Langmuir 2001, 17, 5711-5713
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Ammonia Adsorption at the Aqueous Interface Norton G. McDuffie 1745 Southwest Whiteside Drive, Corvallis, Oregon 97333 Received April 14, 2001. In Final Form: July 13, 2001 At 20 °C, ammonia is adsorbed in the aqueous interface in a limiting monolayer, unlike less soluble gases that exhibit adsorption moving smoothly from monolayer to multiple layers. The limiting monolayer appears at xNH3 g 0.54. The effective diameter of NH3 in this monolayer is 0.293 nm, indicating that ammonia in the interfacial phase is NH3, not NH3‚H2O. In ammonia concentrations where interfacial water OH groups are available, these findings are consistent with published reports showing orientation of the hydrogenbonded complex to leave the water moiety in the liquid phase adjacent to the interface.
Introduction Ammonia solutions in water exhibit reduced surface tension values. The ubiquity of ammonia as a strategic compound, especially in aqueous solutionssin nature, in agriculture, in living organisms, in heat transfer and heat pump applications, in toxicology, in detergent preparations, in minerals recovery, as a significant product from nuclear wastes,1,2 and in industrial chemical processess makes knowledge of its interfacial behavior extremely important. Here, I show that at 20 °C, ammonia is adsorbed at aqueous interfaces differently than other less-soluble gases: it is adsorbed in a limiting monolayer, whereas the less-soluble gases can generally be adsorbed (or condensed) as multilayers at temperatures below critical.3,4 Furthermore, I show that even though ammonia forms a limiting monolayer like methanol and other polar compounds studied by Schay,5,6 it behaves differently in lower concentration ranges. Results from the present study indicate that NH3 fills the surface monolayer, even though it does participate in hydrogen bonding with exposed (or “dangling”) OH groups at the interface as shown in recent reports. In these reports, ammonia in more dilute solutions [Donaldson7 (mole fraction NH3, xNH3 e ∼0.12) and Simonelli and Shultz8 (xNH3 e 0.3)] is shown to be hydrogen bonded to water at the interface, but with ammonia oriented in the surface layer to give bonding of ammonia N to water OH groups exposed at the inner liquid surface.8 These characteristics of interfacial NH3 are of significance in further establishing that the interfaces of aqueous solutions are essentially separate phases that can contain differing molecular forms, not just transition zones between liquid and gas phases. This possibility was recognized for vapor-liquid interfaces about a century ago by Eo¨tvo¨s,9 Ramsay & Shields,10 and Einstein,11 with further elaboration by Jaeger.12,13 Likewise, it is in (1) McDuffie, N. G. Flammable Gas Generation, Retention, and Release in High-Level Waste Tanks: Physical and Chemical Models; WHC-SA-2129-VA; Westinghouse Hanford Company for U. S. Department of Energy: Richland, WA, 1994. (2) McDuffie, N. G. In WM’94, Proceedings of the Symposium on Waste Management at Tucson, Arizona, February 27-March 3, 1994; Post, R. G., Ed.; Laser Options: Tucson, AZ, 1994; Vol. 1, pp 389-392. (3) Jho, C.; Nealon, D.; Shogbola, S.; King, A. D., Jr. J. Colloid Interface Sci. 1978, 65, 141. (4) Strathdee, G. G.; Given, R. M. J. Phys. Chem. 1976, 80, 1714. (5) Schay, G. In Surface and Colloid Science; Matijevic´, E., Eirich, F. R., Eds.; Wiley-Interscience: New York, 1969; Vol. 2, pp 155-211. (6) Schay, G. In Physical Chemistry: Enriching Topics From Colloid and Surface Science; Olphen, H. van, Mysels, K. J., Eds. for IUPAC Commission I.6, Colloid and Surface Chemistry; Theorex: La Jolla, CA, 1975; pp 229-249. (7) Donaldson, D. J. J. Phys. Chem. A 1999, 103, 62. (8) Simonelli, D.; Shultz, M. J. J. Chem. Phys. 2000, 112, 6804.
concurrence with the indications that aqueous salts are excluded from the interface.13,14 Recent findings show that the aqueous vapor-liquid interface for water15 and of some alcohols and soluble surfactants adsorbed at the interfaces of aqueous solutions16 are of molecular thickness as suggested by Einstein11,12 early in his career. The findings herein, along with recent related reports discussed, provide new insight for investigation of aqueous ammonia vaporliquid transport, nucleation and interfacial behavior in growing and contracting bubbles, chemical reactions in ammonia interfaces, detergent action of aqueous ammonia, and interactions of ammonia with colloidal and suspended particles, and especially the large reported reductions of surface tensions of concentrated salt solutions and slurries by dissolved ammonia.17 Gases with low solubilities in water have the effect of reducing the surface, or interfacial, tension of water when in contact with the aqueous phase, notably at elevated pressures.3,4,13,14 This decrease in surface tension can conveniently be expressed as an effective two-dimensional surface pressure
π ) γ0 - γ
(1)
where γ is the interfacial tension of water in contact with the gas at system pressure and γ0 is interfacial tension of water in contact with its vapor (or, to the nearest approximation, air at low pressures). These quantities have units of N/m or J/m2; thus, the relationship to bulkphase pressure volumetric energy content as J/m3 is often expressed. However, unlike in a single-component system, neither surface pressure nor interfacial tension is numerically equivalent to total surface free energy for a multicomponent system.5,13,20 For the less-soluble gases, (9) Eo¨tvo¨s, R. Ann. Phys. Chem., Neue Folge, G. Wiedemann 1886, 27, 448. (10) Ramsay, W.; Shields, J. Z. Phys. Chem. 1893, 12, 433. (11) Einstein, A. Ann. Phys., 4th series 1911, 34, 165. (12) Jaeger, F. M. Proc. Sect. Sci., Koninklijke Akad. Van Wetenschappen, Amsterdam 1914, 17, 416. (13) Freundlich, H. Colloid & Capillary Chemistry, 3rd ed.; Engl. transl. by H. S. Hatfield; E. P. Dutton and Co.: New York, 1922. (14) Adamson, A. W. Physical Chemistry of Surfaces, 2nd ed.; Interscience Publishers: New York, 1967. (15) Fradin, C.; Braslau, A.; Luzet, D.; Smilgles, D.; Alba, M.; Boudet, N.; Mecke, K.; Dalliant, J. Nature 2000, 403, 871. (16) Abramzon, A. A.; Pesin, Ya. A. Russ. J. Gen. Chem. 1993, 63, 1012. (17) Norton, J. D.; Pederson, L. R. Ammonia in Simulated Hanford Double-Shell Tank Wastes: Solubility and Effects on Surface Tension; PNL-10173; Pacific Northwest Laboratory: Richland, WA, 1994. (18) King, H. H.; Hall, J. L.; Ware, G. C. J. Am. Chem. Soc. 1930, 52, 5128.
10.1021/la010549r CCC: $20.00 © 2001 American Chemical Society Published on Web 08/23/2001
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Langmuir, Vol. 17, No. 19, 2001
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Figure 1. Surface pressure, πNH3, relationship at 20 °C to NH3 partial pressure, pNH3, from surface tension data of King et al. (ref 18) and Wilson (ref 19) vapor-liquid equilibrium data.
at temperatures below their critical temperatures, even including those as soluble as CO2 and H2S, π is an upsloping function of fugacity, f, (or pressure, p, where gas deviation from ideality is negligible).3,4 However, I have found that this is not the case for NH3 (Figure 1). I have analyzed the 20 °C isotherm (Figure 1) following the methods of Schay5,6 using a form of the Gibbs equation,
Γ(n) 2 ) -
x1 ∂γ RT ∂ ln(f2/f°)
(2)
which becomes
Γ(n) 2 )
x1 ∂π RT ∂ ln(f2/f°)
(3)
Water is component 1, ammonia is component 2, and xi represents mole fraction of component i. Γ(n) 2 is the surface excess (mol/m2) of NH3 in an interface where the Gibbs dividing surface is set so as to maintain a constant total number, n, of molecules in the interface. The term (f2/f°) is dimensionless fugacity of NH3, which can be approximated by dimensionless partial pressure of NH3, (p2/ p°), if deviation from ideality is negligible. I have determined, by use of the method of Lee & Chen,21 that the maximum deviation from ideality for NH3 is less than 7% in the pressure range for the 20 °C isotherm of NH3/ H2O. The method generally used to analyze surface excess concentrations in water-gas interfaces involves determination of the excess, Γ(1) 2 , defined as the surface excess concentration of component 2, mol/m2, in an interface limited by a dividing surface positioned to provide a surface excess concentration of component 1 of zero.3,5,6,7,14,22 The excess concentration, Γ(1) 2 ,
Γ2(1) )
1 ∂π RT ∂ ln(f2/f°)
(4)
differs from Γ(n) 2 only by the factor x1: (1) Γ(n) 2 ) x1Γ2
(5)
For a gas that forms a multilayered aqueous interface, (19) Liley, P. E.; Reid, R. C.; Buck, E. In Perry’s Chemical Engineers’ Handbook, 6th ed.; Perry, R. H., Green, D. W., Maloney, J. O., Eds.; McGraw-Hill Book Co.: New York, 1984; 3-71. (20) Ip, S. W.; Toguri, J. M. J. Mater. Sci. 1994, 29, 688. (21) Lee, M.-J.; Chen, J.-T. J. Chem. Eng. Jpn. 1998, 31, 518. (22) Lyklema, J. Fundamentals of Interface and Colloid Science, Vol. 1, Fundamentals; Academic Press Ltd.: London, 1991.
Figure 2. Equilibrium surface excess concentration of NH3 in aqueous vapor-liquid interfaces at 20 °C, with reference surface tension curve. The linear portion of surface excess concentration indicates monolayer, with the extrapolated ordinate intercept giving total NH3 concentration in the monolayer (method after Schay, refs 5 and 6).
Γ(1) 2 increases from zero with increasing aqueous gas concentrations and continues to rise as gas partial pressure is increased beyond that required for the formation of a complete monolayer.3 However, for the NH3/H2O 20 °C isotherm, I find that Γ(1) 2 does increase from zero with increasing ammonia concentrations but then becomes constant at a value of 1.23 × 10-5 mol/m2 beyond pNH3 ) 0.42 MPa and xNH3 ) 0.54. This behavior is probably better represented by Figure 2. This plot illustrates the application of Schay’s method to the analysis. The formation of a distinct monolayer is represented in the linear portion of Γ(n) 2 versus x2. This behavior is similar to that reported for methanol/water5 for the monolayer portion of the adsorption isotherm; however, it is quite different otherwise. The methanol/water adsorption isotherm increases in a steady convex parabolic fashion until it reaches the limiting monolayer at a methanol mole fraction of about 0.4. The ammonia/water isotherm (Figure 2), however, increases slowly up to an ammonia mole fraction of about 0.5 and then suddenly jumps to form a cusp up to the beginning of the linear monolayer portion at an ammonia mole fraction of 0.54. Thus, the ammonia-water isotherm is different from those for the less-soluble gases with water and for miscible methanol-water. The value of Γ(1) 2 is found as the ordinate intercept of the back-extrapolated (1) linear portion of the isotherm (Figure 2), since Γ(n) 2 and Γ2 become identical for the monolayer at x1 ) 1.0. The value of 1.23 × 10-5 mol/m2 or 1.35 × 10-19 m2/molecule for the monolayer gives a molecular diameter of 0.293 nm using the surface hexagonal-close-packing correlation of de Boer23
b2 ) 1.57d2
(6)
in which b2 is minimum surface area covered per molecule and d is the molecular diameter. The value of 0.293 nm agrees very well with the diameter of 0.308 nm calculated for NH3 from a two-dimensional form of the van der Waals equation.23 This indicates that the adsorbed layer (1) is a monolayer and (2) is NH3, not NH3‚H2O or NH3‚NH3 (other than of NH3 in association with internal liquid molecular forms). The strong influence of molecular association to form ammonia monohydrate in the liquid bulk phase is evidenced by the effect on ammonia activity (23) Boer, J. H. de The Dynamical Character of Adsorption, 2nd ed.; Oxford at the Clarendon Press: London, 1968.
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Langmuir, Vol. 17, No. 19, 2001 5713
Figure 4. Depiction of distribution of NH3 (dark) and H2O (light) molecules at the vapor-liquid interface with a NH3 surface monolayer. Figure 3. Surface pressure and partial pressure for NH3 in NH3/H2O mixtures at 20 °C. Note pronounced negative deviation from Raoult’s law for NH3 partial pressure for NH3 concentrations in the 0-0.5 mole fraction range and parallel tracking of the two curves in the 0.5-1.0 NH3 mole fraction range.
in the aqueous phase in the 0-0.5 mole fraction range. This is illustrated in Figure 3. In the 0-0.5 mole fraction range, NH3 activity is shown to have a pronounced negative deviation from ideality, as exhibited by the effect on NH3 partial pressures. The effect is not evident in the surface pressure, or surface tension, relationship to NH3 concentration. For NH3 concentrations above a mole fraction of 0.5, the surface pressure and partial pressure curves are more nearly parallel. The negative deviation from ideality in the lower concentrations can be explained by the strong association of NH3 with H2O in the liquid phase. The surface tension reduction in the dilute ammonia solutions (xNH3 e ∼0.12) has been shown by Donaldson7 to be related to a typical Langmuir relationship for surface adsorption as a function of ammonia liquid activity. The Donaldson numbers for surface excess concentrations at 298 K are essentially the same as those reported herein at 273 K for the dilute solutions. Donaldson, however, fitted his results to a Langmuir adsorption model giving a saturated surface excess (Γ(1) 2 ) of 1.98 µmol/m2, considerably less than a saturated monolayer. The reports of sophisticated spectroscopic methods showing hydrogen bonding of NH3 to water OH at the interface7,8 lead to the inference that surface accumulation of ammonia proceeds by some other mechanism at higher bulk con-
centrations, simply due to the lower concentration of available OH groups at the surface. The finding that NH3 forms a monolayer on aqueous solutions leads to a number of interesting questions. Is this monolayer actually pure NH3 as depicted in Figure 4, or is there a defect in the standard methods of thermodynamic analysis? If the monolayer is pure NH3, what is the mechanism for transport across the interface, and what is the mobility of NH3 relative to that in the liquid? Can the propensity of a gas or vapor to form a monolayer on a liquid without further condensation molecular layers be predicted by solute and solvent properties such as solubility parameters or cohesion parameters? Accounting for the nature of the interface is important in many instances. Specifically for NH3/H2O, in capillary condensation in soil or pores of adsorbents, the surface tension-partial pressure relationships can determine relative uptake of NH3. In very thin condensation films on solid surfaces or in porous media or foams, the presence of a vapor/liquid surface richer in NH3 (or other vapor solutes in general) leads to increasing solute concentration as the mixed films become incrementally thinner. The apparent exclusion, or directed orientation, of ammonia hydrates from the surface films adds further basis for treating the vapor-liquid interface as a very different physical and chemical environment. Interactions of ammonia with particulate (or colloidal) material at the interface can be expected to differ from those within the liquid phase. LA010549R
