Determination of equilibrium constants:
In order to compute the equilibrium constant, the initial concentrations and extent of the reaction should be known. However, to utilize any of the equations for finding out the equilibrium constant, we must be sure that the equilibrium has certainly been reached. There are two criteria for validating the presence of equilibrium condition. At first, the value of equilibrium constant must be similar whenever the equilibrium is approached from either side. Secondly, the value of equilibrium constant must be similar whenever the concentrations of the reacting substances are varied very much.
Once it is established that equilibrium has been reached, the subsequent step is to determine the equilibrium concentration of at least one of the reactants or products, in order to compute the extent of the reaction. This can be accomplished either via chemical analysis or through measurement of certain physical property. The complexity with chemical analysis is that the concentration will change throughout the course of the analysis. Therefore, this process can mere be used whenever the reaction can be stopped at equilibrium by several means, example: sudden cooling. Physical methods are more convenient as they don't need stopping the reaction. The physical properties generally employed for this aim are refractive index, density, electrical conductivity and light absorption.
Introduction to Le Chatelier's principle:
An action which changes the pressure, temperature or concentrations of reactants in the system at equilibrium stimulates a response which partially offsets the change as a new equilibrium condition is established. Therefore, Le Chatelier's principle defines that any change to a system at equilibrium will adjust to compensate for that change. In the year 1884 the French chemist and engineer Henry-Louis Le Chatelier stated one of the central theories of chemical equilibria, that illustrates what happens to a system when something briefly eliminates it from a state of equilibrium.
This is significant to understand that Le Chatelier's principle is only a helpful guide to recognize what happens if the conditions are changed in the reaction in dynamic equilibrium; it doesn't give reasons for the changes at the molecular level (example: timescale of change and underlying reaction method).
Le Chatelier's principle defines that if a dynamic equilibrium is disturbed via changing the conditions, the position of equilibrium shifts to counteract the change to re-establish equilibrium. Whenever a chemical reaction is at equilibrium and undergoes a change in temperature, pressure or concentration of products or reactants, the equilibrium shifts in the opposite direction to make up for the change.
Le Chatelier's Principle and the Shift in Equilibrium Position:
Le Chatelier's principle defines that whenever factors which affect equilibrium are changed, the equilibrium will shift to the latest position which tends to minimize such changes. Understanding the factors that control the equilibrium positions of a chemical reaction is very significant in chemical manufacturing. Chemists and chemical engineers in charge of production wish for to select conditions that would provide an optimum yield of the desired products. They would wish for the equilibrium to lie as far to the product side as possible. Three significant factors which affect equilibrium are: Concentration, pressure and temperature.
1) The Effect of Concentration Change:
Let's take the reaction: N2 (g) + 3 H2 (g) ↔ 2 NH3 (g);
KC = [NH]2/[N2][H2]3
Increasing the concentration of either N2 or H2 (together) after an equilibrium has been established, will make Qc < Kc, and a total forward (left to right) reaction takes place. Alike result is obtained if some of NH3 is eliminated from the mixture. On the other hand, if some of N2 or H2 is eliminated or NH3 is added to the equilibrium mixture, this will make Qc > Kc, and a total reaction in the opposite direction will take place till a new equilibrium position is established.
Via periodically adding new batches of N2 and H2 gases to the reactor and eliminating the product (NH3), the above reaction will continuously carry on in the forward direction.
2) The Effect of Changing Pressure by changing Volume:
a) For a reaction like: N2 (g) + 3 H2 (g) ↔ 2 NH3 (g),
The forward reaction would yield in a decrease in pressure due to the less number of gaseous molecules. As the reverse reaction increases the number of molecules and yield in an increase in the net pressure. If the reaction mixture is forced to a smaller volume, the pressure will raise. According to Le Chatelier's principle, the system will carry on in the forward reaction to decrease the pressure. For this kind of reactions, the forward reaction favors a high pressure condition. Whenever the reaction mixture is transferred to a bigger container, the net gas pressure will drop and a total reaction will carry on in the reverse direction to reach latest equilibrium position.
b) For a reaction like: CH4 (g) + H2O (g) ↔ CO (g) + 3H2 (g),
The products include more gaseous molecules. The equilibrium will shift in the direction of the reactants side if pressure is increased via compression of the reaction mixture. This kind of reactions favors the low pressure conditions.
c) Reactions like: CO (g) + H2O (g) ↔ CO2 (g) + H2 (g) or H2 (g) + Cl2 (g) ↔ 2HCl (g),
Has equivalent number of gaseous molecules on both the sides. Changing the overall pressure will not influence the state of equilibrium. Any changes in the pressure imposed on the system encompass equivalent effects on both the sides.
d) It will be noted that a pressure change which is not due to the volume change consists of no effect on the equilibrium of a reaction. Equilibrium is not affected by a pressure change that is due to the addition of gases not comprised in the equilibrium system. For illustration, introducing helium gas to an equilibrium mixture having H2, N2 and NH3, will not influence its equilibrium?
3) The Effect of changing Temperature on Equilibrium:
To comprehend the effect of temperature on equilibrium, we have to recognize whether a reaction is exothermic or endothermic. For illustration, the reaction:
N2 (g) + 3H2 (g) ↔ 2NH3 (g) + 92 kJ,
Is an exothermic reaction - heat is a product. Increasing the temperature signifies that heat is added to the system, and the system responses via going in the opposite direction (making the reactants), which is the direction that absorbs heat. Whenever the temperature is lowered (that is, heat is taken out), a total forward reaction will take place to produce more heat. Therefore, exothermic reactions favor the low temperature condition.
The given is an endothermic reaction:
CH4 (g) + H2O (g) + 205 kJ ↔ CO (g) + 3H2 (g),
Increasing the temperature reasons the reaction to go forward (that is, in direction which eliminates heat from the system) to decrease the stress. This kind of reactions favors a high temperature condition. Therefore, endothermic reactions favor a high temperature condition and an exothermic reaction favors low temperatures.
In chemical kinetics we find out that catalysts speed up the reaction rates, however catalysts do not influence equilibrium.
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