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"f‘ - :- 3-" a”? is“? in by? “ r 5.3% 5.515 " ..‘ um, 35:; l. ‘1 fit; “f“! 433%“??? ‘ < ~14 I; [tn-“1'.“ 2g»: xiii!" \"‘l'.1:i-Sj:"‘ m ‘7 ‘ w :1. ‘ 31555234.;‘1fi-J {'1 w'gjfid ‘ .é‘ I‘m“?! ‘. ,. I, 7 ‘ .Xf’iuél 1‘ f 4' (‘l' ‘ , . z ; V? #51,, It ‘1'./‘ view“ i " , '- . r 7 9; ~ v1" . ." ». 9 3' MiChigan State University This is to certify that the dissertation entitled Effect of Various Oxyanions of Nitrogen on Corrosion Inhibition of Aluminum Alloy presented by Hyung—Joon Kim has been accepted towards fulfillment of the requirements for _Ph_-lD_-_’_degree in W “A Mgfiim Date) iMM l9P7 0-12771 MSU is an Affirmative Action/Equal Oppanuniry Institution MSU RETURNING MATERIALS: Place in book drop to ngRARiEs remove this checkout from .2353. your record. FINES will be charged if book is returned after the date stamped below. MAR 2 1 1996 W EFFECT OF VARIOUS OXYANIONS OF NITROGEN ON CORROSION INHIBITION OF ALUMINUH.ALLOY BY Hyunngoon Kim A DISSERTATION Submitted to Michigan State University in partial fulfillment of the requirements for the degree of DOCTOR OF PHILOSOPHY Department of Metallurgy, Mechanics and Materials Science 1987 ABSTRACT EFFECT OF VARIOUS OXYANIONS OF NITROGEN ON CORROSION INHIBITION OF ALUMINUM ALLOY BY Hyung-Joon Kim Potentiodynamic polarization was used to evaluate inhibitor effectiveness for AA 7075-T6 in solutions containing chloride, at varying temperatures. Inhibitors included various anions of nitrogen and boron, and ammonium ion. Generally, passivation was found in distilled water. In the presence of chloride, however, passivity depends on inhibitor concentration. Inhibitive efficiency of nitrogen oxyanions, based on the pitting potential, is in the decreasing order of: Nzoz- > NH‘}+ > - . . - . 2- - N03 > Combination of NO2 Wlth B407 > N02 . Current density depends logarithmically on temperature in solutions of sodium nitrite and sodium nitrate; the activation energies are 2.68 kcal/mol and 4.88 kcal/mol, respectively. Acknowledgments I would like to express my deepest gratitute and appreciation to Dr. Robert Summitt who as major professor provided direction and encouragement throughout the course of this study and in preparation of this manuscript. I also wish to thank the other members of my guidance committee, Dr. K. Mukherjee, Dr. G. Gottstein and Dr. H. Eick for their inspiration. Thanks are extended to U.S. Air Force and U.S. Office of Naval Research and the Department of Metallurgy, Mechanics and Materials Science at Michigan State University for financial support during the doctoral program. Finally, I would like to especially thank my parents, uncle and wife (Young-Joo) for their encouragement and understanding. ii TABLE OF CONTENTS LIST OF TABLES 0.0.0.0...0.0.0....OOOOOOOOOOOOOOOOOOO V LIST OF FIGURES OOOOOOOOOOOOOOOOOOOOOOOOOOOOOOOOOOOOO. Vi I. INTRODUCTION 0...... ....... .OOOOOOOOOOOOOOOOOO... 1 II. THEORETICAL BACKGROUND AND LITERATURE REVIEW .... 3 A. THEORY OF PASSIVITY ..... ........... ........ 3 B. ELECTROCHEMISTRY OF PASSIVATION ............. 6 C. INHIBITION 00.00.0000...OOOOOOOOOOOOOOOOOOOO 12 D O NITROGEN OXYANIONS O O O O O O O ..... O O O O O O O O O O O O O 3 1 E. RESEARCH PROGRAM .. ....... .................. 43 III. EXPERIMENTAL TEST PROCEDURE AND APPARATUS ....... 44 A. EXPERIMENTAL SYSTEM ........................ 44 B. EXPERIMENTAL TECHNIQUE ............... ...... 51 C. 7075-T6 AL ALLOY HANDLING, CLEANING AND TEST SOLUTION ......... O O O O O O O O O O O O O O O O 0 O O O 0 O O O O O 54 IV. EXPERIMENTAL RESULTS 0 O O O O O O O C O O O O O O O O O O O O ...... O 61 A. DISTILLED WATER AND SODIUM CHLORIDE ........ 61 B. EFFECTS OF INHIBITORS ...................... 61 C. EFFECTS OF TEMPERATURE ..................... 82 iii V. VI. REFERENCES DISCUSSION CONCLUSION OOOOOOOOOOOOOOOOOO0...... ...... 0...... iv 88 109 LIST OF TABLES Page The Oxyanions of Nitrogen ............. ...... 32 Oxidation Potential Data for Reduction of Nitrate Ion ................................. 37 Chemical Composition ............... . ...... .. 55 Mechanical Properties of 7075-T6 Al Alloys .. 55 The Test Solution ...................... ..... 57 Lynch Formulation ......... .................. 58 Pitting Potential of Different Test Solutions With 0.1 wt.% NaCl .......................... 100 10. 11. LIST OF FIGURES Schematic Structure of Passive Film ........... The Corrosion Diagram of Evans ................ Electronic Structure of NO - .................. 2 Electronic Structure of N03- .................. Electronic Structure of NZOZZ— ................ The Experimental System ....................... The Corrosion Cell .......... . ........ . ........ The Anodic Polarization of 7075-T6 A1 Alloy in Dilute Solution ............. ...... .. .......... The Anodic Polarization of 7075—T6 A1 Alloy in 0.1 wt.% Sodium Chloride ...................... The Anodic Polarization of 7075-T6 Al Alloy in Sodium Nitrite ................................ The Effect of Concentrations of Nitrite on Current Density at Constant Potential ......... vi 34 _..__,. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21. The Anodic Polarization of 7075-T6 Al Alloy in Sodium Nitrite with Sodium Chloride ........... 67 The Anodic Polarization of 7075-T6 Al Alloy in Sodium Nitrite and Borax .. ...... .............. 68 The Anodic Polarization of 7075-T6 Al Alloy in Sodium Nitrite and Borax with Sodium Chloride .. 69 The Anodic Polarization of 7075-T6 Al Alloy in Lynch Formulation ....... . ..................... 71 The Anodic Polarization of 7075-T6 A1 Alloy in Lynch Formulation with Sodium Chloride ........ 72 The Cathodic Polarization of 7075-T6 Al Alloy in Sodium Nitrite .............. ..... ............. 73 The Cathodic Polarization of 7075—T6 Al Alloy in Sodium Nitrite with Sodium Chloride ........... 74 The Anodic Polarization of 7075-T6 Al Alloy in Sodium Nitrate ................................ 76 The Anodic Polarization of 7075-T6 A1 Alloy in Sodium Nitrate with Sodium Chloride ........ ... 77 The Effect of Concentrations of Nitrate on Current Density at Constant Potential ......... 78 vii 22. 23. 24. 25. 26. 27. 28. 29. 30. 31. The Anodic Polarization of 7075—T6 Al Alloy in Sodium Hyponitrite ............................ 80 The Anodic Polarization of 7075-T6 Al Alloy in Sodium Hyponitrite with Sodium Chloride ....... 81 The Anodic Polarization of 7075-T6 Al Alloy in Ammonium Hydroxide .... ..... . ..... ............. 83 The Anodic Polarization of 7075-T6 Al Alloy in Ammonium Hydroxide with Sodium Chloride ....... 84 The Anodic Polarization of 7075-T6 Al Alloy in Sodium Nitrite vs. Temperature ................ 85 The Anodic Polarization of 7075-T6 Al Alloy in Sodium Nitrate vs. Temperature ................ 86 Schematic Models of the Passive Film .......... 94 Schematic Polarization Diagram for Oxidizing Inhibitor ................... ........... ....... 97 Dependence of the Current Density on Temperature in Sodium Nitrate ....... ........ .............. 105 Dependence of the Current Density on Temperature in Sodium Nitrite ......................... .... 106 viii I. Introduction This study is an experimental exploration of the relations between the inhibitor/accelerant properties and the electrochemistry in 7075-T6 Al alloy. A potentiodynamic polarization technique is used to evaluate the effective- ness of inhibitor when 7075-T6 high-strength aluminum alloy is exposed to solutions with and without chloride ions. ltrtrite, nitrate, hyponitrite and ammonium hydroxide are used as the inhibitors. ,Aluminum.has a notable resistance to corrosion, which is remarkable from the viewpoint of its high free-energy content which would suggest a powerful tendency to corro- sion. In fact, a free aluminum surface corrodes rapidly and soon is covered by a thin protective oxide film that makes the metal immune to further corrosion. The corrosion resis- tance, therefore, depends upon the resistance of the oxide film to attack, that is, upon the degree of passivity. In practice, the corrosion behavior of aluminum is determined mainly by the behavior of the oxide-covered metal surface towards the corroding medium, which behavior is very de- pendent on the nature of the anion in solution. In.near~ neutral air-saturated solutions, the corrosion of aluminum is generally inhibited by the same anions which are inhibi- tive for iron, e.g., chromate, phosphate, and acetate. Aggressive anions for aluminum include the halide ions, i.e., F , Br , Cl-, and 1-, which cause pitting attack; and anions which form soluble complexes with aluminum, e.g. , citrate and tartrate which cause general attack. In the presence of chloride ions, the breakdown of passivating oxide films results from the severe pitting of the underly- ing metal. Competitive effects are observed in the action of mixtures of inhibitive anions and chloride ions on aluminum. Generally, the results show that the corrosion of aluminum depends on the nature of passivity, which in turn is determined by the specific anions present in the solu- tion. II. Theoretical Background and Literature Review A. Theory of Passivity The protection of metals against corrosion in the environment can be enhanced with inhibitors. When the inhibition.is obtained with the compounds altering mainly ‘the kinetics of anionic reaction -- this mechanism being most effective in neutral electrolytes --, inhibition is closely related to the passivation. The major theories of passivity are divided into four' categories: a) metal modification; b) reaction rate; c) oxide film; d) adsorption. a) Metal-Modification Theories 1) Schonbein( considered that the passive and active states of iron might represent allotropic modifications of the metal. Finklestein(2) supported a metal-modification theory by proposing the valence theory of passivity, where 2+ 3+ the metallic iron is made up of a mixture of Fe and Fe 3+ ions, with passive iron.consisting mostly of Fe and active iron consisting mostly of Fe2+, i.e., equilibrium exists between allotropic forms of the metal, i.e., Fe2+, 3+ Fe , and electrons(3). Metal-modification theories are no longer of serious interest except in the modern counterpart longer of serious interest except in the modern counterpart as a guide to establish the kind of surface films that form application of electronic configuration to transition metals. b) Reaction-Velocity Theory According to this theory(4), the passive state results from a slow rate of metal dissolution independent of any surface filnu The.slow reactivity leads to sluggish hydra- tion of anhydrous ions in the metal lattice, hydration being necessary to ion solubility in an aqueous environ- ment. This slow reactivity easily accounts for the marked anodic polarization characteristic of passivated metal. This theory, however, is expressed in part by the present adsorption theory. The chemisorbed oxygen film results from a displacement of adsorbed H O molecules with a consequent 2 reductdxn1.in rate of metal-ion hydration, or metal-complex formation necessary for the solubility of such ions in ‘water; or, the surface electrical double layer is altered in composition and structure by adsorbed oxygen resulting similarly in a reduced reaction rate. c) Oxide Film Theory An oxide model for passive film was proposed by Gmelin(5), Haber(6), Heathcote(7), Evans(8). Flade sug- gested that an oxide filnlor an oxygen alloy layer at the metal surface was a possible cause of passivity. Also, it is known that a protective oxide film is the passive film on metals such as Cr, Ni, Ti, Al, Mo, etc.. Often, no distinction is made between oxide films and adsorbed films, both kinds being considered to isolate similarly the metal frem its environment. But the distinction should be made because the mechanism of protection by adsorbed films relates to retarded surface reaction rates rather than to retarded diffusion rates through oxides supposedly making up the passive layer as explained with the oxide film theory. d) Adsorption Theory It is proposed by Fredenhagen(9) that the passive film on Fe consists of oxygen that satisfies chemical affinities of the surface and, hence, results in reduced reactivity. Langmuiruo) suggested that his measurements showed that oxygen adsorbed on tungsten exhibited reduced chemical reactivity compared with oxygen in the oxide W0 and that 3: the.well-known passive properties of Cr were probably related to a similar filnlof'adsorbed oxygen. On progres- sive anodic polarization of a metal to more noble potentials, multilayer adsorption of oxygen builds up the passive film thickness. The pronounced negative charge of adsorbed oxygen attracts and slowly incorporates into the film in nonstoichiometric amounts positively charged metal ions, as well as protons from the aqueous environment(1l). The proton comes from the presence of H20 or H in passive (12). (13) (14). films on Fe and on stainless steels It is reported that the mixed adsorbed oxygen-metal-hydrogen (O- M-H) passive film can account for the observation that jpassivity results from less than monolayer amounts of oxygen on the surface<15). In summary, the most promising model for the passive film on metals is one made up of adsorbed oxygen containing metal ions and protons, in indefinite proportions, and including the category of nonstoichiometric oxides (Figure 1). The latter designation is applicable to passive films on such metals as Al, Ta, and Zr, etc. B. ectroc emist of ass vat'on Figure 2 shows the corrosion diagram of Evans repre- sented by two curves, a cathodic-polarization curve CB and an anodic polarization curve AB, that characterize the dependence of the cathode and anode potentials on the current strength. The intersection point of the curves x' The measure of corresponds to the maximum current Ima control for the corrosion rate of the given electrode reaction is determined by the slope of the curves, i.e., the tangents of angles a andpfor the cathodic and anodic .reactions, respectively. Thus the cathodic and anodic potentials are given by WV Initial Film 0,0 0, ,0 . 0,0, 0 ,0‘, 1,,30.0,000 :0 0. .0. .0. .0. .0. .0 // M-O-H Film 0 Oxygen ion 0 Metal ion H Proton Figure 1. Schematic Structure of Initial Passive Film containing less than monolayer amounts of adsorbed oxygen, and of a thicker passive film containing additional metal ions and protons in nonstoichiometric amounts. IR F has ._ o_co:e . . _ _ _ . . _ . . _ _ _ _ _ _ _ ir ifi. . . . . . _ _ PP) L i4. ii; a._ ocms ._ O o_e;o o_oocreo ARC .orucOO Co mrzmooz Imax II The Corrosion Diagram of Evans. Figure 2. 9 o _ ° . w. ‘ '1’. 'Ma“ ° ‘1’. + Po 1 (1) ¢a — ¢a° + I-tan B - $30 + Pa°1: (2) where Pc and Pa are the cathodic and anodic polarization resistances. The cathode and anode potential are equal at the point of intersection of the curves corresponding to the maximum current: 0 . _ ° . 3 I’bc ' Pc Imax ¢a + Pa Imax ( ) o o (4) Imax - ($c - ¢a )/(Pa + PC). Taking into account the ohmic potential drop in the electrolyte (I'R), the current 1' corresponding to the ohmic resistance R is I'-(¢C'- $5W/R- (5) Thus, ¢a' - ¢a° + PaoI', (6) wc' - ¢C' - PcoI'. (7) Consequently, 1' - (¢C° - ¢a°>/- (8) 10 Equation (8) is the fundamental equation describing the corrosion process, taking into account both the polarization resistance and the ohmic resistance. The current can be determined by the initial potential dif- ference and by the sum of the polarization and ohmic resistance. From Equations (1) and (2), the equations of the cathodic and anodic polarization curve are modified as O ¢c _ ¢c - Kl.(1/FC), and (9) o 10 ¢ - ¢ + K2~, ( ) a a where K1 and K2 are the polarizability constants of the cathode and anode at unit current density, and Fc and Fa are the areas of the passive (cathode) and active (anode) parts of the electrodes. The intensity of the corrosion current of the system, defined as the rate of the corro- sion process calculated per unit surface area subjected to corrosion damage, is determined by the equation 0 O I - <¢c - $8 )/[R + (Kl/Fe) + (Kz/Fa)]. (11) When an inhibitor impedes only the anodic process, the corrosion rate may decrease either because of a lower rate of metal ion transfer into solution as a result of adsorp— tion of positively charged particles in the double layer, 11 or because of a contraction of the active part of the electrode as a result of passivation. It can be shown that an anodic inhibitor contracts the active part of the electrode considerably more than it reduces the corrosion current. For a completely polarized system, R is small in comparison with the polarization resistance. From Equation (11). I - (w ° C - ¢a°>/[(K1/Fc) + (Kz/Fa)i - A¢/[(Kl/FC) + (Kz/Fa)i. (12) The corrosion current in the initial electrolyte is The corrosion current in the electrolyte containing the inhibitor is On the condition that A¢1-A¢2 and K1 = K2, the following can be obtained, from Fa1 + Fcl = 1 and Fa2 + Fc2 = 1, Pal/Faz = (11/12) (Fcz/Fcl) ‘ (15) Since FcZ > Fcl' 12 Fal/Faz > 11/12. (16) From Equation (16), it can be seen that the active part of the electrode in the presence of an inhibitor contracts more strongly than the corrosion current is reduced. C. Inhibition (16) notes that organic compounds, particularly Shreir those containing elements of Group V and VI of the Periodic Table, such as nitrogen, phosphorus, arsenic, oxygen, sulfur and selenium are effective inhibitors in aqueous acid solutions. Generally, the primary step in the action iof inhibitors in acid solutions is its adsorption onto the metal surface, which usually is oxide free. Then the adv sorbed inhibdtor acts to retard the cathodic and/or anodic electrochemical process of the corrosion. The corrosion of metals in neutral solutions differs from that in acid solutions in two important respects. In air-saturated solutions, the main cathodic reaction in neutral solutions is the reduction of dissolved oxygen, whereas in acid solution, it is hydrogen evolution. Corroding metal surfaces have no oxide, whereas in neutral solutions metal surfaces are covered with films of oxides, hydroxides or salts, because the solubility of these species is reduced. Inhibition in neutral solutions is 13 caused by compounds which can form or stabilise protective surface films. Because of these differences, substances which inhibit corrosion in acid solution by adsorption on oxide—free surfaces, do not generally inhibit corrosion in neutral solution. Another type of inhibitors in neutral solutions act by stabilising oxide films on the metals to form thin protec— ‘tive, passivating films. Such inhibitors are the anions of weak acids, some of the most important in practice being chromate, nitrite, benzoate, silicate, phosphate and borate. Because of the resistance to the diffusion of metal ions due to passivating oxide films, the anodic reaction of metal dissolution is inhibited; thus these inhibitive ianions often are referred to as anodic inhibitors and they are more generally used than cathode inhibitors to inhibit the corrosion of iron, zinc, aluminum, copper and their alloys, in near neutral solutions. (17) identified the Sheldon, Derby and Van Dem Bussche four major categories of corrosion inhibitors with the mechanism of operation, that is, barrier layer formers, neutralizers, scavengers, and miscellaneous types. The barrier formers were as classified into oxidizers, adsorbed layer formers and conversion layer former. These materials have the ability to deposit on the metal surface and inter- face with the corrosion reaction and thereby lower its rate ‘to an acceptable value. The neutralizing inhibitors reduce 14 the corrosivity of the environment by removing hydrogen ions from the environment. Some scavengers operate on a similar principle. Miscellaneous inhibitors include scale inhibitors and biological growths which prevent corrosion by interfering with other processes. (18) notes that the inorganic anodic inhibitors Cohen for iron usually are either oxidizing anions such as nitrite and chromate, or buffering agents such as borate or phosphate. Oxidizing anions are effective in the presence of oxygen to prevent corrosion, but in general, the effec- tive concentration is lower in the presence of oxygen. Nitrite ion reacts at.the metal surface to form iron oxide and is reducted as far as ammonia(19). The protective film formed on the metal surface is mainly cubic oxide of 20 21 the Fe3O4 - Fe203- 7 type( )'( ). In the absence of an oxidizing agent, the protective film tends to revert to Fe3O4 by reaction with the underlying metal..At a slightly higher potential than the primary passive potential, the protective film breaks down locally in the presence of chloride to form pits, although the film formed reduces corrosion. In the absence of chloride, the film becomes thick with a raising of the potential. (22) Cohen earlier had found that weight loss decreased as both concentration.of nitrite and oxygen increased..At 15 intermediate concentrations, pitting was observed. During ‘the reaction/exposure of nitrite solution to the specimen, the nitrite ion concentration was found to decrease. Cohen suggests that a surface film is formed by ad- sorption of nitrite on the metal surface, followed by a reaction which forms oxide and ammonia. The same adsorption-reaction mechanism was proposed for other oxidizing inhibitors, e.g., chromate and molybdate. Although nonoxidizing inhibitors require the presence of oxygen, a much higher concentration of non-oxidizing oxyanion is required to inhibit corrosion. Rozenfe1d(23)’(24) reviewed the effects of nitrite, nitrate and other simple nitrogen-containing compounds for ferrous alloys in both neutral aqueous solutions and in the presence of accelerants, specifically chloride and sulfate ions. The inhibiting properties of sodium nitrite depend on the concentration of accelerant ions in the electrolyte. In solutions containing sodium chloride and sodium sulfate, sodium nitrite accelerates corrosion up to 9;. 0.08 g/l concentration, but at higher concentrations, it exerts a strongly inhibitive effect, being maximum at 0.2 g/l con- centration (2 wt.%). Rozenfeld suggests that in case of incomplete protection, sodium nitrite diverts a larger part of the surface from the anodic reaction and hence shifts the potential more strongly toward positive values than other inhibitors. The corrosion-intensifying mechanism 16 involving low concentrations of sodium nitrite is the same as observed for other anodic inhibitors (e.g., chromate and bichromate), i.e., the corrosion rate increases as in- Ihibitor increases up to critical value and then decreases. In the absence of accelerant ions, the protective con- 5 4 centration of sodium nitrite is between 10- and 10- mol/l or a. 6.9 x 10-4 The protection concentrations of sodium nitrite are a fuction of the temperature. In the absence of an inhibitor, ‘the corrosion rate of iron increases as the temperature is raised to 50 - 60°C, but at higher temperatures, the corro- sion rate decreases because of the lower oxygen solubility. If sodium nitrite is added, but not sufficient to completely suppress the corrosion process, the corrosion rate increases steadily with a rise in temperature. Apparently at high temperatures, the overvoltage for the reduction reaction of sodium nitrite decreases and the inhibitor begins to be reduced at a low rate, hence the limited oxygen solubility can no longer be the factor which restricts the corrosion rate. At high temperature, however, the increase in the concentration of sodium nitrite decreases the corrosion rate. At 25°C, the protective property of steel in distilled water is obtained at 2 x 4 10- mole/l, while at 70°C, it is obtained at 5 x 10-3 mole/l. 17 When accelerant ions are present, the logarithm of the protective sodium nitrite concentrations is linear with respect to the concentration of accelerant. The protective properties of nitrite are suppressed most strongly by sulfate ions and less strongly by nitrate ions. Therefore, for the same concentration of accelerant ion, a higher sodium nitrite concentration is needed in the presence of sulfate than in the presence of nitrate for protection. To suppress corrosion in the presence of chloride, a lower sodium nitrite concentration is needed than in the presence of sulfate. As the concentration of accelerant ions in- creases, the difference between the effects of sulfate and chloride decreases, whereas the difference between chloride and nitrate remains. If high concentrations of the ions are excluded, then, in terms of aggressiveness, accelerants are ordered as follows: sulfate > chloride > nitrate. In the presence of both accelerant and inhibitive ions, the passivating ion is adsorbed preferentially. Thus, it is harder for nitrite ion to dislodge already-adsorbed chloride ions from the electrode surface than to prevent its initial adsorption. The adsortion of both the passivat— ing ions and the accelerating ions should be facilitated as the potential shifts in the positive direction. Evidently at more positive potentials, the adsortion of chloride is greater than the adsortion of passivating ions, hence the electrode is activated. When various aggressive ions (e.g., chlorides and sulfates) are present, protection is achieved at a sodium nitrite concentration generally greater than 18 the total concentration of accelerant ions. Corrosion occurs when the ratio of inhibitor concentration to total accelerant-ion concentration is less than one. Sodium nitrite also is more effective in suppressing the accelerant properties of chloride than are benzoate and chromate. In the presence of sulfate, nitrite is about as effective as chromate. With respect to nitrate, the effec- ‘tiveness of the different inhibitors decreases in the following order: chromate > benzoate > nitrite. Other compounds containing the NO group and used as corrosion 2 inhibitors include dicyclohexylamine nitrite (DAN) and salts of nitrobenzoic acids (nitrobenzoates). Also accord- ing to Rozenfeld and Marshakov<25) , that concentration of inhibitors required to inhibit crevice corrosion is mudh higher than for general corrosion. Marshall(26) found that a combination of nitrite with N, N-di-(phosphonomethyl) methylamine was an effective corrosion inhibition for ferrous alloys when compared with nitrite alone. The passive film formed in the presence of a combination of nitrite.with N, N-di-(phosphonomethyl) methylamine was more protective and less susceptible to jpitting corrosion than.the passive film formed in the presence of nitrite alone. Ledovskikh(27) also has reported that combinations of sodium nitrite with various organic amines inhibit the corrosion of steel in a neutral medium (0.1 M NaCl) and 19 finds that the protective effect correlates well with physicochemical parameter of the amines (e.g., electron density at the nitrogen atom, ionization constant). The effect of sodium nitrite in chloride and sulfate media for mild steel has been discussed by Legault and Walker(28). Passivity in these systems resulted from an oxide film and was dependent upon the rate of protective oxide film formation exceeding the rate of the film (29) showed the effectiveness deterioration. Also, Wachter of sodium nitrite under the conditions which obtained in (30) concluded that the function petroleum pipelines. Cohen of nitrite was to maintain the oxide film which was formed by the reaction of dissolved oxygen with the steel surface. Also, it was reported that the sulfate ion interferes with the fuctioning of sodium nitrite as a corrosion inhibitor for mild stee1(31). Mercer, Jenkins and Rhoades-Brown(32) investigated the actitn1 20 chloride > nitrate; the order changes with increase in nitrite concentration. (33) describes the role of chloride in iron Foley corrosion as functioning through its property of penetrate ing oxide films that otherwise are protective, and adsorption on the metal surface or the thin oxide film to produce a strong electric field which draws ions from the meta1(34). The halide ion forms a surface complex with Fe, the stability of which determines the corrosion kinetics and catalyzes the reaction by forming some sort of inter- mediate bridging structure. Corrosion rates of iron in chromate and nitrite solur tions were measured as weight loss by Matsuda and Uhlig(35). The critical concentration below which corrosion 5 M for NaNO and 5 x increases or pitting occurs is 1 x 10— 2 10-4 M for NaZCrO4. In the presence of Cl- and SO 42-, these values are increased. Sulfates break down passivity in nitrite solutions more than chlorides do, the reverse is true in chromate solution. Oxyanions of various elements (e.g., chromium and :molybdenum) are corrosion inhibitors for ferrous alloys in neutral aqueous solutions. In acid solutions, these oxidants become cathodic depolarizers and.they influence the anodic reaction of iron, as discussed by Mikhailovskii, 21 Popova and Sokolva(36). Inhibiting properties are at- tributed either to a capacity to "repair" the oxide film formed on the metal in an electrolyte or to adsorption of the oxyanion, resulting in a change in structure of the metal-electrolyte boundary. The influence of oxidizing agents usually is attributed to their enhanced adsorption capacity favoring formation of surface complexes which retard the anodic ionization of the metal. McCafferty(37) classes inhibitors as (a) adsorption (chemisorption), (b) film-forming/passivating, and (c) precipitation. Crevice corrosion of iron accelerated by chloride and inhibited by chromate has been discussed further by McCafferty<38) , finding that the process can be inhibited successfully at chloride concentrations up to 0.6 M/l (3.5 wt.%). The logarithms of each ion's activities are linearly related at the minimum chromate activity necces- sary to inhibit crevice corrosion. McCafferty explained this on the basis of competitive adsorption between the aggressive and inhibitive ions, each of which adsorbs. Steady-state electrode potentials of crevice-corroding iron were found to be - 620 to - 660 mV vs. Ag/AgCl reference electrode, and independent of bulk solution composition. (39) Samuel, Sotoudeh and Foley proposed steps by which aggressive anions act on aluminum: 22 1. .Adsorption of‘Cl- on aluminum oxide (A1203) surface. 2. Complexing of aluminum cation (Al+3) in the oxide lattice with halide ion (e.g., C1.) to form soluble A1Cl4-. 3. Soluble species (AlCl4—) diffuse away from the surface resulting in a thinning of the protective oxide film. 4. At sufficiently thinned sites aluminum reacts directly with the electrolyte. Soluble species diffuse away from the reaction site and the oxide film is thinned to a point where aluminum ions can pass directly from the metal to the solution interface. The critical step then is the formation of soluble complexes at specific sites. For this reason pitting and general corrosion of aluminum depend on the nature of the anion. The preferential adsorption sites may be defects or flaws in the oxide film. An inhibitor may participate in either of the first two steps: - Competing for adsorption sites, hence retarding' formation of soluble species. 23 3 - Competing in the reactionAl+ + 4Cl- = A1C14- to prevent formation of AlCl4-; hence the competing ion must form an insoluble complex. In chloride solution, the passivity of inhibition should be viewed as interfering with the reaction A1 + 2Cl + 20H" A1(0H)2c12', whereby the oxide film is thinned. A species may form a complex ion which may be stable and soluble A13+ + nR’ = A1R3'n, and accelerate corrosion. On the other hand, the species may react as A13+ + 3R- = AlR wherein the species AlR is stable, unshared, very slightly 3 ionized, and resides at or near the aluminum surface. (40)'(41) showed the relation Vedder and Vermilyea between aluminum and water mechanism which involves forma— dissolution of the Al O and tion of amorphous A1203, 2 3, precipitation of AlOOH from solution. The AlOOH layer 24 provides deposition sites near the surface for dissolved species hence allows the reaction to proceed at a high rate despite the very low solubility of Al203 in neutral solu- tions. Inorganic inhibitors function by becoming adsorbed on the oxide and preventing the nucleation and growth of the hydroxide layer(41). 36 The adsorption of Cl on an aluminum surface was found to be a function of time and applied voltage, and highly localized to oxide imperfections and grain bound- aries, as discussed by Verkerk(42). Belov and coworkers(43) established the following order for the degree of adsorp— tion onto an anodic aluminum film: _ - 2- 2- .. H20 < NO3 < C1 < Cr204 < 804 < H2PO4 A similar order also was reported by Herczynska and c1’ and so 2' Prozynska(44) for N03 , 4 . Berzins, Lowson and (45) Mirams measured the adsorption isotherms for chloride 36 on a corroding aluminum surface using Cl as a radioactive tracer. The amount of chloride adsorbed, W was a function C1 of chloride concentration, [Cl_], and time, t, according to; log w = 0.64(log[C1-] + log t) - 78, Cl 25 2 was expressed as g cm- , [Cl-J as mol l-1 where W and t Cl as minutes. The adsortion was localized to corroding pit sites and there was no threshold for the chloride con- centration below which pitting would never occur and the addition of inhibitor (e.g., nitrate or sulfate) would delay but not prevent the onset of pitting. The delay time before the onset of pitting varied a logarithmically with inhibition concentration according to log D = 1.4 log C1 + K1, where K1 was constant and was a function of chloride con— centration and inhibitor species, D as delay time and C as 1 inhibitor concentration. Foroulis and Thubrikar<46) investigated the kinetics of passivity breakdown and nucleation of pitting of preanodized aluminum by chloride. The rate of passivity breakdown at steady-state, critical pitting potential was dependent on chloride concentration, temperature, and oxide film thickness, but is independent of solution pH in the range 5 - 10. It is likely that in the process of pit initiation, adsorption of chloride from solution on the hydrated aluminum oxide surface is the rate-determining step in low chloride concentration, whereas the reaction step resulting in formation and dissolution of the hydroxylchloride aluminum salt on the hydrated oxide sure face becomes the rate-determining step in higher chloride concentration. 26 The influence of anions on the initiation of pitting and the kinetics of pit growth on aluminum alloy Type 7075 (47). The order of has been investigated by Dallek and Foley reaction (i.e., the number of aggressive anions per Al surface reaction site) and the energy of activation for pitting initiation in aggressive anion-containing solution. were measured. The value of the order of reaction (n) and the activation energy (Ea) in 1 N H2S04 varies with the 1 aggressive anions as n = 8, E = 18 kcal mole- for 01-, n a 1 for Br- and n = 2, E = 6.6 kcal = 4, E = 26 kcal mole- a a mole"1 for I-. Nguyen and Foley<48) investigated the dissolution reaction at a bare aluminum surface in.aqueous solutions and concluded that the controlling step was a complexation :reaction of hydrated.cation with the aggressive anion present, which accounted for the strong anion dependence in metal pitting and general corrosion, although hydrolysis of complexes lead to the formation of stable hydrous oxides whose composition was independent of the anion in solution. Samuel, Sotoudeh and Foley(39) also have reported that a chelate, an ion or molecule with two or more atoms having unshared pairs of electrons, might inhibit corrosion or accelerate corrosion. The hydroxy carboxylic acids form soluble chelate compounds with aluminum and corrosion in 27 chloride solution is increased by an order of mag- nitude.Carboxylic acids form stable compounds which function in the same manner as do precipitating inhibitors, although the behavior of either may depend on the con- centration of the coordinating compound. The role of complex ions in the corrosion reaction with potential energy'has been clarified by Foley and Nguyen(49) . Generally speaking, one or more complex ions are involved in a multistep sequence. The relative stability of the complex is the critical factor. The forma- tion of an uncharged, stable, basic salt in a deep energy valley will prevent the reaction from progressing along the reaction coordinate, which was illustrated by comparison of the Cl- and $042- systems. In both cases, initial formation of the complex species was quite rapid, but the sulfate species lay at a low energy level resisting forward move- ment along the reaction coordinate. Thus chloride was aggressive, but sulfate was not in all corrosion reactions such as pitting and stress corrosion. Locking and Mayne(50) investigated the corrosion of aluminum at room temperature in solutions over pH range 2.0 - 12.0 and.found that corrosion was dependent more on the nature of the anion than on the pH of the solution. 28 Bohni and Uhlig(51) investigated the environmental factors affecting the critical pitting potential of pure aluminum, - 0.40 V (SHE). This value was not greatly sensi- tive to temperature (0 - 40°C), to small alloying additions of Mn or Mg, nor to thickness of oxide film produced by anodizing, which was more active in the high Cl- concentra- tion, but became more noble with additions to NaCl of nitrates, chromates, acetates, benzonates or sulfates which were effective as pitting inhibitors in increasing order. The pitting is relative to competitive adsortion on C1- with oxygen for sites on the metal surfaces. Inhibitors competed in turn with Cl- ions, making it necessary to shift the potential in the positive direction in order for C1. to adsorb followed by pit initiation. On the other hand, for 18-8 stainless steel, increas- ing C1- concentration shifts the critical potential to more active values as discussed by Leckie and Uhlig(52).The potential is shifted to more noble values by the presence .. 2... .— of other anions, e.g., ClO , SO , NO , OH-, at suffi- 4 cient concentrations of which act as pitting inhibitors. Lowering the temperature produces a similar effect of enobling the critical potential. It was known that borax-nitrite formulations were effective inhibitors for aqueous corrosion of iron, steel, 29 and aluminum alloys(53). It was reported that formulations (54),(55) such as molybdate-nitrite and benzoate- (56).(57) nitrite also provided effective inhibition for these alloys in tap water. Khobaib, Quakenbush and Lynch(59) have studied corro- sion inhibition effects of nitrite-borax mixtures ( as well as other combinations) on both ferrous and aluminum alloys. These were found to be effective in both distilled water and.tap water, but in chloride concentrations comparable with sea water, they were ineffective at low concentra- tions. Certain surfactants, however, significantly enhanced corrosion inhibition in water of high chloride concentra- tion. In a further study, Khobaib(59) evaluated the inhibit- ing efficacy of more than 200 formulations of organic and inorganic compounds in the presence of urine for use as airplane bilge inhibitors. The most effective combinations again involved boraxg sodium nitrite and a variety of surfactants. These formulations were developed by sys- tematic exploration of compositions suggested by Green and Boies(6o). Nguyen and Foley(61) implied that the corrosive or non-corrosive state of a metal in a given environment depended on the stability of the complex species formed on the film and on the affinity of the anion present, that is, 30 support the first step among steps established by Samuel et al(39), to react with the exposed metal. Nitrate ion was reduced to ammonia by elementary aluminum but not by alumina. Mckissick, Adam and Foley<62) reported that in specific electrolyte mixtures, i.e., specific concentra- tions and specific NaCl - NaNO ratios, the corrosion of 3 aluminum alloy Types 7075 and 2024 was increased by an order of maginitude over that in NaCl solutions of the same concentration. .Anderson and Hocking(63) had examined the behavior of aluminum during anodic polarization in electrolytes con- 2 4 -) and chloride ion. ‘taining anodic inhibitors (e.g., CrO They had established the existence of a critical potential of film breakdown above which pitting corrosion occurred. This potential was dependent upon electrolyte composition, ‘which meant that competitive adsorption of anions at the film-solution interface determined the anodic reaction. In summary, it is suggested that a variety of physicochemical properties of accelerant/inhibitors may be involved in the corrosion process. Many published studies contain hints as to which properties may be critical, e.g., ionic charge, size, mobility, ionization state, adsorption, etc., but systematic exploration of accelerant/inhibitor 31 chemistry is needed for elucidating the crucial properties in corrosion. The nitrogen system, consisting of the oxyanions and ammonium ion, presents an excellent system for systematic evaluation of these factor. D. The Oxyanions of Nitrogen The oxyanions of nitrogen are listed in Table 1 together with apparent oxidation state. Although all of them can form free acids, only nitric acid can be con- sidered stable in aqueous solution or as the pure acid; all the rest decompose with varying degrees of violence. All can be prepared as relatively stable salts, particularly the sodium salt, although the preparations of the salts themselves are.somewhat hazardous. The oxidation/reduction and inhibitive properties of these ions are related to their stability characteristics. 1) Nitrite Sodium nitrite is a readily available laboratory reagent. Nitrous acid and the nitrites are most commonly employed as oxidizing agents, although strong oxidants (e.g., electric current, MnOz, and C12) convert nitrous to nitric acid in alkaline solution. 32 Table 1. The Oxyanions of Nitrogen<64> Formula Name Apparent Oxidation State of N NH4+1 ammonium -3 N202-2 hyponitrite +1 N203-2 nitrohydroxylamate +2 1402"2 hydronitrite +2 NOZ-l nitrite +3 NO3-1 nitrate +5 NO4 peroxynitrate +5 33 structurally, the nitrite ion belongs to the point group C2v which is not linear, but bent. Two resonance structures are postulated (Figure 3). The ONO bond angle is consistent with sp2 hybridization at the nitrogen atom, but the N-O bond lengths are closer to a bond order of 2 than 1050 Important couples describing that nitrites are employed as oxidizing agents are in — + - N20 + 3H20 — ZHNO2 + 4H + 4e NO + H20 = HNO2 + e HNO + H o = NO ‘ + 3H+ + 2e- acidic solutions, and N20 + 60H = 2N02 + BHZO + 4e NO + 20H” = N02“ +HZO + e' NO ‘ + on' = NO ' + H o + 2e" 2 3 2 298 298 -1.29 volts -0.99 volt -0.94 volt -O.15 volt 0.46 volt -0.01 volt 34 0 Z \ O O / z / \ O | Figure 3. Electronic Structure of N02”. 35 in alkaline solutions. The potential in acidic solutions is positively greater than that in alkaline solutions, and the free energy for reduction in the acidic solutions is smaller than in the alkaline solutions. Thus the greater oxidizing strength in acidic solutions is apparent. 2) Nitrate Sodium nitrate is readily available. In general, the nitrate ion is planar (D symmetry) with all N-O bond 3h distances close to 1.22A. This can be represented in 'valence bond terms by resonance structures based on those illustrated by Figure 4. Molecular orbitals also can be constructed for the nitrate ion on the basis of three bonds using sp2 hybrid orbitals, the pz orbitals of nitrogen, and three oxygens combining to form a n-orbital containing two electrons. The data in the Table 2 show the wide variety of reduction reactions available and their relations to other nitrogen oxyanions and compounds. Also shown are stepwise reduction processes of nitrogen CV) to nitrogen (III). As with the case of nitrite, comparing the magnitude of poten- tials and driving force, the effect of acid media on the oxidizing power of nitrate is clearly evident. 3) Hyponitrite 36 +1: / \ +7 / \ +7 / \ Figure 4. Electronic Structure of NO3-. 37 3.9 2.0 and «Tel mod 3.0... 3.0 ~o.o| 20> mwé IIIIIII « v a a \02/86- 0286.. 0 28.6- 02 2.61 2.6.. «luOuz 0:4 841/ \361 .2 ilimofiz 5.2 .22 «as 3.6.. Eel ....w + 0.26 + ...02 ...I. -208 + .22 low; + 0.32 + ...OZmflu 1302.: + .EJ’H inc + 043 + ...OZ ...ii. iflOh + 30an -8 + 0.5 + -.02 ...... -208 + .22 ....w + 0.2.. + -.02m ...». .28 + ...0.2 ...m + 0.2... + .02.“ .ii. 202 + 0.2 -.m + 0.5 + -.02 .ii. -20.. + 02 -..a + 0.2 + -.02 H .28 + ...02 ..“m ...u + 0.28 + .028 ...... -20.. + .0.2 a , e a In ooél as. T «0.25 0.2 351/ \R. T. .2 +2052 +5.2 +52 $4 3.2.. $4.. $6.. .1. -...w + +22 + -.02 ...» 0.2..” + ..mz ewel ... ...E + +2: 1. -.02m H 0.28 + +5.2 23 3.6.. ... -..e + +5 + -.02 n... 0.3 + +2052 36> en. T. n -3 + +5 + -.02 H 0.5 + .22 26> 5 .ol n ....m 1. +22 + .02?» 0.5. + 5.2.: 33> 2.7. .u -.w + :2: + ...02m .1... 0.2... + 0.2 8.6.. a -.m + Li. + -.02 ..H 0.5 + 02 $6.. a -..m + :5 + ...02 H 0.: + .022 :3 5.6.. u 0.28 + .0.2 ...“. -.m + +2.. + -.02a ...... .8228 .532 20:28 use... $8 COH mumuuflz mo sofluoscmm you sumo Hafiuzmuom COfiwmmofixO .N OHQMB 38 Hyponitrite is known chiefly in the form of its salts,although free hyponitrite acid, H2N202 can be prepared as white deliquescent plates; however, it is explosively unstable. Aqueous solutions of the acid are relatively stable: At pH 1 - 3 and 25°C, the acid has a half-life of 16 days, but at pH 4 - 14, it is quite stable. The acid ionizes as a weak dibasic acid, k k2 + 2- 2 2 2 2 2 2H + N202 :1: Z 0 ll 3 + + a: Z O I ll 8 11 k = 9 x 10' 1 x 10' at 25°C and k2 It is reported that the first and second heats of ioniza- tion for hyponitrous acid are 3.7 and 6.1 kcal/mol(67), respectively, the sum of which equals 9.8 kcal/mol, a value lower than that obtained calorimetrically by Latimer and (68) Zimmerman (11.1 kcal/mol) . The enthalphy of formation 0 - and H N O is -12.4 and -15.4 kcal/mol of aqueous HN2 2 2 2 2 respectively(69). In higher pH solution, the hydrogen hyponitrite ion is unstable, HN o ' N o + OH". 39 Thus, the characteristics of the hyponitrite ion related to corrosion can be examined readily, particularly in mild acid to alkaline solutions. (70) in the ultraviolet region at The strong absorption 248 mp (6:3980) is characteristic of compounds containing the N=N bond, which suggests that the bonding in the hyponitrite ion is adequately described by Figure 5 with trans or cis configuration. It has been shown by Raman( 72)'(73) and infrared studies<74)’(75) that the product obtained in the usual preparative reactions is the trans isomer. Such a planar trans centrosymmetric structure belongs to the C2h point group. The hyponitrite metal salts are prepared generally by reduction of the nitrate ion, commonly by sodium amalgam, 2NaNO3 + 8Na/Hg + 4H20 ----> NazNZO2 + 8NaOH + 8Hg. Sodium hyponitrite decomposes on heating to 260°C in vacuo or 335°C at atmospheric pressure to give sodium oxide, nitrite and nitrate together with nitrogen. It is partially hydrolyzed in water or dilute acids to N20, N04, and N2 but it is oxidized by permanganate. Hyponitrites possess both reducing and oxidizing properties as indicated by the couples 40 Figure 5. Electronic Structure of N20 2 (b) cis (a) trans 2— 41 _ + - __ H2N202 — 2N0 + 2H + 2e E298 — 0.712 volt _ + ‘ _ _ HZNZOZ + 2H20 — 2HN02 + 4H + 4e — 0.86 volts + - —- N2 + 2H20 — H2N202 + 2H + 2e — 2.75 volts 2NH 0H+ — H N 0 + 6H+ + 4e“ = -0 44 volts 3 ‘ 2 2 2 ' in acidic solution, and the couples N 0 '2 = 2N0 + 2e‘ E° = -0 10 volt 2 2 298 ° -2 __ _ _ N202 + 40H — 2N02 + 2H20 + 4e — 0.18 volt __ -2 _ __ N2 + 40H — N202 + 2H20 + 2e — 1.60 volts 2NH20H + soH‘ = 5N202‘2 + 6H20 + 4e' = 0.73 volt in alkaline solution. However, the reducing properties predominate. Strong oxidizing agents will convert hyponitrites to nitrates. d) Ammonium Hydroxide 42 The extent to which aqueous ammonia exists as NH40H or as hydrated NH has long been discussed(76) . Ammonium ion 3 is very soluble in water in which it is involved in the equilibrium NH + H 0 = NH + + 0H" 3 2 4 [NH4+] [0H'] K = ---------------- = 1.65 x 10’5 at 25°C. [NH3] Solution of ammonia react with acids, taking up protons to form ammonium ions H H + .0 .. H:N: + H:x ----- > H:N:H + x'. on .0 H H For this reason, ammonium solutions give only a weakly basic reaction or, alternatively, behave like solutions of weak base. The equilibrium is displaced strongly to the left because of the extremely slight dissociation of water. There are, accordingly, only a few ammonium ions present in an aqueous solution of ammonium and correspond- ingly few hydroxyl ions. 43 E. Research Program The object of this research program was‘an experimen- tal exploration of the relations between accelerant/inhibitor properties and the electrochemistry and electrokinectics of corrosion in aluminum. A variety of inhibitor/accelerants were under study. These included (a) ordinary reagent-grade chemicals, e.g., NaCl, NaNO NaNO 2’ 3' and NH4OH, (b) the Lynch formulation mentioned in Section III-C and NaZB4O7, and (c) other compounds of nitrogen, specially sodium hyponitrite ( Na2N202 ). A potentiodynamic polarization technique was used to investigate the effectiveness of inhibitor when a 7075 - T6 high-strength aluminum was exposed to solutions with and without chloride ions. III. Experimental Test Procedure And Apparatus lMany corrosion phenomena can be explained in terms of electrochemical reactions, hence electrochemical techniques can be used to study these phenomena. Measurements of current-potential relations under controlled conditions can yield information on corrosion rates, coating and films, passivity, pitting tendencies and other important data. Potentiodynamic anodic polarization is the charac- terization of a metal specimen by its current-potential relationship. This is fast and reproducible, and clearly indicates the open circuit potential, the corrosion current density in the passive range and pitting potential. The specimen potential is scanned slowly in the positive- or negative- direction. These measurements are used to deter- ndne corrosion characteristics of metal specimens in aqueous environments. A complete current-potential plot of a specimen can be measured. A. Experimental System The experimental system consisted of: (1) Potentiostat and Logarithmic Current Converter (2) Waveform Generator and Recorder (3) Corrosion Cell (4) Electrometer Probe 44 45 (5) Heat Source The experimental arrangement used in this study is shown in Figure 6. (1) Potentiostat and Logarithmic Current Converter The potentiostat was a Princeton Applied Research (PAR) Model 173, an electrochemical measuring instrument. Operated as a potentiostat, the potential between electrodes of an elctrochemical cell is held constant: operated as a galvanostat, the current through a cell is held constant. The instrument featured a current capability of one ampere with compliance voltage as high as 100 V in either polarity and a slew rate of 10 V per microsecond. It incor- porated two independent built-in potential/current sources, each adjustable to any voltage in range of i4.999V as well as logic and switching circuity for controlling the sources from the front panel or by externally derived triggers. Two additional external potential/current programming signals might be added to those provided by the instrument, and a wide variety of triggering and switching waveforms might be employed to control the applied potential/current programs. The logarithmic current converter was an EC & G PAR Model 376. The Model 376 was ideally suited to use in those applications where the current varied over many orders of 46 FUNCTlON GENERATOR (WAVETEK) MODEL 175 RECORDER ‘ Y AXIS |NPUT (SOLTEC) X AXIS INPUT VP-6414S I MODEL 173 376 EXT. ELEC. EXT. suc. IN\ CELL PROBE (A or B) \ i Y BLACK (NOT USED)—> REO——> GREEN CORROSION CELL SAMPLE Figure 6. The Experimental System. 47 magnitude in the course of the experiment. Like the Model 176, the Model 376 provided an output voltage proportional to the cell current. However, it additionally provided an output voltage proportional to the log of the cell current. There was provision for displaying the current of the log of the current on the Model 173 panel meter. In displaying the current, full-scale meter deflection indicated a cur- rent of 1, 2, or 5 times the selected Current Range, the same as for the Model 176. In the Log Meter mode, full- scale deflection indicated a current 1, 2, or 5 decades below the selected current range. (2) Waveform Generator and Recorder Waveform generator used was Wavetek Model 175. The Model 175, Arbtrary Waveform Generator generates any waveform that can be expressed as a function of time. The working Random Access Memory (RAM) had four sets of 256 ad- dresses and each address accepted an 8 bit word. This corresponded to a 1024 (time) by 255 (amplitude) matrix in which to draw waveforms. These storage addressess can be loaded manually by front panel controls or remotely loaded v_ia_ the GPIB interface. There are four fixed waveforms: sine, square, triangle and ramp. In this study, the ramp waveform was used. The recorder used was Soltec VP-6414S X-Y Recorder. This recorded a relation of the form f(x) = y between two 48 variables on A 4 chart paper. The range was from 2 mV/cm to 1 V/cm. (3) Corrosion Cell The Model K47 Corrosion Cell System of PAR was used in this study. Figure 7 shows a corrosion cell for polariza- tion measurements. (a) Working Electrode The working electrode for this investigation was a . . 2 square shaped spec1men w1th surface area of one cm . (b) Reference Electrode .A saturated calomel electrode (SCE) was used as the reference electrode. The SCE is not easily poisoned or con- taminated and is insensitive to electrolyte composition because of its design. A solution bridge from the Luggin capillary to another beaker containing the reference electrode made the electrical connection between this electrode and the solution. This liquid junction was used to avoid the contamination of the solution. The SCE incor~ R tip designed specially to provide porated an unfired Vycor ultra-low liquid leakage rates and minimum IR iii-1) Sodium amalgam The Fieser procedure<81) gave especially pure material: A clean sodium piece (6.9 g for an amalgam con- taining 2% Na, 10g for one containing 3% Na) was placed in a 250 m1 three-neck round-bottom flask. The two side necks carried nitrogen inlet and outlet tubes, while the center 57 Table 5. The Test Solutions Solutions Concentrations(wt.%) Distilled Water NaCl 0.1 NaNO2 0.05, 0.1, 0.5, 1.0, 3.0 NaNO2 + NaCl (0.05,0.1), (0.1,0.1), (O.5,0.1) (l.0,0.1), (8.0,0.1) NaNOz + Na2B4O7 (0.1,0.1), (0.1,0.5), (0.1,3.0) NaNO2 + NazB4O7 + NaCl (0.1,0.1,0.1), (0.1,0.5,0.1) (2.0,2.0,0.1), (3.0,3.0,3.0) (52) Lynch Formulation 0.513 Lynch Formulation + NaCl (0.513,0.1), (0.513,3.5) NaNO 0.05, 0.1, 0.5, 1.0, NaNO + NaCl (0.05,0.1), (0.1,0.1), (O.5,0.1) (l.0,0.1), NazNZOZ 0.05, 0.1, 0.5 Na2N202 + NaCl (0.05,0.1), (0.1,0.1), (O.5,0.1) NH4OH 0.05, 0.1, 0.5, 1.0 NH4OH + NaCl (0.05,0.1), (0.1,0.1), (O.5,0.1) (1.0,0.1) 58 Table 6. Lynch Formulation(52) Formular Recommended Weight Aqueous wt% Borax NaZB4O7'10H20 381.37 Sodium Nitrate NaNO3 84.80 Sodium Nitrite NaNO2 69.00 Sodium Metasilicate Na2SiO3v5H20 212.74 Pentahydrate Sodium Hexameta (NaPO3)XNa20 Phosphate MBT, Mercapto- 167.25 benzothiazole Total 0.002 0.001 0.513 wt.% 59 neck carried a dropping funnel containing 340 g(25 m1) of Hg. The flask was thoroughly flushed with N2, and 10 ml of Hg then was added. The flask was heated on an open flame until the start of the reaction. Additional Hg then was slowly added, with minimum additional heating. After the addition, the hot molten amalgam was poured onto a clean plate and broken up into pieces while still hot and brittle. Amalgams containing less than 3 wt.% Na are not too sensitive to air: however they must be stored in an air- free atmosphere. Complete liquefaction occurs at following 0 O (liquidus) temperatures: 0.5% Na, 0 C; 1.0% Na, 50 C; O O O 1.5% Na, 100 C; 2.0% Na, 130 C; 2.5% Na, 156 c; 3% Na, 0 O 250 C; 4.0% Na, 250 C; 4.0% Na, 320 °C. iii-2) Sodium Hyponitrite, NaZNZOZ' 9H20 The metal salts are generally prepared in variable yield by reduction methods, sodium amalgam being a favorite reducting agent: 2NaNO3 + 8Na + 4H20 = NaZNZOZ + 8NaOH. 170.0 184.0 72 106.0 320.0 60 Sodium amalgam was added with stirring or shaking to 3 in 250 ml of H20. an ice-cooled solution of 85 g of NaNO The amalgam was prepared by dissolving 58 g of Na in 4000 g of Hg (1.43 wt.% Na). When three quarters of the amalgam had been added, the cooling was discontinued and the remainder of the amalgam was added at once. Shaking was continued for 10—15 mimutes, during which the temperature increased tol40 0C. When the temperature began to fall the mixture was poured into a closed flask with a narrow neckq then the first flask was rinsed with 2-3 ml of H20. The washings were combined with the main solution and the whole Shaken vigorously for about ten minutes. The solution was decanted from the Hg and placed over H2804 in a vacuum desiccator at 35°C to 40°C in order to remove all the NH3. The Na N 0 9H 2 2 2 2O separated during this operation. It was suction-filtered on a fritted glass filter, washed with al- cohol at a temperature above 10 0C to remove traces of NaOH, and dried in a desiccator. Sodium Hyponitrite (Na2N202°9H20, 268.14) was small granular crystals or plates. It lost water of crystal- lization under vacuum. IV. Experimental Results A. Distilled Water and Sodium Chloride Figure 8 shows the anodic polarization behavior of Al 7075-T6 in distilled water at room temperature. The profile is continuous with no discontinuity and may be adaptable to Tafel analysis and the corrosion current density can be found. In Figure 9, anodic profile is presented for 0.1 wt.% NaCl at room temperature. The figure shows that this profile is continuous with no discontinuity. It is known that the corrosion potential of corroding 7075-T6 Al alloy depends strongly on the solution concentration, i.e, in- creasing corrosion potentials correspond to decreasing Cl- concentration(82) . With comparison of anodic polarization behavior between the distilled water and sodium chloride, the current density is increased by three orders of mag- nitudes. B. Effects of Inhibitors The effectiveness of inhibitors are based on the corrosion current density and pitting potential values obtained from the anodic polarization curves. The measure of an effective inhibition, therefore, is its ability to reduce the anodic reaction (corrosion) rate to an accept- ably low level, and to raise the pitting potential to more 61 62 ANE6\— O C C'— r- P- ‘- r- " f' ._ ___O E i A .. .. N h—- l—-.-—- < P p I— '- < - r- _ r E .. .. v FF E _ -l u) I Z O r- P- '— Z C 4 DJ E P- O )— N LLJ C.‘ L—IO [I : :1. D: : _ : 3 .. 0 r- L) D D- '- Z C5 7 3 Po h F- r— h- >— - r- v- I I I I I 7 I I I I I r I T I I I T I '0 O O O O O '— | I (303 SA A) 'lVLlNBlOd Figure 9. The Anodic Polarization of 7075-T6 Al Alloy in 0.1 wt.% Sodium Chloride. 64 positive value than the open circuit potential. a) Sodium Nitrite, Borax and Lynch Formulation Figure 10 shows the anodic polarization behavior of 7075-T6 A1 alloy in inhibitor solutions for concentrations between 0.05 wt.% and 3.0 wt.% NaNOZ. Sodium nitrite alone is an effective inhibitor but clearly exhibits a reverse effect at concentrations above 0.1 wt.% where corrosion currents increase. Figure 11 shows the effect of concentra- tion.of nitrite on- L]_ )- - r- r I )- b- 0 )- >Z_ N/‘N - —] p- E r :3 :"- 0 E m : \ - c; I; <( - 1- E - F' V )- TI E - — In : :O - -.. 2: Z : UJ - - C) .. t Z Cr DU r-— "' m E 59 Cr t Z 3 L . L) h F' \ "1’ E' E’o t :r r- )- v I I I I [Tfifi—I I ITTj Tor I I '0 O. 0. 0. O. 0.5 N r- O ‘— CT] (308 SA A) 'IVllNI-llOd Figure 15. The Anodic Polarization of 7075-T6 Al Alloy in Lynch Formulation. 72 Nov -PPPp h k pl 0— r-ppPP P 14Eo\ 40H . (1) Sodium nitrite was found to be an excellent inhibitor (94) while it was for aluminum in engine cooling systems found to accelerate the corrosion of aluminum in alkaline media(95). Sodium nitrite protects 7075-T6 Al alloy in 96 distilled water with good passivity at a tested concentra- tion although the passive current density is increasing with the concentration (Figure 10). The protec- tion given to aluminum alloy in distilled water by sodium nitrite is a direct result of action on the anodic dissolution/passivation procedures. Figure 29 shows a polarization diagram in aqueous system showing the anodic polarization curve with the development of passivity. The addition of an oxidizing inhibitor such as sodium nitrite shifts the cathodic polarization-curve in the positive direction, thus stabilizing the corrosion potential in the passive region. The inhibition species are adsorbed on the metal surface and these adsorbed species develop the pas- sivity, that is, insoluble oxide which protects the metal surface. The presence of sodium chloride affects the inhibition of corrosion of 7075-T6 Al alloy. In the presence of sodium chloride, the inhibition of the corrosion depends on the ratio of nitrite to aggressive ion (Figure 12). Actually, the inhibition is not seen at low concentration of sodium nitrite but it is inhibited at higher concentration. It may be suggested that the passive film formed by the low con- centration of sodium nitrite is still weak and ineffective against chloride attack without high concentration of inhibitor. The breakdown in the presence of insufficient nitrite is localized and severe, which can give rise to intense pitting and perforation. If no passivation results 9T7 COY‘ \ CO!" n * a 1cor 1cor LOG i Figure 29. Schematic Polarization Diagram for Oxidizing Inhibitor. 98 at low concentration, as shown by the experimental results, metal dissolution increases and then the passive film becomes weak. The inhibition process of sodium nitrite is accom- panied by a shift in the corrosion potential in the positive direction, thus sodium nitrite is an anodic in- hibitor. The inhibitor as sodium nitrite has no effect on the cathodic reaction where the current density increases with the concentration of sodium nitrite below -1.0 V as shown Figure 17. The cathodic reaction is No ' + 8H + 6e ----- > NH + 2H 0. (J) The corrosion in non-inhibited solutions of nitrite is usually accompanied by a marked rise in pH (from 6-7 to 9- 10) during the experiment. This also is shown in the results of Mercer et al(32) who investigated the corrosion inhibitors for mild steel in neutral solution. The rise in pH is presumed to be due to cathodic reduction of nitrite ion to ammonia such as reaction (J) during the corrosion process. Cathodic reduction of low concentration of nitrite in chloride-containing solution could lead to eventual disappearance of nitrite from the solution: corrosion then should be of the type occurring in solutions free from inhibition. Increased corrosion rates in non-inhibiting 99 :nitrites solutions have been reported by Sussman et a1(31) (96), (98). , Patterson and Jones and Conoby et al The passive current density of a combination of sodium nitrite and borax is independent of concentration of borax (Figure 13). With the comparison in the case of sodium nitrite, it is increasing with the concentration. In the presence of sodium chloride, the effectiveness of the inhibition is eliminated even in the high concentration, although the pitting potential is increasing as the con- centration is increasing (Figure 14 and Table 7). It is suggested that the film formed by these inhibitors still is unstable in the presence of sodium chloride as in the case of sodium nitrite. Bhansali et al(83) showed that crack growth rate is lowered by an order of magnitude after the addition of a combination of sodium nitrite and borax to chloride solution.'rhese inhibitors have essentially eliminated environmentally assisted crack growth in chloride solutions. The Lynch formulation shows results similar to those of the nitrite and borax combination. The pitting potential of the Lynch formulation is between that in the 0.5 wt.% borax and 0.1 wt.% borax (Figure 15). This results from the fact that 0.35 wt.% borax and 0.05 sodium nitrite of Lynch formulation is between the above concentration of borax. (58) According to Khobaib et al the borax-nitrite base 100 Table 7. Pitting potential of different test solutions with 0.1 wt.% NaCl Concentration (wt.%) Pitting potential (V vs SCE) NaNO 8.0 1.25 NaNo2 + Na2B4O7 (0.1,0.1) -o.53 (0.1,0.5) -o.14 LYNCH 0.513 -0.48 NaNo3 0.5 -o.01 1.0 1.08 8.0 1.41 Na2N202 0.05 -o.50 0.1 -0.63 0.5 0.21 NH4OH 0.05 -o.77 0.1 -o.59 0.5 -o.24 1.0 -O.18 101 formulation provides protection to aluminum and steel, whereas this protection is lost with increasing concentra— tion of chloride. They suggested that the passive film formed by the inhibition is weak and ineffective in the presence of chloride of high concentrations. But this inhibitor can give protection to aluminum and steel in high concentration of chloride if some surface active agents are added even in low concentration to this inhibitor formula- tion. It may be suggested that small additions of surface active agents interfere with the dissolution reaction by interacting synergistically with the passive film already ‘by the inhibition formulation, which leads to a stronger and possibly thicker adsorbed protective film. Thus, the protective film formed by this inhibitor formulation with surface active agents slows down the dissolution reaction, as well as blocks aggressive anions from attacking the metal surface. Once certain anions are adsorbed, they react chemi- (98) cally as shown by Augustynski from the X-ray Photoelectron Spectroscopy(XPS). The results shows that N03- is reduced to NH4+ and the presence of nitrate retards the adsorption of chloride. It also is reported that nitrate is reduced by elemental aluminum in alkaline solution(99) 102 8A1(0H)4' + 3NH3, and the reduction potential for the couple N03-/NH4+ is 0.47 v<1°°2 It has been confirmed (61) that the formation of NH3 results from the chemical reaction between N03- and the aluminum metal. When a solution of NaNO3 was brought into contact with aluminum metal, the formation of ammonia could be recognized after 16 hours by ammonia odor. Upon shaking the suspension vigorously, the formation of NH3 was quickly accelerated within a few minutes. It was reported(62)’(101)that NH was formed on 3 rapidly corroding aluminum alloys 2024 and 7075 in certain mixtures of chloride and nitrate, whereas aluminum alloy 1199, essentially pure aluminum, didn't show the synergis- ‘tic effect of chloride accelerating film dissolution. This resulted from the presence of Cu and Zn as alloying ele- ments which acted as local alkaline cathodic site. Chemical reactions between NH3 produced by reduction of NO3 and intermetallic compound are CuAl2 + 4NH3 = Cu(NH3)42+ in alloy 2024 103 _ -13 Kinstability— 2'4 X 10 2+ . 14an2 + 4NH3 = Zn(NH3)4 in alloy 7075 _ -1o Kinstability_ 3'46 X 10 ° Both complexes are very soluble and will accelerate the attack. The existence of a specific concentration of mix- ture of chloride and nitrate in which corrosion was accelerated was because of the competition between inhibi- tion with nitrate and the acceleration with chloride. Figure 20 shows the effectiveness depending on the concentration of sodium nitrate in the chloride-containing solution. As the concentration of nitrate is increasing, the inhibition among the competition between inhibition and acceleration becomes dominated and the pitting potential is increasing (Table 7). The concentration of 0.1 wt.% NaNO3 in the chloride-containing solution shows the synergistic effect as explained above. The pitting potential has been explained in terms of Cl- penetrating an oxide film which (102),(103) covers the metal surface or in terms of competi- tive adsorption of Cl- and oxygen for sites on the metal (52)1(104) surface which is dependent upon the oxide film. The thicker oxide film, the smaller is the electric film impelling C1“ to penetrate it and the longer is the diffusion path of Cl- from electrolyte to metal surface. 104 (51) also noted that NaNO3 addition to NaCl Bohni and Uhlig solutions moved the pitting potential in noble direction corresponding to improved corrosion resistance and the relationship between chloride activity and inhibiting nitrate activity followed the equation log(Cl-) = 0.6510g(No3') - 0.78. Challender‘los) also showed that the rate of reduction of nitrate is related directly to the concentration required to achieve passivation for Al alloy. There seems to be no threshold for chloride concentra- tion below which pitting would never occur and the addition of inhibition, NO3- would delay but not prevent the.onset of pitting as discussed by Berzins, et al(45). They showed that the delay time before the onset of pitting is logarithmically related to inhibitor concentration. Temperature as might be expected has considerable influence on corrosion rate. Figures 30 and 31 show that the current density logarithmically is dependent upon the temperature in the solution of sodium nitrate and sodium :nitrite although the temperature dependence of sodium nitrite is more scattered than that of sodium nitrates. These figures give the slopes which are related to the 105 Io IITI WIIITI I 'IIIIIII I PIIIIIjtT IIIITIITEI [HUTIT I r5 ID E I '— on 0 0 b3). 2: O . Z _Q 1’ pr) _ o o r") -0. r- "'1 a) I— C 0 Phi a) t m I\ W ! w r //00 E03 3 A/ .. / L_In ‘- IIIII IT IIIIIITI I WIIII I 7 WWII I I PIIIII I I PIIITW IEI 01 “o 0 '- T I '1' 7 .. '- o o o o (,wO/vw) MISNEG lNEHHflO Figure 30. Dependence of the Current Density on Temperature in Sodium Nitrate. INVERSE TEMPERATURE x 1000 106 ,. (o IIITII r PIIIIII I IIIIIII I I 'IIIIIII I lIIIIITT I IIIWIII I {(5 - 10. I'm 2: U I! r m. 8 C) N I" L— . I.” X UJ .... — a on . [I 1"") :3 I— < _ 1.0. DC F) DJ 0. 2 HJ m [n EN 0: 0J . i I- I—N' — (D -- O 0 I—N In I- *0} ¢ IIIIII I I ITTIITI I I 'IIIIII I I IIIIIIII j IIIIIIII T PIIIII I I N “Q 0 " T I II" T .. "' O O O O F P \- P (2000/1101) 11151130 1113111100 Figure 31. Dependence of the Current Density on Temperature in Sodium Nitrite. 107 Arrehenius activation energy 2.68 kcal/mol for nitrite and 4.88 kcal/mol for nitrate. Hyponitrite has both reducing and oxidizing properties but the reducing properties predominate. As indicated by Figure 22, the hyponitrite has a strong passivating effect on 7075—T6 Al alloy. Although the passive current density of hyponitrite is larger than that of nitrite, the region of passivity is wider and the pitting potential is higher (Table 7). This results from the larger reduction potential with comparison between potential of nitrite and hyponitrite. The passivity breaks down in the chloride- containing solution (Figure 23), but hyponitrite has the strongest effectiveness of inhibition among nitrite, nitrate and hyponitrite when the inhibition effectiveness is compared at the same concentration, i.e., the hyponitrite has the strongest passivity film against chloride attack. It has been reported(106)that dry ammonia has no action on aluminum even at elevated temperature. When ammonia is moist or in solution as ammonium hydroxide, the corrosion rate is low for all concentrations at tempera- ‘tures up to 120°F(107I Ammonium hydroxide shows the strong effectiveness in passivation of 7075-T6 Al alloy, although the passive current density is increasing with the con- centration because a thin protective oxide film forms after exposure to ammonium ion and prevents further attack (Figure 24). The region of passivity is as wide as the 108 sodium hyponitrite (Table 7) . Results similar to those of sodium hyponitrite are obtained in chloride-containing solutixni (Figure 25). Thus, aluminum is typically used for ammonia refrigeration tubing, storage vessels, sprays and molds in contact with ammoniacal rubber latex. Aluminum equipment also is used widely in contact with ammonium bicarbonate, ammonium carbonate, urea and ammonium nitrate. VI . Conclusion Generally, the inhibitors used in this study i.e., and Na B O NaNOz, Lynch Formulation, combination of NaNO2 2 4 7, NaNOB, NaZNZOZ' NH4OH show good passivation on 7075-T6 Al alloy in distilled water. In the presence of chloride, the passivity is dependent on the concentration of the inhibitor in the solution which determines the competitive action between the inhibitor and the aggressive ion. The inhibitive efficiency of oxyanions of nitrogen based on the pitting potential is in the decreasing order of: N202- > + -. _ NH4 > N03 > Lynch Formulation > combination of NO2 and The current density depends logarithmically upon the temperature in the solution of sodium nitrite and sodium nitrate, the acivation energies of which is 2.68 kcal/mol and 4.88 kcal/mol, respectively. 109 1. 2. 10. 11. 12. 13. 14. 15. 16. LIST OF REFERENCES C. Schonbein, Pogo. Ann.. 44, 73(1838) A. Finklestein, Z. Physik. Chem., 39, 91(1902) W.J. Muller, Z. Elektrochem., 10, 518(1904) M. Le Blanc, Z. 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