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Understanding Chemical Equilibrium in General Chemistry

Learn about equilibrium constants, Le Châtelier’s Principle, and predicting changes in chemical equilibria in this comprehensive guide. Explore examples, equilibrium concentrations, and the impact of temperature changes.

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Understanding Chemical Equilibrium in General Chemistry

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  1. Ch. 15: Chemical Equilibrium Dr. Namphol Sinkaset Chem 152: Introduction to General Chemistry

  2. I. Chapter Outline • Introduction • The Equilibrium Constant (K) • Le Châtelier’s Principle

  3. I. Introduction • Most reactions are reversible, meaning they can proceed in both forward and reverse directions. • This means that as products build up, they will react and reform reactants. • At equilibrium, the forward and backward reaction rates are equal.

  4. I. Example Equilibrium

  5. II. Equilibrium Concentrations • Equilibrium does not mean that concentrations are all equal!! • However, we can quantify concentrations at equilibrium. • Every equilibrium has its own equilibrium constant.

  6. II. The Equilibrium Constant • equilibrium constant: the ratio at equilibrium of the [ ]’s of products raised to their stoichiometric coefficients divided by the [ ]’s of reactants raised to their stoichiometric coefficients. • The relationship between a balanced equation and equilibrium constant expression is the law of mass action.

  7. II. The Equilibrium Constant • For a general equilibrium aA + bB  cC + dD, the equilibrium expression is:

  8. II. Example Equilibrium Constant

  9. II. Sample Problem • Write the equilibrium constant expression for the reaction: 2H2(g) + O2(g) 2H2O(g).

  10. II. Physical Meaning of K • Large values of K mean that the equilibrium favors products, i.e. there are high [ ]’s of products and low [ ]’s of reactants at equilibrium. • Small values of K mean that the equilibrium favors reactants, i.e. there are low [ ]’s of products and high [ ]’s of reactants at equilibrium.

  11. II. Values of K • Values of K are most easily calculated by allowing a system to come to equilibrium and measuring [ ]’s of the components. • For the equilibrium H2(g) + I2(g) 2HI(g), let’s say equilibrium [ ]’s at 445 °C were found to be 0.11 M, 0.11 M, and 0.78 M for molecular hydrogen, molecular iodine, and hydrogen iodide, respectively.

  12. II. Kc for a H2/I2 Mixture • Note that units are not included when calculating K’s. • Thus, equilibrium constants are unitless.

  13. II. Equilibrium [ ]’s Vs. K • For any reaction, the equilibrium [ ]’s will depend on the initial [ ]’s of reactants or products. • However, no matter how you set up the reaction, the value of the equilibrium constant will be the same if the temperature is the same.

  14. II. Equilibrium [ ]’s Vs. K

  15. III. Predicting Qualitative Changes to Equilibrium • If a system is at equilibrium, what happens when it is disturbed? • Le Châtelier’s Principle allow us to make qualitative predictions about changes in chemical equilibria: • When a chemical system at equilibrium is disturbed, the system shifts in the direction that minimizes the disturbance.

  16. III. Three Ways to Disturb an Equilibrium • We look at three ways that disturb a system at equilibrium: • Adding/removing a reactant or product. • Changing the volume/pressure in gaseous reactions. • Changing the temperature.

  17. III. Adding Reactant/Product • If we add or remove a reactant or product, we’re changing the [ ]’s. • The equilibrium will shift in a direction that will partially consume a reactant or product that is added, or partially replace a reactant or product that has been removed.

  18. III. Adding Reactant

  19. III. Adding Product

  20. III. Changing Volume • Via PV = nRT, changing volume is related to changing pressure and vice versa. • For example, decreasing volume is equivalent to increasing pressure. • Reducing volume of a gaseous reaction mixture shifts the equilibrium in the direction that will, if possible, decrease the # of moles of gas.

  21. III. Changing Pressure

  22. III. Changing Temperature • To determine the effect of changing the temperature, we need to know the heat of reaction, ΔHrxn. • Once we know if a reaction is exo or endo, we can write “heat” as a product or reactant. • We then apply Le Châtelier’s Principle in the same manner as when we considered the add product/reactant case.

  23. III. Changing Temperature

  24. III. Sample Problem • For the reaction N2(g) + 3H2(g) 2NH3(g), ΔHrxn = -46.19 kJ/mole. Which way will the equilibrium shift when each of the following occurs? • NH3 is removed. • The reaction vessel is opened. • The reaction vessel is cooled by 25 °C.

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