Principle LeChatelier

From these two reactions it can be gathered that at higher temperatures less CH4 and more CO will be present in the equilibrium gas. However, by LeChatelier principle, increasing the pressure will increase the methane equilibrium content. Reaction 2.3 represents the steam reforming of higher hydrocarbons, which are present in small quantities in natural gas. 

If a reaction is at equilibrium and we alter the conditions so as to create a new equilibrium state, then the composition of the system will tend to change until that new equilibrium state is attained. (We say tend to change because if the reaction is kinetically inhibited, the change may be too slow to observe or it may never take place.) In 1884, the French chemical engineer and teacher Henri LeChatelier (1850-1936) showed that in every such case, the new equilibrium state is one that partially reduces the effect of the change that brought it about. This law is known to every Chemistry student as the LeChatelier principle. His original formulation was somewhat more complicated, but a reasonably useful paraphrase of it reads as follows ... 

LeChatelier principle If a system at equilibrium is subjected to a change of pressure, temperature, or the number of moles of a substance, there will be a tendency for a net reaction in the direction that tends to reduce the effect of this change. 

Consider an arbitrary mixture of these components at equilibrium, and assume that we inject more hydrogen gas into the container. Because the H2 concentration now exceeds its new equilibrium value, the system is no longer in its equilibrium state, so a net reaction now ensues as the system moves to the new state. The LeChatelier principle states that the net reaction will be in a direction that tends to reduce the effect of the added H2. This can occur if some of the H2 is consumed by reacting with I2 to form more HI in other words, a net reaction occurs in the reverse direction. Chemists usually simply say that the equilibrium shifts to the left . 

Virtually all chemical reactions are accompanied by the liberation or uptake of heat. If we regard heat as a reactant or product in an endothermic or exothermic reaction respectively, we can use the LeChatelier principle to predict the direction in which an increase or decrease in temperature will shift the equilibrium state. Thus for the oxidation of nitrogen, an endothermic process, we can write... 

Suppose this reaction is at equilibrium at some temperature T1 and we raise the temperature to T%. The LeChatelier principle tells us that a net reaction will occur in the direction that will partially counteract this change, meaning that the system must absorb some of this additional heat, and the equilibrium will shift to the right. 

The LeChatelier principle tells us that in order to maximize the amount of product in the reaction mixture, it should be carried out at high pressure and low temperature. However, the lower the temperature, the slower the reaction (this is true of virtually all chemical reactions.) As long as the choice had to be made between a low yield of ammonia quickly or a high yield over a long period of time, this reaction was infeasible economically. 

The value of Q in relation to K serves as an index how the composition of the reaction system compares to that of the equilibrium state, and thus it indicates the direction in which any net reaction must proceed. For example, if we combine the two reactants A and B at concentrations of 1 mol L-1 each, the value of Q will indeterminately large(l/0). If instead our mixture consists only of the two products C and D, Q = 0-s-l = 0. It is easy to see (by simply application of the LeChatelier principle) that the ratio of Q/K immediately tells us whether, and in which direction, a net reaction will occur as the system moves toward its equilibrium state ... 

To see if you really understand this, try explaining to yourself how the LeChatelier Principle as it applies to concentrations of reaction components follows from the idea of opposing reaction steps. 

Comment This net reaction describes the dissolution of limestone by acid it is responsible for the eroding effect of acid rain on buildings and statues. This is an example of a reaction that has practically no tendency to take place by itself (the dissolution of calcium carbonate) begin driven by a second reaction having a large equilibrium constant. From the standpoint of the LeChatelier principle, the first reaction is pulled to the right by the removal of carbonate by the hydrogen ion. Coupled reactions of this type are widely encountered in all areas of chemistry, and especially in biochemistry, in which a dozen or so reactions may be linked in this way. 

In the section that introduced the LeChatelier principle, it was mentioned that diluting a weak acid such as acetic acid CH3COOH ( HAc ) will shift the dissociation equilibrium to the right ... 

If the reaction is exothermic, AH° is negative, and the equilibrium constant decreases with increase in temperature. If the reaction is endothermic, AH° is positive then Kp increases with increase in temperature. Since an increase in the equilibrium constant implies an increase in the yield of products, Eq. (11.57) is the mathematical expression of one aspect of the LeChatelier principle. 

It must be noted here that there are certain types of systems that do not obey the LeChatelier principle in all circumstances (for example, open systems). A very general... 

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