Equilibrium
Equilibrium
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Equilibrium occurs when you have a REVERSIBLE REACTION in a CLOSED SYSTEM.
A CLOSED SYSTEM is a system where nothing can ENTER or LEAVE. E.g. a conical flask with a rubber bung closing it.
At the start of any reaction, only REACTANT particles are present. As these collide and REACT, they form PRODUCT particles.
In normal reactions, the amount of reactants would DECREASE to ZERO at the end of the reaction, and the amount of products would INCREASE to a MAXIMUM.
However this does NOT happen in a reversible reaction.
The PRODUCTS can also react to reform the REACTANTS in the REVERSE reaction.
This means there are TWO reactions occurring in the closed system:
When the FORWARD and REVERSE reactions both occur at the SAME RATE, we say the reaction is in EQUILIBRIUM.
At this point, the AMOUNTS of reactants and products STAY THE SAME.
You can change the RELATIVE AMOUNTS of reactants and products in a reaction at equilibrium by changing the CONDITIONS of the reaction. This is known as SHIFTING the POSITION of equilibrium and you can shift it in TWO WAYS:
Shifting the equilibrium to the LEFT will cause the relative amounts of the REACTANTS to INCREASE and the PRODUCTS to DECREASE.
Shifting the equilibrium to the RIGHT will cause the relative amounts of the REACTANTS to DECREASE and the PRODUCTS to INCREASE.
There are THREE factors that can shift the equilibrium and affect its POSITION:
To work out the effect that CHANGING one of these FACTORS has on a system at equilibrium, you can use LE CHATELIER'S PRINCIPLE.
Le Chatelier’s principle states that if a CHANGE is made to the conditions of a system at EQUILIBRIUM, the system will respond and SHIFT to COUNTERACT the change.
For example, if you have a system at equilibrium and you heat it to INCREASE the temperature, the equilibrium will respond in a way to DECREASE the temperature back to normal.
Here are some more examples:
If you change TEMPERATURE, the direction that the equilibrium shifts depends on whether the forward reaction is EXOTHERMIC or ENDOTHERMIC.
ENDOTHERMIC reactions show a DECREASE in temperature, while EXOTHERMIC reactions show an INCREASE in temperature.
This means that if you INCREASE the temperature of a system at equilibrium, it will try to COUNTERACT the change by DECREASING the temperature, so it will shift to the ENDOTHERMIC side of the reaction. The opposite will happen if you DECREASE the temperature.
This shifting of equilibrium can affect the amount of reactants and products in the system.
For an example where the FORWARD reaction is EXOTHERMIC:
If the FORWARD reaction is ENDOTHERMIC, the opposite would happen:
For gaseous reactions, if you change PRESSURE, the direction that the equilibrium shifts depends on the NUMBER OF MOLECULES (moles) on either side of the equation. These are the THREE scenarios you can have:
The more MOLECULES you have in a closed system, the GREATER the PRESSURE.
This means if you INCREASE the pressure of a system at equilibrium, it will try to COUNTERACT the change by DECREASING the pressure, so it will shift to the side with FEWER MOLECULES. The opposite will happen if you DECREASE the pressure.
This shifting of equilibrium can affect the amount of reactants and products in the system.
For an example where there are MORE MOLES on the LEFT:
If there are MORE MOLES on the RIGHT, the opposite would happen:
If there are the SAME number of moles on the left and right, NOTHING HAPPENS.
If you were to change the CONCENTRATION, you will have disrupted the equilibrium balance, meaning the reaction is no longer in equilibrium for a short amount of time.
When this happens, the reaction RESPONDS and gets itself back into equilibrium and COUNTERACTS the change that was made to it.
You can predict the direction the equilibrium shifts for the short amount of time by knowing whether the concentration that changes is a REACTANT or PRODUCT.
Here are all the possible scenarios:
Equilibrium occurs when you have a REVERSIBLE REACTION in a CLOSED SYSTEM.
A CLOSED SYSTEM is a system where nothing can ENTER or LEAVE. E.g. a conical flask with a rubber bung closing it.
At the start of any reaction, only REACTANT particles are present. As these collide and REACT, they form PRODUCT particles.
In normal reactions, the amount of reactants would DECREASE to ZERO at the end of the reaction, and the amount of products would INCREASE to a MAXIMUM.
However this does NOT happen in a reversible reaction.
The PRODUCTS can also react to reform the REACTANTS in the REVERSE reaction.
This means there are TWO reactions occurring in the closed system:
When the FORWARD and REVERSE reactions both occur at the SAME RATE, we say the reaction is in EQUILIBRIUM.
At this point, the AMOUNTS of reactants and products STAY THE SAME.
You can change the RELATIVE AMOUNTS of reactants and products in a reaction at equilibrium by changing the CONDITIONS of the reaction. This is known as SHIFTING the POSITION of equilibrium and you can shift it in TWO WAYS:
Shifting the equilibrium to the LEFT will cause the relative amounts of the REACTANTS to INCREASE and the PRODUCTS to DECREASE.
Shifting the equilibrium to the RIGHT will cause the relative amounts of the REACTANTS to DECREASE and the PRODUCTS to INCREASE.
There are THREE factors that can shift the equilibrium and affect its POSITION:
To work out the effect that CHANGING one of these FACTORS has on a system at equilibrium, you can use LE CHATELIER'S PRINCIPLE.
Le Chatelier’s principle states that if a CHANGE is made to the conditions of a system at EQUILIBRIUM, the system will respond and SHIFT to COUNTERACT the change.
For example, if you have a system at equilibrium and you heat it to INCREASE the temperature, the equilibrium will respond in a way to DECREASE the temperature back to normal.
Here are some more examples:
If you change TEMPERATURE, the direction that the equilibrium shifts depends on whether the forward reaction is EXOTHERMIC or ENDOTHERMIC.
ENDOTHERMIC reactions show a DECREASE in temperature, while EXOTHERMIC reactions show an INCREASE in temperature.
This means that if you INCREASE the temperature of a system at equilibrium, it will try to COUNTERACT the change by DECREASING the temperature, so it will shift to the ENDOTHERMIC side of the reaction. The opposite will happen if you DECREASE the temperature.
This shifting of equilibrium can affect the amount of reactants and products in the system.
For an example where the FORWARD reaction is EXOTHERMIC:
If the FORWARD reaction is ENDOTHERMIC, the opposite would happen:
For gaseous reactions, if you change PRESSURE, the direction that the equilibrium shifts depends on the NUMBER OF MOLECULES (moles) on either side of the equation. These are the THREE scenarios you can have:
The more MOLECULES you have in a closed system, the GREATER the PRESSURE.
This means if you INCREASE the pressure of a system at equilibrium, it will try to COUNTERACT the change by DECREASING the pressure, so it will shift to the side with FEWER MOLECULES. The opposite will happen if you DECREASE the pressure.
This shifting of equilibrium can affect the amount of reactants and products in the system.
For an example where there are MORE MOLES on the LEFT:
If there are MORE MOLES on the RIGHT, the opposite would happen:
If there are the SAME number of moles on the left and right, NOTHING HAPPENS.
If you were to change the CONCENTRATION, you will have disrupted the equilibrium balance, meaning the reaction is no longer in equilibrium for a short amount of time.
When this happens, the reaction RESPONDS and gets itself back into equilibrium and COUNTERACTS the change that was made to it.
You can predict the direction the equilibrium shifts for the short amount of time by knowing whether the concentration that changes is a REACTANT or PRODUCT.
Here are all the possible scenarios:
