Reaction predicting the direction

Knowing the value of the equilibrium constant for a chemical reaction lets us judge the extent of the reaction, predict the direction of the reaction, and calculate equilibrium concentrations from any initial concentrations. Let s look at each possibility. 

The value of the equilibrium constant for a reaction makes it possible to judge the extent of reaction, predict the direction of reaction, and calculate equilibrium concentrations (or partial pressures) from initial concentrations (or partial pressures). The farther the reaction proceeds toward completion, the larger the value of Kc. The direction of a reaction not at equilibrium depends on the relative values of Kc and the reaction quotient Qc, which is defined in the same way as Kc except that the concentrations in the equilibrium constant expression are not necessarily equilibrium concentrations. If Qc Kcr net reaction goes from left to right to attain equilibrium if Qc > Kc/ net reaction goes from right to left if Qc = Kc/ the system is at equilibrium. 

For each of the following reactions, predict the direction in which the equilibrium will shift when the volume of the container is increased. 

Practice Problem A For each reaction, predict the direction of shift caused by increasing the volume of the reaction vessel. 

The sign of AG can be used to predict the direction in which a reaction moves to reach its equilibrium position. A reaction is always thermodynamically favored when enthalpy decreases and entropy increases. Substituting the inequalities AH < 0 and AS > 0 into equation 6.2 shows that AG is negative when a reaction is thermodynamically favored. When AG is positive, the reaction is unfavorable as written (although the reverse reaction is favorable). Systems at equilibrium have a AG of zero. 

All of the calculated values are in very good agreement with experiment. The theoretical calculations correctly predict the direction of each reaction only the first one is exothermic. 

Predict the direction in which reaction will occur to reach equilibrium, starting with 0.10 mol of N204 and 0.20 mol of N02 in a 2.0-L container. 

Strategy First calculate the partial pressures of N204 and N02, using the ideal gas law as applied to mixtures P, = tiiRT/V. Then calculate Q. Finally, compare Q and K to predict the direction of reaction. 

Predict the direction of a reaction, given K and the concentrations of reactants and products (Example 9.5). 

Sometimes it is important to know under what conditions a precipitate will form. For example, if we are analyzing a mixture of ions, we may want to precipitate only one type of ion to separate it from the mixture. In Section 9.5, we saw how to predict the direction in which a reaction will take place by comparing the values of J, the reaction quotient, and K, the equilibrium constant. Exactly the same techniques can be used to decide whether a precipitate is likely to form when two electrolyte solutions are mixed. In this case, the equilibrium constant is the solubility product, Ksp, and the reaction quotient is denoted Qsp. Precipitation occurs when Qsp is greater than Ksp (Fig. 11.17). 

Although thermodynamics can be used to predict the direction and extent of chemical change, it does not tell us how the reaction takes place or how fast. We have seen that some spontaneous reactions—such as the decomposition of benzene into carbon and hydrogen—do not seem to proceed at all, whereas other reactions—such as proton transfer reactions—reach equilibrium very rapidly. In this chapter, we examine the intimate details of how reactions proceed, what determines their rates, and how to control those rates. The study of the rates of chemical reactions is called chemical kinetics. When studying thermodynamics, we consider only the initial and final states of a chemical process (its origin and destination) and ignore what happens between them (the journey itself, with all its obstacles). In chemical kinetics, we are interested only in the journey—the changes that take place in the course of reactions. 

The orbital mixing theory was developed by Inagaki and Fukui [1] to predict the direction of nonequivalent orbital extension of plane-asymmetric olefins and to understand the n facial selectivity. The orbital mixing rules were successfully apphed to understand diverse chemical phenomena [2] and to design n facial selective Diels-Alder reactions [28-34], The applications to the n facial selectivities of Diels-Alder reactions are reviewed by Ishida and Inagaki elesewhere in this volume. Ohwada [26, 27, 35, 36] proposed that the orbital phase relation between the reaction sites and the groups in their environment could control the n facial selectivities and review the orbital phase environments and the selectivities elsewhere in this volume. Here, we review applications of the orbital mixing rules to the n facial selectivities of reactions other than the Diels-Alder reactions. 

So far in this chapter, you have worked with reactions that have reached equilihrium. What if a reaction has not yet reached equilihrium, however How can you predict the direction in which the reaction must proceed to reach equilihrium To do this, you substitute the concentrations of reactants and products into an expression that is identical to the equilihrium expression. Because these concentrations may not he the concentrations that the equilihrium system would have, the expression is given a different name the reaction quotient. The reaction quotient, Qc, is an expression that is identical to the equilihrium constant expression, but its value is calculated using concentrations that are not necessarily those at equilihrium. 

Various samples were analyzed. The concentrations are given in the table below. Decide whether each sample is at equilibrium. If it is not at equilibrium, predict the direction in which the reaction will proceed to establish equilibrium. 

The entropy of reaction by itself, however, is not sufficient to predict the direction of a reaction. At 25° C, you know that H2O (1) is the stable phase, not H20(g). Moreover, the second reaction... 

Thermodynamics is a powerful tool. It states that at constant temperature and pressure, the system always moves to a state of lower Gibbs free energy. Equilibrium is achieved when the lowest Gibbs free energy of the system is attained. Given an initial state, thermodynamics can predict the direction of a chemical reaction, and the maximum extent of the reaction. Macroscopically, reactions... 

The thermodynamics of these reaction systems have been investigated, resulting in methods to predict the direction of a typical reaction a priori. Furthermore, studies on kinetics, enzyme concentration, pH/temperature effects, mixing, and solvent selection have opened up new perspectives for the understanding, modeling, optimization, and possible large-scale application of such a strategy. 

Then, an absolute electrode potential was defined, Z fabs). It was established that the absolute electrode potential for the reference hydrogen electrode has a value between -4.44 V and -4.78 V, and a scale of absolute potentials lor different reactions was obtained. This was an important step because knowledge of this scale allows one to predict the direction of electron flow when two electrodes are brought into electrical contact. 

In order to determine the products of decomposition for equilibrium reactions the Kistiakowsky-Wilson or the Springall Roberts rules can be applied as a starting point. From the products of decomposition the heat and temperature of explosion can then be calculated. The temperature of explosion can then be used to calculate the products of decomposition. In practice, this process is repeated many times until there is agreement between the answers obtained. Equilibria of complex reactions and of multi-component systems are today obtained by computer however, the ability to use tabulated data is useful in predicting the direction and extent of the reaction. 

The change in free energy, AG, can be used to predict the direction of a reaction at constant temperature and pressure. Consider the reaction ... 

AG° is predictive only under standard conditions Under standard conditions, AG° can be used to predict the direction a reaction proceeds because, under these conditions, AG° is equal to AG. However, AG° cannot predict the direction of a reaction under physiologic conditions, because it is composed solely of constants (R, T, and Keq) and is, therefore, not altered by changes in product or substrate concentrations. 

Definition of enthalpy and entropy Definition of free energy Enthalpy (a measure of the change in heat content of the reactants and products) and entropy (a measure of the change in the randomness or disorder of reactants and products) determine the direction and extent to which a chemical reaction will proceed. When combined mathematically, they can be used to define a third quantity, free energy, which predicts the direction in which a reaction will spontaneously proceed. 

An equilibrium constant tells us—virtually at a glance—about the composition of the reaction mixture at equilibrium and, specifically, whether we can expect a high or low concentration of product. The constant also allows us to predict the direction in which the reaction will proceed when the reactants and products are present at arbitrary concentrations. It is important to realize, however, that an equilibrium constant tells us nothing about the rate at which equilibrium is reached. Thermodynamics, as always, is silent about rates. 

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