Determining the Effect of Environmental Conditions on Iron

Laboratory Experiment pubs.acs.org/jchemeduc

Determining the Effect of Environmental Conditions on Iron Corrosion by Atomic Absorption Esteban Malel†,‡ and Deborah E. Shalev*,†,§ †

The Jerusalem College of Engineering, 26 Shriebom Street, Jerusalem, Israel The Institute of Chemistry and the §Wolfson Centre for Applied Structural Biology, The Hebrew University of Jerusalem, Safra Campus Givat Ram, Jerusalem, Israel

‡

S Supporting Information *

ABSTRACT: Iron corrosion is a complex process that occurs when iron is exposed to oxygen and humidity and is exacerbated by the presence of chloride ions. The deterioration of iron structures or other components can be costly to society and is usually evaluated by following the properties of the corroding material. Here, the iron ions released into solution due to corrosion were detected directly by atomic absorption and their concentration was determined using a calibration curve. Iron corrosion was measured in samples immersed in aqueous solutions that differed in salinity (increasing NaCl concentrations), pH, temperature and presence of oxygen, and under the cathodic protection of a zinc ingot. The corrosion of the iron samples in solution was accelerated by high salinity and temperatures, low pH, the presence of chloride ions and oxygen, and the absence of cathodic protection. Material deterioration due to exposure may be arrested or enhanced by understanding the conditions that expedite the reactions. The experiment was performed by third-year material engineering students and would also be appropriate for an upper-level analytical lab. KEYWORDS: Second-Year Undergraduate, Upper-Division Undergraduate, Analytical Chemistry, Chemical Engineering, Laboratory Instruction, Hands-On Learning/Manipulatives, Atomic Spectroscopy, Electrochemistry, Metals, Oxidation/Reduction

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coating layer is usually porous, allowing corrosion to continue at a slower rate. Corrosion of iron may take place in the atmosphere, splash zones, and while immersed14 (see Supporting Information for the reduction−oxidation reactions of corrosion). Iron is often exposed to environmental conditions that vary according to weather, proximity to the sea, humidity, and pollution. The conditions of atmospheric corrosion change dramatically from winter to summer and from rural to urban areas. Industrial areas have more corrosive environments due to acid rain. Marine surroundings have increased chloride concentrations in the air causing pitting corrosion (see Supporting Information and supporting Figure S1). Hara and co-workers showed how deicing with salt impacts bridge corrosion during the summer.9 The splash zone situated along the seawater line is a harsh environment combining humidity, high salt concentrations, and oxygen. Moreover, corrosion may be accelerated by mechanical erosion in these zones. Iron structures immersed in water or buried in soil are also subject to corrosion, albeit under more stable conditions that depend on the soil or water composition. The importance of measuring corrosion of iron and other metals under different conditions has been addressed in this Journal using anodic polarization,19,20 weight loss,21,22 and through observation,23 and a laboratory experiment using atomic absorption to quantitatively analyze iron in multivitamin tablets was developed.24 The biocompatibility of corrosion of different metal alloys in biological systems has been studied

orrosion and environmental behavior are important subjects in material sciences as these govern the change in mechanical properties and appearance of materials over time. The deterioration of iron girders or other iron components due to exposure can be evaluated by following the corrosion process1,2 closely and may be arrested or enhanced by understanding the conditions under which corrosion occurs. Many studies have addressed the chemical processes involved in iron corrosion.3−6 Most methods have been aimed at detecting signs of corrosion on the iron itself,7,8 in the form of changes in properties, scaling, or pitting. The experiments presented here measure the concentration of iron released into solution and are complementary to stress and fracture experiments that measure the effect of corrosion on the iron material itself. Iron corrosion is a complex process1,2 that occurs when iron is exposed to oxygen and humidity. The process is accelerated7,9,10 in the presence of chloride ions that react to give a number of intermediate iron chloride products, which subsequently oxidize.11−13 Corrosion processes diminish the tensile strength of iron by making it more brittle and reducing its mass due to changes in the chemical composition of the metal.1,14 Corrosion depends on a number of parameters including temperature, pH, and salt composition. The corrosion process eventually yields a layer of oxide that provides the iron with a degree of protection and partial passivation occurs. Effort is made to prevent or retard the corrosion process through chemical coating.15−18 However, the © 2012 American Chemical Society and Division of Chemical Education, Inc.

Published: December 6, 2012 490

dx.doi.org/10.1021/ed3001107 | J. Chem. Educ. 2013, 90, 490−494

Journal of Chemical Education

Laboratory Experiment

Figure 1. Student results for dependence of corrosion on environmental conditions: (A) the sodium chloride salt content of the solution, showing triple distilled water and salt concentrations equal to those of drinking water, irrigation water, and seawater; (B) the temperature of the solution, showing salt concentrations equal to those of irrigation water at 2 and 19 °C; (C) the pH of the solution, showing salt concentrations equal to those of irrigation water at pH 2, pH 7, and pH 12; (D) deaeration, showing salt concentrations equal to those of a deaerated solution and irrigation water; (E) zinc galvanization on corrosion, showing salt concentrations equal to those of irrigation water with iron contact with a zinc rod and with none; and (F) reproducibility of corrosion measurement as a function of time showing results for 10 nails with ∼2% weight differences in irrigation water solution at 19 °C. All error bars represent two standard deviations.

degree of corrosion under the chosen conditions was measured by atomic absorption spectrometry and the dissolved iron concentration was calculated as a function of time. At the end of the 4-h lab period, results from all student groups were shared and compared, and a general discussion was held. The students were introduced to the effects of corrosion and to atomic absorption spectrometry as an analytical tool, and their results matched theoretical expectations.

using atomic absorption spectroscopy to detect metals that were released to solution by different alloys.25,26 Atomic absorption27,28 is widely used to detect minute quantities of metals in corrosion processes,29 biological fluids,30 pharmaceuticals,30 food,31 industrial processes,32 and pollutants,33 among others. In this experiment, iron ions corroded from a solid sample were directly detected using atomic absorption and their concentration was determined using a calibration curve. Iron samples were completely immersed in solutions in which the electric potentials depended on the composition of the solution. Solutions differed in salinity, pH, oxygenation, and temperature. Finally, an iron sample was exposed to a corrosive environment under the cathodic protection of a zinc ingot.

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Measurements

Pure, clean, nonanodized pieces of iron were placed in 100 mL of solution simulating different corrosion conditions and gently rocked (see Supporting Information and Figure S2, therein). The solutions were sampled for approximately 2 h every 10 min and measured at the end of the experiment by atomic absorption spectroscopy (Perkin-Elmer, AAnalyst 400, Fe wavelength 248.33 nm) against a calibration curve prepared for a range of 0.01−10 mg/L (ppm) from ferric chips dissolved in 1:1 HCl/triple-distilled water (TDW). All solutions were prepared in TDW.

EXPERIMENTAL DETAILS

Procedure

The experiment was performed in an advanced chemistry course by classes of third-year material engineering students. Each pair of students picked a number of experimental conditions to examine. The students made five solutions of known iron ion concentrations to make a calibration curve. The 491




Saline Solutions

temperature was found to increase the degree of corrosion (Figure 1B), whereas corrosion in neutral and basic solutions was negligible compared to corrosion under acidic conditions (Figure 1C). The rate of corrosion slightly diminished in deaerated water relative to regular water (Figure 1D), and corrosion was essentially arrested when the iron sample was in contact with a zinc rod (Figure 1E). The standard deviation of the measured results was measured in 10 repeat experiments and was found to change with time between approximately 0.03 ppm at the initial times up to 0.23 ppm after 2.5 h (Figure 1F). The time-dependent value was used to present an error of two standard deviations shown in this pane and all others.

NaCl solutions were made to give 8.5 mM (upper advisable salt concentration in tap water), 34 mM (upper advisable salt concentration for irrigation water), and 600 mM (approximate concentration of salt in seawater). Temperature Dependence

Sodium chloride solutions, 600 mM, were equilibrated at 2 °C and room temperature prior to immersing iron samples. pH Solutions

Phosphate buffered solutions with pH values of 2, 7, and 12 with identical ionic strength values of 600 mM, were prepared to study the influence of pH in iron corrosion. As ionic strength can also affect corrosion, this variable was kept constant in solutions of different pH.

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DISCUSSION The corrosive action of chloride ions has been the focus of numerous studies.4,7,9,12,13,34 Corrosion of iron in marine surroundings presents a major problem for industry. Whereas iron passivation can arrest the process, passivation does not occur in solutions with high chloride concentrations where corrosion is intensified.7,10 The mechanism of chloride and chlorine oxidation in aqueous solution is still debated. Some theories suggest that chloride adsorbs to the metal surface7,13,14 and penetrates the passivation layers, thereby enhancing corrosion.1,7,15 Chloride ions have been shown to integrate into the Fe3O4 crystal lattice, causing the well-known green rust.4,7 Although the mechanism is not fully understood, environments rich in chloride ions increase pitting corrosion (see the Supporting Information) and prevent the formation of a perfect passivation layer. Temperature increases the rate of all chemical reactions as seen from the Arrhenius equation. Increasing temperature can induce different behavior of the corrosion rate of metals (Figure 1B). For example, nitric acid causes a rapid exponential increase in the rate of iron corrosion in response to increasing temperature. At room temperature, iron exposed to nitric acid is in the immune state very close to the transition region. A small change in temperature increases the oxidative power of the acid, causing a rapid increase in the corrosion rate.14 The chemical equilibrium depends on both the electrode potential of the reaction and the H+ concentration; their combined influence can be plotted in a single diagram (a Pourbaix diagram) where the pH is plotted on the x axis and the potential on the y axis. A simplified Pourbaix diagram of the Fe/H2O system is shown in Figure 2. The stable phases in the diagram are separated by solid lines. The dotted lines A and B represent the region of stability of water. Below line B, water is

Deaerated Solution

Two methods were compared and found to give essentially identical results. In the more rigorous method, a solution of 34 mM sodium chloride was prepared where the water was boiled for 10 min, cooled to room temperature in a closed container, the salt was dissolved, and the solution was covered with a layer of impermeable silicone oil. The easier method consisted of bubbling nitrogen into the preprepared salt solution for 10 min. The latter method may suffer from a slight increase in salt concentration due to evaporation; however, this effect was not observed. The deaerated solution and one prepared with regular TDW were used to investigate the effect of dissolved oxygen on corrosion. Cathodic Protection Effect

Iron samples were placed in direct contact with a zinc rod in a 600 mM salt solution. Reproducibility of the Measurement

The reproducibility of the measurement was assessed by measuring the corrosion of 10 iron nails weighing within 2% of each other, in irrigation water solutions (34 mM NaCl solutions). The values were measured over 2.5 h and the standard deviations as a function of time were used to provide error estimations for all measurements.

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HAZARDS Hydrochloric acid, phosphoric acid, and ferrous chloride are corrosive and must be handled appropriately. RESULTS The concentration of iron ions was measured by the students by atomic absorption as a function of time in differing solutions and under differing conditions. The concentrations of iron were calculated using calibration curves prepared by the students (e.g., Supporting Figure S4) and are presented as a function of time. Some of the measurements were within experimental error of each other. These values would differentiate significantly with time; however, the expected trend is already evident within the limited duration of the measurement. An aggregate of student results measured under different experimental conditions is presented in Figure 1. The rate of iron corrosion was found to be dependent on the degree of salinity of water (Figure 1A). Little corrosion was seen in TDW, and drinking water and irrigation water showed increased corrosion, although average values in each set were within experimental error of each other. As expected, the seawater showed the highest degree of corrosion. Elevated

Figure 2. Simplified Pourbaix diagram of Fe−H2O system.14 492



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unstable and decomposes to H2. Above line A, water is also unstable, but decomposes to O2. In the area between lines A and B, water is stable. Corrosion of iron immersed in water is possible in two regions that correspond to the transition of Fe to Fe2+. According to the Pourbaix diagram, Fe(s) is thermodynamically stable, “immune to corrosion”, at potentials below −0.7 V (vs a normal hydrogen electrode, NHE) and pH values between −2 and 10.1,14−17 If the potential of the system is raised over approximately 0.6 V and the pH increased to alkaline values, the metal itself will not be stable but its surface will become covered by a stable oxide film (Fe2O3, Fe3O4). The degree of protection that the oxide layer provides governs the degree to which the film shields the metal from the solution and further corrosion.2,14,35 The rate of corrosion was expected to diminish as the pH increased and eventually stop due to passivation above pH 8 for an aerated solution. There was only a slight difference between the neutral and basic solutions, although the basic solution showed less corrosion. The presence of oxygen in the solution raises the potential of the iron, thereby increasing the rate of corrosion (Figure 1D) of the samples that were at pH 7. The solubility of oxides formed by iron corrosion in aerated solution is also pH dependent. At pH values above 8, passivation of the metal is achieved by forming Fe3O4 and Fe2O3.1,14 Galvanic protection is used to prevent corrosion in corrosive environments.16,17,36 Coupling iron to zinc (Supporting Figure S5) lowered its potential, by shifting it to the immune zone of the Pourbaix diagram (see Figure 2), which depicts the species present in an aqueous electrochemical system. Because Zn has a lower standard reduction potential than Fe, its oxidation is thermodynamically more favorable, as in Galvanic cells. When coupled, Fe and Zn have the same reduction potential, lowering the Fe potential and increasing the Zn potential. Therefore, Zn acts as a sacrificial anode. It is oxidized and produces a flow of electrons from the Zn to the Fe where H+ is reduced.15 The effect of the actual surface area of the sample available for corrosion is difficult to estimate. Iron nails, of the same size and shape, showed weights within 2% of each other. Figure 1F shows the iron concentrations resulting from corrosion with error bars of ±2 standard deviation units from the average value. The geometry of the sample and uniformity of the surface are likely to be responsible for variations in rate. The results measured by the students for the various corrosion conditions were larger than the variability shown among samples. The students succeeded in measuring the dependencies in a robust manner. The vast majority of the fit calibration curves had R2 values above 0.990; lower values were invariably due to errors in dilution calculations, most of which could be traced. The resulting graphs clearly showed differences in corrosion, which could not be detected by visual observation. The students liked being able to choose the studied experimental conditions and discussed real-world applications.


ASSOCIATED CONTENT

S Supporting Information *

Information on materials used; detailed experimental procedures and atomic absorption acquisition parameters; students handouts including theoretical background on oxidation− reduction reactions that occur during iron corrosion and the process of pitting, particularly in the presence of chloride ions; prelab questions and answers for students; notes to instructors and a discussion of the results. This material is available via the Internet at http://pubs.acs.org.

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS We thank Daniel Mandler for the idea that grew into this experiment, Eitan Manor for constructive discussions, and Michael Mizrahi for his help in choosing and preparing the iron samples. We extend our heartfelt thanks to our classes of 2010 and 2011 material sciences students and their instructors for their cooperation and participation in developing this experiment.

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REFERENCES

(1) Pourbaix, M. Lectures on Electrochemical Corrosion; Plenum Press: New York, 1973, p 361. (2) Vernon, W. H. J. Trans. Soc. 1935, 31, 1668. (3) Graedel, T. E.; Frankenthal, R. P. J. Electrochem. Soc. 1990, 137, 2385. (4) Refait, P.; Genin, J. M. R. Corros. Sci. 1997, 39, 539. (5) Sagoecrentsil, K. K.; Glasser, F. P. Corrosion 1993, 49, 457. (6) Stratmann, M. Ber. Bunsen−Ges. Phys. Chem. 1990, 94, 626. (7) Forsberg, J.; Hedberg, J.; Leygraf, C.; Nordgren, J.; Duda, L. C. J. Electrochem. Soc. 2010, 157, C110. (8) Wang, Z. S.; Xu, C. C.; Dong, X. Q. Chin. J. Chem. Eng. 2008, 16, 299. (9) Hara, S.; Miura, M.; Uchiumi, Y.; Fujiwara, T.; Yamamoto, M. Corros. Sci. 2005, 47, 2419. (10) Elkot, A. M.; Elhaleem, S. A.; Mohammed, S. Monatsh. Chem. 1992, 123, 965. (11) Misawa, T.; Hashimot., K.; Shimodai., S. Corros. Sci. 1974, 14, 131. (12) Refait, P.; Drissi, S. H.; Pytkiewicz, J.; Genin, J. M. R. Corros. Sci. 1997, 39, 1699. (13) Refait, P. H.; Abdelmoula, M.; Genin, J. M. R. Corros. Sci. 1998, 40, 1547. (14) Fontana, M. G. Corrosion Engineering; 3rd ed.; McGraw-Hill Book Co.: Singapore, 1987, Chapter 9, pp 435−449. (15) Sorensen, P. A.; Kiil, S.; Dam-Johansen, K.; Weinell, C. E. J. Coat. Technol. Res. 2009, 6, 135. (16) Baldwin, K. R.; Robinson, M. J.; Smith, C. J. E. Corros. Sci. 1993, 35, 1267. (17) Wilcox, G. D.; Gabe, D. R. Corros. Sci. 1993, 35, 1251. (18) Kraljic, M.; Mandic, Z.; Duic, L. Corros. Sci. 2003, 45, 181. (19) Solorza, O.; Olivares, L.; Ibanez, J. G. J. Chem. Educ. 1991, 68, 175. (20) Arce, E. M.; Ramirez, R.; Cortes, F.; Ibanez, J. G. J. Chem. Educ. 1991, 68, 351. (21) Onuchukwu, A. I. J. Chem. Educ. 1989, 66, 681. (22) Higa, C. O.; Dacosta, S.; Agostinho, S. M. L. J. Chem. Educ. 1989, 66, 441. (23) Meloan, C. E. J. Chem. Educ. 1986, 63, 456.

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SUMMARY This experiment addresses the important subject of iron corrosion and allows the students to both directly measure the effect under different conditions and to learn about the analytical sensitivity of atomic absorption spectroscopy. The results were very reproducible and correlated the theory well. The experiments are easily amenable to checking other conditions. 493




(24) Pinnell, R. P.; Zanella, A. W. J. Chem. Educ. 1981, 58, 444. (25) Wever, D. J.; Veldhuizen, A. G.; de Vries, J.; Busscher, H. J.; Uges, D. R. A.; van Horn, J. R. Biomaterials 1998, 19, 761. (26) Bumgardner, J. D.; Johansson, B. I. J. Biomed. Mater. Res. 1998, 43, 184. (27) Kirchhoff, G.; Bunsen, R. Ann. Phys. 1860, 186, 161. (28) Walsh, A. Spectrochim. Acta 1955, 7, 108. (29) Barteneva, O. I.; Bartenev, V. V. Prot. Met. Phys. Chem. Surf. 2010, 46, 129. (30) Ivanenko, N. B.; Ganeev, A. A.; Solovyev, N. D.; Moskvin, L. N. J. Anal. Chem. 2011, 66, 784. (31) Grindlay, G.; Mora, J.; Gras, L.; de Loos-Vollebregt, M. T. C. Anal. Chim. Acta 2011, 691, 18. (32) Monkhouse, P. Prog. Energy Combust. Sci. 2011, 37, 125. (33) Korzhova, E. N.; Kuznetsova, O. V.; Smagunova, A. N.; Stavitskaya, M. V. J. Anal. Chem. 2011, 66, 222. (34) McNallan, M. Mater. Perform. 1994, 33, 54. (35) Skorchelletti, V. V. Theory of Metal Corrosion; Keter Publishing House Ltd.: Jerusalem, 1973, Chapter 6, pp 201−205. (36) Chang, Y. N.; Wei, F. I. J. Mater. Sci. 1991, 26, 3693.

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Determining the Effect of Environmental Conditions on Iron

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