Friday, March 18, 2011

Chemical Equilibrium

Definition of Chemical Equilibrium

Chemical equilibrium applies to reactions that can occur in both directions. In a reaction such as:
CH4(g) + H2O(g) <--> CO(g) + 3H2(g)
The reaction can happen both ways. So after some of the products are created the products begin to react to form the reactants. At the beginning of the reaction, the rate that the reactants are changing into the products is higher than the rate that the products are changing into the reactants. Therefore, the net change is a higher number of products.
Even though the reactants are constantly forming products and vice-versa the amount of reactants and products does become steady. When the net change of the products and reactants is zero the reaction has reached equilibrium. The equilibrium is a dynamic equilibrium. The definition for a dynamic equilibrium is when the amount of products and reactants are constant. (They are not equal but constant. Also, both reactions are still occurring.)

Equilibrium Constant

To determine the amount of each compound that will be present at equilibrium you must know the equilibrium constant. To determine the equilibrium constant you must consider the generic equation:
aA + bB <--> cC + dD
The upper case letters are the molar concentrations of the reactants and products. The lower case letters are the coefficients that balance the equation. Use the following equation to determine the equilibrium constant (Kc). Kc equation
For example, determining the equilibrium constant of the following equation can be accomplished by using the Kc equation.
Using the following equation, calculate the equilibrium constant.
N2(g) + 3H2(g) <--> 2NH3(g)
A one-liter vessel contains 1.60 moles NH3, .800 moles N2, and 1.20 moles of H2. What is the equilibrium constant?
example equilibrium constant
Answer: 1.85

Le Chatelier's Principle

Le Chatelier's principle states that when a system in chemical equilibrium is disturbed by a change of temperature, pressure, or a concentration, the system shifts in equilibrium composition in a way that tends to counteract this change of variable. The three ways that Le Chatelier's principle says you can affect the outcome of the equilibrium are as follows:
  • Changing concentrations by adding or removing products or reactants to the reaction vessel.
  • Changing partial pressure of gaseous reactants and products.
  • Changing the temperature.
These actions change each equilibrium differently, therefore you must determine what needs to happen for the reaction to get back in equilibrium.

Example involving change of concentration:

In the equation
2NO(g) + O2(g) <--> 2NO2(g)
If you add more NO(g) the equilibrium shifts to the right producing more NO2(g)
If you add more O2(g) the equilibrium shifts to the right producing more NO2(g)
If you add more NO2(g) the equilibrium shifts to the left producing more NO(g) and O2(g)

Example involving pressure change:

In the equation
2SO2(g) + O2(g) <--> 2SO3(g),
an increase in pressure will cause the reaction to shift in the direction that reduces pressure, that is the side with the fewer number of gas molecules. Therefore an increase in pressure will cause a shift to the right, producing more product. (A decrease in volume is one way of increasing pressure.)

Example involving temperature change:

In the equation
N2(g) + 3H2(g) <--> 2NH3 + 91.8 kJ,
an increase in temperature will cause a shift to the left because the reverse reaction uses the excess heat. An increase in forward reaction would produce even more heat since the forward reaction is exothermic. Therefore the shift caused by a change in temperature depends upon whether the reaction is exothermic or endothermic.

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