Le Chatelier’s Principle
Sadika explains about Le Chatelier’s Principle.
Le Chatelier’s Principle
For any given chemical process, the correct raw materials need to be brought together in the right place, at the right time, and under the right conditions. Chemists need to carefully monitor processes to ensure optimum conditions. Important conditions include:
- concentration of reactants and products
- presence of a catalyst
- pressure (only if gases are involved)
If a system is in equilibrium and a change is made that upsets the equilibrium, then the system alters in such a way as to counteract the change and a new equilibrium is established.
The Haber process involves an equilibrium reaction, and knowledge of Le Chatelier’s principle is needed in order to predict how reaction conditions will impact on the production of ammonia by this process
Rule 1: Temperature
If the forward reaction is endothermic, increasing the temperature favours the formation of the product.The yield of product will be increased.
If the forward reaction is exothermic, reducing the temperature favours the formation of the product. The yield of the product will be increased.
These rules apply to all reactions.
Rule 2: Pressure
applies only to reactions with one or more gaseous reactants or products.
Increasing the system pressure (by reducing the volume) causes the equilibrium to shift to the side of the equation with the least number of gaseous molecules.
Decreasing the system pressure (by increasing the volume) causes the equilibrium to shift to the side of the equation with the most number of gaseous molecules.
Where there are equal numbers of molecules on the reactant and product sides, pressure has no effect on the position of the equilibrium.
Rule 3: Concentration
Increasing the reactant concentration shifts the equilibrium to the right to make more products and to reduce the concentration of the added reactant.
Decreasing the concentration of a reactant shifts the equilibrium to the left to make more of that reactant. Consequently the concentration of products decreases.
Rule 4: Catalysts
Catalysts make reactions go faster.
They do not affect the position of equilibrium.
Catalysts cause the forward and reverse reactions to speed up by equal amounts.
Catalysts are useful because they lower the activation energy, and equilibrium is achieved faster.
This is important in industry as there are considerable savings to be made in time and energy as reactions can be conducted (where appropriate) at lower temperatures.
Synthesis of Ammonia
Explain why the synthesis of ammonia has high activation energy.
To form ammonia, nitrogen atoms and hydrogen atoms must combine.
However, nitrogen and hydrogen do not exist as single atom, they both exist as diatomic molecules held together by strong covalent bonds.
Nitrogen molecules are held together by very strong triple covalent bonds and the single covalent bond between hydrogen atoms is also quite strong.
A lot of energy is needed to break these bonds and form atoms that can rearrange to form ammonia.
This energy is called the activation energy, the energy needed to start the chemical reaction.
Therefore the activation energy is high.
Use Le Chatelier’s Principle to explain why the yield of product in the Haber process is reduced at higher temperatures.
Le Chatelier’s principle states that, if a system already in equilibrium is disturbed by change in conditions, such as temperature, pressure, or concentrations, the equilibrium will shift to attempt to compensate for that change.
Since the forward equilibrium reaction to form ammonia is exothermic, higher temperatures will favour the backward, endothermic reaction, which absorbs some of the added heat. Hence the proportion of ammonia in the equilibrium mixture will be decreased by increasing the temperature..