, as in some nuclear reactors, impurities are measured in very small units such as parts per billion (ppb) or 1 mg of contaminant per liter of water. Water analysis for drinking water supplies is concerned mainly with pollution and bacteriological tests. For industrial supplies a mineral analysis is of more interest. The important constituents can be classified as follows:

• Dissolved gases (oxygen, nitrogen, carbon dioxide, ammonia, sulfurous gases)• Mineral constituents, including hardness salts, sodium salts (chloride, sulfate,

nitrate, bicarbonate, etc.), salts of heavy metals, and silica• Organic matter, including that of both animal and vegetable origin, oil, trade

waste (including agricultural) constituents and synthetic detergents• Microbiological forms, including various types of algae and slime forming

bacteria• The pH of natural waters is rarely outside the fairly narrow range of 4.5 to 8.5.

High values, at which corrosion of steel may be suppressed, and low values, at which gaseous hydrogen evolution occurs, are not often found in natural waters. Copper is affected to a marked extent by pH value. In acidic waters, slight corrosion occurs and the small amount of copper in solution causes green staining of fabrics and sanitary ware. In addition redeposition of copper on aluminum or galvanized surfaces sets up corrosion cells resulting in severe pitting of the metals.

• From a corrosion standpoint, the most significant contaminant is dissolved oxygen (DO) from ambient air. This is illustrated in the following Figure. Oxygen is a cathodic depolarizer that reacts with and removes hydrogen from the cathode during electrochemical corrosion, thereby permitting corrosion attack to continue.

Effect of oxygen concentration on the corrosion of low-carbon steel in tap water at different temperatures

• Dissolved oxygen can destroy the protective hydrogen film that can form of many metals and oxidize dissolved ions into insoluble forms. Deposits of rust in a plumbing system is such an example of differential aeration cells and accelerate corrosion.

• Dissolved oxygen (DO) refers to the volume of oxygen that is contained in water. Oxygen enters the water by photosynthesis of aquatic biota and by the transfer of oxygen across the air-water interface. The amount of oxygen that can be held by the water depends on the water temperature, salinity, and pressure. Gas solubility increases with decreasing temperature (colder water holds more oxygen). Gas solubility increases with decreasing salinity (freshwater holds more oxygen than does saltwater). Both the partial pressure and the degree of saturation of oxygen will change with altitude . Finally, gas solubility decreases as pressure decreases. Thus, the amount of oxygen absorbed in water decreases as altitude increases because of the decrease in relative pressure.

Oxygen dissolved in water is probably the most troublesome corrosion producing substance. The product of corrosion of iron by oxygen containing water is a mixture of iron oxides, usually hydrated, and generally referred to as rust. The following equations illustrate this in the simplest form and in water containing only dissolved oxygen.

Fe + 2H+ Fe++ + 2H

2H + ½ O2 H2O

2Fe++ + ½ O2 + H2O 2Fe+++ + 2OH-

The action of oxygen is two fold - it depolarizes the cathode, and it oxidizes the ferrous ions to ferric ions, which form the insoluble (above pH 3) ferric hydroxide.

• If the formation of ferric hydroxide occurs out of contact with the metal, there is no stifling of the reaction by this corrosion product. In a closed system, this reaction will continue until the dissolved oxygen is used up and the ferrous hydroxide smothers the reaction.

• However, in a system in contact with the air, the oxygen supply is continually replenished. The rate of corrosion, in this case, is generally restricted by the transport of oxygen from the air through the water to the metal. While the rate of corrosion may not be as high as that produced by the attack of acids, the inexhaustible supply of air assures that the reaction may continue for a long period of time.

• The corrosion rate of steel in water has been found to be approximately proportional to oxygen content up to 5.5 cc/litre. The corrosion rate is lower at higher oxygen concentrations. This may be explained by the fact that at low oxygen concentrations, the corrosion product formed is not an impermeable to oxygen diffusion as that formed at higher oxygen concentrations.

• The effect of oxygen on corrosion with increasing temperature is also shown in the following Figure, indicating the results in a closed versus an open container (the latter permitting natural deaeration by ebullition). In a closed vessel, corrosion continues to increase with temperature, whence the requirement for removing DO from hot water systems and boilers.

Effect of oxygen on corrosion of steel

• In waters containing high salt concentrations, corrosion is proportional to the amount of oxygen dissolved in the water. As the salt concentration in the water increases, the solubility of oxygen decreases and, consequently, the corrosion rate is reduced.

• Carbon dioxide dissolved in water can contribute to the corrosion of steel. Corrosion caused by water containing dissolved carbon dioxide is characterized by clean, uniformly thinned surfaces below the water line.

• The rate and amount of corrosion is dependent upon the salts dissolved in the water, the carbon dioxide content, the oxygen content, the temperature, and the composition of the steel.

• Carbon dioxide is present in water as: (1) the carbon dioxide in carbonate ions; (2) the carbon dioxide necessary to convert the carbonates to bicarbonates ; (3) the amount of carbon dioxide necessary to keep the bicarbonates in solution; and (4) any excess carbon dioxide. This excess carbon dioxide is referred to as “aggressive” carbon dioxide, and is the most corrosive form.

• For equal concentrations, carbon dioxide dissolved in water is not as corrosive as oxygen dissolved in water. At 60 oC, a solution containing 4 mL of dissolved oxygen is seven times as corrosive as one containing the same volume of dissolved carbon dioxide.

• In distilled water with only oxygen present, the corrosion rate of mild steel (0.15% carbon) increases twofold between 60 and 90 oC; whereas the corrosion rate increases by 2.6 fold in the same temperature range with only carbon dioxide dissolved in distilled water. The increase in corrosion rate can be attributed to increased diffusion rates of the dissolved gases with increased temperature.

• The greater increase in corrosion rate in the carbon dioxide solution with increased temperature is caused by a decrease in the acidity - or finite carbon dioxide concentration - required at the metal surface for hydrogen evolution. This makes the carbon dioxide more corrosive.

• In corrosion caused by carbon dioxide dissolved in water, the following reaction applies where carbon dioxide reacts with water to form bicarbonate but not carbonate:

2CO2 + 2H2O + 2e- 2HCO3- + H2

• The depolarizing reaction of oxygen at the cathode is:

2O2 + 4H2O + 8e- 8OH-

• Comparing of these reactions indicates that, on the basis of electrons, oxygen dissolved in water should be roughly four times corrosive as an equal molar amount of carbon dioxide.

• When dissolved in water, carbon dioxide acts as an acid, so that the acidity of the solution and the corrosion rate are increased by increasing partial pressure of the carbon dioxide.

• Water that contains both dissolved oxygen and carbon dioxide is more corrosive to steel than water which contains only an equal total concentration of one of these gases. For example, in a solution containing 10 cc of carbon dioxide per litre, the corrosion rate is 8 mils per year (mpy); in a solution containing 0.67 cc of dissolved oxygen per litre, the corrosion rate is 4 mpy. A solution containing the above amounts of each of these selected gases has a corrosion rate of 17 mpy, which is greater than the sum of the corrosion rates of the individual gases.

• In waters containing high ratios of carbon dioxide to oxygen, corrosion is accompanied by the evolution of hydrogen; the corrosion rate is governed by the acidity and by the composition and physical structure of the steel. At low ratios, oxygen depolarization controls the reaction and all low alloy composition steels corrode the same.

• Hydrogen sulfide, like carbon dioxide, is not corrosive in the absence of moisture. Water containing dissolved hydrogen sulfide is corrosive.

• Hydrogen sulfide is very soluble in water and, when dissolved, behaves as a very weak dibasic acid, yielding an acid solution in distilled water.

• Generally, when hydrogen sulfide exists at the bottom of an oil or gas well dissolved in a brine, there is no oxygen or other oxidizing agent with it. Under these conditions, the dissolved hydrogen sulfide will attack iron and nonacid resistant alloys.

• When the brine is pumped to the surface, oxygen from the air dissolves in the brine giving a water which may even be corrosive to acid resistant alloys. Dissolved oxygen will slowly oxidize the hydrogen sulfide to give water and free sulfur according to the equation:

H2S + ½ O2 H2O + S

• The corrosion rates in the brine solutions are higher than those in distilled water, and the brine containing hydrogen sulfide gives the highest corrosion rate.

• The principal ions found in water are calcium, magnesium, sodium, bicarbonate, sulfate, chloride and nitrate. A few parts per million of iron or manganese may sometimes be present and there may be traces of potassium salts, whose behavior is very similar to that of sodium salts. From the corrosion point of view the small quantities of other acid radicals present, e.g. nitrite, phosphate, iodide, bromide and fluoride generally have little significance. Larger concentrations of some of these ions, notably nitrite and phosphate, may act as corrosion inhibitors, but the small quantities present in natural waters will usually have little effect.

• Chlorides have probably received the most study in relation to their effect on corrosion. Like other ions, they increase the electrical conductivity of the water so that the flow of corrosion currents will be facilitated. They also reduce the effectiveness of natural protective films, which may be permeable to small ions.

• Nitrate is very similar in its effects to chloride but is usually present in much smaller concentrations. Sulfate in general appears to behave very similarly, at least on carbon steel materials. In practice, high sulfate waters may attack concrete, and the performance of some inhibitors appears to be adversely affected by the presence of sulfate. Sulfates have also a special role in bacterial corrosion under anaerobic conditions.

• Another mineral constituent of water is silica, present both as a colloidal suspension and dissolved in the form of silicates. The concentration varies very widely and, as silicates are sometimes applied as corrosion inhibitors, it might be thought that the silica content would affect the corrosive properties of a water. In general, the effect appears to be trivial; the fact that silicate inhibitors are used in waters with a high initial silica content suggests that the form in which silica is present is important.

• Salts dissolved in water have a marked influence on the corrosivity of the water. At extremely low concentrations of dissolved salts, different anions and cations show varying degrees of influence on the corrosivity of the water.

• The anions most commonly found in water are chloride, sulfate, and bicarbonate. The sulfate ion has a greater effect on the corrosivity of the water than the chloride ion, and the bicarbonate ion shows inhibitive tendencies.

• In solution containing bicarbonate ion along with either chloride or sulfate ion at concentrations up to 100 ppm, the bicarbonate ion showed increasing inhibition with increasing concentration but did not entirely prevent corrosion.

• The influence of anion on corrosivity also depends upon the metal in question. The corrosion rate of stainless steel is greater in 0.1 mole potassium chloride than in 0.1 mole potassium sulfate because of the greater penetration of the protective oxygen or oxide coating by the chloride ion.

• The order of decreasing penetration power of anions has been given as:

chloride > bromide > iodide > fluoride > sulfate > nitrate > monohydrogen phosphate.

• The order of decreasing corrosiveness of cations has been given as:

ferric > chromic > ammonium > potassium > sodium > lithium> barium > strontium > Calcium > Manganese > Cadmium > Magnesium.

• A relative dilute alkaline solution affords corrosion protection to iron, but a highly alkaline solution does not. In a water solution of 4% sodium hydroxide, the corrosion rate of iron is very low, and the potential of the iron on the hydrogen scale is approximately 0.1 volt. At extremely high concentrations of sodium hydroxide, the potential drops to an active value, and iron corrodes, forming soluble sodium ferrite (NaFeO2).

• Other constituents that contribute to corrosion are chlorides and carbon dioxide, possibly calcium, as will be explained later, and sulfides or ammonia from industrial or natural sources. Of course, many other man-made contaminants can be found in local areas if industries are permitted to discharge their waste products into water resources. As with other chemical reactions, corrosion increases with elevated temperature, unless stifled by insoluble scales, the removal of corrosive gases, or the addition of corrosion inhibitors.

• The formation of scale on a surface can be positive by providing an excellent protection of the substrate or negative by accentuating pitting at pores, cracks, or other voids in the film. If the film attains any significant thickness, the loss of heat transfer through the metal and deposited scale can also be a problem in certain applications. Thus, the development of scales on metal surfaces is an important consideration when using metals in waters. The effect of oxygen and pH on the corrosion rate of steel at two temperatures is shown in the following Figure.

Corrosion of steel in water containing 5 ppm of dissolved oxygen at two different temperatures as a function of the water pH.

• In a broad range of about pH 5 to 9, the corrosion rate can be expressed simply in terms of the amount of DO present (e.g., µm/y per ml. DO per liter of water). At about pH 4.5, acid corrosion is initiated, overwhelming the oxygen control. At about pH 9.5 and above, deposition of insoluble ferric hydroxide tends to stifle the corrosion attack. Amphoteric metals also show an increase in the corrosion rate in alkaline environments. Aluminum and lead are examples of amphoteric metals. The following Figure shows the behavior of steel and aluminum as a function of pH.

Corrosion of steel and aluminum as a function of pH at the same temperature (22oC)

• One of the common ways of generating hydrogen in a laboratory is to place zinc into a dilute acid, such as hydrochloric or sulfuric. When this is done, there is a rapid reaction in which the zinc is attacked or “dissolved” and hydrogen is evolved as a gas.

Rapid evolution of hydrogen bubbles during the corrosion of a zinc strip in a 1 M HCl acid solution

• These reactions are described in the following equations:

Zn + 2HCl ZnCl2 + H2

Zn + 2H+ + 2 Cl. Zn2+ + 2 Cl- + H2

• These equations are the chemical shorthand for the statement: One zinc atom + two hydrochloric acid molecules dissociated as ions H+ and Cl- and becomes one molecule of zinc chloride in the first equation and written as a soluble salt in the form of Zn2+ and Cl- ions in the second equation + one molecule of hydrogen gas which is given off as indicated by the vertical arrow. It should be noted that the chloride ions do not participate directly in this reaction, although they could play an important role in real corrosion situations.

• Similarly, zinc combines with sulfuric acid to form zinc sulfate (a salt) and hydrogen gas as shown in the following equations:

Zn + H2SO4 ZnSO4 + H2

Zn + 2H+ + SO42- Zn2+ + SO4

2- + H2

• Note that each atom of a substance that appears on the left-hand side of these equations must also appear on the right-hand side. There are also some rules that denote in what proportion different atoms combine with each other. As in the preceding reaction, the sulfate ions that are an integral part of sulfuric acid do not participate directly to the corrosion attack and therefore one could write these equations in a simpler form:

Zn + 2H+ Zn2+ + H2

• Many other metals are also corroded by acids often yielding soluble salts and hydrogen gas as shown in Equations and for respectively iron and aluminum:

Fe + 2H+ Fe2+ + H2

2Al + 6H+ 2Al3+ + 3H2

• Note that zinc and iron react with two H+ ions, whereas aluminum reacts with three. This is due to the fact that both zinc and iron, when corroding, each lose two electrons and display two positive charges in their ionic form. They are said to have a valence of +2 or II, whereas aluminum loses three electrons when leaving an anodic surface and hence displays three positive charges and is said to have a valence of +3 or III. Some metals have several common valences, others only one. The following Figure shows Some of the oxidation states found in compounds of the transition-metal elements.

Oxidation states found in compounds of the metallic elements. A solid circle represents a common oxidation state, and a ring represents a less common (less energetically favorable)

oxidation state

• The corrosion of metals can also occur in fresh water, seawater, salt solutions, and alkaline or basic media. In almost all of these environments, corrosion occurs importantly only if dissolved oxygen is also present. Water solutions rapidly dissolve oxygen from the air, and this is the source of the oxygen required in the corrosion process. The most familiar corrosion of this type is the rusting of iron when exposed to a moist atmosphere.

4Fe + 6H2O + 3O2 4Fe(OH)3

• In this equation, iron combines with water and oxygen to produce an insoluble reddish-brown corrosion product that falls out of the solution, as shown by the downward pointing arrow. During rusting in the atmosphere, there is an opportunity for drying, and this ferric hydroxide dehydrates and forms the familiar red-brown ferric oxide (rust) or Fe2O3, as shown below:

2Fe(OH)3 Fe2O3 + 3H2O

Similar reactions occur when zinc is exposed to water or moist air followed by natural drying.

2Zn + 2H2O + O2 Zn(OH)2

The resulting zinc oxide is the whitish deposit seen on galvanized pails, rain gutters, and imperfectly chrome-plated bathroom faucets. It also familiarly called 'white rust' a non-protective and even destructive form of corrosion that attacks incompletely passivated galvanized steel material or galvanized components subjected to marine atmospheres.

White rust on seaside road railing

• As discussed previously, the iron that took part in the reaction with hydrochloric acid in had a valence of 2, whereas the iron that takes part in the reaction shown in the previous equation has a valence of 3. The clue to this lies in the examination of the equation for the corrosion product Fe(OH)3. Note that water ionized into H+ and OH-. It is further known that hydrogen ion has a valence of 1 (it has only one electron to lose). It would require three hydrogen ions with the corresponding three positive charges to combine with the three OH- ions held by the iron. It can thus be concluded that the iron ion must have been Fe3+ or a ferric ion.

• Also note that there is no oxidation or reduction (electron transfer) during either reaction. In both cases the valences of the elements on the left of each reaction remain what it is on the right. The valences of iron, zinc, hydrogen, and oxygen elements remain unchanged throughout the course of these reactions, and it is consequently not possible to divide these reactions into individual oxidation and reduction reactions.

• In a broad range of about pH 5 to 9, the corrosion rate can be expressed simply in terms of the amount of DO present (e.g., µm/y per ml. DO per liter of water). At about pH 4.5, acid corrosion is initiated, overwhelming the oxygen control. At about pH 9.5 and above, deposition of insoluble ferric hydroxide tends to stifle the corrosion attack. Amphoteric metals also show an increase in the corrosion rate in alkaline environments. Aluminum and lead are examples of amphoteric metals. The following Figure shows the behavior of steel and aluminum as a function of pH.

Corrosion of steel and aluminum as a function of pH at the same temperature (22oC)

• The effect of temperature changes on the corrosion rate in waters is more complex than the simple chemical principle that an increase in temperature increases the reaction rate.

• An increase in the temperature of a corroding system has four main effects: (1) the rate of chemical reaction is increased; (2) the solubility of gases in the water is decreased; (3) the solubility of some of the reaction products may change, resulting in different corrosion reaction products; and (4) viscosity is decreased, and any thermal differences will result in increased circulation.

• Generally, corrosion rate increases with increasing temperature. This is particularly true when corrosion is due to the presence of mineral acids in water, resulting in hydrogen evolution. However, in waters which are corrosive due to the presence of dissolved oxygen, the corrosion rate increases with increasing temperature only until the temperature becomes high enough to cause an appreciable decrease in the oxygen solubility. Further temperature increase beyond this value results in a decrease in the corrosion rate for open systems, where the oxygen is free to escape.

• In a closed system, the oxygen can not escape and the corrosion rate continues to increase with increasing temperature.

• In waters containing calcium or magnesium bicarbonates, temperature increases result in the evolution of carbon dioxide and increased corrosion. At the same time, calcium or magnesium carbonates may deposit on the metal surface, resulting in the formation of a protective coating. Some of the carbonate ion may also combine with ferrous ion forming ferrous carbonate in the corrosion product.

• If the temperature increase is not uniform over the metal system, the hotter areas tend to be anodic to the colder areas. This may result in pitting corrosion.

• Zinc is normally anodic to steel in water and is sometimes used as a sacrificial material to the cathodic protection of steel. However, this may change with increased temperature. In the temperature range 60 to 90 oC, the zinc may become the cathode, resulting in corrosion of the steel.

• The flow of water over a metal surface influences the corrosion rate chiefly through the effects of water movement on the other factors governing corrosion. In stagnant waters or waters at zero velocity, the general corrosion rate is usually low, but localized or pitting corrosion may occur.

• Generally, some motion in a corrosive system causes greater uniformity and results in a thinning type of corrosion rather than pitting. Some flow or motion is also desirable when corrosion inhibitors are used, so that the inhibitors may be effectively distributed.

• Turbulence may occur at high velocities, and the turbulence may result in nonuniform conditions that lead to pitting corrosion. At high velocities, the film of corrosion product may be removed as it forms, resulting in corrosion.

• Systems which contain areas of high and low velocity may experience deposition of sludges or suspended solids in areas of low velocity. These sludge deposits can restrict oxygen diffusion, resulting in corrosion under the deposit.

• In oxygen free systems, the area subject to the highest velocity becomes anodic to the area subject to lowest velocity, and corrodes. When dissolved oxygen present, an oxygen concentration cell is formed, and the area of low velocity (receiving less oxygen) becomes the anodic area.

• Extremely high velocities may give rise to low pressure areas where vapor bubbles may form; on collapse, these can cause cavitation erosion at areas of higher pressures. Impingement attack may occur, under turbulent flow conditions, if the water carries debris and air bubbles. The forward ends of the corrosion pits may be undercut because of the impingement of the air bubbles. Mechanical erosion can result from waters carrying suspended sand or other particles at high velocities.