Factors that Affect Chemical Equilibrium

Chemical equilibrium means that there is a balance between the rate of forward and backward reaction.  But this condition is so sensitive that whatever changes in some factors or variables there is a change in the production of products, it can either produce more or less products.  These variables or factors are concentration, pressure, volume, and temperature.

Le Chatelier's Principle

Now in chemical equilibrium there is a general rule that will guide us in predicting the direction in which a chemical reaction will move in case concentration, pressure, volume and temperature are changed, this is called the Le Chatelier's Principle.  This was formulated by a French Chemist Henri Le Chatelier.  Le Chateliers Principle states that if an external stress is applied to a system at equilibrium the system adjust in such a way that the stress is partially offset or balance as it tries to reestablish the equilibrium.  Stress means the variables or factors that are changed, the concentration, pressure, volume and temperature.

Concentration

Concentration is the amount of solute that is present in a given solvent.  The greater the amount of the solute in the solution the higher is the concentration.  Now what happens when there is a change in the concentration of the reactants and products? Let us give an example:

FeSCN⁺²(aq)        ⇄        Fe⁺³(aq)    +   SCN^-(aq)

                                                 red                       pale yellow          colorless

The above reaction shows what happens with Iron (III) thiocyanate [Fe(SCN)3] when dissolve in water.  The solution gives a red color due to the presence of hydrated FeSCN⁺² ion.  The substance when dissolve in water gives Fe⁺³ ion and SCN^- ion.  The system of chemical equilibrium occurs.

Now, what will happen if we add substance into the solution?  Let us say sodium thiocyanate (NaSCN) is added to the solution.  In this case, the stress added to the system is the  increase in the concentration of SCN^- ion from the dissociation of NaSCN.  To offset the stress the Fe⁺³ ions will react with SCN^- ions, and the equilibrium will shift from right to left.

FeSCN⁺²(aq)     ←     Fe⁺³(aq)    +   SCN^-(aq)

The above reaction will deepen the red color.  What if iron (III) nitrate [Fe(NO₃)₃] is added?  The red color will also deepen because the concentration of Fe⁺³ ions will increase from the dissociation of Fe(NO₃)₃, and it will react with SCN^- ions, and therefore the equilibrium will shift from right to left.  Both Na⁺ ions and NO₃^- ions are colorless spectator ions.

Suppose we add oxalic acid (H₂C₂O₄) to the original solution.   Oxalic acid ionizes in water to form oxalate ion (C₂O₄⁺²), which binds strongly with Fe⁺³ ions forming a yellow ion Fe(C₂O₄)₃^-3 which removes free Fe⁺³ ions in solution. And therefore, more FeSCN⁺² units dissociate and the equilibrium shifts from left to right.  The solution will now turn to yellow because of the formation of Fe(C₂O₄)₃^-3.

FeSCN⁺²(aq)     →    Fe⁺³(aq)    +   SCN^-(aq)

Changes in Pressure and Volume

Not all solutions are affected of the change in pressure and volume.  Only gases  are the one affected because of the ability of gases to be compressed.  While liquids and gases are virtually incompressible.   Let us recall the formula for ideal gas equation:

PV = nRT

P = (n/V)RT  

The equation tells that P is inversely related to V, which means as the volume is decreased the pressure increases and vice versa.  On the other hand, the concentration, n, of the gas is directly proportional to the pressure of gas, which means if the concentration is increased the pressure also increases.

Let us have an example, suppose we have the reaction below:

N₂O₄(g)    ⇄    2NO₂(g)

The equilibrium system is in a cylinder fitted with a movable piston.  What will happen if the pressure on the gases is increased?  Increasing the pressure will result in decrease in volume and so concentration of N₂O₄ and NO₂ will increase.  Therefore, increasing pressure will cause the shift of equilibrium to the left, more N₂O₄ will form because of the decrease in volume.  Why N₂O₄?  Simply because N₂O₄ has less number moles.  In other words if pressure is decreased the equilibrium will shift to the right in favor of more number of moles NO₂.

Therefore, an increase in pressure (decrease in volume) favors the net reaction which has less number of moles, in the above reaction, it favors the backward reaction; and the decrease in pressure (increase in volume) favors the the net reaction which has more number of moles.  In the above reaction, it favors the forward reaction.

Let us have another examples:
Predict the direction of the net reaction of the following equilibrium system if pressure will be increased.
a.   2PbS(s)  +   3O₂     ⇄      2PbO(s)   +   2SO₂(g)
b.  PCl₅(g)     ⇄      PCl₃(g)   +   Cl₂(g)

Solution:  
We have to be reminded that only gases are affected by pressure because solids and liquids are incompressible.
a. There are 3 moles of reactants and 2 moles of the products not including the solids, therefore increasing pressure or decreasing the volume will cause the system to shift to the right  in favor of less than number of moles or to the product side.
b.  Here only one mole in the reactant side and two moles in the product side increasing the pressure will shift the equilibrium to the left to the reactant side with less number of moles.

Changes in Temperature

A change on concentration, pressure and volume does not alter the equilibrium constant it only alters the relative amounts of reactants and products, only change in temperature can alter the equilibrium constant.  There are two reactions possible the exothermic reaction and the endothermic reaction.  Exothermic reaction is a chemical reaction that releases heat or energy while endothermic reaction is the chemical reaction that absorbs heat.

Let us have an example,

N₂O₄(g)      ⇄     2NO₂(g)

The above reaction is endothermic in the forward reaction (absorbs heat, △H^o > 0),

heat  +  N₂O₄(g)    →   2NO₂(g)          △H^o = 58.0 kJ/mole

so the reverse reaction is exothermic (releases heat, △H^o < 0)

2NO₂(g)   →    N₂O₄(g)  +  heat      △H^o = -58.0 kJ/mole

At equilibrium, the heat effect is zero because there is not net reaction.  Increasing the temperature will favor the endothermic reaction, the forward reaction (the shift is from left to right), the amount of N₂O₄ decreases and an increase in the amount of NO₂. On the other hand a decrease in the temperature will favor the exothermic reaction reaction, the backward reaction (the shift is from right to left), which decreases the NO₂ and increases N₂O₄.

Therefore we can say that an increase in temperature will favor the endothermic reaction and the decrease in temperature favors the exothermic reaction.

The Effect of Catalyst

Catalyst enhances the chemical reaction. Adding catalyst to the reaction will not alter the chemical equilibrium constant or will not shift the equilibrium to any direction, it just increases the reaction rate by lowering the activation energy of the reacting molecules.

Adding a catalyst to a reaction that is not at equilibrium will only help the reaction to achieve equilibrium at the soonest possible time.  The same equilibrium can be obtained without a catalyst only that it takes longer time to occur.