Guldberg Waage equilibrium constant

Within experimental error, Guldberg and Waage obtained the same value of K whatever the initial composition of the reaction mixture. This remarkable result shows that K is characteristic of the composition of the reaction mixture at equilibrium at a given temperature. It is known as the equilibrium constant for the reaction. The law of mass action summarizes this result it states that, at equilibrium, the composition of the reaction mixture can be expressed in terms of an equilibrium constant where, for any reaction between gases that can be treated as ideal,... 

It was shown by Guldberg and Waage that when solids are present in a system, their active masses may be taken as constant and included in the equilibrium constant, K. For example, in the reaction ... 

The last equation, one of the most important physicochemical equations, expresses exactly the law of mass action, formulated for the first time by Guldberg and Waage in a less exact form. The equation enables the calculation of the equilibrium composition of a reaction mixture or determination of theoretically possible yields of chemical processes starting from the known value of the equilibrium constant K which can be determined by thermodynamic methods. 

In this formula K m is the dissociation constant expressed solely by the equilibrium concentrations, according to the classical Guldberg-Waage interpretation of the law of mass action. This value is identical with the true thermodynamical dissociation constant Km in highly diluted solutions only, for which the mean activity coefficient y+w very nearly equals unity. In all other solutions K m is not a true constant, but it depends on the actual concentration and on the presence of additional electrolytes therefore, it is called the apparent dissociation constant, in contradistinction of the true dissociation constant. For concentration expressed in terms of molarity, a similar equation is valid-... 

Guldberg and Waage found that the equilibrium concentrations for every reaction system that they studied obeyed this relationship. That is, when the observed equilibrium concentrations are inserted into the equilibrium expression constructed from the law of mass action for a given reaction, the result is a constant (at a given temperature and assuming ideal behavior). Thus the value of the equilibrium constant for a given reaction system can be calculated from the measured concentrations of reactants and products present at equilibrium, a procedure illustrated in Example 6.1. 

This simple equation, which interrelates the mole fractions of the components in the true equilibrium state, is essentially an expression of Guldberg and Waage s law of mass action. The quantity K T,p) is called the equilibrium constant of the reaction considered. 

The term constant is justified, for no numerical relation whatsoever was assumed to exist between C s and ys beyond the fact that they corresponded m each reservoir to equilibrium at the same temperature Hence we obtain the Guldberg Waage expression, or the Law of Mass Action for a system obeying the gas laws... 

Another way in which heterogeneous equilibria differ from homogeneous equilibria is the manner in which the different constituents offset the equilibrium. Guldberg and Waage showed that when a solid is a component of a reversible chemical process, its active mass can be considered constant, regardless of how much of the solid is present. That is, adding more solid does not bring about a shift in the equihbrium. So the expression for the equilibrium constant need not contain any concentration terms for substances present as solids. That is, the standard state of a solid is taken as that of the solid itself, or unity. Thus, for the equilibrium... 

Our derivation of the equilibrium constant in Section 17.1 was based on kinetics. But the fundamental observation of equilibrium studies was stated many years before the principles of kinetics were developed. In 1864, two Norwegian chemists, Cato Guldberg and Peter Waage, observed that at a given temperature, a chemical system reaches a state in which a particular ratio of reactant and... 

Guldberg and Waage proposed the definition of the equilibrium constant as a certain ratio of concentrations. What relationship allows us to use a particular ratio of partial pressures (for a gaseous reaction) to express an equilibrium constant Explain. 

Equation 10.1 is an example of the law of mass action, which holds that for a reversible reaction at equilibrium and at a constant temperature, a certain ratio of reactant and product concentrations has a constant value K (called the equilibrium constant). This law was first formulated by two Norwegian chemists, Cato Guldberg and Peter Waage, in 1864. 

The kinetic law of mass action, first developed by CM Guldberg and P Waage in 1864, says that reaction rates should depend on stoichiometry in the same way that equilibrium constants do. According to this law, the initial rate of product formation, d[P]/dt for the reaction in Equation (19.11), depends on reactant concentrations ... 

We see that in determining the equilibrium the concentrations of the solids do not appear at all. This important result was first stated by Guldberg and Waage in 1867 in the form that the active mass of a solid is constant. It is true only when the solids are of unvarying composition. 

Quantitative measurements of simple and enzyme-catalyzed reaction rates were under way by the 1850s. In that year Wilhelmy derived first order equations for acid-catalyzed hydrolysis of sucrose which he could follow by the inversion of rotation of plane polarized light. Berthellot (1862) derived second-order equations for the rates of ester formation and, shortly after, Harcourt observed that rates of reaction doubled for each 10 °C rise in temperature. Guldberg and Waage (1864-67) demonstrated that the equilibrium of the reaction was affected by the concentration ) of the reacting substance(s). By 1877 Arrhenius had derived the definition of the equilbrium constant for a reaction from the rate constants of the forward and backward reactions. Ostwald in 1884 showed that sucrose and ester hydrolyses were affected by H+ concentration (pH). 

Chemical equilibria for many reactions have been studied. In each case it has been found that at equilibrium the concentrations of reactants and products remained constant. As long ago as 1864, Guldberg and Waage were working on equilibrium reactions they claimed that each chemical taking part had an active mass , which was the force that controlled the progress of a reaction. They concluded that the force was proportional to the masses of the chemicals involved. 

Though the concentrations of A, B, C, and D at equilibrium would shift as one or more of these components were added to or taken from the mixture, Guldberg and Waage found they could cling to one unchanging factor. The ratio of the product of the concentrations of the substances on one side of the double arrow to the product of the concentrations on the other side of the double arrow, at equilibrium, remains constant. 

---