Le ChStelier’s principle

Le ChStelier s Principle If an equilibrium system is stressed, the equilibrium will shift in the direction that relieves the stress. 

Even though the acyclic form of a monosaccharide may be present in only trace amounts, the equilibrium can be tipped in its favor by Le ChStelier s principle (Section 9.8). Suppose, for example, that the carbonyl group of the acyclic form reacts with a reagent, thus depleting its equilibrium concentration. The equilibrium will then shift to compensate for the loss, thus producing more of the acyclic form, which can react further. 

Combining this observation with Le ChStelier s principle gives us a handle on manipulating reactions with heat. If a reaction requires heat, adding heat will cause the reaction to shift to the product side. 

From Le ChStelier s principle we know that for exothermie reaetions, the equilibrium shifts to the left (i.e., K and deerease) as the temperature increases. Figmes C-1 and C-2 show how the equilibrimn eonstant varies with temperature for an exothermie reaction and for an endothermic reaction, respectively. 

Le ChStelier s principle states that chemical equilibria adjust to relieve applied stresses. 

Could you have predicted this result using Le ChStelier s principle Yes. Think of the increased concentration of CO as a stress on the equilibrium. The equilibrium system reacts to the stress by consuming CO at an increased rate. This response, called a shift to the right, forms more CH4 and H2O. Any increase in the concentration of a reactant results in a shift to the right and additional product. 

How could an industrial chemist regulate the temperature to increase the amount of methane in the equilibrium mixture According to Le ChStelier s principle, if heat is added, the reaction shifts in the direction in which heat is used up that is, the reaction shifts to the left. A shift to the left means a decrease in the concentration of methane because methane is a reactant in the reverse reaction. However, lowering the temperature shifts the equilibrium to the right because the forward reaction liberates heat and relieves the stress. In shifting to the right, the equilibrium produces more methane. 

The product of [H+] and [OH ] always equals 1.0 X 10 at 298 K. This means that if the concentration of H+ ion increases, the concentration of OH ion must decrease. Similarly, an increase in the concentration of OH ion causes a decrease in the concentration of H+ ion. You can think about these changes in terms of Le ChStelier s principle, which you learned about in Chapter 18. Adding extra hydrogen ions to the self-ionization of water at equilibrium is a stress on the system. The system reacts in a way to relieve the stress. The added H+ ions react with OH ions to form more water molecules. Thus, the concentration of OH ion decreases. Example Problem 19-1 shows how you can use to calculate the concentration of either the hydrogen ion or the hydroxide ion if you know the concentration of the other ion. 

It is apparent from the above discussion that Le ChStelier s principle is the dominant concept behind most chemical reactions in the real world. It is particularly important in biochemical reactions, and external factors such as temperature can have a significant effect on biological equilibria. Catalysts (enzymes) are also key players in many biological and physiological reactions, as we shall see in Chapter 22. 

Given the negative AH° , Le ChStelier s principle predicts that low temperature should favor NH3 formation, and so the answer is reasonable. 

Le ChStelier s principle predicts that when we add energy to this system at equilibrium by heating it, the shift will be in the direction that consumes energy—that is, to the left. 

Plan (a) We can use standard enthalpies of formation to calculate AfP for the reaction, (b) We can then use Le ChStelier s principle to determine what effect temperature will have on the equilibrium constant... 

Thermodynamic data on these gases are given in Appendix C. You may assume that AH" and AS do not vary with temperature. (a) At what temperature will an equilibrium mixture contain equal amounts of the two gases (b) At what temperature will an equilibrium mixture of 1 atm total pressure contain twice as much NO2 as N2O4 (c) At what temperature will an equilibrium mixture of 10 atm total pressure contain twice as much NO2 as N2O4 (d) Rationalize the results from parts (b) and (c) by using Le ChStelier s principle. [Section 15.7]... 

Competitive inhibition can be reversed by increasing the concentration of substrate and letting Le ChStelier s principle operate. This is illustrated by the following equilibria that would exist in a solution containing enzyme (E), substrate (S), and competitive inhibitor (1) ... 

Figure 15.13 Temperature and Le ChStelier s principle. In the molecular level views, only... 

Use Le ChStelier s principle to explain why the equilibrium vapor pressure of a liquid increases with increasing temperature. 

Le ChStelier s principle simply summarizes the observed behavior of equilibrium systems therefore, it is incorrect to say that a given equilibrium sNft occurs "because oT Le Ch elier s principle. 

The extent to which a weak acid ionizes depends on the initial concentration of the acid. The more dilute the solution, the greater the percentage ionization (Figure 15.4). hi qualitative terms, when an acid is diluted, the concentration of the particles in the solution is reduced. According to Le ChStelier s principle (see Section 14.5), this reduction in particle concentration (the stress) is counteracted by shifting the reaction to the side with more particles that is, the equilibrium shifts from the nonionized acid side (one particle) to the side containing ions and the conjugate base (two particles) HA H+ + A. Consequently, the concentration of particles increases in the solution. 

The common ion effect is the shift in equilibrium caused by the addition of a compound having an ion in common with the dissolved substance. The common ion effect plays an important role in determining the pH of a solution and the solubility of a slightly soluble salt (to be discussed later in this chapter). Here we will study the common ion effect as it relates to the pH of a solution. Keep in mind that despite its distinctive name, the common ion effect is simply a special case of Le ChStelier s principle. 

To reestablish equilibrium, some AgCl will precipitate out of the solution, as Le ChStelier s principle would predict, until the ion product is once again equal to Kgp. The effect of adding a common ion, then, is a decrease in the solubility of the salt (AgCl) in solution. Note that in this case [Ag ] is no longer equal to [CU] at equilibrium rather, [Ag ] > [CU]. 

When heat is applied to a system in equilibrium, the reaction that absorbs heat is favored. When the process, as written, is endothermic, the forward reaction is increased. When the reaction is exothermic, the reverse reaction is favored. In this sense heat may be treated as a reactant in endothermic reactions or as a product in exothermic reactions. Therefore temperature is analogous to concentration when applying Le ChStelier s principle to heat effects on a chemical reaction. 

An ion added to a solution already containing that ion is called a common ion. When a common ion is added to an equilibrium solution of a weak electrolyte or a slightly soluble salt, the equilibrium shifts according to Le ChStelier s principle. For example, when silver nitrate (AgNOs) is added to a saturated solution of AgCI,... 

The transport of oxygen and carbon dioxide between the lungs and tissues is a complex process that involves several reversible reactions, each of which behaves in accordance with Le ChStelier s principle. 

Equation 10.16 is a mathematical statement of Le ChateUer s principle for the change in the equilibrium constant with changes in temperature. Suppose that T2>Ti. If the reaction is endothermic AH° > 0), the right-hand side of the equation will be positive and K2 > Ki, that is, the reaction shifts toward the products in an endothermic reaction, consistent with Le ChStelier s principle. If the reaction is exothermic (AH° 0), the right-hand side of the equation will be negative and K2 1, that is, the reaction shifts toward the reactants in an exothermic reaction. 

However, the basis of Le ChStelier s principle holds for any system at equilibriuiiL... 

Now we add enough CI2 to increase its concentration by 0.075 M. Before any reaction occurs, this addition creates a new set of initial concentrations. Then the system reacts and comes to a new equilibrium position. From Le ChStelier s principle, we predict that adding more reactant will produce more product, that is, shift the equilibrium position to the right. Experiment shows that the new [PCI5] at equilibrium is 0.637 M. 

The hydronium ion concentration can have a profound effect on the solubility of an ionic compound. If the compound contains the anion of a weak add, addition of HjO (from a strong add) increases its solubility. Once again, Le ChStelier s principle explains why. An especially interesting case occurs with calcium carbonate. In a saturated solution of CaC03, we have... 

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