Brief Summary
This session covers chemical equilibrium, focusing on key concepts, types of reactions, and factors affecting equilibrium. It includes explanations of homogeneous and heterogeneous reactions, endothermic and exothermic reactions, and the relationship between Kp and Kc. The session also discusses Le Chatelier's principle and its applications.
- Chemical equilibrium is a state where the rate of forward and backward reactions are equal, and measurable properties remain constant.
- Types of chemical reactions include homogeneous (same phase) and heterogeneous (different phases).
- Factors affecting equilibrium include temperature, pressure, and the addition of inert gases.
Session Intro
The session begins with a welcome and encouragement for viewers to like, share, and subscribe. The instructor expresses excitement for the session on chemical equilibrium, emphasizing its importance for exams.
Equilibrium Definition
Equilibrium is defined as a state where properties do not change, using examples like a water tank with constant inflow and outflow. Chemical equilibrium is introduced, highlighting that it involves a balance where changes are not constant.
Chemical Equilibrium Examples
The discussion covers examples of chemical reactions reaching equilibrium, such as the reversible reaction A ⇌ B. It explains that equilibrium is reached when the concentrations of reactants and products become constant, though not necessarily equal.
Representing Chemical Reactions
A chemical reaction is defined as a symbolic representation of a chemical change, which may or may not be visible. The session touches on types of chemical reactions, including homogeneous and heterogeneous reactions.
Types of Chemical Reactions: Homogeneous vs. Heterogeneous
Homogeneous reactions involve reactants and products in the same phase, exemplified by H2(g) + I2(g) ⇌ 2HI(g). Heterogeneous reactions involve different phases, such as CaCO3(s) ⇌ CaO(s) + CO2(g). The distinction is based on the state or phase of the reactants and products.
Energy Basis: Endothermic vs. Exothermic Reactions
Reactions are classified based on heat absorption (endothermic) or release (exothermic). Endothermic reactions absorb heat, resulting in a positive ΔH, while exothermic reactions release heat, resulting in a negative ΔH. Stability of reactants and products is also discussed in relation to these reactions.
Equilibrium and Irreversible Reactions
Equilibrium exists only in reversible reactions, which can proceed in both directions. Irreversible reactions proceed in one direction until completion. Reversible reactions are generally carried out in closed containers, while irreversible reactions occur in open containers.
Equilibrium State and Properties
The equilibrium state is defined by constant measurable properties such as pressure, temperature, moles, concentration, and volume. The rate of forward and backward reactions must be equal at equilibrium. Graphs illustrating rate versus time and concentration versus time are used to explain these concepts.
Equilibrium Constant: Kc and Kp
The equilibrium constant can be expressed in terms of concentration (Kc) or partial pressure (Kp). Kc is defined for both gases and solutions, while Kp is defined only for gases. Dalton's Law of Partial Pressure is relevant for understanding Kp.
Dalton's Law of Partial Pressure
Dalton's Law states that the total pressure exerted by a mixture of non-reacting gases is equal to the sum of the partial pressures of each individual gas. The partial pressure of a gas is the pressure it would exert if it occupied the volume alone.
Kp and Kc Relationship
The relationship between Kp and Kc is given by the equation Kp = Kc(RT)^Δng, where Δng is the change in the number of moles of gas in the reaction. The session explains how to apply this equation and interpret the value of Δng.
Extent of Reaction and Equilibrium Constant
The extent of a reaction is related to the value of the equilibrium constant K. A high K indicates a greater concentration of products at equilibrium, while a low K indicates a greater concentration of reactants. Specific values of K (e.g., K > 10^3, K < 10^-3) are used to describe the extent of the reaction.
Variation of Equilibrium Constant
The session discusses how the equilibrium constant changes when a reaction is reversed, multiplied by a factor, or divided by a factor. These changes are important for solving problems involving equilibrium.
Degree of Dissociation
The degree of dissociation (α) is defined as the number of moles dissociated divided by the initial number of moles. The percentage dissociation is α multiplied by 100. The session provides an example of calculating α and Kp for the reaction PCl5(g) ⇌ PCl3(g) + Cl2(g).
Reaction Quotient
The reaction quotient (Qc) is a measure of the relative amount of products and reactants present in a reaction at any given time. It is used to determine if a reaction is at equilibrium and, if not, which direction it must shift to reach equilibrium.
Le Chatelier's Principle
Le Chatelier's principle states that if a change of condition (e.g., change in concentration, pressure, or temperature) is applied to a system in equilibrium, the system will shift in a direction that relieves the stress. The session explains how to apply this principle to predict the effect of changes on equilibrium.
Factors Affecting Equilibrium: Temperature, Pressure, and Inert Gases
The factors affecting equilibrium are temperature, pressure, and the addition of inert gases. Temperature changes affect endothermic and exothermic reactions differently. Pressure changes shift the equilibrium towards fewer moles of gas. Adding inert gases at constant volume has no effect, while adding them at constant pressure shifts the equilibrium towards more moles of gas.
Session Outro
The session concludes with a thank you and encouragement for viewers to like, subscribe, share, and comment. The instructor mentions upcoming biology sessions and expresses gratitude for the viewers' participation.
