Ammonia equilibrium concentration

Essential for synthesis considerations is the abiUty to determine the amount of ammonia present ia an equiUbrium mixture at various temperatures and pressures. ReHable data on equiUbrium mixtures for pressures ranging from 1,000 to 101,000 kPa (10 —1000 atm) were developed early on (6—8) and resulted ia the determination of the reaction equiUbrium constant (9). Experimental data iadicates that is dependent not only on temperature and pressure, but also upon the ratio of hydrogen and nitrogen present. Table 3 fists values for the ammonia equilibrium concentration calculated for a feed usiag a 3 1 hydrogen to nitrogen ratio and either 0 or 10% iaerts (10). 

First, the outlet concentration of ammonia from reactor cannot be increased by 10-20 times due to the limitation of the ammonia equilibrium concentration under the same reaction conditions and at normal pressures and low temperatures, it does not have much practical significance. 

Various amines find application for pH control. The most commonly used are ammonia, morpholine, cyclohexylamine, and, more recently AMP (2-amino-2-methyl-l-propanol). The amount of each needed to produce a given pH depends upon the basicity constant, and values of this are given in Table 17.4. The volatility also influences their utility and their selection for any particular application. Like other substances, amines tend towards equilibrium concentrations in each phase of the steam/water mixture, the equilibrium being temperature dependent. Values of the distribution coefficient, Kp, are also given in Table 17.4. These factors need to be taken into account when estimating the pH attainable at any given point in a circuit so as to provide appropriate protection for each location. 

FIGURE 9.2 (a) In the synthesis of ammonia, the concentrations of N, and H, decrease with time and that of NH1 increases until they finally settle into values corresponding to a mixture in which all three are present and there is no further net change, (bi If the experiment is repeated with pure ammonia, it decomposes, and the composition settles down into a mixture of ammonia, nitrogen, and hydrogen. (The two graphs correspond to experiments at two different temperatures, and so they correspond to different equilibrium compositions.)... 

Now suppose we add ammonium chloride to aqueous ammonia until the solution contains similar concentrations of NH3(aq) and NH4+(aq). The ammonia equilibrium is... 

Use the seven-step procedure. We want to determine pH, for which we need to know the equilibrium concentration of either H3 or OH. The major species present in aqueous ammonia are molecules of NH3 and H2 O. Both of these compounds produce hydroxide ions as minor species in solution NH3(ts q) + H2 0(/) NH4+((2 q) -b OH (a q) = 1.8 x 10" ... 

The strong base is a soluble hydroxide that ionizes completely in water, so the concentration of OH matches the 0.25 M concentration of the base. For the weak base, in contrast, the equilibrium concentration of OH is substantially smaller than the 0.25 M concentration of the base. At any instant, only 0.8% of the ammonia molecules have accepted protons from water molecules, producing a much less basic solution in which OH is a minor species. The equilibrium concentration of unproton-ated ammonia is nearly equal to the Initial concentration. Figure 17-7 summarizes these differences. 

The water equilibrium always exists in aqueous solution. In general, we can focus our initial attention on the equilibria involving other major species (NH3 in this example). Nevertheless, the water equilibrium does exert its effect on the concentrations of OH and H3 O. In this example, the concentration of hydroxide anion is established by the ammonia equilibrium, but the concentration of hydronium cations must be found by applying the water equilibrium. We use this feature in several of our examples in this chapter. 

In principle, the calculation of concentrations of species of a complexation equilibrium is no different from any other calculation involving equilibrium constant expressions. In practice, we have to consider multiple equilibria whenever a complex is present. This is because each ligand associates with the complex in a separate process with its own equilibrium expression. For instance, the silver-ammonia equilibrium is composed of two steps ... 

Figure 2.1. Equilibrium concentration of ammonia in a mixture of initially 1 3 N2 H2 as a function of temperature for several total pressures. Note the slight deviation due to non-ideality of the gases.

Study of the aminodeethoxylation with N-labeled liquid ammonia shows that in the 4-imino compound no incorporation of the label has taken place, proving that in the replacement of the ethoxy group no ring opening is involved. It is unknown whether the aminodeethoxylation occurs according to routes (a) and (b) in the a-adduct 22 or in the starting material 21, which is present in only a small equilibrium concentration with 22 (Scheme III.13). One can expect, however, that despite its low concentration, the aminodeethoxylation reaction takes place in the pyrimidinium salt 21, being more reactive towards to nucleophiles than the neutral adduct 22. 

Figure 17.20. Control of temperature in multibed reactors so as to utilize the high rates of reaction at high temperatures and the more favorable equilibrium conversion at lower temperatures, (a) Adiabatic and isothermal reaction lines on the equilibrium diagram for ammonia synthesis, (b) Oxidation of SOz in a four-bed reactor at essentially atmospheric pressure, (c) Methanol synthesis in a four bed reactor by the ICI process at 50 atm not to scale 35% methanol at 250°C, 8.2% at 300°C, equilibrium concentrations.

The overall strategy for this calculation is to replace the partial pressures that appear in K by the molar concentrations, and thereby generate Kc. We need to keep track of the units so we write activities as Pj/bar and molar concentrations as Q]/(mol-L 1), as explained in the earlier side-notes. We consider a specific case the relation between K and Kc for the ammonia equilibrium, reaction C. 

A mixture of reactants and products in the equilibrium state is called an equilibrium mixture. In this chapter, we ll address a number of important questions about the composition of equilibrium mixtures What is the relationship between the concentrations of reactants and products in an equilibrium mixture How can we determine equilibrium concentrations from initial concentrations What factors can be exploited to alter the composition of an equilibrium mixture This last question is particularly important when choosing conditions for the synthesis of industrial chemicals such as hydrogen, ammonia, and lime (CaO). 

In the early 1900s, the German chemist Fritz Haber discovered that a catalyst consisting of iron mixed with certain metal oxides causes the reaction to occur at a satisfactory rate at temperatures where the equilibrium concentration of NH3 is reasonably favorable. The yield of NH3 can be improved further by running the reaction at high pressures. Typical reaction conditions for the industrial synthesis of ammonia are 400-500°C and 130-300 atm. 

Tabus XII.—Effect of Pressure and Temperature on the Equilibrium Concentration op Ammonia. 

The steam requirements in an ammonia unit can be reduced by lowering the steam-to-carbon ratio to the primary reformer. However a number of drawbacks can exist downstream in the I I I S and LTS reactors. The drawbacks include By-product formation in the HTS, Pressure drop buildup in the HTS, Reversible poisoning of the LTS catalyst, and Higher CO equilibrium concentrations exiting the HTS and LTS reactors. 

Equilibrium concentrations of carbon or ammonia are not found in short combustion chambers used in rocket motors. The reason for this non-equilibrium situation is that the rate of formation of soot is very slow and carbon does not have time to form. Similarly the dissociation of NH3 is very slow. Thus in ethylene oxide monopropellant rocket motors one finds very little carbon, whereas equilibrium considerations predict carbon as a predominant product and in hydrazine decomposition chambers one finds an excess of NH3 over that predicted by equilibrium considerations. In ethylene oxide motors carbon forms from the decomposition of methane, not the reaction represented above, thus both non-equilibrium situations give higher performance than expected, since the endothermic reactions do not have time to take place. Of course, carbon also could form in cool reactions which take place in boundary layers along the walls where velocities are slow. 

Greater than equilibrium concentrations of intermediate species have been observed in the combustion products of several reactant systems. Examples are the concentrations of ammonia in the products of the decomposition of hydrazine (32), the concentration of CH4 in ethylene oxide decomposition (33), nitric oxide and ammonia in the products of the reaction of hydrazine and nitrogen tetroxide (34), and chlorine monofluoride in the products of the reaction of hydrazine and chlorine pentafluorlde (35). 

Equilibrium concentrations of carbon or ammonia are not found in short combustion chambers used in rocket motors. 

When the reactants and products of a given chemical reaction are mixed, it is useful to know whether the mixture is at equilibrium and, if it is not, in which direction the system will shift to reach equilibrium. If the concentration of one of the reactants or products is zero, the system will shift in the direction that produces the missing component. However, if all the initial concentrations are not zero, it is more difficult to determine the direction of the move toward equilibrium. To determine the shift in such cases, we use the reaction quotient (Q). The reaction quotient is obtained by applying the law of mass action, but using initial concentrations instead of equilibrium concentrations. For example, for the synthesis of ammonia,... 

It is important to understand the factors that control the position of a chemical equilibrium. For example, when a chemical is manufactured, the chemists and chemical engineers in charge of production want to choose conditions that favor the desired product as much as possible. In other words, they want the equilibrium to lie far to the right. When Fritz Haber was developing the process for the synthesis of ammonia, he did extensive studies on how the temperature and pressure affect the equilibrium concentration of ammonia. Some of his results are given in Table 6.2. Note that the amount of NH3 at equilibrium increases with an increase in pressure but decreases with an increase in temperature. Thus the amount of NH3 present at equilibrium is favored by conditions of low temperature and high pressure. 

Around 1900 Fritz Haber began to investigate the ammonia equilibrium [11] at atmospheric pressure and found minimal ammonia concentrations at around 1000 °C (0.012 %). Apart from Haber, Ostwald and Nernst were also closely involved in the ammonia synthesis problem, but a series of mistakes and misunderstandings occurred during the research. For example, Ostwald withdrew a patent application for an iron ammonia synthesis catalyst because of an erroneous experiment, while Nernst concluded that commercial ammonia synthesis was not feasible in view of the low conversion he found when he first measured the equilibrium at 50 - 70 bar [12] - [14],... 

---