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The ammonium nitrate formed in reaction (54) can exist either as a solid particle or in solution, and since this reaction is an equilibrium, it can redissociate to form the reactants. The deliquescence point for NH4N03 at 25°C is 62% RH; i.e., at a water vapor concentration corresponding to 62% RH, the solid particle dissolves to form a concentrated liquid solution.

Mozurkewich (1993) has treated this system in detail and recommends that the equilibrium constant for reactions (54, —54) when the ammonium nitrate is a solid can be calculated as a function of temperature from ln(^54 _54) = 118.87 - 24084/7 - 6.025 ln(T), (F)

where T is in Kelvin and K54_54 is in (nanobar)2. Since 1 bar = 0.987 atm, this is within a few percent of the value expressed in units of (ppb)2.

At water vapor concentrations above the deliquescence point, the equilibrium is that between the reac-tant gases and aqueous ammonium nitrate. As treated in detail by Mozurkewich, the equilibrium constant, K*4_54, then depends on the solution concentrations or activities:

where the a parameters are the solution-phase activities for NH3 and HN03, respectively, which clearly depend on the solution concentrations, and the H factors are the Henry's law constants for these gases. The activity coefficients needed to obtain the activities in solutions of varying concentrations can be obtained by using standard thermodynamic approaches outlined by Mozurkewich (1993).

Figure 7.18 gives the ratio (K*/K)54_i4 of the calculated equilibrium constants for solution-phase ammonium nitrate compared to the solid salt product at various temperatures and water activities. As the water activity, i.e., water vapor pressure above the solution, increases, the equilibrium constant falls. That is, at higher relative humidities, relatively less HN03 and NH3 are found in the vapor phase at equilibrium. This may be why relatively more ammonium nitrate in particles collected on filters evaporates at lower RHs compared to higher ones.

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