Introduction
In chemistry, equilibrium refers to the state in which both reactants and products are present in concentrations that have no further tendency to change with time. It occurs when the forward and reverse reactions occur at equal rates. Understanding equilibrium is crucial for predicting the concentrations of substances in chemical reactions, which is fundamental for solving problems in NEET Chemistry.
Types of Equilibrium
Chemical Equilibrium
Chemical equilibrium is the state in which the rate of the forward reaction equals the rate of the reverse reaction in a closed system. At this point, the concentrations of reactants and products remain constant over time.
Dynamic Nature of Chemical Equilibrium
Even though the concentrations of reactants and products remain constant, the reactions continue to occur at the molecular level. This is known as dynamic equilibrium.
ExampleConsider the synthesis of ammonia: $$ \text{N}_2(g) + 3\text{H}_2(g) \rightleftharpoons 2\text{NH}_3(g) $$ At equilibrium, the rate of formation of ammonia equals the rate of its decomposition.
Physical Equilibrium
Physical equilibrium involves changes in the physical state rather than chemical composition. Examples include phase transitions like vaporization, condensation, melting, and freezing.
ExampleWater in a closed container: $$ \text{H}_2\text{O}(l) \rightleftharpoons \text{H}_2\text{O}(g) $$ At equilibrium, the rate of evaporation equals the rate of condensation.
Law of Mass Action
The Law of Mass Action states that the rate of a chemical reaction is proportional to the product of the concentrations of the reactants, each raised to a power equal to the stoichiometric coefficient.
Equilibrium Constant ($K_c$)
For a general reaction: $$ aA + bB \rightleftharpoons cC + dD $$ The equilibrium constant ($K_c$) is given by: $$ K_c = \frac{[\text{C}]^c [\text{D}]^d}{[\text{A}]^a [\text{B}]^b} $$
Note$K_c$ is dimensionless and depends only on temperature.
Equilibrium Constant in Terms of Partial Pressure ($K_p$)
For reactions involving gases, the equilibrium constant can also be expressed in terms of partial pressures: $$ K_p = \frac{(P_C)^c (P_D)^d}{(P_A)^a (P_B)^b} $$
Tip$K_p$ and $K_c$ are related by the equation: $$ K_p = K_c (RT)^{\Delta n} $$ where $\Delta n = (c + d) - (a + b)$ and $R$ is the gas constant.
Le Chatelier's Principle
Le Chatelier's Principle states that if a dynamic equilibrium is disturbed by changing the conditions, the position of equilibrium moves to counteract the change.
Changes in Concentration
Increasing the concentration of reactants shifts the equilibrium to the right, favoring the formation of products, and vice versa.
Changes in Pressure
For gaseous equilibria, increasing the pressure shifts the equilibrium toward the side with fewer moles of gas.
ExampleFor the reaction: $$ \text{N}_2(g) + 3\text{H}_2(g) \rightleftharpoons 2\text{NH}_3(g) $$ Increasing pressure favors the formation of ammonia because there are fewer moles of gas on the product side.
Changes in Temperature
The effect of temperature change depends on whether the reaction is exothermic or endothermic.
- For exothermic reactions, increasing temperature shifts the equilibrium to the left.
- For endothermic reactions, increasing temperature shifts the equilibrium to the right.
Do not confuse the effect of temperature on $K_c$ and $K_p$ with other changes. Temperature changes can alter the values of the equilibrium constants.
Applications of Equilibrium Concepts
Solubility Product ($K_{sp}$)
The solubility product constant is used for sparingly soluble salts. For a salt $AB$ that dissociates as: $$ AB(s) \rightleftharpoons A^+(aq) + B^-(aq) $$ The solubility product $K_{sp}$ is given by: $$ K_{sp} = [A^+][B^-] $$
Common Ion Effect
The common ion effect describes the decrease in solubility of an ionic compound due to the presence of a common ion.
ExampleAdding NaCl to a solution of AgCl decreases the solubility of AgCl because of the common ion $Cl^-$.
Buffer Solutions
Buffers are solutions that resist changes in pH upon addition of small amounts of acid or base. They are typically made from a weak acid and its conjugate base or a weak base and its conjugate acid.
TipThe pH of a buffer solution can be calculated using the Henderson-Hasselbalch equation: $$ \text{pH} = \text{p}K_a + \log \left( \frac{[\text{A}^-]}{[\text{HA}]} \right) $$
Conclusion
Understanding equilibrium is essential for solving many problems in NEET Chemistry. By mastering the concepts of dynamic equilibrium, the Law of Mass Action, and Le Chatelier's Principle, you can predict the behavior of chemical systems under various conditions. Additionally, applications like the solubility product, common ion effect, and buffer solutions are crucial for practical problem-solving.
NoteRemember that equilibrium is a dynamic process, and the system continuously adjusts to maintain balance.
