5.2.1.1 Gas generator—water gas method

This method, developed by Bosch in 1920, is characterized by two distinct reactions for processing coal into syngas. On the one hand, the partial oxidation of the coal, and on the other hand, its processing with steam. Each of the main reactions is also affected by the chemical equilibrium among the species involved, Boudouard in the case of carbon oxidation, and water gas shift (WGS) in the case of steam processing.

5.2.1.1.1 Partial oxidation (gas generator)

This stage consists of flowing oxygen through a bed of incandescent coal. Limited air and a certain bed height are required for the production of CO and CO₂ together. In the first 15 cm of the bed C is combusted into CO and CO₂. When oxygen is no longer available, Boudouard equilibrium occurs so that the CO₂ reacts with the bed to produce CO (see the reactions below). The global process is exothermic:

$\begin{array}{l}2{\mathrm{C}}_{(s)}+1.5{\mathrm{O}}_{2(\mathrm{g})}\to {\mathrm{CO}}_{(g)}+{\mathrm{CO}}_{2(g)}\hfill \\ {\mathrm{C}}_{(s)}+{\mathrm{CO}}_{2(g)}\leftrightarrow 2{\mathrm{CO}}_{(g)}(\text{Boudouard})\hfill \end{array}$

$\begin{array}{l}3{\mathrm{C}}_{(s)}+1.5{\mathrm{O}}_2\to 3{\mathrm{CO}}_{(g)}\hfill \\ {\mathrm{C}}_{(s)}+\frac{1}{2}{\mathrm{O}}_{2(g)}\to \mathrm{CO}+26.41\text{kcal}/\text{mol}\hfill \end{array}$

The Boudouard reaction is of particular interest. It is endothermic,

${\mathrm{C}}_{(s)}+{\mathrm{CO}}_{2(g)}\leftrightarrow 2{\mathrm{CO}}_{(g)}41\text{kcal}/\text{mol}$

with an equilibrium constant given by the following expression: Ortuño (1999)

$K_{\mathrm{p}}=\frac{P_{\mathrm{CO}}^2}{P_{{\mathrm{CO}}_2}}\Rightarrow \log (K_{\mathrm{p}})=9.1106-\frac{8841}{T},T[=]\mathrm{K}$ (5.1)

(5.1)

Pressure and temperature have an effect on the composition of the gas. High temperatures favor the production of CO (it is an endothermic reaction), while pressure reverts the equilibrium into CO₂; the number of moles is higher on the right-hand side. Let n_i be the moles of species i, X the conversion, and P_T the total pressure. Note that the solid coal does not take part in the equilibrium constant given by Eq. (5.2). Fig. 5.4 shows the effect of pressure and temperature in the equilibrium.

$K_{\mathrm{p}}=\frac{P_{\mathrm{CO}}^2}{P_{{\mathrm{CO}}_2}}=\frac{n_{\mathrm{CO}}^2·P_{\mathrm{T}}}{n_T\cdot n_{{\mathrm{CO}}_2}}=\frac{{(2X)}^2\cdot P_{\mathrm{T}}}{(1+X)·(1-X)}=\exp \left[9.1106-\frac{8841}{T(\mathrm{K})}\right]$ (5.2)

(5.2)

Thus, for the production of CO, we need limited availability of oxygen, a certain bed height, and high temperature (at the risk of melting the ash).

5.2.1.1.2 Steam processing (water gas)

The other stage in the process for syngas production from coal using this method is the use of steam. The main reaction produces CO and H₂. It is an endothermic reaction ($31.35\text{kcal}/\text{mol}$). Thus, CO is favored at high temperature:

${\mathrm{C}}_{(s)}+{\mathrm{H}}_2{\mathrm{O}}_{(g)}\to {\mathrm{CO}}_{(g)}+{\mathrm{H}}_{2(g)}$

The reaction is an equilibrium whose constant can be written as follows:

${\exp }^{\left[\frac{-131,000}{(8.314\cdot T(K)}+\frac{134}{8.314}\right]}={\displaystyle \frac{(P_{\mathrm{T}}·n_{\mathrm{CO}}·n_{{\mathrm{H}}_2})}{(n_{{\mathrm{H}}_2\mathrm{O}}·(n_{\mathrm{Total}}))}}$ (5.3)

(5.3)

The second reaction that occurs is the well-known water gas shift reaction (WGSR), given by the following equation:

$\mathrm{CO}{}_{(g)}+{\mathrm{H}}_2{\mathrm{O}}_{(g)}\leftrightarrow {\mathrm{CO}}_{2(g)}+{\mathrm{H}}_{2(g)}$

It is an exothermic reaction affected by temperature. The wide application of this reaction deserves more attention. Thus, a more detailed discussion is needed.

The equilibrium between the four species is as follows:

$\mathrm{CO}{}_{(g)}+{\mathrm{H}}_2{\mathrm{O}}_{(g)}\leftrightarrow {\mathrm{CO}}_{2(g)}+{\mathrm{H}}_{2(g)}\mathrm{\Delta }H=-41\text{kJ}/\text{mol}$

It is widely used for modulating the concentration of CO and hydrogen in a stream, in the production of hydrogen from steam reforming of hydrocarbons, or in carbonous mass gasification. The equilibrium is reversible, and the reverse WGS can also be used if needed. The reaction is an equilibrium whose constant has been widely evaluated in the literature. The following are several correlations by Graaf et al. (1986), eq. (5.4) and by Susanto and Beenackers (1996), eqs. (5.5) and (5.6):

$\log K_{\mathrm{p}}=\log \left(\frac{P_{{\mathrm{CO}}_2}{P}_{{\mathrm{H}}_2}}{P_{{\mathrm{H}}_2\mathrm{O}}{P}_{\mathrm{CO}}}\right)=\log \left(\frac{n_{{\mathrm{CO}}_2}{n}_{{\mathrm{H}}_2}}{n_{{\mathrm{H}}_2\mathrm{O}}{n}_{\mathrm{CO}}}\right)=\left(\frac{2073}{T(\mathrm{K})}-2.029\right)$ (5.4)

(5.4)

$\log (K_{\mathrm{p}})=\frac{2232}{T(\mathrm{K})}+0.0836\log (T(\mathrm{K}))+0.00022T(\mathrm{K})+2.508$ (5.5)

(5.5)

$K_{\mathrm{p}}=0.0265e^{\left[\frac{33,010}{8.314·T(\mathrm{K})}\right]}$ (5.6)

(5.6)

There is some flexibility for locating the WGS reactor within the processes, depending on the presence of sulfur compounds in the gas. It can be located either before the sulfur removal step (sour shift) or after sulfur removal (sweet shift).

Sweet operation of the WGSR occurs in two steps at high and low temperature, respectively. They are known as high-temperature shift (HTS) and low-temperature shift (LTS).

A conventional high-temperature sweet shifting operates between 315°C and 530°C and uses chromium- or copper-promoted, iron-based catalysts. The catalysts are calcinated over 500°C so that the Fe reaches α phase. Next, they are reduced at temperatures between 315°C and 460°C. Their half-time life is around three years. The difficulty in removing the energy generated in the reaction makes them work adiabatically. Thus, the feed is at around 310–360°C and 10–60 atm, and it is allowed to heat up. The conversion of CO is over 96%. The spatial velocity of the gas is typically 300–4000/h.

Low-temperature sweet shifting (up to 370°C) uses copper–zinc–aluminum catalysts. For instance, ICI catalyst contains 30% by weight of CuO, 45% of ZnO, and 13% of Al₂O₃. Furthermore, the use of supported Cu over chromium oxide (15–20% of CuO, 68–73% of ZnO, and 9–14% of Cr₂O₃) constitutes the Haldor Tøpsoe commercial catalyst that operates from 220°C to 320°C. The aim of this catalyst is to achieve 99% conversion of CO operating at 3–40 atm. The feed enters the reactor at 200–220°C and operates adiabatically. The spatial velocity of the gas is typically 300–4000/h. The half-life time of the catalyst ranges from two to four years.

A second type of WGSR is sour WGSR. For processing gases containing H₂S, a catalyst based on cobalt–molybdenum is used. It is possible to add Li, Na, Cs, or K. The reaction is typically located after the water scrubber, where syngas is saturated with water at about 200–250°C. Furthermore, this catalyst can also convert carbonyl sulfide and other organic sulfur compounds into H₂S. The reactor operates adiabatically, reaching conversions of CO of around 97%. Apart from packed beds, lately membrane reactors are gaining attention. They are based on the particularly high diffusivity of hydrogen through metals; see Fig. 5.5. Thus, the reaction occurs in the inner pipe and the hydrogen crosses the palladium-based wall. The purity of the hydrogen produced reaches 99.99% (Osenwengie Uyi, 2007).

Let us illustrate the operation and yield of WGS reactors in a numerical example.

Example 5.1

A WGS reactor works at 600K and the feed consists of 36% CO, 30% CO₂, 20% steam, and 14% hydrogen.

a. Compute the conversion of the reaction.

b. Plot the profile of the conversion with the temperature.

Hint: use the correlation for K_p given by the following equation:

$K_{\mathrm{p}}=0.0265e^{\left[\frac{33,010}{8.314T}\right]}$

Solution

a. We write the equation of the constant (see Eq. 5.4) as follows:

$K_{\mathrm{p}}=\frac{(n_{{\mathrm{CO}}_2}+X)(n+X)}{(n_{\mathrm{CO}}-X)(n_{{\mathrm{H}}_2\mathrm{O}}-X)}=0.0265e^{\left[\frac{33,010}{8.314T}\right]}$

Solving for X we choose the lowest solution: $X=16.4$.
With respect to the limiting reactant, H₂O, the conversion is $X=\frac{16.4}{20}=0.82$.

b. For a range of temperatures, we compute the conversion following the same procedure as before (Fig. 5E1.1).

Now that both reactions, the gas generator, and the water gas path have been evaluated, the production of syngas consists of using both to process the coal and adjust the composition of the syngas.

$\begin{array}{l}\mathrm{Air}\mathrm{gas}3.76{\mathrm{N}}_2+{\mathrm{O}}_2+2\mathrm{C}\to 2\mathrm{CO}+3.76{\mathrm{N}}_2\mathrm{Exothermic}\\ \mathrm{Water}\mathrm{gas}{\mathrm{H}}_2\mathrm{O}+\mathrm{C}\to \mathrm{CO}+{\mathrm{H}}_2\mathrm{Endothermic}\end{array}$

Fig. 5.6 shows the scheme for the process dividing the use of coal between the two processes. In Example 5.2 we evaluate the performance of such a process.

Unfortunately, at pressures above atmosphere and low temperatures, methane is produced following this reaction:

${\mathrm{C}}_{(s)}+2{\mathrm{H}}_{2(g)}\leftrightarrow {\mathrm{CH}}_{4(g)}$

When the temperature rises, the equilibrium reverses. The equilibrium constants for the reaction for a couple of temperature values are Log₁₀ K_p (25°C)=9 and Log₁₀ K_p (100°C)=−2. Table 5.1 shows the effect of temperature on methane production.

Table 5.1

Methane Formation in the Flue Gas/Water Gas Method

T (°C)	%CO2	%CO	%H2	%CH4	%H2O
400	25.2	0.50	11.3	20.0	43.0
600	19.5	15.5	37.5	8.50	19.0
800	2.50	46.0	48.0	1.50	2.00
1000	0.20	50.0	49.0	0.50	0.30

5.2.1.2 Coal distillation

Coal distillation is the chemical process that decomposes a carbonous material by heating it up to high temperatures in the absence of air. It is basically a pyrolysis that allows the production of a number of species, from coke to coal gas, gas carbon, Buckmisterfullerene, and hydrogen—when operating from 800°C to 1000°C. As the coal is heated, several processes take place:

Typically, from 100 kg of coal we obtain 72 kg of coke, 22 kg of gas (hydrogen, ethylene, ethane, methane, CO, CO₂, and nitrogen), and 6 kg of tar containing mainly benzene, toluene, xylene, phenol, naphthalene, anthracene, and liquor.

5.2.1.3 Gasification

Gasification is the partial oxidation of a carbon substrate into a gas with a low heating value; it proceeds via a series of reactions that occur in the presence of a gasification agent such as air, oxygen, and/or steam. To produce hydrogen, gas reforming after gasification is required. There are a number of designs for the gasifiers. They are first classified as either fixed or fluidized beds. Fig. 5.7 shows four designs: (a) a fixed bed with countercurrent solids and gas circulation; (b) a downdraft-type fixed bed where the carbonous material and the gases exit the unit from the bottom; (c) a bubbling-type fluidized bed where only one unit is involved; and (d) a circulating-type fluidized bed that requires solid gas separation to recycle the solids. The product gas composition is highly dependent on the configuration, the raw materials, and the operating conditions. Brigdwater (1995) presented some tables that gathered the yields of different designs. The operating temperatures range from 700°C to 1000°C and the pressure from slightly above atmospheric to 30 atm. Typically, the largest production capacity is obtained using fluidized beds.

Lately, gasification has been used to process biomass to obtain syngas. This has been used for the production of different biofuels such as ethanol, Fischer–Tropsch (FT) liquids, or hydrogen (Martín and Grossmann, 2013). The two basic designs are either the one based on the FERCO/Battelle design or the one by the Gas Technology Institute (GTI). Fig. 5.8 shows the scheme for the GTI gasifier. It is a bubbling-type gasifier that operates above atmospheric pressure (5–20 atm) and 800–900°C. It consists of one single unit, and in order to avoid gas dilution, oxygen (instead of air) and steam are used as fluidification agents. The gas produced has a low concentration of hydrocarbons, but the fraction of CO₂ is large. It is called Renugas. Fig. 5.9 shows the scheme of the FERCO/Battelle gasifier. The system consists of two units: the gasifier and the combustor. In the gasifier, steam is mixed with the carbonous material. The energy for the gasification is provided by hot sand. The solids, sand and char, are separated in the cyclone, and both are sent to a combustor where air is used to burn the char and reheat the sand. A cyclone separates the flue gas from the solids (the sand), which are recycled to the gasifier. The syngas produced in this system has a higher concentration of hydrocarbons but typically lower concentration of CO₂. Furthermore, since the combustion of the char takes place in a second unit, air can be used without diluting the syngas.

The main reactions can be divided into several types:

– Oxygen-based reactions:

$\begin{array}{l}\mathrm{C}+{\mathrm{O}}_2\to {\mathrm{CO}}_2\hfill \\ \mathrm{C}+0.5{\mathrm{O}}_2\to \mathrm{CO}\hfill \\ \mathrm{CO}+0.5{\mathrm{O}}_2\to {\mathrm{CO}}_2\hfill \end{array}$

– Steam-based reactions:

$\begin{array}{l}\mathrm{C}+{\mathrm{H}}_2\mathrm{O}\to \mathrm{CO}+{\mathrm{H}}_2\hfill \\ \mathrm{C}+2{\mathrm{H}}_2\mathrm{O}\to {\mathrm{CO}}_2+2{\mathrm{H}}_2\hfill \\ \mathrm{CO}+{\mathrm{H}}_2\mathrm{O}\to {\mathrm{CO}}_2+{\mathrm{H}}_2\hfill \end{array}$

– Methane formation:

$\begin{array}{l}\mathrm{C}+2{\mathrm{H}}_2\to {\mathrm{CH}}_4\hfill \\ 2\mathrm{C}+2{\mathrm{H}}_2\mathrm{O}\to {\mathrm{CH}}_4+{\mathrm{CO}}_2\hfill \\ \mathrm{CO}+3{\mathrm{H}}_2\to {\mathrm{CH}}_4+{\mathrm{H}}_2\mathrm{O}\hfill \\ {\mathrm{CO}}_2+4{\mathrm{H}}_2\to {\mathrm{CH}}_4+2{\mathrm{H}}_2\mathrm{O}\hfill \end{array}$

– The Boudouard reaction:

$\mathrm{C}+{\mathrm{CO}}_2\to 2\mathrm{CO}$

– Other reactions (eg, hydrocarbon production):

$\begin{array}{l}{\mathrm{H}}_2+0.5{\mathrm{O}}_2\to {\mathrm{H}}_2\mathrm{O}\hfill \\ {\mathrm{CH}}_4+2{\mathrm{O}}_2\to {\mathrm{CO}}_2+2{\mathrm{H}}_2\mathrm{O}\hfill \\ {\mathrm{C}}_2{\mathrm{H}}_4+3{\mathrm{O}}_2\to 2{\mathrm{CO}}_2+2{\mathrm{H}}_2\mathrm{O}\hfill \\ {\mathrm{C}}_n{\mathrm{H}}_{2n}+{\mathrm{H}}_2\to {\mathrm{C}}_n{\mathrm{H}}_{2n+2}\hfill \\ {\mathrm{C}}_n{\mathrm{H}}_m+n{\mathrm{H}}_2\mathrm{O}\to n\mathrm{CO}+(n+m/2){\mathrm{H}}_2\hfill \end{array}$

If gasification is the first stage to hydrogen production, after gas purification, a WGS stage is needed.

5.2.1.4 Thermal decomposition/partial oxidation

1. Hydrocarbon + Heat → H₂ + Lamp Black
This reaction is of interest for the production of lamp black more than for hydrogen; the hydrogen is not pure enough.

2. Hydrocarbon + Heat → H₂ + Unsaturated Hydrocarbon
For instance, the production of ethylene from methane:

$\begin{array}{l}2{\mathrm{CH}}_4\leftrightarrow {\mathrm{C}}_2{\mathrm{H}}_2+2{\mathrm{H}}_2\hfill \\ {\mathrm{C}}_2{\mathrm{H}}_2\leftrightarrow 2\mathrm{C}+{\mathrm{H}}_2\hfill \end{array}$

3. Hydrocarbon + O₂ + Cooling → H₂ + CO  (partial oxidation)

Operating under limited availability of oxygen, this reaction not only produces energy, but also syngas. The yield to hydrogen is lower than when using steam. The oxygen needed can be fed as pure oxygen if the syngas is due to produce fuels, or air if nitrogen is required (eg, in the production of ammonia). The presence of nitrogen typically results in an increased size of all the units downstream, and it is not recommended.

5.2.1.5 Steam reforming of hydrocarbons

In the section on gasification, reforming was mentioned as a means to enhance the proportion of hydrogen by decomposing the hydrocarbons into CO and hydrogen. If the feedstock is already made of hydrocarbons, reforming is the first stage towards syngas:

$\begin{array}{l}\mathrm{Hydrocarbon}+\mathrm{Steam}\to {\mathrm{H}}_2+\mathrm{CO}\hfill \\ {\mathrm{C}}_2{\mathrm{H}}_2+2{\mathrm{H}}_2\mathrm{O}\to 2\mathrm{CO}+3{\mathrm{H}}_2\hfill \\ {\mathrm{CH}}_4+{\mathrm{H}}_2\mathrm{O}\to \mathrm{CO}+3{\mathrm{H}}_2\hfill \end{array}$

The most representative case is methane steam reforming. The process involves a coupled equilibria, that of the decomposition of the methane and that of the WGS. The constants are presented below Davies and Lihou (1971):

${\mathrm{CH}}_4+{\mathrm{H}}_2\mathrm{O}\leftrightarrow \mathrm{CO}+3{\mathrm{H}}_2{K}_{\mathrm{p}}={10}^{\left[\frac{-11,650}{T}+13.076\right]}T[=]\mathrm{K}$ (5.7)

(5.7)

$\mathrm{CO}{}_{(\mathrm{g})}+{\mathrm{H}}_2{\mathrm{O}}_{(\mathrm{g})}\leftrightarrow {\mathrm{CO}}_{2(\mathrm{g})}+{\mathrm{H}}_{2(\mathrm{g})}{K}_{\mathrm{p}}={10}^{\left[\frac{1910}{T}-1.784\right]}T[=]K$ (5.8)

(5.8)

Steam reforming is an endothermic process that requires a large amount of energy. Furthermore, the products show a larger number of moles. Therefore, pressure drives the equilibrium to the reactants, not to the product. The WGSR is exothermic and increases the proportion of hydrogen. Here we present the kinetics for this particular packed bed reactor, where y represents the molar fraction:

$\begin{array}{l}{r}_1=\frac{{\rho }_{\mathrm{b}}·A_1·e^{\left[-E{a}_1/RT\right]}·P_{{\mathrm{CH}}_4}}{(1+K_{\mathrm{a}}·P_{{\mathrm{H}}_2})}\hfill \\ r_2={\rho }_{\mathrm{b}}·A_2·e^{\left[-E{a}_2/RT\right]}\left(y_{{\mathrm{CH}}_4}{y}_{{\mathrm{H}}_2}-\frac{y_{{\mathrm{CO}}_4}{y}_{{\mathrm{H}}_2}}{K_{{\mathrm{eq}}_2}}\right)\hfill \end{array}$ (5.9)

(5.9)

The reforming of natural gas for the production of hydrogen consists of four stages, namely, purification, reforming, WGS, and gas purification. Fig. 5.10 shows the scheme of the process. The natural gas must be purified to remove sulfur compounds that are poisonous to the catalysts downstream. Typically a hydrodesulfurization stage transforms the organic sulfur compounds into H₂S using as catalyst an alumina base impregnated with cobalt and molybdenum (usually called a CoMo catalyst) operating at 300–400°C and 30–130 atm in a fixed bed reactor. For example, ethanethiol (C₂H₅SH):

${\mathrm{C}}_2{\mathrm{H}}_5\mathrm{SH}+{\mathrm{H}}_2\to {\mathrm{H}}_2\mathrm{S}+{\mathrm{C}}_2{\mathrm{H}}_6$

After that, a bed of ZnO, operating at 150–200°C, is used to remove H₂S:

$\mathrm{ZnO}+{\mathrm{H}}_2\mathrm{S}\to {\mathrm{H}}_2\mathrm{O}+\mathrm{ZnS}$

Alternatively, iron-based oxides (Fe₂O₃) operating at 25–50°C can also be used:

$\begin{array}{l}{\mathrm{Fe}}_2{\mathrm{O}}_3+3{\mathrm{H}}_2\mathrm{S}\to {\mathrm{Fe}}_2{\mathrm{S}}_3+3{\mathrm{H}}_2\mathrm{O}\hfill \\ 2{\mathrm{Fe}}_2{\mathrm{S}}_3+3{\mathrm{O}}_2\to 2{\mathrm{FeO}}_3+6\mathrm{S}\hfill \end{array}$

At this point, the gas is fed to the reformer. Remember that hydrocarbon reforming is endothermic and a fraction of the feed is used to provide that energy. The primary reformer uses Ni/Al₂O₃ as a catalyst, and the secondary one, if used, is Ni supported on ceramics. Around 10% of the feed is typically used as fuel. In the reforming reaction, CO is produced, as seen above. Thus WGS is carried out in two stages: an HTS at 350°C using Fe₃O₄ as a catalyst, and an LTS at 200°C using CuO as catalyst, operating adiabatically, to achieve more than 97% conversion. Finally, the CO₂ generated is eliminated. Several technologies that will be presented in the following section can be used, such as absorption in ethanolamines or adsorption in fixed beds, pressure swing adsorption (PSA) systems. Finally, the traces of CO and CO₂ are eliminated by methanation. The reaction consumes a small amount of hydrogen using Ni/Al₂O₃ as catalyst:

$\mathrm{CO}+{\mathrm{CO}}_2+7{\mathrm{H}}_2\to 2{\mathrm{CH}}_4+{\mathrm{3H}}_2\mathrm{O}+\mathrm{Heat}$

The purity achieved is above 99.95 mol%. Fig. 5.11 shows a scheme of the reformer furnace. The process typically requires 1.3 Mg/d of steam at 2.6 MPa, and produces 1.9 Mg/d at 4.8 MPa.

In the production of ammonia, the reforming system consists of two reformers (see Fig. 5.11). A primary reformer is fed with desulfurized gas that has been preheated to 400–600°C. The furnace consists of a number of high-nickel chromium alloy tubes filled with a nickel-based catalyst. The heat to the process is provided by burning fuel, typically a part of the supply, requiring from 7 to 8 GJ per ton of ammonia. In the secondary reformer, a packed bed reactor, the nitrogen needed for preparation of the synthesis gas is added. Air is compressed and heated on the convection section of the primary reformer and fed to the secondary reformer. Next, HTS and LTS are carried out to produce more hydrogen by reducing the CO generated in the reforming stage. Throughout the chapter, the complete process will be described.

In Table 5.2 the comparisons between fired furnaces and gas-heated reformers are presented, justifying the use of the second ones.

5.2.1.6 Dry reforming and autoreforming

So far steam reforming and partial oxidation have been considered to produce syngas or hydrogen. There are two more alternatives. The first one is the combination of the previous two. Autoreforming uses the energy generated in the partial oxidation of a fraction of the feed to provide the energy required for the steam reforming reactions. In this way, the operation is adiabatic. The stoichiometry depends on the energy balances. For example:

$3{\mathrm{CH}}_4+{\mathrm{O}}_2+{\mathrm{H}}_2\mathrm{O}\to 3\mathrm{CO}+7{\mathrm{H}}_2$

Dry reforming is a novel technology in which CO₂ is used to treat hydrocarbons to produce syngas. The reaction is endothermic and the yield to hydrogen is low, but it represents a process where CO₂ becomes a carbon source to be reintegrated in the process:

${\mathrm{C}}_2{\mathrm{H}}_2+2{\mathrm{CO}}_2\to 4\mathrm{CO}+{\mathrm{H}}_2$

5.3.2.1 Sour gas removal

Sour gas removal is a more general case of carbon capture. The technologies have been used at industrial scale over the last few decades, but only now are they attracting more attention Plasynski and Chen (1997), GPSA (2004).

5.3.2.1.1 Alkali solutions

These are used for the simultaneous removal of CO₂ and H₂S via absorption; Fig. 5.12 shows a scheme. The use of ethanolamines is based on a series of consecutive reactions between the sour gases and the alkali solution. The absorption occurs at low temperature (29°C) and moderate pressure (29 atm). To avoid dilution of the solution by the condensed water, it is separated upon cooling. There are a number of rules of thumb to determine the flow of amine as a function of the type. Typically, for monoethanolamine, a solution of 15–25% by weight with a correction factor of 0.25–0.4 based on a mol-to-mol-based reaction is used. The reaction is exothermic, and therefore the exit liquid heats up. The reactions for capturing the CO₂ are as follows:

$\begin{array}{l}2{\mathrm{H}}_2\mathrm{O}\leftrightarrow {\mathrm{OH}}^{-}+{\mathrm{H}}_3{\mathrm{O}}^{+}\hfill \\ {\mathrm{CO}}_{2(g)}\leftrightarrow {\mathrm{CO}}_{2(\mathrm{aq})}\hfill \\ {\mathrm{CO}}_{2(\mathrm{aq})}+2{\mathrm{H}}_2\mathrm{O}\leftrightarrow {{\mathrm{HCO}}_3}^{-}+{\mathrm{H}}_3{\mathrm{O}}^{+}\hfill \\ {{\mathrm{HCO}}_3}^{-}+{\mathrm{H}}_2\mathrm{O}\leftrightarrow {{\mathrm{CO}}_3}^{2-}+{\mathrm{H}}_3{\mathrm{O}}^{+}\hfill \\ {\mathrm{CO}}_{2(\mathrm{aq})}+{\mathrm{RNH}}_2\leftrightarrow {\mathrm{RNH}}_2^{+}{\mathrm{COO}}^{-}\hfill \\ {\mathrm{RNH}}_2^{+}{\mathrm{COO}}^{-}+{\mathrm{RNH}}_2\leftrightarrow {\mathrm{RNH}}_3^{+}+{\mathrm{RNHCOO}}^{-}\hfill \\ {\mathrm{RNH}}_2^{+}{\mathrm{COO}}^{-}+{\mathrm{H}}_2\mathrm{O}\leftrightarrow {\mathrm{H}}_3{\mathrm{O}}^{+}+{\mathrm{RNHCOO}}^{-}\hfill \\ {\mathrm{RNH}}_2^{+}{\mathrm{COO}}^{-}+{\mathrm{OH}}^{-}\leftrightarrow {\mathrm{H}}_2\mathrm{O}+{\mathrm{RNHCOO}}^{-}\hfill \end{array}$

where the absorption kinetics are given by the equations below (Kierzkowska-Pawlak and Chacuk, 2010):

$\begin{array}{l}r=k_2\left[\mathrm{Amine}\right]{[\text{CO}}_2]=k_{\mathrm{ob}}{[\text{CO}}_2]\hfill \\ k_{\mathrm{ob}}=k_2{\left[\text{Amine}\right]}_0=2.07·{10}^9\exp \left(\frac{-5912.7}{T(\mathrm{K})}\right){\left[\mathrm{Amine}\right]}_0\hfill \\ k_2[=]{\mathrm{m}}^3{\mathrm{kmol}}^{-1}{\mathrm{s}}^{-1}\hfill \end{array}$ (5.10)

(5.10)

As long as:

$\begin{array}{l}3<Ha<<E_i\hfill \\ 3<\frac{\sqrt{k_{\mathrm{ob}}{D}_{{\mathrm{CO}}_2}}}{k_{\mathrm{L}}}<<1+\frac{D_{\mathrm{amine}}\left[\mathrm{Amine}\right]}{{\nu }_{\mathrm{amine}}{D}_{{\mathrm{CO}}_2}{\left[{\mathrm{CO}}_2\right]}_i}\hfill \\ {\left[{\mathrm{CO}}_2\right]}_i=p_{{\mathrm{CO}}_2}/H;{\nu }_{\mathrm{amine}}={\mathrm{stoichiometry}}_{{\mathrm{CO}}_2-\mathrm{amine}};D_i=\mathrm{Diffusivities}(m/s);k_{\mathrm{L}}=Sh·D_{{\mathrm{CO}}_2}/d_{\mathrm{b}}\hfill \\ \log (H)=5.3-0.035°{\mathrm{C}}_{\mathrm{amine}}-1140/T(\mathrm{K})\hfill \end{array}$ (5.11)

(5.11)

The solution is regenerated by distillation. The column operates at 1.7 atm and is fed with the solution at 93°C. The distillate exits at 54°C and the bottoms at 125°C. The energy involved in the different heat exchangers can be computed based on rules of thumb as a function of the flow of amine in gallons per minute, see Figure 5.12 (GPSA, 2004)

Example 5.4

A stream of 40 kmol/s at 34.5 atm and 298K of syngas with 10% CO₂ is processed to remove 99.5% of it before further stages. The equilibrium data for the system CO₂–amine are given in the table below. The L/G ratio is 1.5, the minimum. Determine the number of stages in a gas–liquid absorption tower for the removal of CO₂ from syngas assuming 25% efficiency per stage (Table 5E4.1).

Table 5E4.1

Equilibrium Data

mol CO2/mol Amine	PCO2(Pa)
0.025	65
0.048	197
0.145	1204
0.276	3728
0.327	4815
0.45	8253
0.9	26,325
10	1,480,620
15	2,918,315
20	4,722,992

Solution

We first draw the XY curve where X represents the moles of gas per mol of solution free of gas and Y is the moles of CO₂ in the gas per mol of stream free of CO₂:

$\begin{array}{l}X(\mathrm{free}\mathrm{of}x)=\frac{x}{1-x}\hfill \\ Y(\mathrm{free}\mathrm{of}y)=\frac{y}{1-y}\hfill \\ G\prime =G(1-y)\hfill \\ G\prime (Y{}_{n+1}-Y_1)=L\prime (X_n-X_{\mathrm{o}})\hfill \end{array}$

Since y_n+1=0.1 and y₁=0.005, the corresponding Ys are as follows (Table 5E4.2):

$Y_{n+1}=0.11\mathrm{and}{Y}_1=0.005$

Table 5E4.2

Equilibrium Data

mol CO2/mol Amine	PCO2(Pa)	mol CO2/mol Inert
0.025	65.4479554	1.89705E-06
0.048	197	5.71018E-06
0.145	1204	3.48998E-05
0.276	3728	0.00010807
0.327	4815	0.000139585
0.45	8252.67269	0.000239265
0.9	26,324.9963	0.000763626
10	1,480,619.55	0.044840925
15	2,918,315.23	0.092405305
20	4,722,991.64	0.158612026

If we plot the equilibrium data for Y_n+1, we have X_max=16.5, and thus:

$\begin{array}{l}{{\displaystyle \frac{L\prime }{G\prime }}|}_{\min }=\frac{Y_{n+1}-Y_1}{X_{\max }}=0.00641\hfill \\ G\prime =G(1-y)=36\text{kmol}/s\hfill \\ L\prime /G\prime =0.0096\hfill \\ L\prime =4156\to X_{\max }=11.02\text{mol}/\text{mol}\hfill \end{array}$

We also plot the operating line (triangles) and the minimum operating line (squares) (Fig. 5E4.1).

Therefore, we need three stages. With an efficiency of 0.25, 12 trays are needed for the operation.

5.3.2.1.2 Physical absorption

For operating pressures above 50 atm, physical absorption is widely extended using different solvents that define the processes, such as Selexol (using polyethylene glycol dimethyl ether, UOP) or Rectisol (using methanol, Lurgi). In general, these processes operate at very low temperatures and high pressures. Thus, only if the gas is already at high pressure is it interesting to use them. The flowsheet is similar to the one presented for the use of alkali solutions, an absorber and a desorber column.

Currently, the Selexol process operates on a license by UOP LLC. The solvent, a mixture of dimethyl esters of polyethylene glycol, absorbs or dissolves the sour gases at 20–140 atm. The solvent is regenerated by reducing the pressure. The Selexol process can selectively separate CO₂ and H₂S so that the latter can be sent to the Claus process to obtain elementary sulfur and from it sulfuric acid; the CO₂ can be further used. Below 2 MPa the absorption capacity of the solvent is greatly reduced and alkali-based processes are more interesting.

The Rectisol process (Lurgi) uses methanol under cryogenic conditions (−40°C) and from 27.6 atm to 68.9 atm for the removal of H₂S and CO₂. It is also selective with the gases.

Fig. 5.13 shows the comparison between physical and chemical sorbents. Physical sorbents show a linear absorption capacity with the partial pressure of the gas, and for them to be competitive, the pressure must be higher than when using alkali solutions. However, the energy involved in the regeneration of the solvent is lower.

5.3.2.1.3 Pressure swing adsorption (PSA)

The operation of PSA systems for the removal of sour gases is similar to the one presented in Chapter 3, Air, where they were used to separate nitrogen and oxygen. The difference is the adsorbent bed. The operating conditions are low temperatures (25°C) and moderate pressures (4–6 atm). Typically, CO₂ is the one removed. In Fig. 5.14 the adsorption capacity for molecular sieve carbon is shown. Zeolites can also be used.

5.3.2.2 Removal of H₂S

This was briefly depicted for the reforming of natural gas. This stage takes place from 150°C to 400°C following the reaction below in an adsorbent bed or in a scrubber:

${\mathrm{H}}_2\mathrm{S}+\mathrm{ZnO}\to \mathrm{ZnS}+{\mathrm{H}}_2\mathrm{O}$

Instead of using a bed of ZnO, it is possible to use iron compounds. They operate optimally at 25–50°C. Condensation of water on the catalyst should be avoided since it reduces the contact area.

${\mathrm{Fe}}_2{\mathrm{O}}_3+{\mathrm{3H}}_2\mathrm{S}\to {\mathrm{Fe}}_2{\mathrm{S}}_3+3{\mathrm{H}}_2\mathrm{O}$

$2{\text{Fe}(\text{OH})}_3+{\mathrm{3H}}_2\mathrm{S}\to {\mathrm{Fe}}_2{\mathrm{S}}_3+6{\mathrm{H}}_2\mathrm{O}$

In the Claus process the recovered H₂S can be further processed to recover the S in it and produce H₂SO₄. This process was discovered by Carl Friedrich Claus, who obtained a patent in 1883. The process has a thermal and a catalytic stage.

In the thermal stage, the H₂S is burned to produce SO₂ that reacts with the feed of H₂S to produce sulfur, as in the following reactions:

$\begin{array}{l}2{\mathrm{H}}_2\mathrm{S}+3{\mathrm{O}}_2\to 2{\mathrm{SO}}_2+{\mathrm{2H}}_2\mathrm{O}\hfill \\ 2{\mathrm{H}}_2\mathrm{S}+{\mathrm{SO}}_2\to 3\mathrm{S}+{\mathrm{2H}}_2\mathrm{O}\hfill \\ 10{\mathrm{H}}_2\mathrm{S}+5{\mathrm{O}}_2\to 2{\mathrm{H}}_2\mathrm{S}+{\mathrm{SO}}_2+7/2{\mathrm{S}}_2+8{\mathrm{H}}_2\mathrm{O}\hfill \end{array}$

Two thirds of the H₂S are converted to elementary sulfur.

In the catalytic stage, the Al₂O₃ or titanium oxide IV is used as catalyst. The reaction progresses as follows:

$2{\mathrm{H}}_2\mathrm{S}+{\mathrm{SO}}_2\to 3\mathrm{S}+2{\mathrm{H}}_2\mathrm{O}$

The sulfur obtained can be in the form of S₆, S₇, S₈, or S₉. This stage can be subdivided into three phases: heating up, catalytic reaction, and cooling with condensation. Heating up avoids condensation of sulfur on the catalytic bed. The first catalytic stage operates at 315–330°C. For the next two stages, the bottom of the bed is at 240°C and 200°C, respectively, since the conversion is favored. Finally, the gas is cooled down to 130–150°C and the excess energy is used to produce steam. Fig. 5.15 shows the scheme of the process.

5.3.2.3 Removal of CO (ammonia synthesis)

The WGSR is used to convert CO into CO₂. A two-step method for sweet gases is employed. First the gas stream is passed over a Cr/Fe₃O₄ catalyst at 360°C, and later over a Cu/ZnO/Cr catalyst at 210°C.

The remaining CO can be separated at low temperatures or by absorption in copper solutions. In the first case, after the removal of CO₂, methane and CO are cooled down to −180°C at 40 atm, condensing both. The mixture is expanded to 2.5 atm and distilled to obtain CO from the top and methane from the bottom. The second method can use several copper aqueous solutions such as CuCl in HCl (at 300 atm, recovering the CO by expansion at 50°C), and solutions of CuCl and AlCl₃ in toluene operating at 25°C and 20 atm. The CO is recovered at 100°C and 1–4 atm.

5.3.2.4 Water removal (ammonia synthesis)

The gas mixture is cooled to 40°C, at which temperature the water condenses out and is removed.

5.3.2.5 Removal of carbon oxides: methanation (ammonia synthesis)

The remaining CO₂ and CO after carbon capture are converted to methane (methanation) using an Ni/Al₂O₃ catalyst at 325°C:

$\begin{array}{c}2\mathrm{CO}+3{\mathrm{H}}_2\to {\mathrm{CH}}_4+{\mathrm{H}}_2\mathrm{O}\\ {\mathrm{CO}}_2+4{\mathrm{H}}_2\to {\mathrm{CH}}_4+2{\mathrm{H}}_2\mathrm{O}\end{array}$

The water produced in the reactions is removed by condensation at 40°C.

5.4.1.1 Introduction

History: Ammonia was already known back in ancient times. There are up to three possible origins for the word ammonia. It is typically related to the discovery of sal ammoniac (ammonium chloride) near the Temple of Zeus Ammon, in the Siwar Oasis, Libya. The word means “sand” in Greek. It was reported that the salt was found beneath the sand. Pliny also reported the existence of another hammoniacum close to the oracle of Ammon. It was a plant secretion deposited in the sand as droplets. The last possible origin is also related to the Temple of Zeus Ammon since the priest, while burning camel dung for fuel, observed the formation of some white crystals. It was not until 1716 that ammonia was linked to rotten food. A few years later, in 1727, it was obtained from lime and ammonium chloride, as reported in the Solvay process:

$\begin{array}{l}\mathrm{CaO}+{\mathrm{H}}_2\mathrm{O}\to \mathrm{Ca}{(\mathrm{OH})}_2\hfill \\ \mathrm{Ca}{(\mathrm{OH})}_2+{\mathrm{NH}}_4\mathrm{Cl}\to {\mathrm{CaCl}}_2+{\mathrm{NH}}_4\mathrm{OH}\hfill \\ {\mathrm{NH}}_4\mathrm{OH}\leftrightarrow {\mathrm{NH}}_3+{\mathrm{H}}_2\mathrm{O}\hfill \end{array}$

Free ammonia was first prepared in 1774 by J. Priestly using ammonium carbonate, and in 1785, the chemical formula was discovered by C. Berthollet. In the 1840s ammonia was a byproduct in coke production. The current process for the production of ammonia is based on the research by Fritz Haber from 1880 to 1900 on the equilibrium of the reaction between nitrogen and hydrogen. In 1908 BASF and Haber came to an agreement to design an industrial plant. In the 5 years working with Karl Bosch, they developed the so-called Haber–Bosch process. The first plant dates back to 1913 in Ludwigshafen, with a capacity of 30 t/day, a process that displaced the use of coal and charcoal as sources for ammonia. In 1953, 143 facilities were already in operation (The HaberBosch Heritage. 1997, www.science.uva.nl).

Properties: Ammonia is a colorless gas with a characteristic odor whose density is 0.7714 g/cm³. However, it can easily be liquefied at moderate pressure. In Table 5.3 we can see some values for the vapor pressure as a function of the temperature. The heat of vaporization is 5580 cal/mol and the critical temperature is 132.9°C. It is highly soluble in water, in particular at high pressure and low temperature. At 20°C and 1 atm, a solution of 28% can be obtained; it is sold commercially. It solidifies at −78°C in the form of colorless crystals.

$\begin{array}{l}{\mathrm{NH}}_3+{\mathrm{H}}_2\mathrm{O}\to {\mathrm{NH}}_4\mathrm{OH}\hfill \\ {\mathrm{NH}}_4\mathrm{OH}\to {\mathrm{NH}}_4^{+}+{\mathrm{OH}}^{-}\hfill \end{array}$

From a health and safety point of view, a feature worth mentioning is the fact that ammonia–air mixtures may explode with a hot point. Furthermore, in the gas phase it is toxic since it increases the blood pressure if breathed in. The eyes are sensitive to it, and it burns the skin.

Production: Ammonia is naturally produced from rotten, nitrogenated organic species. For decades it was obtained from the distillation of carbonous materials. It is currently produced from the reaction of nitrogen and hydrogen. This is the process that will be analyzed below.

Usage: Ammonia is the basic raw material for the nitrogen industry. Nitrogen fertilizers are the most widely used fertilizers, which constitute 80% of the ammonia market. The increase in consumption is 18% over the several last years. Urea, caprolactama, and explosives, as well as nitric acid, are other well-known products.

Storage: Ammonia is stored as a compressed liquid in sphere tanks made of chromium–nickel stainless steel; the use of Zn or Cu would produce soluble amines.

5.4.1.2 Stage III reaction (H₂+N₂)

The reaction is shown below. We are going to evaluate its thermodynamics and kinetics as well as the design of the converter Appl (1998). Finally, we will describe different commercial processes for the production of ammonia.

$\frac{1}{2}{\mathrm{N}}_2+\frac{3}{2}{\mathrm{H}}_2\leftrightarrow {\mathrm{NH}}_3;\mathrm{\Delta }{H}_{\mathrm{r}}=-48{\text{kcal}/\text{mol}}_{{\mathrm{NH}}_3}$

5.4.1.2.1 Thermodynamics of the reaction

Ammonia production is an exothermic reaction and thus low temperatures favor conversion. However, the kinetics sets the lower bound of the operating temperature. On the other hand, there is a decrease in the number of moles as the reaction progresses. Therefore, the higher the pressure, the larger the conversion. The ideal operation would be isothermal, since high conversions could be achieved. However, removing the energy generated as the reaction advances is a technical challenge and the design of the reactors look for an optimal operating line. The typical pressures and temperatures are 100–1000 atm and 400–600°C (Vancini, 1961). In Table 5.4 we see some values of the equilibrium constant as a function of both pressure and temperature. Alternatively, in the literature we can find correlations (Eqs. 5.12 and 5.13) to compute the equilibrium constant either for atmospheric pressure, or as a function of it; the parameters b and u can be seen in Table 5.5.

${\log }_{10}(K_{\mathrm{p}})=\frac{2250.322}{T}-0.85430-1.51049{\log }_{10}T-2.58987\times {10}^{-4}T+1.48961\times {10}^{-7}{T}^2$ (5.12)

(5.12)

${\log }_{10}(K_{\mathrm{p}})=\frac{2074.8}{T}-2.4943{\log }_{10}T-b·T+1.8564\times {10}^{-7}{T}^2+uT[=]\mathrm{K}$ (5.13)

(5.13)

Table 5.4

Values for the Equilibrium Constant

	Temperature (°C)
P (atm)	325	350	400	450	500
1.0	0.0401	0.0266	0.0129	0.0060	0.0038
100			0.0137	0.0072	0.0040
300				0.0088	0.0049
600				0.0130	0.0065
1000				0.0233	0.0099

Table 5.5

Values for Parameters b and u

P (atm)	b	u
300	1.256×10−4	2.206
600	1.0856×10−4	3.059
1000	2.6833	4.473

The equilibrium constant is a pseudoconstant since it is computed with partial pressures instead of fugacities. Thus, for the stoichiometry presented, the equilibrium is as follows Reklaitis (1983):

$K_{\mathrm{p}}=\frac{P_{{\mathrm{NH}}_3}}{P_{{\mathrm{N}}_2}^{0.5}{P}_{{\mathrm{H}}_2}^{1.5}}{K}_{\mathrm{f}}=\frac{f_{{\mathrm{NH}}_3}}{f_{{\mathrm{N}}_2}^{0.5}{f}_{{\mathrm{H}}_2}^{1.5}}$ (5.14)

(5.14)

$f_{{\mathrm{NH}}_3}=u_{{\mathrm{NH}}_3}{P}_{{\mathrm{NH}}_3}=u_{{\mathrm{NH}}_3}P{x}_{{\mathrm{NH}}_3}$ (5.15)

(5.15)

$K_{\mathrm{f}}=\frac{f_{{\mathrm{NH}}_3}}{f_{{\mathrm{N}}_2}^{0.5}{f}_{{\mathrm{H}}_2}^{1.5}}=\frac{{\nu }_{{\mathrm{NH}}_3}}{{\nu }_{{\mathrm{N}}_2}^{0.5}{\nu }_{{\mathrm{H}}_2}^{1.5}}\frac{x_{{\mathrm{NH}}_3}}{x_{{\mathrm{N}}_2}^{0.5}{x}_{{\mathrm{H}}_2}^{1.5}}=K_{\mathrm{u}}{K}_{\mathrm{p}}\Rightarrow K_{\mathrm{p}}=\frac{K_{\mathrm{f}}}{K_{\mathrm{u}}}$ (5.16)

(5.16)

Example 5.5

Compute the ammonia fraction at equilibrium for an operating pressure equal to 1 atm. The reaction for the production of ammonia is given as follows:

$0.5{\mathrm{N}}_2+1{\mathrm{.5H}}_2\leftarrow \to {\mathrm{NH}}_3$

The feed enters the reaction in stoichiometric proportions. Assume that no inerts accompany the feed and that the correlation for the equilibrium constant is given by the following equation:

$\log (K_{\mathrm{p}})=\frac{2250.322}{T}-0.85340-1.51049\log (T)-2.58987\times {10}^{-4}T+1.48961\times {10}^{-7}{T}^2$

Solution

For this example, we use the following nomenclature: T=temperature, P_T=P, and the molar fraction for the species involved is detoned as x, y and z for NH₃, N₂ and H₂ respectively. Therefore:

$\begin{array}{l}{P}_{{\mathrm{NH}}_3}=x{P}_{\mathrm{T}};P_{{\mathrm{N}}_2}=y{P}_{\mathrm{T}};P_{{\mathrm{H}}_2}=z{P}_{\mathrm{T}};\hfill \\ K_{\mathrm{p}}=\frac{P_{{\mathrm{NH}}_3}}{P_{{\mathrm{N}}_2}^{0.5}·P_{{\mathrm{H}}_2}^{1.5}}=\frac{x·P_{\mathrm{T}}}{{(y·P_{\mathrm{T}})}^{0.5}{(z·P_{\mathrm{T}})}^{1.5}}=\frac{1}{P_{\mathrm{T}}}\frac{x}{y^{0.5}·z^{1.5}}\hfill \\ K_{\mathrm{p}}·P_{\mathrm{total}}=\frac{x}{y^{0.5}{z}^{1.5}}\hfill \end{array}$

$\begin{array}{l}z=3·y\hfill \\ x+y+z=1\hfill \\ x+y+3·y=1\hfill \\ x+4·y=1\Rightarrow y=\frac{1-x}{4}\hfill \\ z=3·y=\frac{3}{4}(1-x)\hfill \end{array}$

Thus:

$K_{\mathrm{p}}=\frac{x}{P_{\mathrm{T}}{\left(\frac{1-x}{4}\right)}^{0.5}{\left(\frac{3}{4}(1-x)\right)}^{1.5}}=\frac{4^2}{3^{1.5}}\frac{x}{P_{\mathrm{T}}{(1-x)}^2}=3.0792\frac{x}{P_{\mathrm{T}}{(1-x)}^2}$

We can solve the equation for x as follows:

$x=\frac{\left(2+\frac{3.0792}{K_{\mathrm{p}}·P_{\mathrm{T}}}\right)-\sqrt{\left({\left(2+\frac{3.0792}{K_{\mathrm{p}}·P_{\mathrm{T}}}\right)}^2-4\right)}}{2}$

We compute, for different temperatures, the fraction of ammonia given by the variable x (Table 5E5.1).

Table 5E5.1

Equilibrium Constant Values

P (atm)	T (K)	Kp	x
1	473	0.59515055	0.14221511
1	523	0.17715127	0.051733
1	573	0.06429944	0.02005278
1	623	0.02717032	0.00867146
1	673	0.01293599	0.00416616
1	723	0.00677637	0.00219106

Assuming that the correlation is valid for different pressures, Fig. 5E5.1 presents the ammonia fraction of the product gas as a function of the pressure and temperature. As predicted by Le Chatelier’s principle, the conversion increases with the pressure and for lower temperatures. As described above, the optimal operation in the reaction would be isothermal, since high conversion could be obtained in one stage. Adiabatic operation leads us to find the limit given by the equilibrium line. The reader can also use the values for K_p given in Table 5.4 to compute the ammonia fraction obtained.