Of Le Chatelier

When esterification is the objective water is removed from the reaction mixture to encourage ester formation When ester hydrolysis is the objective the reaction is carried out m the presence of a generous excess of water Both reactions illustrate the applica tion of Le Chatelier s principle (Section 6 10) to organic synthesis... 

The principle of Le Chatelier shows that when the pressure applied to a gaseous system is increased, dre equilibrium composition will chairge in order to reduce tire number of gaseous molecules. In the case of tire steam reforming of metlrane, the partial pressures of methane and steam will increase as the pressure is increased. In the water-gas reaction, where tire number of molecules is the same on both sides of the equation, the effect of increasing... 

Both LFL and UFL valnes for mixtnres can be estimated by nse of the Le Chatelier eqnation (Growl and Lonvar 1990). However, the methods have some limitations with respect to calcnlating the UFL for certain mixtnres. Britton (1996) determined that the eqnation does not apply to the UFL of mixtnres containing decomposable components snch as ethylene oxide or to mixtnres containing ethyl ether. Mashnga and Growl (2000) discnss the derivation of Le Chatelier s mixing rnle for flammable limits. 

Knowledge of chemical principles pays rewards in technological progress. Control of chemical reactions is the key. The large scale commercial production of nitrogen compounds provides a practical example of the beneficial application of Le Chatelier s Principle. 

Can we predict the optimum conditions for a high yield of NH3 Should the system be allowed to attain equilibrium at a low or a high temperature Application of Le Chatelier s Principle suggests that the lower the temperature the more the equilibrium state will favor the production of NHS. Should we use a low or a high pressure The production of NH3 represents a decrease in total moles present from 4 to 2. Again Le Chatelier s Principle suggests use of pressure to increase concentration. But what about practicality At low temperatures reaction rates are slow. Therefore a compromise is necessary. Low temperature is required for a desirable equilibrium state and high temperature is necessary for a satisfactory rate. The compromise used industrially involves an intermediate temperature around 500°C and even then the success of the... 

In our view, such students difficulties as described above originate, mainly, from a failure to recognise and teach chemical eqniUbrinm as a process. These problems are the outcome of a quantitative approach being taken to the theme to the detriment of understanding chemical eqnilibrinm qnalitatively. As a conseqnence, students have difficulty understanding these and other issnes in terms of Le Chatelier s Principle, which demands knowledge abont how the processes occnr. 

The buffer solution works on the basis of Le Chatelier s principle. Consider the equation for the reaction of acetic acid with water ... 

The effect of external pressure on the rates of liquid phase reactions is normally quite small and, unless one goes to pressures of several hundred atmospheres, the effect is difficult to observe. In terms of the transition state approach to reactions in solution, the equilibrium existing between reactants and activated complexes may be analyzed in terms of Le Chatelier s principle or other theorems of moderation. The concentration of activated complex species (and hence the reaction rate) will be increased by an increase in hydrostatic pressure if the volume of the activated complex is less than the sum of the volumes of the reactant molecules. The rate of reaction will be decreased by an increase in external pressure if the volume of the activated complex molecules is greater than the sum of the volumes of the reactant molecules. For a decrease in external pressure, the opposite would be true. In most cases the rates of liquid phase reactions are enhanced by increased pressure, but there are also many cases where the converse situation prevails. 

Although it is not an explanation, Le Chatelier s principle is used to predict the effect of changes in conditions on the position of equilibrium. One statement of Le Chatelier s principle is If a system in equilibrium is subjected to a change which disturbs the equilibrium, the system responds in such a way as to counteract the effect of the change . The factors that may change the position of an equilibrium are concentration, temperature and pressure. 

Proper usage of Le Chatelier s rule requires flammability limit data at the same temperature and pressure. Also, flammability data reported in the literature may be from disparate sources, with wide variability in the data. Combining data from these disparate sources may cause unsatisfactory results, which may not be obvious to the user. 

C. V. Mashuga and D. A. Crowl, Derivation of Le Chatelier s Mixing Rule for Flammable Limits, Process Safety Progress, (2000), 19(2) 112-117. 

Each species within a buffer solution participates in an equilibrium reaction, as characterized by an equilibrium constant K. Adding an acid (or base) to a buffer solution causes the equilibrium to shift, thereby preventing the number of protons from changing, itself preventing changes in the pH. The change in the reaction s position of equilibrium is another manifestation of Le Chatelier s principle (see p. 166). 

Suppose we add a solution of Na2S04 to this equilibrium system. The additional sulfate ion will disrupt the equilibrium by Le Chatclier s principle and shift it to the left. This decreases the solubility. The same would be true if you tried to dissolve PbS04 in a solution of Na2S04 instead of pure water—the solubility would be less. This application of Le Chatelier s principle to equilibrium systems of a slightly soluble salt is the common-ion effect. 

At a given temperature, a reaction will reach equilibrium with the production of a certain amount of product. If the equilibrium constant is small, that means that not much product will be formed. But is there anything that can be done to produce more Yes, there is— through the application of Le Chatelier s principle. Le Chatelier, a French scientist, discovered that if a chemical system at equilibrium is stressed (disturbed) it will reestablish equilibrium by shifting the reactions involved. This means that the amounts of the reactants and products will change, but the final ratio will remain the same. The equilibrium may be stressed in a number of ways changes in concentration, pressure, and temperature. Many times the use of a catalyst is mentioned. However, a catalyst will have no effect on the equilibrium amounts, because it affects both the forward and reverse reactions equally. It will, however, cause the reaction to reach equilibrium faster. 

Your Chemistry 12 Electronic Learning Partner can help you reinforce your understanding of Le Chatelier s principle. 

Draw diagrams to show the above examples of Le Chatelier s principle in human physiology, ecology, and economics. Show how different conditions affect the equilibrium and how the systems react to establish a new equilibrium. 

In this section, you determined the solubility product constant, Kgp, based on solubility data. You obtained your own solubility data and used these data to calculate a value for Kgp. You determined the molar solubility of ionic solutions in pure water and in solutions of common ions, based on their Ksp values. In section 9.3, you will further explore the implications of Le Chatelier s principle. You will use a reaction quotient, Qsp, to predict whether a precipitate forms. As well, you will learn how selective precipitation can be used to identify ions in solution. 

How is the common ion effect an application of Le Chatelier s principle Illustrate your answer with an example, including diagrams and chemical equations. 

The inflammability of gas mixtures were investigated by early workers, such as Davy, Bunsen and particularly the French School of Le Chatelier and M. Berthelot. 

From (8.34c) or (8.35), it is easy to see that if the chosen reaction is endothermic (AH° > 0), then a T increase tends to promote product formation (the reaction shifts right ). Conversely, if the reaction is exothermic (AH° < 0), a temperature increase promotes formation of reactants (the equilibrium shifts left ). Such conclusions appear intuitive from the perspective of Le Chatelier s principle, and indeed we shall show in Section 8.6 that such Le Chatelier-like conclusions arise from deep theoretical roots that permeate the Van t Hoff equation and many other thermodynamic relationships. 

The shift of Keq with P change can therefore be predicted from the sign of AV° if reaction volume / creases (A gas > 0), then a P increase shifts Keq toward the reactant side, whereas if reaction volume decreases (A gas < 0), a P increase will promote product formation. These inferences are consistent with the expectations of Le Chatelier s principle. 

The underlying idea is the restorative tendency of equilibrium, tending to counteract the effects of attempted changes on an original equilibrium system. This restorative tendency is associated with the stability of chemical equilibrium, and we therefore use the rigorous stability condition (8.13) to prove the above statement of Le Chatelier s principle in a general form. 

To see how (8.43) corresponds to the usual statement of Le Chatelier s principle, let us consider the specific case of a pressure change (Y = P) under isothermal conditions. 

According to the principle of Le Chatelier. the reaction should go back to the left to partially offset the increase in dichromate, which appears on the right side of Reaction 6-7. We can verify this algebraically by setting up the reaction quotient, Q. which has the same form as the equilibrium constant. The only difference is that Q is evaluated with whatever concentrations happen to exist, even though the solution is not at equilibrium. When the system reaches equilibrium, Q = K. For Reaction 6-7. 

These statements can be understood in terms of Le Chatelier s principle as follows. Consider an endothermic reaction ... 

This application of Le Chatelier s principle is called the common ion effect. A salt will be less soluble if one of its constituent ions is already present in the solution. 

To understand why this should be so, look at the Ka and Kb reactions in terms of Le Chatelier s principle. Consider an acid with pKa = 4.00 and its conjugate base with p/fb = 10.00. Let s calculate the fraction of acid that dissociates in a 0.10 M solution of HA. 

The solubility of most ionic compounds increases with temperature, despite the fact that the standard heat of solution (AH°) is negative for about half of them. Discussions of this seeming contradiction can be found in G. M. Bodner, On the Misuse of Le Chatelier s Principle for the Prediction of the Temperature Dependence of the Solubility of Salts, J. Chem. Ed. 1980,57, 117, and R. S. Treptow, Le Chatelier s Principle Applied to the Temperature Dependence of Solubility, J. Chem. Ed. 1984,61, 499. 

In Chapter 16, we apply the fundamental general equilibrium expression to gaseous equilibrium reactions. In this chapter, we apply the same expression to the equilibria that involve weak acids and bases in aqueous solution, the principal difference being that all concentrations are expressed in moles/liter (rather than in atmospheres as for gases). All the general conclusions given in Chapter 16, and summarized in the principle of Le Chatelier, apply to equilibria in solutions as well as to those in gases. 

---