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Ind. Eng. Chem. Res. 2002, 41, 914-921
Corrosion at Metal InterfacessA Study of Corrosion Rate and Solution Properties, Including Electrical Conductance, Viscosity, and Density Ashvin K. Dewan Hightower High School, 3333 Hurricane Lane, Missouri City, Texas 77459
Diego P. Valenzuela, Shiela T. Dubey, and Ashok K. Dewan* Equilon Enterprises LLC, Westhollow Technology Center, Houston, Texas 77251
In this paper, we attempt to correlate the rate of corrosion at the metal surface with bulk solution properties such as electrical conductance, viscosity, and density. The premise is that the process solution environment within the pipe or refinery equipment controls the severity of the corrosion at the metal surface. Furthermore, the migration of a protonated amine to the metal surface (instead of a proton in acid solutions) is the rate-determining step. For this study, the corrosion rate and solution properties of a single neutralizer salt (DEA‚HCl) plus ammonium chloride, were examined more closely at fixed temperature, pressure, and composition. The metallurgy was restricted to carbon steel. Metal coupons of the same geometry and size were used to collect corrosion rate data in a nonflowing static environment. Under such conditions, the experimental data support the hypothesis that the corrosion rate at the metal surface is strongly influenced by the diffusion of the protonated amine. Diffusion, like all transport properties (i.e., electrical conductivity, viscosity, density, and diffusion coefficients), is a function of temperature, pressure, and solution composition (Dillon, C. P. Materials Selector for Hazardous Chemicals; Materials Technology Institute: St. Louis, MO, 1997; Vol. 1, p 69). The diffusion coefficient plus the Reynolds and Schmidt numbers determine mass-transfer-limited electrochemical processes (Eisenberg, M.; Tobias, C. W.; Wilke, C. R. J. Electrochem. Soc. 1954, 101, 306). The practical significance of this study lies in its ability to define, select, and/or screen neutralizer candidates via analytical measurements of their solution properties. Any mitigation in the corrosion rate resulting from a judicious choice of the neutralizer amine can have a significant economic impact. Introduction Corrosion at metal surfaces is a severe industrial problem that costs the U.S. industry alone in excess of $170 billion/year.1 Minimizing this corrosion can save substantial money and prevent accidents due to equipment and/or pipe failure. Proven methods for corrosion control include the selection of appropriate metallurgy, the use of filming inhibitors, and the addition of neutralizing amines.2 In crude unit overheads, the addition of neutralizing amines is a very common industrial practice. The primary corrosive agent in such instances is hydrogen chloride, which makes its way into the overhead system from the crude tower. In the presence of water, this causes acid corrosion at metal surfaces. Above the aqueous dew point, an undesirable side reaction of the neutralizing amine to amine hydrochloride is probable. The amine hydrochloride salt can be very corrosive. The presence of alkaline earth (magnesium, calcium, etc.) chloride impurities in the crude oil supplied to refineries introduces hydrochloric acid into the tops of crude columns. In a typical crude column, because of the presence of stripping steam in the column, or water wash in an upstream desalter unit, or water in the supplied crude oil, hydrolysis of these † Paper 83e presented at the American Institute of Chemical Engineers National Meeting, Nov 14, 2000, Los Angeles, CA * Author to whom correspondence should be addressed (
[email protected]).
alkaline earth chlorides occurs and releases hydrochloric acid.5 For example
MgCl2 + 2H2O ) 2HCl + Mg(OH)2 Whereas magnesium and calcium chloride salts tend to hydrolyze in the feed heater and bottom section of the crude column, the sodium chloride salt resists hydrolysis and is stable. The HCl vapors are volatile and exit the crude column along with the column overhead vapor.6 If untreated, the HCl vapor will condense in the vicinity of the aqueous dew point in the crude column overhead system, causing active acid corrosion by the overall reaction
Fe + 2HCl ) FeCl2 + H2 Atomic hydrogen can form blisters in the metal wall, and iron can dissolve in the free water phase produced in the crude column overhead system.1 In the presence of organic amines, the overall reaction for corrosion is altered because of the availability of the protonated amine (instead of the proton in acid corrosion) to the following
Fe + 2Am‚HCl ) FeCl2 + 2Am + H2 Thus, the diffusion of the protonated amine (in the amine hydrochloride) can be expected to strongly influ-
10.1021/ie001027u CCC: $22.00 © 2002 American Chemical Society Published on Web 02/27/2002
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ence the rate of corrosion at the metal surface.7 Highly conductive, nonviscous solutions are more likely to promote rapid corrosion than poorly conducting, highly viscous solutions. At the atomic level, corrosion occurs when the protonated amine accepts electrons from the surface of the metal. The elemental metal is oxidized (iron goes to ferrous ion after donating its electrons to the protonated amine) and dissolves in the solution. The hydrogen ions receiving the electrons combine to liberate hydrogen gas. The number of protonated amine molecules that can access the metal surface depends on (1) the conductivity of the solution (i.e., availability of ionic reactants), (2) the viscosity of the solution media (highly viscous solutions will impede the motion of the protonated amine toward the metal surface), (3) the density of the solution (how close the molecules are to each other), (4) the solution temperature (higher temperatures lead to faster diffusion) (5) the metal (its ability to donate electrons and the ease with which metal ions dissolve into the solution), (6) the amine salt structure (effectiveness in adhering to the metal surface and receiving electrons from the metal), (7) the accumulation of corrosion products on the metal surface (insoluble products inhibit the diffusion of protonated molecules from the solution to the metal surface and of metal ions into the solution), and (8) the velocity of the solution media (velocity disturbs the solution film on the metal surface). In a static cell, the corrosion rate will depend on the physical properties of the solution and the geometry of the metal surface (available metal surface area).4 The corrosion rate (CR) can be defined as follows
CR ) KTaFbµcΛdGe where K is a constant, T is the temperature, F is the density, µ is the viscosity, Λ is the molar conductance, G represents the geometry. At a fixed temperature and constant metal coupon geometry, the expression simplifies to
CR ) KFxµyΛz Thus, measuring the physical properties of the solution (F, µ, Λ) should be sufficient for predicting the corrosion rate in a static system (see Appendix A for more details of the model heuristics). Solution Properties Electrical Conductivity. Electrical conductivity is a measure of the mobility of an ion in solution. Positive ions, or cations, move toward the cathode, and negative ions, or anions, move toward the anode. The conductivity of a solution depends on the concentration and mobility of all ions present in solution. The ion mobility, in turn, depends on the ion size and charge, the dielectric constant of the solvent, and the solution temperature.8 The determination of electrical conductivity (EC) consists basically of measuring the AC resistance of a column of solution. With alternating current (AC), any concentration changes due to oxidation/reduction reactions at the electrodes are minimized, and polarization effects reduced.9 Ohm’s law and the Wheatstone bridge are used to measure the resistance (R).10,11 Electrical conductivity (EC) in Ω-1 (or siemens, S) is defined as the reciprocal of the resistance in ohms (Ω) measured
between opposite faces of a cube 1 cm on each edge. The resistance of a conductivity cell of constant cross section immersed in a solution is proportional to the distance (L) between electrodes, inversely proportional to the cross-sectional area (A) of the electrolyte, and inversely proportional to the specific electrical conductance of the solution.12 Thus
SC ) (L/A)(R-1) ) K × EC where K is the cell constant in cm-1, SC is the specific conductance in S/cm, R is the resistance in Ωx, EC is the electrical conductivity in siemens, L is the length in cm, and A is the cross-sectional area in cm2. The cell constant is determined from measurements using a calibration fluid of known specific conductance. Viscosity. Viscosity is a measure of the internal friction of a fluid. This friction becomes apparent when a layer of fluid is made to move in relation to another layer. The greater the friction, the greater the amount of force required to cause this movement, which is called shear.5 Shearing occurs whenever the fluid is physically moved or distributed, as in pouring, spreading, spraying, mixing, etc. Viscosity can be Newtonian or non-Newtonian. A Newtonian fluid is one in which the ratio of the shear stress to the shear rate is constant. The ratio is variable for a non-Newtonian fluid.5 Density. The density of a solution is a measure of the mass per unit volume, or the heaviness (or concentration) of the molecules that make up a unit volume of solution. Typically, the density is measured in grams per cubic centimeter. The greater the density, the more concentrated the solution. For this study, the density of the solutions used varies from 1 to 1.5 g/cm3. Experimental Procedures Sparging. All water used for this study was made oxygen-free by sparging. To do this, a flask containing DI (deionized) water was boiled on a hot plate under a nitrogen blanket. The water was allowed to boil for 1 h and was then allowed to cool under a nitrogen blanket. An oxygen kit was used to test for oxygen in solution. The water was deemed acceptable only if less than 1 ppm of oxygen was detected in solution. Solution Preparations. All solution preparations were performed in a drybox under a nitrogen blanket. This step was taken to minimize any contamination due to oxygen. The same philosophy was used in preparing the desired mixture solutions and in transferring smaller samples for individual tests of the solution properties. The original mixture solutions were stored in the drybox until measurements for that solution mixture were completed. Coupon Preparation. For this study, the selection of the coupon and coupon preparation for the corrosion experimentssmetal loss, iron in solution, and pH change during corrosionswere deemed very important for reproducibility and experimental consistency. All coupons were carbon steel, obtained from the same source, with identical geometries and similar dimensions. The coupons were rinsed in an acetone bath, followed by a toluene dip, to clean the surface chemically. They were then scrubbed with pumice soap, sponged, and stored in acetone. An “erasing” machine was used to wipe off the metal surfaces. To avoid any oxidation of the cleaned surfaces, the coupons were transported in containers filled with acetone. The coupons were dried with nitro-
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gen and weighed immediately to minimize exposure to air/oxygen. Corrosion Experiments. The pH readings (before and after the corrosion experiments) were taken on a pH meter that was calibrated with standard buffer solutions before each measurement. The calibration step before the actual reading avoided instrument drift and ensured that the pH probe head was clean for measurement. The ambient-temperature (25 °C) corrosion tests were conducted in the drybox by allowing the coupons to sit in test tubes containing the mixture solution. The test duration was standardized to be 9 days for all tests. The 50 °C tests were conducted in specially designed automated oven boxes with the ability to maintain the set temperature over the extended period of testing. For safety reasons (for example, explosion of test tube), the samples were placed in capped aluminum cylinders before the ovens were heated. Upon completion of the tests, the caps were unscrewed very slowly to allow any gases to escape in the event that a test tube had accidentally broken on heating. The difference in weight of an “erased” coupon before and after the experiment was used as a measure of metal loss. To determine the iron concentration in solution, a spectrophotometer calibrated for the expected range of iron in solution was used. An indicator reagent powder from Hach was used to dope one cell while the other cell contained a blank (i.e., without reagent). Hach method 265 for total Fe and a wavelength of 510 nm were typically used. Conductivity Measurements. A Metrohm 712 conductivity meter was used to measure the electrical conductance of various mixture solutions at different temperatures. The known mixture solution was placed in a cylindrical cell, which was then immersed in a constant-temperature water bath. A flow of nitrogen was maintained above the cell to minimize oxygen/air contamination. A Pt(1000) sensor was used to record the temperature of the water bath as close to the cell as feasible. A conductivity probe immersed in the cell recorded the electrical conductance. Readings were taken at 25, 30, 40, 50, and 60 °C, after the water bath had been allowed to equilibrate for a minimum of 30-45 min at each temperature setting. Viscosity Measurements. A Brookfield DV-II+ programmable viscometer with a UL adaptor and a small cell adaptor was used to measure mixture solution viscosities at various temperatures/compositions. The higher viscosity readings were taken on a standard cell. Expected viscosities of less than 15 cP were measured in a small cell adaptor. Readings of less than 4 cP required special handling and calibration with standard solutions. For our readings, we chose to maximize the torque in the cell. A thermostat was used to control the temperature. Data were recorded at 25, 30, 40, 50, and 60 °C. Density Measurements. Density measurements were made from the hottest (70 °C) to the coldest temperature (ambient 25 °C) in 10 °C intervals. The instrument used for density measurements was an Anton Paar DMA density meter. Data were recorded at 25, 30, 40, 50, 60, and 70 °C. The instrument was calibrated with known fluids. Results and Observations This section describes the solution properties of DEA‚ HCl and NH4Cl. There are two basic differences in the
Figure 1. Corrosion mechanism in pure water.
Figure 2. Corrosion mechanism in the presence of amine ‚HCl.
Figure 3. Densities of DEA‚HCl solutions at fixed temperature.
properties of these two salts. First, diethanolamine hydrochloride salt is completely soluble in water throughout the temperature range covered in this work, 25-70 °C. Therefore, it was possible to obtain properties of this salt up to the pure salt point. On the other hand, the ammonium chloride salt is relatively less soluble in water, about 11 mol %.13 In the latter case, this limited the maximum solution concentration that could be prepared for the experiments. Second, the melting point of the DEA‚HCl salt is very low (we were unable to crystallize this salt even at temperatures as low as -10 °C).14 On the other hand, the melting point of NH4Cl is about 521 °C.15 Therefore, under the experimental conditions explored in this work, DEA‚HCl was a molten salt, whereas the NH4Cl salt was a solid. Figures 1 and 2 illustrate the corrosion mechanisms in pure water and in amine‚HCl solution. Figures 3 and 4 (also Tables 1 and 2) depict the density values of aqueous solutions of DEA‚HCl and NH4Cl, respectively. The temperature ranged from 25 to 70 °C. In both cases, the density was strongly dependent on the salt concentration and a weak function of the temperature. Figures 5 and 6 (also Tables 3 and 4) show the measured viscosities for these two salts. As expected, for both systems, the viscosity was found to be a strong function of temperature and composition. Again, the aqueous solubility determined the experimentally ac-
Ind. Eng. Chem. Res., Vol. 41, No. 5, 2002 917
Figure 4. Densities of NH4Cl solutions at fixed temperature.
Figure 7. Conductances of DEA‚HCl solutions at fixed temperature. Table 3. Viscosities of Various DEA‚HCl Solutions at Different Temperatures
Figure 5. Viscosities of DEA‚HCl solutions at fixed temperature.
viscosity (cP) 40 °C 50 °C
[DEAHCI] (mol %)
25 °C
30 °C
0.00 2.00 5.05 10.00 25.01 49.94 74.97 100.00
0.91 1.03 2.81 3.48 15.55 116.19 519.70 2220.03
0.82 0.91 2.67 3.28 12.94 87.62 372.92 1577.94
0.67 0.71 2.43 3.03 9.41 54.31 204.76 819.88
0.56 0.57 2.25 2.90 7.26 37.35 121.53 441.03
60 °C 0.47 0.46 2.11 2.87 5.90 28.17 77.34 244.89
Table 4. Viscosities of Various NH4Cl Solutions at Different Temperatures [NH4Cl (mol %)
25 °C
30 °C
0.00 0.42 4.57 7.50 9.90
0.91 0.92 0.92 0.92 0.96
0.82 0.79 0.83 0.81 0.85
viscosity (cP) 40 °C 50 °C 0.67 0.62 0.68 0.65 0.69
0.56 0.51 0.57 0.53 0.58
60 °C 0.47 0.44 0.49 0.45 0.50
Table 5. Electrical Conductances of Various DEA‚HCl Solutions at Different Temperatures Figure 6. Viscosities of NH4Cl solutions at fixed temperature. Table 1. Densities of Various DEA‚HCl Solutions at Different Temperatures [DEAHCI] (mol %)
25 °C
0.00 2.00 5.05 10.00 25.01 49.94 74.97 100.00
0.9969 1.0305 1.0696 1.1146 1.1759 1.2152 1.2313 1.2462
30 °C
density (g/cm3) 40 °C 50 °C
60 °C
70 °C
0.9955 1.0288 1.0675 1.1123 1.1732 1.2125 1.2287 1.2435
0.9924 1.0252 1.0635 1.1078 1.1686 1.2075 1.2238 1.2386
0.9839 1.0164 1.0533 1.0983 1.1588 1.1978 1.2140 1.2289
0.9775 1.0100 1.0480 1.0919 1.1525 1.1914 1.2076
0.9885 1.0211 1.0590 1.1032 1.1637 1.2027 1.2189 1.2336
Table 2. Densities of Various NH4Cl Solutions at Different Temperatures [NH4Cl] (mol %)
25 °C
0.00 0.42 4.57 7.50 9.90
0.9969 1.0008 1.0337 1.0532 1.0666
30 °C
density (g/cm3) 40 °C 50 °C
60 °C
70 °C
0.9955 0.9994 1.0319 1.0514 1.0644
0.9924 0.9962 1.0285 1.0479 1.0602
0.9839 0.9878 1.0203 1.0398 1.0510
0.9775 0.9813 1.0146 1.0340 1.0451
0.9885 0.9924 1.0246 1.0440 1.0559
cessible range of concentrations for solutions of ammonium chloride. Probably the most interesting property measured was the conductance of the aqueous solutions of these two salts. Figure 7 (and Table 5) shows that the conductance
[DEAHC] (mol %)
∼25 °C
2.00 5.05 10.00 25.01 49.94 74.97 100.00
52.07 72.65 67.89 30.10 6.38 1.67 0.40
electrical conductance (mS/cm) ∼30 °C ∼40 °C ∼50 °C 57.45 80.55 75.37 34.62 8.16 2.29 0.57
69.04 96.42 91.84 45.52 12.56 2.97 1.13
74.80 112.60 109.20 56.87 18.14 6.05 2.00
∼60 °C 83.60 126.50 126.30 68.58 24.51 9.15 3.36
for DEA‚HCl initially increases with concentration, passes through a maximum, and then decreases to a very low value corresponding to the conductance of the pure molten salt. The fact that this molten salt has a finite conductance implies the presence of ionic groups. Figure 8 (and Table 6) shows the electrical conductance of ammonium chloride in solution up to its maximum solubility in water. These data suggest that the maximum conductance occurs at concentrations close to the solubility limit at that temperature. Figures 9 and 10 (also Tables 7 and 8) show the molar conductances of DEA‚HCl and NH4Cl, respectively. In both cases, the molar conductance decreases with concentration and increases with temperature. This appears to be the opposite of the behavior for viscosity, which suggests a relationship between these two properties. Figures 11 and 12 (also Tables 9 and 10) show the weight loss of coupons immersed in solutions of DEA‚ HCl and NH4Cl, respectively. Weight loss in solutions
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Figure 8. Conductances of NH4Cl solutions at fixed temperature.
Figure 11. Corrosion rates of DEA‚HCl solutions at 25 and 50 °C.
Figure 9. Molar conductances of DEA‚HCl solutions at fixed temperatures.
Figure 12. Corrosion rates of NH4Cl solutions at 25 and 50 °C. Table 7. Molar Conductances of Various DEA‚HCl Solutions at Different Temperatures
Figure 10. Molar conductances of NH4Cl solutions at fixed temperatures. Table 6. Electrical Conductances of Various NH4Cl Solutions at Different Temperatures [NH4Cl] (mol %)
∼25 °C
0.42 4.57 7.50 9.90
24.52 217.20 309.80 366.60
electrical conductance (mS/cm) ∼30 °C ∼40 °C ∼50 °C 26.98 230.70 334.30 393.60
32.61 264.80 380.50 447.10
38.69 294.60 426.90 495.60
∼60 °C 45.95 326.10 470.80 549.20
of DEA‚HCl goes through a maximum located between 5 and 15 mol %, and then the corrosion rate decays to a minimum at the molten-salt limit. Corrosion in ammonium chloride solutions behaves in a more unpredictable way. It appears that the proximity to the solubility limit might be affecting corrosion in this case. Notice that a similar erratic behavior was found for the viscosity of ammonium chloride in the same temperature and concentration range. Figures 13 and 14 (also Tables 9 and 10) show the increase of pH with corrosion as a result of consumption of protons and release of unprotonated amines. Figure 15 is an attempt to show the strong relationship between the solution properties and the corrosion rate. Our proposed model (see Appendix A) was compared to the weight loss observed in the DEA‚HCl solutions. There is a remarkable similarity in the shapes
[DEACHl] (mol %)
25 °C
2.00 5.05 10.00 25.01 49.94 74.97 100.00
52.20 32.56 18.44 4.98 0.84 0.21 0.05
molar conductance (mS cm2 mol-1) 30 °C 40 °C 50 °C 60 °C 57.74 36.63 20.69 5.82 1.08 0.24 0.06
67.81 43.75 25.30 7.61 1.67 0.40 0.13
76.51 60.90 30.07 9.54 2.38 0.70 0.24
83.85 57.79 35.00 11.61 3.24 1.12 0.39
Table 8. Molar Conductances of Various NH4Cl Solutions at Different Temperatures [NH4Cl] (mol %) 0.42 4.57 7.50 9.90
molar conductance (mS cm2 mol-1) 25 °C 30 °C 40 °C 50 °C 60 °C 106.23 89.92 81.38 74.87
116.73 96.79 87.68 80.42
140.74 110.47 100.22 91.49
168.77 124.06 112.64 102.52
200.81 137.55 124.97 113.52
of the curves and a close match of the plots at two different temperatures. This clearly illustrates that the corrosion rate in a static cell at fixed temperature, for known geometry and known metallurgy of the coupon, is governed by the properties of the solution. Figure 16 presents a comparison of the weight loss observed in various amine hydrochloride solutions of equimolar (10%) strength under identical conditions at 25 and 50 °C. The objective of the comparison was to determine whether screening of corrosion inhibitors was feasible and in agreement with field observations. Laboratory measurements for competitive amines are still incomplete. Future work will focus on the screening of alternate inhibitors on the basis of their solution properties at fixed temperatures. Discussion The data presented here indicate that corrosion from amine chloride solutions is strongly controlled by bulk
Ind. Eng. Chem. Res., Vol. 41, No. 5, 2002 919 Table 9. Corrosion Rates of Various DEA‚HCl Solutions at 25 and 50 °C Using Metal Loss, Iron in Solution, and pH Change wt of coupon after scrubbing (g)
spectrowt photometer loss wt loss pH (mg) (mg) before
[DEAHCl] (mol %)
wt of coupon (g)
0.00 2.00 5.05 10.00 25.01 49.94 74.97 100.00
13.49496 13.92257 13.60666 13.65858 13.83922 14.26055 13.77998 13.97659
25 °C 13.419248 2.48 13.906398 15.59 13.58976 16.90 13.64552 13.06 13.83187 7.35 14.25619 4.36 13.77659 3.39 13.97600 0.59
0.13 11.78 13.72 11.15 5.25 2.09 1.26 0.91
7.064 5.048 5.016 4.627 3.881 3.473 3.3693 7.124
7.182 7.262 7.164 6.807 5.993 5.320 4.803 7.233
0.00 2.00 5.05 10.00 25.01 49.94 74.97 100.00
13.98107 14.02360 13.65338 13.72125 13.74405 14.01705 14.01010 13.95297
50 °C 13.97773 3.34 14.00428 19.32 13.62912 24.26 13.70809 13.16 13.71975 24.30 13.99668 20.37 13.99571 14.39 13.94808 4.89
0.15 13.50 19.19 10.05 18.68 16.76 10.53 1.44
7.064 5.048 5.016 4.627 3.881 3.473 3.363 7.124
8.944 7.478 7.402 6.878 6.446 6.266 6.151 7.274
pH after
Figure 13. pH’s of DEA‚HCl solutions at 25 and 50 °C.
Table 10. Corrosion Rates of Various NH4Cl Solutions at 25 and 50 °C Using Metal Loss, Iron in Solution, and pH Change wt of coupon after scrubbing (g)
Figure 14. pH’s of NH4Cl solutions at 25 and 50 °C. spectrowt photometer loss wt loss pH pH (mg) (mg) before after
[NH4Cl] (mol %)
wt of coupon (g)
0.00 0.42 4.57 4.57 7.50 9.90 9.90
13.49496 13.71914 13.60088 14.09232 13.92790 13.61731 13.84464
25 °C 13.49248 2.48 13.71048 8.66 13.57900 21.88 14.07142 20.90 13.92126 6.64 13.61039* 6.92 18.83564 9.00
0.13 4.74 19.55 17.83 3.87 5.39 7.39
7.064 6.316 5.624 5.624 5.470 5.220 5.220
9.182 7.827 7.420 7.318 6.780 6.594 6.779
0.00 0.42 4.57 7.50 9.90
13.98107 13.89985 13.97697 13.81816 13.93662
13.97773 13.89180 13.96859 13.80144 13.92176
50 °C 3.34 8.05 8.38 16.72 14.86
0.15 3.73 4.75 13.27 11.22
7.064 6.316 5.624 5.470 5.220
8.944 8.024 7.140 7.284 7.074
transport properties. The second observation is that corrosion in the amine hydrochloride environment is, for the most part, independent of the concentration of the free protons in the bulk solution. Acid corrosion involves the following steps: (1) diffusion of the proton from the bulk to the metal surface, (2) adsorption of the proton onto the metal surface, and (3) redox reaction to produce elemental/molecular hydrogen and ferrous ions. Amine salt corrosion includes the following steps: (1) diffusion of the protonated amine to the metal surface, (2) adsorption of the protonated amine onto the metal surface, (3) dissociation of the protonated amine to produce a proton and an unprotonated amine, and (4) redox reaction to produce elemental/molecular hydrogen and ferrous ions. A necessary step of the corrosion process is migration of the protonated amine to the metal surface. Because our experiments were performed in a static cell (no shear stress) without any external electrostatic gradient, diffusional migration appears to have been controlling metal corrosion. The driving force was generated by a decrease, with respect to bulk phase, of the concentration of protonated amines at the metal surface
Figure 15. Solution model vs measured corrosion rates for DEA‚ HCl solutions at 25 and 50 °C.
Figure 16. Corrosion rates of 10 mol % amine‚HCl solutions at 25 and 50 °C.
as a result of corrosion. Diffusion coefficients were not measured in this work. However, for the completely dissociated electrolyte solution, transport properties such as diffusion, conductance, and viscosity are strongly coupled. This is because, in electrolyte solutions, resistance to external forces, such as shear stress, chemical potential, and electrostatic potentials, are reactions to similar stresses of ionic structures. The functional relationships among these transport properties are very complex, especially at the high concentrations used in this work, and beyond the intended experimental character of this work. However, at least in theory, there
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should be a function relating diffusion coefficients to viscosity and conductance, and if corrosion is determined by diffusion, there should also be a function relating corrosion to viscosity and conductance. The evidence presented in this work indicates that such a function exists. We hope that the data presented in this work will aid in the development/validation of a more rigorous relationship. Conclusions and Future Work The following conclusions can be drawn from this study: (1) There is a strong relationship (0.9 e correlation coefficient e 1.0) between metal loss and solution properties for corrosion occurring in DEA‚HCl and NH4Cl solutions. The proposed corrosion rate model based on solution properties was tested to demonstrate the existence of such a relationship, at 25 and 50 °C, for any aqueous mixture solution of DEA‚HCl. The model successfully mirrored the true corrosion rate, defined in terms of metal loss. (2) The corrosion rate increased with increasing temperature, with the 50 °C data showing more rapid metal loss than the ambient-temperature (25 °C) data. There is a distinct maximum in the corrosion rate versus composition data for DEA‚HCl at both temperatures. A maximum was also observed for the NH4Cl corrosion data at 25 °C, which dampened out at 50 °C. This seems to suggest that, at elevated temperatures, the maximum in the corrosion rate might shift so far to the right on the composition scale that it would appear that dilute solutions exhibit a monotonic increase in corrosion rate with increasing composition. This observation might also explain some of the observed confusion in the open literature on this subject. Systematic experimentation (such as this study) is one way to elucidate this behavior. (3) Viscosity decreased with temperature but increased with concentration for all systems. The change in viscosity was more pronounced in the more viscous systems. Very dilute solutions required viscosity corrections. Calibration fluids with known viscosities were used to correct the data. (4) Molar conductance increased with temperature for all systems. Dilute solutions were capable of sustaining more current than concentrated solutions. For DEA‚ HCl, the electrical conductance curve mirrored the corrosion rate curve. (5) Density decreased for the solutions as temperature increased. The more viscous solutions were denser. Density also increased with concentration. (6) The NH4Cl solution had a saturation limit near 11 mol % at ambient temperature and was restricted to a narrow composition region. Nonetheless, it was interesting to observe a maximum in the corrosion rate data under ambient conditions. This maximum seemed to migrate toward lower concentration and rapidly become monotonic with rising temperature. The DEA‚ HCl salt (like most amine hydrochlorides) was completely soluble in water for all compositions; other amine hydrochlorides are soluble up to 90 mol %. (7) The electrical conductivity (in mS/cm) showed a maximum in the dilute region of DEA‚HCl solutions. However, a transformation of the data to molar conductance (in mS cm2 mol-1) simplified the graphs, producing a hyperbola-like curve. Additional solution
products examined (e.g., conductance × viscosity) appeared to smooth the data further. (8) The solution properties can be used to screen amine hydrochloride neutralizers. This aspect of the study is in progress and is showing very promising results. The corrosion rate of MEA‚HCl, a member of the same family of hydrochloride salt as DEA‚HCl, was predicted very well with the existing DEA‚HCl corrosion rate model. Other amine hydrochloride salts belonging to similar chemical families are still being examined. Work to date involves comparisons of various amine hydrochloride equimolar salts with each other and rank orders them by corrosivity and solution properties. Ultimately, model predictions of corrosion rate (i.e., metal loss) for these other amine hydrochloride salts will be compared to laboratory data to validate the corrosion rate model. (9) In this study, the 100 mol % DEA hydrochloride salt was very mildly corrosive. This finding was also borne out by the electrical conductivity measurements. (10) All experiments (except the hydrochloric acid corrosion tests) were conducted in an oxygen-free drybox. The water used for solution was deionized and sparged of O2 for 1-2 h to remove any dissolved oxygen. The pH probe was repeatedly recalibrated to maintain its reliability. Likewise, the conductivity meter and viscometer were tested versus known calibration fluids. (11) The pH measurements (before and after experiments) are a good guide to expected corrosion behavior. OLI Engine software can be used to solve the complex system of equations on a computer and thereby determine the resulting metal loss using the pH change and fluid properties. Some of these ideas require testing and future work (outside the scope of this study). (12) The current study focused on monoamine salts, with only one active site for HCl addition to the amine. Diamine salts (e.g., EDA, ethylene diamine) belong to another family of amines and are expected to behave differently because of the presence of two active sites in the molecule for HCl addition. Consequently, the corrosion rate due to such diamines will be affected by the interplay between the two active sites and the metal surface. This phenomenon might require an alternative prediction model. Appendix A Model Heuristics.
CR ) f[D] The corrosion rate (CR) in a static cell with no velocity factor, constant coupon geometry, and constant metallurgy is hypothesized to be dictated by diffusion of the protonated amine from bulk solution to metal surface.
Do )
ΛoRT ziF2
According to the Nernst equation, diffusivity at infinite dilution (Do) is dependent on the conductance at infinite dilution (Λo), the temperature (T), and the charge of the ion (zi), where F is the Faraday constant and R is the gas constant.
[
D ) Do 1 + m2
(
)]
∂ ln γ( µH2O FH2O ∂m2 µsoln Fsoln
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When a solute is introduced, the Nernst equation must be corrected for composition. By introducing the term [1 + m2(∂ ln γ(/∂m2)] into the equation, the solute concentration can be taken into account. m2 represents the molality of the amine hydrochloride, and γ( is the activity coefficient of the amine hydrochloride. µsoln is the viscosity of the solution, and Fsoln represents the density of the solution.
D ) f[µsoln, Fsoln, Λsoln] Diffusivity, therefore, for any solution composition, is dependent on the solution properties.
CR ) f[D] ∼ k1ΛsolnaµsolnbFsolnc Therefore, the corrosion rate (CR) can be defined as a function of the molar conductance (Λsoln), viscosity (µsoln), and density (Fsoln), according to the equation shown, with regression constants a, b, and c. Literature Cited (1) Korb, L. J.; Olson, D. L. Metal Handbook, 9th ed.; ASM International: Materials Park, Ohio, 1987, pp 17-26, 23-44, 229, 231-233, 242, 283-303, 485-523, 1134-1145. (2) Process Industries Corrosion; National Association of Corrosion Engineers: Houston, TX, 1975; pp 6-15, 92-100. (3) Eisenberg, M.; Tobias, C. W.; Wilke, C. R. J. Electrochem. Soc. 1954, 101, 306. (4) Dillon, C. P. Materials Selector for Hazardous Chemicals; Materials Technology Institute: St. Louis, MO, 1997; Vol. 1, p 69.
(5) Coulson, J. M.; Richardson, J. F. Chemical Engineering; Pergamon Press: Elmsford, NY, 1965; Vol. 1, p 44. (6) Hefter, G. T.; North, N. A.; Tan, S. H. Organic Corrosion Inhibitors in Neutral Solutions. Corrosion 1997, 53 (8), 657. (7) Dubey, S. Structure and Solvation in the Ionization of Selected Amines in Water; University of Calgary: Calgary, Alberta, Canada, 1975; pp 96-134. (8) Balakrishnan, P. V. Liquid-Vapor Distribution of Amines and Acid Ionization Constants of Their Ammonium Salts in Aqueous Systems at High Temperatures. J. Solution Chem. 1988, 17 (9), 825. (9) Glasstone, S.; Lewis, D. Elements of Physical Chemisty; Van Nostrand: Princeton, NJ, 1960; p 25 (10) Fogiel, M. The Essentials of Physical Chemistry I; Research and Education Association: Piscataway, NJ, 1987; p 86. (11) van der Linde, W.; Northcott, D.; Redmond, W.; Robertson, R. Basic Dissociation Constant for Ehtylamine by a Convenient Conductance Methodology. Can. J. Chem. 1969, 47 (2), 279. (12) Kittsley, S. L. Physical Chemistry; College Outline Series; Barnes & Noble, Inc.: New York, 1955; p 147. (13) Cohen-Adad, R., Lorimer, J. W., Eds. Alkali Metal and Ammonium Chlorides in Water and Heavy Water (Binary Systems); Solubility Data Series; Pergamon Press: New York, 1991; pp 413493. (14) Daubert, T. E.; Pamer, R. P. Physical and Thermodynamic Properties of Pure Chemicals; Design Institute for Physical Property Data, AIChE: University Park, PA, 1994. (15) Kraus, C. A. Electrolytes: From Dilute Solutions to Fused Salts. J. Phys. Chem. 1954, 58 (9), 673.
Received for review December 1, 2000 Revised manuscript received December 5, 2001 Accepted December 19, 2001 IE001027U
