IB Chemistry: The Position of Equilibrium
The Position of Equilibrium
Deduce the equilibrium constant expression (Kc) from the equation for a homogeneous reaction.
Homogeneous Reaction: A homogenous reaction is one where all the substances are in the same phase. E.g. A reaction where all the reactants and products are liquids.
Equilibrium Constant
The equilibrium constant, Kc, is a comparison between the concentration of the Products and the reactants.
Take a reaction like the one above. If we form the equation above in words, it will sound something like this:
“a” moles of [A] and “b” moles of [B] react to form “c” moles of [C] and “d” moles of [D].
According to the Equilibrium Law,
E.g. N2 + 3H2 –> 2NH3
There’s no [D] here, so let’s just consider it as [C]/[A][B]
Kc = [NH3]2 / [N2][H2]3
E. g 2SO2 (g) + O2 –> 2SO3
Kc = [SO3]2 / [O2][SO2]2
E.g. H2 (g) + I2 (g) –> 2HI (g)
Kc = [HI]2 / [I2][H2]
Deduce the extent of a reaction from the magnitude of the equilibrium constant.
The value of Kc can tell us the position of the equilibrium.
Kc =1: The concentration of the reactants and the products will be exactly the same.
Kc >> 1, the reaction will go almost to completion
Kc > 1, then the concentration of the products will be greater than that of the reactants, and the equilibrium will lie to the right.
Kc<<1, there is hardly any reaction taking place.
Kc <1, then the concentration of the reactants will be greater than that of the products, and the equilibrium will lie to the left.
Apply Le Chatelier’s principle to predict the qualitative effects of changes of temperature, pressure and concentration on the position of equilibrium and on the value of the equilibrium constant.
Le Chatelier’s Principle: Le Chatelier’s Principle basically dictates that as long as temperature remains constant, the value of Kc, will remain constant. Therefore, if the system is subject to a certain amount of stress, the system will shift in the direction which will favor a relief in stress.
The three types of stress that can occur on a system are:
Temperature
Pressure
Concentration
Concentration
A + B ⇌ C + D
Here is a table you can use to memorize what will happen to the equilibrium as you change the concentration.
Change in concentration
Direction of shift
Increase concentration of a reactant
–>
Increase concentration of a product
<–
Decrease concentration of a reactant
<–
Decrease concentration of a product
–>
Intuitive explanation:
If you increase the concentration of a reactant, the equilibrium will shift to the right to increase concentration of products to obey the Le Chatelier’s Principle.
If you increase the concentration of a product, the equilibrium will shift to the left to increase the concentration of products to obey the Le Chatelier’s Principle.
If you decrease the concentration of a reactant, the equilibrium will shift to the left to increase the concentration of the reactants to achieve equilibrium.
If you decrease the concentration of a product, the equilibrium will shift to the right to increase the concentration of the products to achieve equilibrium.
Pressure
A + B ⇌ C + D
Here is a table to memorize the effects of pressure on what happens to the direction of equilibrium.
Change in pressure of a gaseous reaction
Direction of Shift
Inert gas
No effect.
Increased pressure.
Favor the side with less moles of gas, so equilibrium will shift to the side with less moles of gas.
Temperature
Change in temperature
Direction of Shift
Increase in temperature
In the endothermic direction.
Decrease in temperature
In the exothermic direction.
In an exothermic reaction, heat is released as a product.
If you take the heat away, the equilibrium will be shifted to the right to form more products.
The forward reaction in exothermic reactions is therefore increased by lowering the temperature.
The opposite occurs in an endothermic reaction.
However, the change in temperature will also change the value of Kc.
E.g. In an exothermic reaction, the concentration of the products will decrease as the temperature increases, so the value of Kc decreases.
State and explain the effect of a catalyst on an equilibrium reaction.
Catalysts decreases the activation energy and hence increases the rate at which equilibrium is reached.
However, catalyst’s will speed both the forward and reverse reactions, and will therefore have no effect on the position of equilibrium, and hence no effect on the value of Kc.
Apply the concepts of kinetics and equilibrium to industrial processes
There are two industrial processes that we will focus on in the IB Diploma.
The aim here is to achieve the highest possible yield with the least amount of time and for the least cost.
Why?
To earn money….
These are:
Haber Process
Contact Process
Haber Process
N2 (g) + 3 H2 (g) ⇌ 2 NH3 (g) (ΔH = −92.22 kJ·mol−1)
This is the process of manufacturing Ammonia for Industrial use.
Hydrogen obtained from natural air.
Nitrogen obtained through distillation of liquid air.
Temperature
As ΔH = −92.22 kJ·mol−1 gives you a negative enthalpy change, the reaction is exothermic and thus a lower temperature should be used to give the highest yield of ammonia.
HOWEVER, lowering the temperature will lower the rate of reaction, as experienced in Unit 6, Kinetics.
The solution here is to find the Optimal Temperature, which is usually around 450 °C, give or take a few degrees.
Catalyst
An iron catalyst is used to maximize surface area, and hence the efficiency of the production.
Pressure
4 moles of reactants ⇌ 2 moles of products
The reaction takes place at high pressure (150-250 atm). Why?
At high pressure, Le Chatlier’s Principle dictates that the direction of equilibrium favors the side with fewer moles of gas. There are only 2 moles of products, so the direction of equilibrium will shift to the right, hence increasing rate in which equilibrium is achieved.
Summary:
High pressure
Low temperature, but the temperature has to be optimal.
Iron catalyst used.
Still, only 15% yield.
Contact Process
2 SO2(g) + O2(g) ⇌ 2 SO3(g) : ΔH = −197 kJ mol−1
This is the manufacture of Sulfuric Acid
Most industrially produced chemical, with production levels of over 150 million tonnes annually.
Used for fertilizers, detergents and etc.
Temperature
Similar to the Haber Process , enthalpy change is a negative value, so the reaction is exothermic. As a result, a low temperature will be used to give a high yield of sulfuric acid. However, a high temperature (approx. 450°C) will be used to ensure a reasonable rate of reaction, so the yield will be sacrificed to a certain extent.
Pressure
3 moles of reactants ⇌ 2 moles of products
Manufacture takes place at high pressure, the reasons being similar to that of the Haber Process. Essentially, high pressure is used to shift the reaction to the product side.
Catalyst
Vanadium Oxide is used.
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