This chapter covers the principles of chemical equilibrium, including its significance in biological and environmental processes. It emphasizes understanding dynamic equilibrium, the equilibrium constant, and the factors affecting equilibrium states.
Equilibrium - Quick Look Revision Guide
Your 1-page summary of the most exam-relevant takeaways from Chemistry Part - I.
This compact guide covers 20 must-know concepts from Equilibrium aligned with Class 11 preparation for Chemistry. Ideal for last-minute revision or daily review.
Complete study summary
Essential formulas, key terms, and important concepts for quick reference and revision.
Key Points
Define dynamic equilibrium.
Dynamic equilibrium occurs when the rates of forward and reverse reactions equalize, maintaining constant reactant and product concentrations.
State Le Chatelier's principle.
If a system at equilibrium experiences a change (in concentration, pressure, or temperature), the system shifts to counteract that change, restoring equilibrium.
Write Kc expression for reactions.
Kc = [C]^c[D]^d/[A]^a[B]^b; uses concentrations of products/reactants raised to stoichiometric coefficients at equilibrium.
Relationship between Kp and Kc.
Kp = Kc(RT)^(Δn); where Δn is the change in moles of gas, accounting for pressure in gaseous equilibria.
Effect of concentration changes.
Adding reactants shifts equilibrium towards products; removing reactants shifts it towards reactants, aiming to minimize changes.
Effect of temperature on equilibrium.
For exothermic reactions, increasing temperature shifts equilibrium left (towards reactants); for endothermic, it shifts right (towards products).
Define solubility product constant (Ksp).
Ksp = [cation]^x[anion]^y; represents equilibrium between a sparingly soluble salt and its ions in solution.
Common ion effect.
Adding a common ion decreases solubility of a salt due to Le Chatelier's principle, favoring the formation of the solid.
Define acid and base (Arrhenius).
Acids produce H+ ions in water; bases produce OH- ions in water.
Brønsted-Lowry acid-base theory.
Acids are proton donors; bases are proton acceptors. Conjugate pairs differ by one proton.
Characteristics of weak vs. strong acids.
Strong acids fully ionize in solution, resulting in high concentrations of H+; weak acids only partially ionize.
Describe pH scale.
pH is the negative logarithm of H+ concentration; neutral water has pH = 7, while lower values indicate acidity and higher values indicate basicity.
Buffer solutions and their importance.
Buffers resist pH changes upon dilution or addition of acids/bases, crucial for biological and chemical applications.
Describe ionic equilibrium.
Ionic equilibrium occurs when ions are formed or consumed, establishing dynamic balance between undissociated species and their ions.
Ionization constant of water.
Kw = [H+][OH-] = 1 x 10^-14 at 25°C; temperature-dependent as it reflects self-ionization of water.
Identify conjugate acid-base pairs.
Conjugate pairs differ by one proton; e.g., HCl/Cl- (acid/base), NH4+/NH3, etc.
Calculate Kc from concentration.
Determine equilibrium constant Kc using concentrations of products and reactants at equilibrium.
Dynamic activity in equilibrium.
Equilibrium is not static; there is constant motion with reactants converting to products and vice versa.
Use of catalysts.
Catalysts speed up the reaction by lowering activation energy, affecting the rate but not the equilibrium position.
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