Equilibrium constants for gaseous reactions

The accuracy of the calculated equilibrium constants for gaseous reactions was good because of the high accuracy of experimentally determined thermodynamic data for gases. As a result, the values obtained for reactions such as NO - V2 N2 + 02, NO + CO - Vz N2 + C02, and CO+ H20 C02 + H2 agreed well with the values reported else-... 

In general, liquids have lower entropies than gases, since the molecules of gas have much more freedom and randomness. Solids, of course, have still lower entropies. Any reaction in which the reactants are all liquids and one or more of the products is a gas is therefore thermodynamically favored by the increased entropy the equilibrium constant for that reaction will be higher than it would otherwise be. Similarly, the entropy of a gaseous substance is higher than that of the same substance dissolved in a solvent. 

Both reactions involve the formation of a vapour-transporting species from four gaseous reactant molecules, but whereas the tetra-iodide of zirconium is a stable molecule, the nickel tetracarbonyl has a relatively small stability. The equilibrium constants for these reactions are derived from the following considerations ... 

The procedure of Beutier and Renon as well as the later on described method of Edwards, Maurer, Newman and Prausnitz ( 3) is an extension of an earlier work by Edwards, Newman and Prausnitz ( ). Beutier and Renon restrict their procedure to ternary systems NH3-CO2-H2O, NH3-H2S-H2O and NH3-S02 H20 but it may be expected that it is also useful for the complete multisolute system built up with these substances. The concentration range should be limited to mole fractions of water xw 0.7 a temperature range from 0 to 100 °C is recommended. Equilibrium constants for chemical reactions 1 to 9 are taken from literature (cf. Appendix II). Henry s constants are assumed to be independent of pressure numerical values were determined from solubility data of pure gaseous electrolytes in water (cf. Appendix II). The vapor phase is considered to behave like an ideal gas. The fugacity of pure water is replaced by the vapor pressure. For any molecular or ionic species i, except for water, the activity is expressed on the scale of molality m ... 

Murad and Hildenbrand (j ) studied the gaseous equilibria involving ZrF, intensities were measured 3 cal K mol above threshold over the temperature range 1665-1747 K, and the equilibrium constants for the reaction 2 Ca(g) + ZrF (g) = 2 CaF(g) + ZrF (g) were calculated. Using the reported equilibrium constants, the enthalpy changes... 

Thermodynamics views a chemical reaction as a process in which atoms flow from reactants to products. If the reaction is spontaneous and is carried out at constant T and P, thermodynamics requires that AG < 0 for the process (see Section 13.7). Consequently, G always decreases during a spontaneous chemical reaction. When a chemical reaction has reached equilibrium, AG = 0 that is, there is no further tendency for the reaction to occur in either the forward or the reverse direction. We will use the condition AG = 0 in the following three subsections to develop the mass action law and the thermodynamic equilibrium constant for gaseous, solution, and heterogeneous reactions. 

The equilibrium constant for this reaction is 1.8 X 10 at 600 K. Gaseous CH4, H2O, and CO are introduced into an evacuated container at 600 K, and their initial partial pressures (before reaction) are 1.40 atm, 2.30 atm, and 1.60 atm, respectively. Determine the partial pressure of H2(g) that will result at equilibrium. 

Write a reaction for the dehydrogenation of gaseous ethane (C2Fdg) to acetylene (C2Fd2). Calculate AG° and the equilibrium constant for this reaction at 25°C, using data from Appendix D. 

Solid carbon can react with gaseous water to form carbon monoxide gas and hydrogen gas. The equilibrium constant for the reaction at 700.0 K is p = 1.60 X 10. If a 1.55 L reaction vessel initially contains 145 torr of water at 700.0 K in contact with excess solid carhon, find the percent by mass of hydrogen gas of the gaseous reaction mixture at equilibrium. 

To proceed fiirther, to evaluate the standard free energy AG , we need infonnation (experimental or theoretical) about the particular reaction. One source of infonnation is the equilibrium constant for a chemical reaction involving gases. Previous sections have shown how the chemical potential for a species in a gaseous mixture or in a dilute solution (and the corresponding activities) can be defined and measured. Thus, if one can detennine (by some kind of analysis)... 

When one of the elements is solid, as in tire case of carbon in the calculation of the partial pressures of tire gaseous species in the reaction between methane and air, CO(g) can be used as a basic element together widr hydrogen and oxygen molecules, and thus the calculation of the final partial pressure of methane must be evaluated using the equilibrium constant for CH4 formation... 

Only one equilibrium constant for gas-phase reactions (Chapter 12), the thermodynamic constant K, often referred to as Kp. This simplifies not only the treatment of gaseous equilibrium, but also the discussion of reaction spontaneity (Chapter 17) and electrochemistry (Chapter 18). 

The equilibrium state is generated by minimizing the Gibbs free energy of the system at a given temperature and pressure. In [57], the method is described as the modified equilibrium constant approach. The reaction products are obtained from a data base that contains information on the enthalpy of formation, the heat capacity, the specific enthalpy, the specific entropy, and the specific volume of substances. The desired gaseous equation of state can be chosen. The conditions of the decomposition reaction are chosen by defining the value of a pair of variables (e.g., p and T, V and T). The requirements for input are ... 

The equilibrium constant, Kc, is calculated by substituting equilibrium concentrations into the equilibrium expression. Experimentally, this means that a reaction mixture must come to equilibrium. Then one or more properties are measured. The properties that are measured depend on the reaction. Common examples for gaseous reactions include colour, pH, and pressure. From these measurements, the concentrations of the reacting substances can be determined. Thus, you do not need to measure all the concentrations in the mixture at equilibrium. You can determine some equilibrium concentrations if you know the initial concentrations of the reactants and the concentration of one product at equilibrium. 

Write the expression for the equilibrium constant for each of the following reactions. Write the pressure of a gaseous molecule, X, as Px. 

Table 16-1 gives equilibrium constants for a number of gaseous reactions. 

When data from the HSC program were used, AHf, AGf°, and S° have been listed for reactants and products as a function of temperature. Gaseous equilibria are listed in separate tables in each section, e.g., S02 + 1/2 02 = S03 for sulfates. In many cases the equilibrium constants for the decomposition reactions can be calculated directly from data in the program. When this was not possible, other methods, using the equations given above, were used. 

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