**2. Bibliographic review**

The possibility of obtaining carbon dioxide from gas mixtures attracts the attention of inventors and investigators toward the end of the nineteenth century, as narrated by Wellford Martin and Killeffer [1]. In 1937, the two authors turn to be among the first ones to recognize the possibility of producing CO2 from the flue gases of power plants.

In the first decades of the twentieth century, the process employing aqueous ammonia for the removal of CO2 and H2 S is used extensively for the purification of coke-oven gas. Carbon dioxide is indeed a major component that must be removed to greatly increase the heating value of that gas. During the following years, amines, specifically alkanolamines, become preferred over ammonia for few reasons [2]. First, the use of amines leads to lower issues of pipe plugging and air polluting. Second, amines are characterized by higher effectiveness in capture H2 S, which can be used as an affordable source of elemental sulfur. Ultimately, ammonia is still an expensive substance because the industrial ammonia production is still to be established. Historically, the first alkanolamine to become commercially viable is triethanolamine (TEA) in the year 1930.

Through the last few decades, the amines that reach commercial maturity for gas purification are monoethanolamine (MEA), diethanolamine (DEA), and methyldiethanol-amine (MDEA). In particular, MEA is taken frequently as the reference process for the carbon capture in post-combustion configuration. By contrast, the aqueous ammonia is reconsidered explicitly for carbon capture only quite recently by both research centers and industrial companies.

The following sections provide an overview of the information about the general absorptionregeneration scheme, the chemistry of solution, the thermodynamic equilibrium models, the kinetics investigations, and the aqueous ammonia process for carbon capture.

#### **2.1. General absorption-regeneration scheme**

end-user cost. Meanwhile, renewable sources are expected to be implemented more and more

Carbon capture is proposed as a viable way of effectively exploiting the conventional resources. It can be implemented in a pre-combustion, a post-combustion, or even an oxycombustion configuration. Among them, the post-combustion option has the large benefit of being readily applicable to the already existing power plants as well as industrial processes

The post-combustion carbon capture can be accomplished by adsorption on solid materials or by chemical absorption in liquid solutions. The chemical absorption in amine aqueous solutions has been in use for decades in a number of chemical and petrochemical areas, such as the Oil & Gas or the urea preparation. Currently, the so-called advanced amines are under research with the goal of reducing the energy demand when applied to power plants and industrial processes. As an alternative to amines, the chemical absorption in ammonia aque-

This chapter covers the chemical absorption of carbon dioxide by an aqueous solution of ammonia. The next sections will present, in sequence, an overview of a number of works retrieved from the open literature, the simulation by either an equilibrium- or a rate-based approach, the environmental as well as economic assessments and, lastly, the future develop-

The possibility of obtaining carbon dioxide from gas mixtures attracts the attention of inventors and investigators toward the end of the nineteenth century, as narrated by Wellford Martin and Killeffer [1]. In 1937, the two authors turn to be among the first ones to recognize

In the first decades of the twentieth century, the process employing aqueous ammonia for

dioxide is indeed a major component that must be removed to greatly increase the heating value of that gas. During the following years, amines, specifically alkanolamines, become preferred over ammonia for few reasons [2]. First, the use of amines leads to lower issues of pipe plugging and air polluting. Second, amines are characterized by higher effectiveness

ammonia is still an expensive substance because the industrial ammonia production is still to be established. Historically, the first alkanolamine to become commercially viable is trietha-

Through the last few decades, the amines that reach commercial maturity for gas purification are monoethanolamine (MEA), diethanolamine (DEA), and methyldiethanol-amine (MDEA). In particular, MEA is taken frequently as the reference process for the carbon capture in

from the flue gases of power plants.

S, which can be used as an affordable source of elemental sulfur. Ultimately,

S is used extensively for the purification of coke-oven gas. Carbon

diffusely to allow independence from fossils in a later future.

ous solution has received great attention during the last decade.

that are fueled by coal and natural gas.

108 Carbon Capture, Utilization and Sequestration

ments of the process itself.

**2. Bibliographic review**

the possibility of producing CO2

nolamine (TEA) in the year 1930.

and H2

the removal of CO2

in capture H2

The carbon capture by aqueous ammonia as well as aqueous amines is based on the conventional absorption-regeneration scheme, which is illustrated in **Figure 1**. In simple words, the gas to be treated flows upward through the absorber, countercurrent to the falling absorbing solution, and purified from CO<sup>2</sup> . The generated rich solution from the bottom of the absorber is heated in a heat exchanger, recovering energy from the lean solution (see subsequent text), and enters the regenerator at a point near to its top. In the regenerator, a heat source (such as steam) releases the captured CO2 , which exits from the top of the column, while the generated lean solution from the bottom. The lean solutions flow, through the mentioned heat exchanger, to the top of the absorber closing the scheme. An exhaustive description of the absorption-regeneration scheme is provided by Kohl and Nielsen [3].

#### **2.2. Chemistry of the solution**

The carbon capture by aqueous ammonia is based on the ternary system CO2 –NH3 –H2 O, which yields an electrolyte solution. At the absorber conditions, the main reactions are [4] as follows:

$$2\,\mathrm{H\_2O} \Leftrightarrow \,\mathrm{H\_3O^+} \star \mathrm{OH^-} \tag{1}$$

$$\rm{CO}\_2 + 2\, H\_2O \rightleftharpoons H\_3O^+ + HCO\_3^- \tag{2}$$

$$HCO\_3^- + H\_2O \quad \text{és} \quad CO\_3^{2-} + H\_3O^+ \tag{3}$$

$$\mathrm{NH}\_{3\mathrm{(aq)}} \star \mathrm{H}\_{2}\mathrm{O}\_{0} \begin{array}{c} \Leftrightarrow \mathrm{NH}\_{4}^{\*} \star \mathrm{OH}^{-} \end{array} \tag{4}$$

$$\mathrm{NH}\_{3\text{(aq)}} \star \mathrm{HCO}\_{3}^{-} \Leftrightarrow \mathrm{NH}\_{2}\mathrm{COO}^{-} \star \mathrm{H}\_{2}\mathrm{O}\_{0} \tag{5}$$

$$NH\_4^+ + HCO\_3^- \leftrightharpoons NH\_4HCO\_{3()}\tag{6}$$

The ternary system is explored by Burrows and Lewis as early as 1912 [5]. In a more recent work, 1982, Pawlikowski et al. [6] investigate vapor-liquid equilibria of many systems, including CO2 –NH3 –H2 O, by way of the gas-liquid chromatography for temperatures ranging from

by Darde et al. [15]. The most recent improvement of the Extended UNIQUAC model [16]

+ *H*<sup>2</sup> *O*(*l*) ⇆ *HCO*<sup>3</sup>

<sup>−</sup> ⇆ *CO*<sup>3</sup>

*CO*2(*g*) ⇆ *CO*2(*aq*) (11)

*NH*3(*g*) ⇆ *NH*3(*aq*) (12)

*H*<sup>2</sup> *O*(*g*) ⇆ *H*<sup>2</sup> *O*(*l*) (13)

<sup>2</sup><sup>−</sup> ⇆ (*NH*4)

<sup>−</sup> ⇆ (*NH*4)

**Figure 2** illustrates a comparison of the experimental data by Kurz et al. [10], indicated by hollow markers, and the computed values by way of the improved extended UNIQUAC model,

Hsu et al. [17] describe the absorption reaction kinetics of amines and ammonia solutions with carbon dioxide in flue gases. The temperature of investigation is 50°C, which is relatively high

2

2

<sup>+</sup> + *HCO*<sup>3</sup>

<sup>+</sup> + *CO*<sup>3</sup>

<sup>2</sup><sup>−</sup> + 2 *HCO*<sup>3</sup>

There are relatively few investigations on the kinetics for the ternary system NH3

<sup>+</sup> + *CO*<sup>3</sup>

indicated by lines. The agreement is generally high.

+ *HCO*<sup>3</sup>

Moreover, three vapor–liquid equilibrium relations are as follows:

+ *H*<sup>+</sup> ⇆ *NH*<sup>4</sup>

<sup>+</sup> (7)

Chemical Absorption by Aqueous Solution of Ammonia http://dx.doi.org/10.5772/intechopen.78545 111

<sup>−</sup> + *H*<sup>+</sup> (8)

<sup>2</sup><sup>−</sup> + *H*<sup>+</sup> (9)

<sup>−</sup> ⇆ *NH*<sup>2</sup> *COO*<sup>−</sup> + *H*<sup>2</sup> *O*(*l*) (10)

<sup>−</sup> ⇆ *N H*<sup>4</sup> *HCO*3(*s*) (14)

*CO*<sup>3</sup> · *H*<sup>2</sup> *O*(*s*) (16)

*CO*<sup>3</sup> · 2 *N H*<sup>4</sup> *HCO*3(*s*) (17)

–CO2 –H2 O.

<sup>+</sup> + *NH*<sup>2</sup> *COO*<sup>−</sup> ⇆ *NH*<sup>2</sup> *COONH*4(*s*) (15)

comprises a full set of equilibrium reactions. First, speciation reactions are as follows:

*NH*3(*aq*)

*HCO*<sup>3</sup>

Lastly, the four solid formations are as follows:

*NH*<sup>4</sup>

*NH*<sup>4</sup>

2 *NH*<sup>4</sup>

2 *NH*<sup>4</sup>

**2.4. Kinetics investigations**

*CO*2(*aq*)

*NH*3(*aq*)

**Figure 1.** Process flow diagram of the absorbtion-regeneration scheme for ammonia and amines [3].

100 to 150°C. Kawazuishi and Prausnitz [7] provide measurements from previous works by other scientists with the scope of calibrating the expressions of dissociation equilibrium constants and Henry's constants for temperatures in the 100–205°C interval and for total liquidphase concentrations to 10 molal. Göppert and Maurer [8] report vapor-liquid equilibrium data between 333.15 and 393.15 K at pressures up to about 7 MPa for water-rich mixtures and concentrations to about 16 molal for ammonia and 13 molal for carbon dioxide. In 1992, Pelkie et al. [9] use conductivity measurements to estimate the ammonium ion, NH4<sup>+</sup> , concentration at a temperature of 25°C and over a wide span of pressures and concentrations. The vaporliquid-solid equilibrium is considered in the work by Kurz et al. [10], which focuses on the solubility of weak electrolyte gases into the aqueous phase in the temperature range from 313 to 353 K at pressures up to 0.7 MPa. In a subsequent study, the enthalpy changes upon partial evaporation of aqueous solutions, including CO2 –NH3 –H2 O, are reported by Rumpf et al. [11] at temperatures from 313 to 393 K. Finally, speciation is measured with 13C NMR by Holmes et al. [12] at 25 and 35°C and by Mani et al. [4] at room temperature.

#### **2.3. Thermodynamic equilibrium models**

As indicated, the ternary system is an electrolyte solution. The thermodynamic model for such a complex system shall account for the electric interactions among the species, including strong and weak forces. The strong forces are described by long-range terms that represent electrostatic interactions between ions. The weak forces instead by short-range terms that represent the ion dipole interactions and the non-electrostatic interactions.

Two common equilibrium descriptions are the Electrolyte Non-Random Two Liquid (e-NRTL) model [13] and the Extended UNIQUAC model [14]. A comparison between them is proposed by Darde et al. [15]. The most recent improvement of the Extended UNIQUAC model [16] comprises a full set of equilibrium reactions. First, speciation reactions are as follows:

$$NH\_{3(aq)} + H^\* \not\Rightarrow NH\_4^\* \tag{7}$$

$$\text{CO}\_{2(aq)} + \text{H}\_2\text{O}\_{0^\circ} \leftrightharpoons \text{HCO}\_3^- + \text{H}^+\tag{8}$$

$$HCO\_3^- \not\cong CO\_3^{2-} + H^+ \tag{9}$$

$$\mathrm{NH}\_{3(aq)} \star \mathrm{HCO}\_3^- \Leftrightarrow \mathrm{NH}\_2\mathrm{COO}^- \star \mathrm{H}\_2\mathrm{O}\_{(l)}\tag{10}$$

Moreover, three vapor–liquid equilibrium relations are as follows:

$$\text{CO}\_{2(l)} \not\simeq \text{CO}\_{2(q)}\tag{11}$$

$$\text{NH}\_{3(l)} \not\cong \text{NH}\_{3(aq)}\tag{12}$$

$$H\_2\big\mathcal{O}\_{\left(\mathcal{O}\right)} \not\cong H\_2\big\mathcal{O}\_{\left(\mathcal{O}\right)}\tag{13}$$

Lastly, the four solid formations are as follows:

100 to 150°C. Kawazuishi and Prausnitz [7] provide measurements from previous works by other scientists with the scope of calibrating the expressions of dissociation equilibrium constants and Henry's constants for temperatures in the 100–205°C interval and for total liquidphase concentrations to 10 molal. Göppert and Maurer [8] report vapor-liquid equilibrium data between 333.15 and 393.15 K at pressures up to about 7 MPa for water-rich mixtures and concentrations to about 16 molal for ammonia and 13 molal for carbon dioxide. In 1992, Pelkie

at a temperature of 25°C and over a wide span of pressures and concentrations. The vaporliquid-solid equilibrium is considered in the work by Kurz et al. [10], which focuses on the solubility of weak electrolyte gases into the aqueous phase in the temperature range from 313 to 353 K at pressures up to 0.7 MPa. In a subsequent study, the enthalpy changes upon partial

at temperatures from 313 to 393 K. Finally, speciation is measured with 13C NMR by Holmes

As indicated, the ternary system is an electrolyte solution. The thermodynamic model for such a complex system shall account for the electric interactions among the species, including strong and weak forces. The strong forces are described by long-range terms that represent electrostatic interactions between ions. The weak forces instead by short-range terms that

Two common equilibrium descriptions are the Electrolyte Non-Random Two Liquid (e-NRTL) model [13] and the Extended UNIQUAC model [14]. A comparison between them is proposed

–NH3 –H2 , concentration

O, are reported by Rumpf et al. [11]

et al. [9] use conductivity measurements to estimate the ammonium ion, NH4<sup>+</sup>

**Figure 1.** Process flow diagram of the absorbtion-regeneration scheme for ammonia and amines [3].

et al. [12] at 25 and 35°C and by Mani et al. [4] at room temperature.

represent the ion dipole interactions and the non-electrostatic interactions.

evaporation of aqueous solutions, including CO2

**2.3. Thermodynamic equilibrium models**

110 Carbon Capture, Utilization and Sequestration

$$NH\_4^+ + HCO\_3^- \Leftrightarrow NH\_4HCO\_{3()}\tag{14}$$

$$\mathrm{NH\_4^\* + NH\_2COO^- \Leftrightarrow \mathrm{NH\_2COONH\_{4\odot}}}\tag{15}$$

$$2\,\mathrm{NH}\_4^+ + \mathrm{CO}\_3^{2-} \Leftrightarrow \left(\mathrm{NH}\_4\right)\_2\mathrm{CO}\_3 \cdot \mathrm{H}\_2\mathrm{O}\_{60} \tag{16}$$

$$2\,\mathrm{NH}\_4^+ + \mathrm{CO}\_3^{2-} + 2\,\mathrm{HCO}\_3^- \Leftrightarrow \,\mathrm{\{NH}\_4\}\_2\mathrm{CO}\_3 \cdot 2\,\mathrm{NH}\_4\,\mathrm{HCO}\_{30} \tag{17}$$

**Figure 2** illustrates a comparison of the experimental data by Kurz et al. [10], indicated by hollow markers, and the computed values by way of the improved extended UNIQUAC model, indicated by lines. The agreement is generally high.

#### **2.4. Kinetics investigations**

There are relatively few investigations on the kinetics for the ternary system NH3 –CO2 –H2 O. Hsu et al. [17] describe the absorption reaction kinetics of amines and ammonia solutions with carbon dioxide in flue gases. The temperature of investigation is 50°C, which is relatively high

**Figure 2.** Comparison of the computed values via the Extended UNIQUAC model [16] (indicated by Ex-UQ) of the partial pressure of CO2 and NH3 against the experimental data by Kurz et al. [10]. Left: at 313 K and 6.3 molal of NH<sup>3</sup> . Right: at 353 K and 11.8 molal of NH<sup>3</sup> .

for the carbon capture process. Similarly, Diao et al. [18] investigate the removal efficiency of the sole ammonia solution in the 25–55°C interval and regress the parameters of the rate constant for the capture reaction in the Arrhenius form.

Among the reactions that describe the system, only a subset is expected to significantly influence the kinetics of the overall process. These kinetics-affecting reactions are as follows:

$$\rm{CO\_2\star OH^-} \rightarrow \rm{HCO\_3^-} \tag{18}$$

for aqueous ammonia between 0.6 and 6 mol/L, temperature between 5 and 20°C, and the

kinetic of reaction (R20) by the stopped flow apparatus in the range of temperature between

3 and 10 mmol/L. Lastly, Jilvero et al. [23] study it by a different perspective. They implement an adsorption column in a commercial code taking the design parameters of an existing pilot

Lillia et al. [24] conduct a comparison of these four investigations and the resulting param-

**Figure 3** visualizes the trends with respect to the (reciprocal of the) temperature and against the experimental data from the investigations. Each Arrhenius law fits the data well for each work. Apparently, though, the data themselves are not in complete agreement. The results from Pinsent et al. and Wang et al. are in mutual agreement, but in disagreement with those by Puxty et al. and Jilvero et al. Noticeably, Pinsent et al. and Wang et al. measured the data at low ammonia concentrations, while Puxty et al. and Jilvero et al. at high concentrations. In short, there is likely a dependence of the kinetic parameters on the

Subsequently, Lillia et al. [25] propose an alternative kinetics based on the two-film theory [26] as represented in **Figure 4**. Their study covers the region typical for the absorption columns:

0.2 to 0.6. The study yields an Arrhenius constant with a pre-exponential factor of 1.41 × 108

The concept on what was going to be referred to as the novel ammonia-scrubbing process for the carbon capture is proposed by Bai and Yeh in 1997 based on experimental data [27].

Their work highlights the remarkable potential of high removal efficiencies, over 95%, and

complete another experimental campaign with the scope of comparing amine and ammonia scrubbing and they confirm the potential of the second over the first solvent. Experiments are conducted at room temperature in their first work and between 10 and 40°C in the later one. In 2005, Yeh et al. [29] publish the results of three-cycle absorption-regeneration tests conducted on MEA and ammonia in a batch reactor maintained at about 25°C. They also reported

an approximate estimate of energy usage that is lower than the one for MEA.

per kg of NH3

s)] and an activation energy of 60,680 [J/mol]. It has a linear dependence on the CO<sup>2</sup>

concentrations from 5 to 15%, and CO2

concentration with an exponent of 1.89.

−*E*\_\_\_\_*<sup>A</sup>*

15 and 45°C, ammonia concentration between 2.0 and 16 mmol/L, and the initial CO<sup>2</sup>

plant and they tune the kinetics parameters against the experimental data.

eters, reported in **Table 1**, for the Arrhenius equation written as follows:

−*d NH*<sup>3</sup>

loading between 0 and 0.8 molCO2/molNH3. Wang et al. [22] assess the

*dt* <sup>=</sup> *<sup>r</sup>* <sup>=</sup> *<sup>k</sup>*<sup>2</sup> <sup>∗</sup> [*NH*3][*CO*2] (3)

*RT* (4)

Chemical Absorption by Aqueous Solution of Ammonia http://dx.doi.org/10.5772/intechopen.78545

between

113

loadings from

. Shortly after, Yeh and Bai [28]

initial thin liquid film CO<sup>2</sup>

ammonia concentration.

[mol/(m<sup>3</sup>

temperatures from 15 to 35°C, NH3

concentration and a dependence on the NH3

absorption capacities, around 0.9 kg of CO2

**2.5. Aqueous ammonia process for carbon capture**

\_\_\_\_\_\_

*k*<sup>2</sup> = *A* ∗ *e*

$$\rm{HCO\_3^-} \rightarrow \rm{CO\_2} + \rm{OH^-} \tag{19}$$

$$\mathrm{NH}\_3 + \mathrm{CO}\_2 \rightarrow \mathrm{NH}\_2\mathrm{COO}^- + \mathrm{H}^\* \tag{20}$$

$$\mathrm{NH}\_{2}\mathrm{COO}^{-}\star H^{\star} \rightarrow \mathrm{NH}\_{3}\star\mathrm{CO}\_{2} \tag{21}$$

Among them, reactions (R18) and (R20) are considered to be the slowest. The first is studied by Pinsent et al. [19], while the second by five different works as discussed subsequently.

Reaction (R18) is investigated by Pinsent et al. [19] via the rapid thermal method in the range 0–40°C. The fitting yields the Arrhenius constant with a second order as follows:

$$\frac{-d\,\text{CO}\_2}{dt} = r = k\_2 \ast \, [\text{CO}\_2] \text{[OH}^-\text{]}\tag{1}$$

$$k\_2 = A \ast e^{\frac{-E\_i}{RT}} \text{ with } A = 4.32 \ast 10^{13} \frac{kmol}{(m^3 \ast s)} \text{ and } E\_A = 13249 \frac{cal}{mol} \tag{2}$$

Moreover, Pinsent et al. [20] assess reaction (R20) by the rapid thermal method in the range of ammonia concentration between 0.027 and 0.19 mol/l. By contrast, Puxty et al. [21] study it by measuring the rate of CO2 absorption into a falling thin film using a wetted wall column for aqueous ammonia between 0.6 and 6 mol/L, temperature between 5 and 20°C, and the initial thin liquid film CO<sup>2</sup> loading between 0 and 0.8 molCO2/molNH3. Wang et al. [22] assess the kinetic of reaction (R20) by the stopped flow apparatus in the range of temperature between 15 and 45°C, ammonia concentration between 2.0 and 16 mmol/L, and the initial CO<sup>2</sup> between 3 and 10 mmol/L. Lastly, Jilvero et al. [23] study it by a different perspective. They implement an adsorption column in a commercial code taking the design parameters of an existing pilot plant and they tune the kinetics parameters against the experimental data.

Lillia et al. [24] conduct a comparison of these four investigations and the resulting parameters, reported in **Table 1**, for the Arrhenius equation written as follows:

$$\frac{-d\,\mathrm{NH}\_3}{dt} = \,r = k\_2 \ast \, [\mathrm{NH}\_3] \, [\mathrm{CO}\_2] \tag{3}$$

$$k\_2 = A \ast e^{\frac{\cdot E\_s}{RT}} \tag{4}$$

**Figure 3** visualizes the trends with respect to the (reciprocal of the) temperature and against the experimental data from the investigations. Each Arrhenius law fits the data well for each work. Apparently, though, the data themselves are not in complete agreement. The results from Pinsent et al. and Wang et al. are in mutual agreement, but in disagreement with those by Puxty et al. and Jilvero et al. Noticeably, Pinsent et al. and Wang et al. measured the data at low ammonia concentrations, while Puxty et al. and Jilvero et al. at high concentrations. In short, there is likely a dependence of the kinetic parameters on the ammonia concentration.

Subsequently, Lillia et al. [25] propose an alternative kinetics based on the two-film theory [26] as represented in **Figure 4**. Their study covers the region typical for the absorption columns: temperatures from 15 to 35°C, NH3 concentrations from 5 to 15%, and CO2 loadings from 0.2 to 0.6. The study yields an Arrhenius constant with a pre-exponential factor of 1.41 × 108 [mol/(m<sup>3</sup> s)] and an activation energy of 60,680 [J/mol]. It has a linear dependence on the CO<sup>2</sup> concentration and a dependence on the NH3 concentration with an exponent of 1.89.

#### **2.5. Aqueous ammonia process for carbon capture**

for the carbon capture process. Similarly, Diao et al. [18] investigate the removal efficiency of the sole ammonia solution in the 25–55°C interval and regress the parameters of the rate

**Figure 2.** Comparison of the computed values via the Extended UNIQUAC model [16] (indicated by Ex-UQ) of the

against the experimental data by Kurz et al. [10]. Left: at 313 K and 6.3 molal of NH<sup>3</sup>

Among the reactions that describe the system, only a subset is expected to significantly influence the kinetics of the overall process. These kinetics-affecting reactions are as follows:

*NH*<sup>3</sup> + *CO*<sup>2</sup> → *NH*<sup>2</sup> *COO*<sup>−</sup> + *H*<sup>+</sup> (20)

*NH*<sup>2</sup> *COO*<sup>−</sup> + *H*<sup>+</sup> → *NH*<sup>3</sup> + *CO*<sup>2</sup> (21)

Among them, reactions (R18) and (R20) are considered to be the slowest. The first is studied by Pinsent et al. [19], while the second by five different works as discussed subsequently.

Reaction (R18) is investigated by Pinsent et al. [19] via the rapid thermal method in the range

*dt* <sup>=</sup> *<sup>r</sup>* <sup>=</sup> *<sup>k</sup>*<sup>2</sup> <sup>∗</sup> [*CO*2]

Moreover, Pinsent et al. [20] assess reaction (R20) by the rapid thermal method in the range of ammonia concentration between 0.027 and 0.19 mol/l. By contrast, Puxty et al. [21] study

0–40°C. The fitting yields the Arrhenius constant with a second order as follows:

*RT* with *<sup>A</sup>* <sup>=</sup> 4.32 <sup>∗</sup> <sup>1013</sup> \_\_\_\_\_\_ *kmol*

−*d CO*<sup>2</sup>

<sup>−</sup> (18)

.

[*OH*<sup>−</sup>] (1)

*mol* (2)

(*m*<sup>3</sup> <sup>∗</sup> *<sup>s</sup>*) and *EA* <sup>=</sup> <sup>13249</sup> \_\_\_\_ *cal*

absorption into a falling thin film using a wetted wall column

<sup>−</sup> → *CO*<sup>2</sup> + *OH*<sup>−</sup> (19)

constant for the capture reaction in the Arrhenius form.

.

*CO*<sup>2</sup> + *OH*<sup>−</sup> → *HCO*<sup>3</sup>

*HCO*<sup>3</sup>

and NH3

partial pressure of CO2

Right: at 353 K and 11.8 molal of NH<sup>3</sup>

112 Carbon Capture, Utilization and Sequestration

\_\_\_\_\_\_

−*E*\_\_\_\_*<sup>A</sup>*

*k*<sup>2</sup> = *A* ∗ *e*

it by measuring the rate of CO2

The concept on what was going to be referred to as the novel ammonia-scrubbing process for the carbon capture is proposed by Bai and Yeh in 1997 based on experimental data [27].

Their work highlights the remarkable potential of high removal efficiencies, over 95%, and absorption capacities, around 0.9 kg of CO2 per kg of NH3 . Shortly after, Yeh and Bai [28] complete another experimental campaign with the scope of comparing amine and ammonia scrubbing and they confirm the potential of the second over the first solvent. Experiments are conducted at room temperature in their first work and between 10 and 40°C in the later one. In 2005, Yeh et al. [29] publish the results of three-cycle absorption-regeneration tests conducted on MEA and ammonia in a batch reactor maintained at about 25°C. They also reported an approximate estimate of energy usage that is lower than the one for MEA.


**Table 1.** Arrhenius parameters of the rate of reaction (R20) from different experimental works.

phase are realized in a dedicated process section, while the packed absorption and desorption columns operate free of solids. A similar approach is proposed by Gao et al. [35], pursuing the decreasing energy consumption by the addition of alcohols to reinforce the crystallization. Bonalumi et al. [36] suggest to operate the process at cool conditions (20–35°C) rather than chilled (5–20°C), to minimize the load on the chillers in favor of the load on air coolers. The two processes are visualized in **Figure 5**. In the cool process, one chilling load is still present

concentration profile in the liquid phase. Right: the CO<sup>2</sup>

–NH3 –H2

Chemical Absorption by Aqueous Solution of Ammonia http://dx.doi.org/10.5772/intechopen.78545

O. Left: the CO<sup>2</sup>

partial pressure profile

partial

115

**Figure 4.** Representation of the two-film theory [26] applied to the ternary system CO2

**Figure 5.** Process flow diagram of the chilled (top) and cooled (bottom) aqueous ammonia process.

pressure profile in the gas and the CO<sup>2</sup>

in both the gas and liquid phases.

**Figure 3.** Comparison among values for k2 and experimental data from cited works: Puxty et al. [21], Wang et al. [22], and Pinsent et al. [20]. The dashed lines are obtained fitting the experimental data. The dashed line for Jilvero et al. [23] is the trend proposed by the authors of that work.

In 2006, EIG Inc. [30] applies for a patent on the chemical absorption of the carbon dioxide into aqueous ammonia at chilled conditions. The company Alstom is engaged in its intensive development, establishing the commercial name Chilled Ammonia Process (CAP). As summarized by Lombardo et al. [31], the chilled process is tested first at bench scale with SRI International. Later, it is verified at pilot scale with the Electric Power Research Institute and two utilities: WE energies in its Pleasant Prairie (WI, USA) and E.ON in its Karlshamm (Sweden) plant. Ultimately, the product validation is executed in the facility of the American Electric Power in Columbus (OH, USA) and in the world's largest test facility of the Technology Center of Mongstad (Norway). The process has evolved during the years and it is still under development by the company General Electric, which has acquired it recently [32].

At the same time, the process is investigated by a number of research centers. Ullah et al. [33] analyze the use of a capacitive deionization in the conventional scheme of the ammonia-based process to reduce the regeneration energy requirement, concluding that the reduction can be as much as 37.5%. Sutter et al. [34] propose instead the controlled solid bicarbonate formation to decrease the energy requirement. Precipitation, separation, and dissolution of the solid

**Figure 4.** Representation of the two-film theory [26] applied to the ternary system CO2 –NH3 –H2 O. Left: the CO<sup>2</sup> partial pressure profile in the gas and the CO<sup>2</sup> concentration profile in the liquid phase. Right: the CO<sup>2</sup> partial pressure profile in both the gas and liquid phases.

phase are realized in a dedicated process section, while the packed absorption and desorption columns operate free of solids. A similar approach is proposed by Gao et al. [35], pursuing the decreasing energy consumption by the addition of alcohols to reinforce the crystallization.

Bonalumi et al. [36] suggest to operate the process at cool conditions (20–35°C) rather than chilled (5–20°C), to minimize the load on the chillers in favor of the load on air coolers. The two processes are visualized in **Figure 5**. In the cool process, one chilling load is still present

**Figure 5.** Process flow diagram of the chilled (top) and cooled (bottom) aqueous ammonia process.

In 2006, EIG Inc. [30] applies for a patent on the chemical absorption of the carbon dioxide into aqueous ammonia at chilled conditions. The company Alstom is engaged in its intensive development, establishing the commercial name Chilled Ammonia Process (CAP). As summarized by Lombardo et al. [31], the chilled process is tested first at bench scale with SRI International. Later, it is verified at pilot scale with the Electric Power Research Institute and two utilities: WE energies in its Pleasant Prairie (WI, USA) and E.ON in its Karlshamm (Sweden) plant. Ultimately, the product validation is executed in the facility of the American Electric Power in Columbus (OH, USA) and in the world's largest test facility of the Technology Center of Mongstad (Norway). The process has evolved during the years and it is still under develop-

**Figure 3.** Comparison among values for k2 and experimental data from cited works: Puxty et al. [21], Wang et al. [22], and Pinsent et al. [20]. The dashed lines are obtained fitting the experimental data. The dashed line for Jilvero et al. [23]

 **in Eq. (4)**

**s)] EA [cal/mol]**

At the same time, the process is investigated by a number of research centers. Ullah et al. [33] analyze the use of a capacitive deionization in the conventional scheme of the ammonia-based process to reduce the regeneration energy requirement, concluding that the reduction can be as much as 37.5%. Sutter et al. [34] propose instead the controlled solid bicarbonate formation to decrease the energy requirement. Precipitation, separation, and dissolution of the solid

ment by the company General Electric, which has acquired it recently [32].

is the trend proposed by the authors of that work.

**Source Arrhenius parameters of** *k***<sup>2</sup>**

114 Carbon Capture, Utilization and Sequestration

**A [kmol/(m3**

Pinsent et al. [20] 1.35 × 10<sup>11</sup> 11,585 Puxty et al. [21] 1.66 × 1014 14,577 Wang et al. [22] 5.01 × 10<sup>11</sup> 12,279 Jilvero et al. [23] 6.51 × 10<sup>13</sup> 14,362

**Table 1.** Arrhenius parameters of the rate of reaction (R20) from different experimental works.

for the water wash on top of the absorber. Water wash is required indeed to minimize the tendency of ammonia to escape from the absorber, which is called ammonia slip.

common performance index, the specific heat duty [MJth/kgCO2] is defined as the ratio of the reboiler heat duty [MWth] and the mass flow rate [kgCO2/s] of effectively captured carbon dioxide. However, this second index does not include the information on the capture efficiency (first index) nor on the temperature at which the heat duty is required (or, in equivalent terms, the loss of electric power generation from the steam turbine due to the steam bled for the

A third index is adopted to solve this issue about the specific heat duty. Consequently, the new index allows to compare consistently plants characterized by different capture efficiencies, regeneration temperatures, and electric efficiency penalties. The Specific Primary Energy

*EREF* <sup>−</sup> *<sup>E</sup>* <sup>≡</sup>

where all parameters refer to either the power plant equipped with the carbon capture or

[−] the net electric efficiency, and *REF* stays for reference.

**Electric power, MWe Chilled Cooled Electric power, MWe Chilled Cooled**

AC11 2.357 2.351 CH24 0.045 0.000 AC12 0.000 0.132 FN21 3.154 3.342 CH11 4.860 0.000 PM21 1.629 1.410 CH12 1.058 0.000 PM22 2.362 1.121 FN11 3.943 4.177 PM23 <0.001 0.003 PM11 0.597 0.592 PM24 0.010 0.010 PM12 0.201 0.142 Subtotal *64.380 17.219*

Subtotal *13.148 7.394* RB21 45.131 57.207 *ABS-RGN-GW (2)* RB22 1.878 15.321 AC21 0.220 0.671 Subtotal *47.009 72.528*

AC23 0.018 1.770 AC31 0.226 0.326 AC24 0.000 0.952 AC32 0.775 0.957 AC25 0.000 0.018 CM31 6.771 15.421 CH21 36.349 2.801 CM32 6.019 14.825 CH22 20.310 0.495 PM31 1.784 0.652 CH23 0.139 0.000 Subtotal *15.575 32.181*

**Table 3.** Predicted electric consumption of the capture island for the chilled and the cooled aqueous ammonia capture

3600( \_\_1 *ηe* − \_\_\_\_ <sup>1</sup> *<sup>η</sup><sup>e</sup>*,*REF*) \_\_\_\_\_\_\_\_\_\_\_

 *Compression (3)*

**Total loss 140.112 129.323**

*EREF* <sup>−</sup> *<sup>E</sup>* (5)

Chemical Absorption by Aqueous Solution of Ammonia http://dx.doi.org/10.5772/intechopen.78545

], *E* the specific CO<sup>2</sup>

emission

117

Consumption for Carbon Avoided (*SPECCA*) [MJth/kgCO2] is defined as

the reference plant without it: *HR* is the heat rate [MJth/MWh<sup>e</sup>

PM13 0.102 0.000 Power island

AC22 0.144 4.626 *CO2*

*SPECCA* <sup>≝</sup> *HR* <sup>−</sup> *HR* \_\_\_\_\_\_\_\_\_*REF*

regenerator).

[kgCO2/MWh<sup>e</sup>

*Exhaust cooling (1)*

], *<sup>η</sup><sup>e</sup>*

computed by Bonalumi et al. [36].

The ammonia-based capture is proposed typically for existing coal- and natural gas-fired power plants. Nonetheless, Bonalumi and Giuffrida [37] consider it for an air-blown integrated gasification combined cycle (IGCC) fired with high-sulfur coal, while Pérez-Calvo et al. [38] for cement plants, both achieving promising indications.
