Thermodynamics of the Chemisorption of Carbon Dioxide and Water Vapor by Metal Oxides
Abstract Generated abstract
This paper evaluates metal oxides as potentially regenerable adsorbents for carbon dioxide and water vapor in gas drying and purification. Using standard thermodynamic data, the authors calculate Gibbs free energies, equilibrium constants, and corresponding equilibrium partial pressures for carbonate and hydroxide decomposition reactions over a range of temperatures. The results indicate that most examined oxides can chemisorb carbon dioxide or water at room temperature, while regeneration should be possible under vacuum or moderate heating for selected metals, with copper, zinc, and nickel oxides offering relatively low-temperature regeneration. The study identifies promising oxide systems but notes uncertainties for beryllium carbonate data and emphasizes that adsorption kinetics must be studied experimentally before practical conclusions can be drawn.
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PROCEEDINGS
of the Tomsk Polytechnic Institute named after S. M. Kirov, Order of the October Revolution and Order of the Red Banner of Labor
Volume 257 1973
THERMODYNAMICS OF CHEMISORPTION OF CARBON DIOXIDE
AND WATER VAPOR BY METAL OXIDES
N. F. Stas, G. G. Savel’ev, V. V. Nakhalov,
A. S. Guzenberg, A. M. Ryabkin
(Presented by the scientific seminar of the Department
of General and Inorganic Chemistry)
In recent years, adsorption methods for drying gases and purifying them from undesirable impurities have become widespread in the chemical industry. In this connection, the search for new adsorbents is of current importance, especially regenerable ones that can be used repeatedly.
Oxides of alkali and of some alkaline-earth metals have long been used for absorbing carbon dioxide and moisture; however, they are single-use adsorbents, since the carbonates and hydroxides formed during adsorption decompose only at high temperatures: the carbonates of lithium, sodium, potassium, barium, and strontium—above \(1000^\circ\mathrm{C}\), and calcium carbonate—at about \(900^\circ\mathrm{C}\).
At the same time, there are metals whose carbonates and hydroxides decompose at relatively low temperatures; therefore, sorbents based on the oxides of these metals can be easily regenerated. In this connection, the study of the adsorption properties of oxides is of both scientific and practical importance.
For a preliminary evaluation of the prospects of oxides as regenerable adsorbents of carbon dioxide and water vapor, thermodynamic calculations were carried out for the equilibrium of the metal oxide–carbon dioxide and metal oxide–water systems, and for the dependence of the position of equilibrium on external conditions: temperature and pressure.
The isobaric-isothermal potentials and equilibrium constants of reversible reactions
\[ \mathrm{Me}_m(\mathrm{CO}_3)_n \rightleftharpoons \mathrm{Me}_m\mathrm{O}_n + n\mathrm{CO}_2, \tag{1} \]
\[ mMe(OH)_n \rightleftharpoons Me_mO_n + nH_2O . \tag{2} \]
at various temperatures were calculated by formulas (1):
\[ \Delta G_T^{0}=\Delta H_{298}^{0}-T\Delta S_{298}^{0} -\Delta C_{p\,298}^{0}T\left(\ln \frac{T}{298}+\frac{298}{T}-1\right), \tag{3} \]
\[ \lg K_p=\Delta G_T^{0}/4.573T, \tag{4} \]
where \(\Delta H_{298}^{0}\), \(\Delta S_{298}^{0}\), \(\Delta C_{p\,298}^{0}\) are the heat effect (cal/mol), the entropy change (cal/mol·deg), and the heat-capacity change (cal/mol·deg) when the reaction proceeds under standard conditions; \(T\) is the temperature, °K. The values of \(K_p\) obtained in this way from formula (4) are numerically equal to the partial pressure of carbon dioxide and water vapor in the corresponding system at the given temperature. In the calculations, standard values of thermodynamic functions taken from handbooks [2,3] were used.
The sequence and interpretation of the calculation results can be shown using the example of the system
\[ Ag_2CO_3 \rightleftharpoons Ag_2O + CO_2 . \tag{5} \]
According to Hess’s law, the changes in enthalpy, entropy, and heat capacity when this reaction proceeds under standard conditions are, respectively, 19.53 kcal/mol, 40.15 cal/mol·deg, and \(-2.26\) cal/mol·deg. The change in the isobaric-isothermal potential at \(25^\circ C\) is 7559 cal/mol, and the equilibrium constant is \(2.9\cdot10^{-6}\).
The positive value of \(\Delta G^{0}\) means that, at a temperature of \(25^\circ C\), this reaction can proceed from right to left and that, at this temperature, absorption of carbon dioxide by silver oxide is possible. At a stoichiometric ratio of the components, chemisorption is possible down to a residual pressure of carbon dioxide of \(2.9\cdot10^{-6}\) atm (\(2.2\cdot10^{-3}\) mm Hg).
At \(100^\circ C\), the equilibrium pressure of \(CO_2\) in this system is 1.6 mm Hg. Consequently, at this temperature silver oxide can absorb carbon dioxide if its partial pressure in the gas being purified is greater than 1.6 mm Hg; conversely, silver carbonate will decompose if the system is evacuated to a carbon dioxide partial pressure lower than this value.
At a temperature of \(215^\circ C\), the equilibrium pressure of carbon dioxide is 760 mm Hg; consequently, at this temperature complete decomposition of silver carbonate is possible.
The results of calculations of the equilibrium of the reactions of carbon dioxide and water with the oxides of twelve metals differing in their properties are summarized in the table.
Table
Partial pressure of carbon dioxide and water vapor in the decomposition reactions of carbonates and hydroxides of certain metals
| Reaction | 25 | 100 | 200 | 300 |
|---|---|---|---|---|
| $\mathrm{CuCO_3 \rightleftharpoons CuO + CO_2}$ | $3{,}8 \cdot 10^{-2}$ | $4{,}23$ | $221$ | $2915$ |
| $\mathrm{Ag_2CO_3 \rightleftharpoons Ag_2O + CO_2}$ | $2{,}9 \cdot 10^{-6}$ | $2{,}1 \cdot 10^{-3}$ | $0{,}51$ | $17{,}5$ |
| $\mathrm{BeCO_3 \rightleftharpoons BeO + CO_2}$ | $1{,}9 \cdot 10^{4}$ | $1{,}3 \cdot 10^{5}$ | $6{,}2 \cdot 10^{5}$ | $1{,}2 \cdot 10^{6}$ |
| $\mathrm{MgCO_3 \rightleftharpoons MgO + CO_2}$ | $3{,}2 \cdot 10^{-9}$ | $1{,}1 \cdot 10^{-5}$ | $1{,}1 \cdot 10^{-2}$ | $0{,}99$ |
| $\mathrm{ZnCO_3 \rightleftharpoons ZnO + CO_2}$ | $5{,}3 \cdot 10^{-4}$ | $0{,}16$ | $20{,}4$ | $477$ |
| $\mathrm{CdCO_3 \rightleftharpoons CdO + CO_2}$ | $1{,}0 \cdot 10^{-9}$ | $4{,}2 \cdot 10^{-6}$ | $4{,}5 \cdot 10^{-3}$ | $0{,}43$ |
| $\mathrm{HgCO_3 \rightleftharpoons HgO + CO_2}$ | $1{,}3 \cdot 10^{-7}$ | $3{,}4 \cdot 10^{-5}$ | $3{,}8 \cdot 10^{-3}$ | $8{,}2 \cdot 10^{-2}$ |
| $\mathrm{MnCO_3 \rightleftharpoons MnO + CO_2}$ | $3{,}4 \cdot 10^{-11}$ | $4{,}0 \cdot 10^{-7}$ | $1{,}1 \cdot 10^{-3}$ | $0{,}18$ |
| $\mathrm{FeCO_3 \rightleftharpoons FeO + CO_2}$ | $5{,}9 \cdot 10^{-8}$ | $1{,}2 \cdot 10^{-4}$ | $7{,}1 \cdot 10^{-2}$ | $4{,}6$ |
| $\mathrm{CoCO_3 \rightleftharpoons CoO + CO_2}$ | $2{,}2 \cdot 10^{-7}$ | $3{,}3 \cdot 10^{-4}$ | $0{,}15$ | $8{,}4$ |
| $\mathrm{NiCO_3 \rightleftharpoons NiO + CO_2}$ | $1{,}1 \cdot 10^{-5}$ | $5{,}6 \cdot 10^{-3}$ | $1{,}08$ | $33{,}1$ |
| $\mathrm{PbCO_3 \rightleftharpoons PbO + CO_2}$ | $2{,}0 \cdot 10^{-8}$ | $2{,}7 \cdot 10^{-5}$ | $1{,}1 \cdot 10^{-2}$ | $0{,}57$ |
| $\mathrm{Mg(OH)_2 \rightleftharpoons MgO + H_2O}$ | $6{,}6 \cdot 10^{-7}$ | $4{,}7 \cdot 10^{-4}$ | $0{,}12$ | $4{,}29$ |
| $\mathrm{Cu(OH)_2 \rightleftharpoons CuO + H_2O}$ | $7{,}0 \cdot 10^{-4}$ | $0{,}12$ | $8{,}95$ | $167$ |
| $\mathrm{Be(OH)_2 \rightleftharpoons BeO + H_2O}$ | $2{,}2 \cdot 10^{-2}$ | $1{,}70$ | $64{,}6$ | $692$ |
| $\mathrm{Zn(OH)_2 \rightleftharpoons ZnO + H_2O}$ | $3{,}3 \cdot 10^{-2}$ | $2{,}33$ | $83{,}1$ | $871$ |
| $\mathrm{Mn(OH)_2 \rightleftharpoons MnO + H_2O}$ | $3{,}3 \cdot 10^{-5}$ | $1{,}5 \cdot 10^{-2}$ | $2{,}55$ | $71{,}9$ |
| $\mathrm{Fe(OH)_2 \rightleftharpoons FeO + H_2O}$ | $4{,}7 \cdot 10^{-4}$ | $9{,}1 \cdot 10^{-2}$ | $7{,}53$ | $134$ |
| $\mathrm{Co(OH)_2 \rightleftharpoons CoO + H_2O}$ | $1{,}0 \cdot 10^{-2}$ | $1{,}39$ | $87{,}5$ | $1300$ |
| $\mathrm{Ni(OH)_2 \rightleftharpoons NiO + H_2O}$ | $8{,}8 \cdot 10^{-3}$ | $0{,}81$ | $35{,}7$ | $573$ |
Partial pressure (atm) at temperature (°C).
It is evident from the table that at a temperature of 25°C the equilibrium of all reactions, with the exception of the reaction between beryllium oxide and carbon dioxide, is shifted toward the absorption of $\mathrm{CO_2}$ and $\mathrm{H_2O}$ by metal oxides with the formation of the corresponding carbonates and hydroxides. At this temperature, thermodynamically the most probable is the absorption of $\mathrm{CO_2}$ by the oxides of manganese, cadmium, magnesium, lead, and iron, and of water vapor by the oxides of magnesium and manganese. Thermal decomposition of car-
...carbonates and hydroxides of the indicated metals—the regeneration of the adsorbents is possible at \(150{-}250^\circ\) in a vacuum of \(10^{-4} - 10^{-5}\) atm.
Copper, zinc, and nickel oxides have a lower chemical affinity for carbon dioxide and water. However, these oxides too can be used at room temperature for drying and purifying gases; their advantage is that regeneration is thermodynamically possible at \(88{-}120^\circ\) under atmospheric pressure, and at still lower temperatures (\(30{-}70^\circ\)) when evacuated to \(10^{-2}\) atm.
As for beryllium oxide, the results of the calculation for the BeO–\(\mathrm{CO_2}\) system should be treated critically. The calculations imply that formation of beryllium carbonate at \(25^\circ\) is thermodynamically impossible; however, the existence of this compound is known from the literature, and methods for its synthesis are available [4, 5]. It is possible that the standard thermodynamic constants of beryllium carbonate that we found in the only handbook [2] were determined inaccurately. On the other hand, the possibility cannot be excluded that the cited methods, which include a number of operations in aqueous solutions, lead to the formation not of the carbonate but of the basic carbonic-acid salt.
Final conclusions regarding the use of metal oxides as regenerable adsorbents of \(\mathrm{CO_2}\) and \(\mathrm{H_2O}\) can be formulated after experimental studies of adsorption kinetics.
References
-
A. N. Krestovnikov, V. N. Vigdorovich. Chemical Thermodynamics. Metallurgizdat, Moscow, 1962.
-
G. B. Naumov, B. N. Ryzhenko, I. L. Khodakovsky. Handbook of Thermodynamic Quantities. Atomizdat, Moscow, 1971.
-
M. Kh. Karapet’yants, M. L. Karapet’yants. Basic Thermodynamic Constants of Inorganic and Organic Substances. “Khimiya,” Moscow, 1968.
-
S. D. Shargorodsky, O. I. Shor. Ukr. Khim. Zh., 20, 357, 1954.
-
Handbook of Preparative Inorganic Chemistry, ed. G. Brauer. IL, Moscow, 1956.