Chemical reaction equilibria in chemical industry
Introduction to Chemical Equilibrium
Chemical equilibrium is a state in the course of a chemical reaction where the concentrations of both products and reactants reach the limit which prevents further process. Le-Chatelier’s principle is one of the important idea to understand the behavior of a system in equilibrium.
it states that “whenever a system at equilibrium is disturbed the position of equilibrium shifts in such a direction so as to minimize or nullify the change”.
Another important idea which leads to further development on this science is the usage of Gibb’s free energy for quantifying the equilibrium. It is stated as “Equilibrium is attained when the Gibbs free energy of the system is at its minimum value (assuming the reaction is carried out at constant pressure &temperature)”. The equilibrium constant for a reaction finds its relation to the Gibb’s free energy as:
Where,
G is the Gibbs free energy
R is the universal gas constant &
T is the temperature.
Le-Chatelier’s principle of equilibrium is used in the industrial applications the reaction scheme involves different parameters like temperature, pressure, concentration of reaction species. if you done a change in any single parameter results in the change of equilibrium lead to undesired product formation. The principle can be used to understand reaction conditions that will favour in increasing product formation. This idea was discovered by Le Chatelier and Karl Ferdinand Braun. The changes done in the parameter will helps in achieving the desired product. Since the high quantity product is more desirable in industrial applications for economic as well as profitability, unless the system yields the loss is effected due to undesired product. A mistake done by the operator for the process will result in costly loss to the process industry means it will effect on their economics. The concentration of reactant and product plays a important role in equilibrium; for example, if the reactant concentration is higher it will lead to forward reaction and similarly a higher concentration of product leads to backward reaction.
Factors that Affect Chemical Equilibrium
Changes in Concentration
According to Le Chatelier’s principle, adding extra quantity of reactant to a system will shift the equilibrium to the right, towards the side of the products stream. & after reducing the concentration of any quantity from product will also shift equilibrium to the right.
The converse is also true. If we add some extra amount of product to a system, the equilibrium will shift to the left side, in order to produce more reactants. Or, if we remove reactants from the system then the equilibrium will also be shifted to the left.
Thus, according to Le Chatelier’s principle, reversible reactions are self balancing or self-correcting. when they are thrown out of balance by a change in temperature, concentration or pressure, the system will naturally shift in such a way as to “re-balance” itself after the change.
Formation of methanol from syngas
CO+2H2⇌CH3OH
Okay so for any industry they want maximum yield of formation of desired product so for getting maximum yield Le-Chatelier’s principle plays an important role. Suppose we want to increase the concentration of CO in the above system then. By Le Chatelier’s principle, we can say that the amount of methanol will increase, thereby decreasing the total change in CO. If we add a species to the reaction, the reaction will direct to the side opposing the addition of the species. Likewise, the removal of a species would cause the reaction to disturb and favour the side where the species was reduced. In simple words we can say that to get higher yield of methanol more & more amount of CO required adding so that maximum amount of CO gets reacted to form methanol, & also another way we can increase the yield of methanol is by the continuous removal of product means reducing the quantity of product.
This observation is supported by the collision theory. As the concentration of CO is increased, the frequency of successful collisions of that reactant would increase as well, allowing for an increase in the forward reaction, and thus the generation of the product. if a desired product is not thermodynamically favoured, then also the end-product can be obtained if it is continuously removed from the solution.
Changes in Temperature
The effect of temperature on equilibrium has to do with the heat of reaction. Recall that for an endothermic reaction, heat is absorbed in the reaction, and the value of ΔH is positive. Thus, for an endothermic reaction, we can take heat as being a reactant:
heat+A⇌B ΔH=+
For an exothermic reaction, the situation is just the opposite. Heat is released, so heat is a product, and the value of ΔH is with a negative sign:
A⇌B+heat ΔH=−
If we take heat as a reactant or a product, by applying Le Chatelier’s principle just like we did in our discussion on increasing or lowering concentrations. Let us suppose, if we increase the temperature on an endothermic reaction, it is essentially like adding more reactant to the system, and therefore, by Le Chatelier’s principle, the equilibrium will shift the right. Conversely, decreasing the temperature on an endothermic reaction will shift the equilibrium to the left, since decreasing the temperature in this case is equivalent to removing a reactant.
For an exothermic reaction, heat is a product. Therefore, increasing the temperature will shift the equilibrium to the left, while decreasing the temperature will shift the equilibrium to the right.
Lime production from Limestone
The reaction scheme is endothermic as it absorbs 178 kJ of heat in the form of energy for conversion to the desired product of CaO whose formation is influenced by high temperature making the reaction feasible till it reaches to the temperature of the reaction scheme to be always favoured on the right hand side.
So in simple words we can say that to get the higher yield of CaO it is required to increase the temperature of reactants so more & more energy gets absorbed and equilibrium shifts in forward direction.
the gaseous carbon di-oxide is evolved from the reaction & there is no trace of gaseous reactant and hence the lowering of pressure favours the formation of more products. When we compare it to the old kilns, modern rotary kiln favours high production of lime by continuous extraction of lime from the process.
Changes in Pressure
A change in pressure or volume will result in an attempt to restore equilibrium by creating more or less moles of gas. For example, if the pressure in a system increases, or the volume decreases, the equilibrium will shift to favour the side of the reaction that involves fewer moles of gas. Similarly, if the volume of a system increases, or the pressure decreases, the production of additional moles of gas will be favoured.
N2+3H2⇌2NH3 ΔH=−92kJ mol−1
take the number of moles of gas on the left-hand side and the number of moles of gas on the right-hand side. When you change the volume of system, the partial pressures of the gases change. If we decrease pressure by increasing volume, the equilibrium of the above reaction would shift to the left, because the reactant side has greater number of moles than the product side. The system tries to react opposite in the decrease in partial pressure of gas molecules by shifting to the side that exerts greater pressure.
Similarly, if increase pressure by decreasing volume, the equilibrium would shift to the right, counteracting the pressure increase by shifting to the side with fewer moles of gas that exert less pressure.
Lastly, for a gas-phase reaction in which the number of moles of gas on both sides of the equation are equal, the system will be unaffected by changes in pressure, since Δn=0.
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