Limestone processes, equilibria

Equations 8 and 9 can be used for values of I up to 1. M. The second term in these equations accounts for the reversal of slope of activity coefficient versus ionic strength from negative to positive as ionic strength increases. Equations 8 and 9 have been widely used in the equilibrium calculations of the lime or limestone processes. With coals of moderate chloride content and for systems without extensive sludge dewatering, the ionic strength is well below 1.0 M, and equations 8 and 9 reasonable. 

For WGS, commercial catalysts are only operated up to 550 °C and no catalysts are available for higher temperatures, because adverse equilibrium conversion makes the process impractical in the absence of a CO2 sorbent. Han and Harrison [38] have shown that, at 550 °C, dolomite and limestone have a sufficiently high WGS activity. For SMR a conventional Ni SMR catalyst is used in a 1 1 ratio with CaO [30]. Meyer et al. [32] have also used a Ni-based catalyst in combination with limestone and dolomite, and achieved CH4 conversions of 95% at 675 °C while the CH4 conversion at equilibrium was 75%. 

The greenhouse effect warms the surface of Earth. The increase of C02(g) in the environment is one of the factors that contributes to this process. Many older buildings are made of limestone, CaC03(s). Consider the following equilibrium. 

Currently, thermal reduction processes have replaced the electrolysis method. The starting material in these methods is limestone, which is calcined to produce calcium oxide. The latter is ground, mixed and compacted with aluminum, and reduced at temperatures between 1,000° to 1,200°C under vacuum. Calcium vapors formed in low yield under such thermodynamic conditions are transferred from the reactor and condensed in cool zones, thus shifting the equilibrium to allow formation of more calcium vapors. The reactions are as follows ... 

Case 2. The raindrop falls on a limestone rock and comes to equilibrium with calcite, while remaining in equilibrium with the PCO2 of the atmosphere. This process results in dissolution of the rock, and addition of calcium and alkalinity to the raindrop. The knowns are Pc02> the initial composition of the solution, and that it must be in equilibrium with calcite. In this case, as in all others, electroneutrality must apply, but mass balance will not because of CO2 exchange with the atmosphere. First the equations will be written for calcium. 

The most actively cycled reservoir of carbon is atmospheric C02 (it constitutes 0.034% of the atmosphere). Carbon dioxide dissolves readily in water and is in direct equilibrium with dissolved inorganic forms of carbon (H2C03, HCO, and CO7-, see Section 6.2.1.3). Once there, it may precipitate as solid calcium carbonate (limestone). Corals and algae encourage this reaction and build up limestone reefs in the process, but a much larger portion in the deep sea equilibrates only at the slow rate of... 

Figure 26-18 shows the carbon cycle. Carbon dioxide in the atmosphere is in equilibrium with an enormous quantity that is dissolved in oceans, lakes, and streams. Some of this dissolved CO2 was once in the form of calcium carbonate (CaC03), the main component of the shells of ancient marine animals. The shells were eventually converted into limestone, which represents a large store of carbon on Earth. When the limestone was exposed to the atmosphere by receding seas, it weathered under the action of rain and surface water, producing carbon dioxide. Some of this CO2 was released into the atmosphere. This process continues today. 

Most analytical techniques require the state of chemical equilibrium. At equilibrium, the rate of a forward process or reaction and that of the reverse process are equal. The photo at left shows the beautiful natural formation called "Frozen Niagra in Mammoth Cave National Park in Kentucky. As water seeps over the limestone surface of the cave, calcium carbonate dissolves in the water according to the chemical equilibrium... 

This reaction is responsible for formation, over thousands of years, of caves and gorges in limestone areas, as CO2-laden rainwater very slowly dissolves the rock. The process is slow because the equilibrium constant is small. This constant can be... 

Even though a natural system may not be at equilibrium overall, there may well be small parts of the system which are at or not very far from equilibrium. The system is then said to have areas of local equilibrium, and thermodynamics can be applied to these smaller parts. For example, a solution flowing through a limestone may be in the process of dissolving calcite - the system is not at equilibrium. However, the calcium and carbonate and other ions in the solution may well be at or not far from equilibrium among themselves. In other words, a portion of the solution, if removed... 

Consider a combination of processes, such as the solution mentioned above dissolving limestone in one area, then flowing to another location where it loses some of its CO2 content, and precipitates calcite. The overall process is, of course, far from equilibrium. Nevertheless, the process can be considered in a number of separate steps, each of which is not far from equilibrium. Thus a state of local undersaturation might be calculated, then the calcium and carbonate content of the solution increased slightly, or the CO2 content decreased slightly, or whatever, and the calculation repeated until some final state is achieved. The overall process is simulated as a number of equilibrium steps. This is a type of geochemical modeling known as titration, and is discussed in Chapter 8. 

Acid rain also dissolves the calcium carbonate in the marble and limestone of buildings and monuments. Ironically, the same process that destroys these structures saves lakes in limestone-rich soil. As we discussed previously, added H30 shifts the following equilibrium to the right to dissolve more limestone and form more bicarbonate ... 

The COt Acceptor Gasification Process is discussed in light of the required properties of the CaO acceptor. Equilibrium data for reactions involving the CO% and sulfur acceptance and for sulfur rejection jit the process requirements. The kinetics of the reactions are also sufficiently rapid. Phase equilibrium data in the binary systems CaO-Ca(OH)t and Ca(OH)jr-CaCOs show the presence of low melting eutectics, which establish operability limits for the process. Data were obtained in a continuous unit which duplicates process conditions which show adequate acceptor life. Physical strength of many acceptors is adequate, and life is limited by chemical deactivation. Contrary to earlier findings both limestones and dolomites are equally usable in the process. Melts in the Ca(OH)2-CaC03 system are used to reactivate spent acceptors. 

The reaction occurs naturally when acid rain-water percolates through limestone strata to form stalactites in underground caves the acid rain dissolves the rock to form soluble bicarbonate, which then decomposes back to carbonate when it deposits on the stalactite. The same process takes place in the sea, both with limestone and with other alkaline sediments washed out by rivers. This serves ultimately to hold the pH of the ocean almost constant, but the equilibrium process between the river sediment and the ocean bulk is very slow. Reaction (3.3) does, however, provide a potential route for neutralizing the acidity of the injected carbon dioxide, via the deliberate addition of appropriate quantities of alkaline mineral. Such a procedure would be most suited to point sources of carbon dioxide (e.g., power stations) that are located close to both the ocean and large deposits of the mineral. 

The success of the process depends on the fact that the solubility product of lime (AT2) is much higher than that of limestone ( Ti). Therefore, the value of K is high, and equilibrium conversions in turn can be rather high. 

Limestone (CaCOs) can be calcined in a kiln to produce solid CaO and CO2 gas. The reaction is endothermic, and so, in every piece of limestone, heat must be supplied to the reaction site. This heat must be transported through the laminar flowing boundary layer in the gas phase, and through the CaO product layer. Simultaneously, the CO2 which is produced must be transported away in the opposite direction. The overall process involves a coupled transport of heat and mass. The relationship between the CO2 partial pressure and the temperature at the reaction site can be determined from thermodynamic data (unless thermodynamic equilibrium is not achieved, in which case additional kinetic data will also be required to determine the relationship). Fuel is burned in the kiln to supply the heat necessary to maintain the reaction. 

The Solvay process [Equation (13.20) in Figure 13.10] for the production of sodium carbonate starts a) with limestone or chalk (CaC03) that is heated to produce quicklime (CaO) and carbon dioxide, b) The carbon dioxide is combined with aqueous ammonia to produce the ammonium and bicarbonate ions. Although sodium bicarbonate is fairly soluble, keeping an excess of sodium ions (from NaCl) keeps the equilibrium of (r) shifted to the right to precipitate the salt. In d ) heating the sodium bicarbonate produces the desired primary product, sodium carbonate. 

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