Chemical Equilibrium Fundamentals

Chemical equilibrium is a fundamental concept in chemistry. It describes a state where the concentrations of reactants and products remain constant over time, as the rates of the forward and reverse reactions balance each other.

Equilibrium in Chemistry

Key Definitions

• Equilibrium: In a chemical reaction, equilibrium is established when the rate of the forward reaction matches the rate of the reverse reaction.

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  Equilibrium: A condition in a chemical reaction wherein the rates of the forward and reverse reactions are equal.

• Dynamic Equilibrium: This occurs when the rates of the forward and reverse reactions are equal, resulting in constant concentrations of reactants and products.

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  Dynamic Equilibrium: A state within a reversible reaction with equal forward and reverse reaction rates, maintaining consistent concentrations.

• Static Equilibrium: A hypothetical state where no reactions occur, and the system remains unchanged.

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  Static Equilibrium: A state of equilibrium with no net change and no motion.

Dynamic vs. Static Equilibrium

Aspect	Dynamic Equilibrium	Static Equilibrium
Nature of System	Continuous molecular reactions	Inactive, stable conditions
Rate of Reactions	Forward and reverse rates are equal	No reactions
Concentration Change	No net concentration change	Constant due to inactivity

Understanding Dynamic Equilibrium

Dynamic equilibrium is achieved in closed systems, preventing the exchange of matter with the surroundings.

• Examples: In a sealed soda bottle, carbon dioxide dissolves and escapes the liquid at the same rate, maintaining a constant concentration.

Equilibrium in the Haber Process

The Haber Process for ammonia synthesis exemplifies dynamic equilibrium:

• Conditions: High pressure, optimal temperature, and catalysts are applied to maintain equilibrium and maximise ammonia yield.

• Mathematical Representation:

  $aA + bB \rightleftharpoons cC + dD: K_c = \frac{[C]^c[D]^d}{[A]^a[B]^b}$

Le Chatelier's Principle

Le Chatelier's Principle describes how equilibrium systems adjust to changes in concentration, temperature, and pressure.

Effects of Changes

• Concentration: Changing the concentration of reactants or products will shift the equilibrium to counteract the alteration.

• Temperature: Increasing temperature favours endothermic reactions, while decreasing temperature favours exothermic reactions.

• Pressure: Increasing pressure favours the side with fewer gas molecules.

Practical Investigations

Experiment 1: Cobalt(II) Chloride Equilibrium

• Observe the effect of temperature changes on equilibrium colours, as explained by Le Chatelier's Principle.

Worked Example: If we have the equilibrium: CoCl₂·6H₂O (pink) ⇌ CoCl₂·4H₂O (blue) + 2H₂O

When we heat this system:

1. The equilibrium shifts to the right (endothermic direction)
2. More water is released
3. The solution becomes more blue in colour

Experiment 2: Iron(III) Nitrate and Potassium Thiocyanate System

• Notice how increasing reactant concentration shifts the equilibrium position.

Conclusion

Understanding Equilibrium: Fully grasping the distinctions between static and dynamic equilibrium is fundamental for understanding chemical reactions, facilitating both practical and theoretical applications in scientific and industrial settings.

Essential Concepts:

Studying equilibrium equips students with the understanding necessary for practical scientific applications.