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Dynamic Equilibrium in Chemical Reactions: Understanding K, Q, and Le Chatelier’s Principle

This resource delves into the fundamental concepts of chemical equilibrium, emphasizing the dynamic nature of bond formation and breaking. Explore equilibrium expressions, the reaction quotient (Q), and the equilibrium constant (K) as they relate to reaction direction and system stability. Key topics include Le Chatelier’s Principle, solubility constants (Ksp), acid-base equilibria, and the common ion effect. Additionally, discover practical applications such as titration curves and how external factors influence equilibrium in chemical systems.

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Dynamic Equilibrium in Chemical Reactions: Understanding K, Q, and Le Chatelier’s Principle

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Presentation Transcript

  1. Big Idea #6 “Any bond or intermolecular attraction that can be formed can be broken. These two processes are in dynamic competition, sensitive to initial concentration and external perturbations.”

  2. Concepts: • Equilibrium expression, magnitude of K • Le Chatelier’s Principle • Reaction Quotient • Multistep process ( relationship of equilibrium constant) • Ksp • Common Ion Effect • Acid/Base Equilbria • Buffers • Titration

  3. Equilibrium Review Key topics: Equilibrium expression Equilibrium Constant (Kc) Gas Equilibrium (Kp) Relationship Kc- Kp ICE charts (molar relationships) Reaction Quotient (Q) Le Chatelier’s Principle Solubility Constant (Ksp)

  4. Equilibrium expression and Kc • Coefficients are used as exponents. • Describe concentrations of (aq) and (g) only. • Kc = [products]/[reactants] • Kc > 1 means equilibrium favors forward direction (formation of product) • Kc < 1 describes an equilibrium that favors reverse reaction (formation of reactants) • Kc =1 system is at equilbrium

  5. Relationship: Kc & Kp • Kc = equilibrium concentrations • Kp = equilibrium pressure (gas systems only) • Kp = Kc(RT) Δn(difference in coefficient gaseous products and reactants)

  6. ICE charts • Stoichiometric relationship of reactants: products • Assume reaction occurs in the forward direction. • Some strategies: • Perfect square • Quadratic equation • 5% rule (used when K is very small-compared to initial concentration)

  7. Reaction Quotient (Q) • Used to determine the direction required for a system to achieve equilibrium. • Q<K - reaction must move forward • Q> K – reaction must move in reverse • Q = K – reaction at equilibrium

  8. Le Chatelier’s Principle Factors that disrupt an equilibrium system: • Change in concentration (adding/removing reactants or products) • Change in volume/pressure (for gas equilibrium) • Change in temperature (endo/exo reactions) Catalyst: allows a system to reach equilibrium more quickly but does not alter the equilibrium.

  9. Basic concepts: Acid-Base chemistry & pH • Recognizing acid/base and conjugate base/acid • Calculation of pH, pOH, [H3O+], [OH-] • Calculating pH for solutions of strong acids/base • Ionization constant: Ka, Kb • Polyprotic acid (and associated Kavalues) • pKa, pKb • Acid-Base properties of salts • Predicting direction of acid-base reaction • Types of acid-base reactions • Calculations with equilibrium constants • Titration of acid/base and characteristic titration curves

  10. Solubility constant (Ksp) • Remember, expression does not include (s) salt(s) cation (aq) + anion(aq) • Common ion effect: presence of an ion at the start of the “reaction”. Alters the solubility (think Le Chatelier) • pH and solubility: role of pH may impact the solubility of an insoluble salt based on the common ion effect (ex. Mg(OH)2 enhances by the presence of H+ ions/acidic) • Formation of a precipitate- again use Q when Q> Ksp a precipitate will form

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