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Equilibrium in chemistry refers to a state where the forward and reverse reactions occur at the same rate. This means that the concentrations of reactants and products remain constant over time. It's important to understand that equilibrium does not mean the amounts are equal, but rather that their rates of production and consumption are balanced.
The concept of equilibrium is crucial in predicting how a chemical system will respond to changes. Whether involving gases, liquids, or solids, each system seeks a state of equilibrium where any disturbance will be counteracted. This is where Le Chatelier's Principle comes into play, helping predict shifts in equilibrium when external conditions such as pressure, temperature, or concentration change.

• Static equilibrium: No net change occurs over time, but microscopic changes can still happen.
• Dynamic equilibrium: Both reactions continue to happen, but at an equal rate, leading to no overall change.

A chemical reaction is a process that leads to the transformation of one set of chemical substances to another. It involves breaking and forming bonds between atoms. During reactions, reactants are converted to products. The reaction between ozone ( $O\_3$) and oxygen ( $O\_2$) is a classic example of decomposition.
In the decomposition reaction:

• The reactants are 2 moles of ozone (O3).
• The products are 3 moles of molecular oxygen (O2).

This reaction highlights how a change in the number of moles on each side of an equation can be significant, especially under varying conditions like pressure.
Understanding the nature of chemical reactions helps predict how chemical systems will behave. When discussing reactions at equilibrium, it is essential to count moles for both sides, as they often determine how a system responds to external changes.

Pressure can significantly affect a gaseous equilibrium. Le Chatelier's Principle helps predict the direction in which an equilibrium will shift in response to an external pressure change.
When pressure increases, systems will shift the equilibrium position to reduce that pressure. This usually means favoring the side of the reaction with fewer moles of gas. In the ozone decomposition reaction:

• Left side (reactants): 2 moles of gas (O3)
• Right side (products): 3 moles of gas (O2)

Hence, if pressure is increased, the system will shift towards the side with fewer moles, which in this case is the reactant side (O3).
This change helps minimize the effect of added pressure and brings the system back towards equilibrium. Thus, understanding pressure effects helps in controlling and optimizing reactions in industrial and laboratory settings.