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Chemical Equilibrium & K C

Chemical Equilibrium & K C. PLN 13 & PLN 14. Important Concepts. Equilibrium Constant Reaction Quotient. Introduction To Chemical Equilibrium. No Reaction “Goes to Completion” Reactants are forming products However, products are also forming reactants

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Chemical Equilibrium & K C

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  1. Chemical Equilibrium & KC PLN 13 & PLN 14

  2. Important Concepts • Equilibrium Constant • Reaction Quotient

  3. Introduction To Chemical Equilibrium • No Reaction “Goes to Completion” • Reactants are forming products • However, products are also forming reactants • Reactants are formed at such a slow rate, it’s barely noticeable usually • When the Reaction Appears to Stop • Product and reactant concentrations level off • An equilibrium is reached

  4. Dynamic Equilibrium • The Reaction Only Appears to Have Stopped • The forward reaction is going just as fast as the reverse reaction • A Balance is Reached, but the Reactions are Still Occurring • Dynamic equilibrium

  5. EquilibriumExamples • Ionization of Weak Acids or Weak Bases • Dissolution/Precipitation • BiogeochemicalCycles • Carbon Cycle

  6. Position Of Equilibrium • The Position is How Far the Reaction Proceeds Before Reaching Equilibrium • An Equilibrium that Lies to the Right: • Reaction almost completed • System is mostly products • Equilibrium Lies to the Left: • Reactants barely reacted • System is mostly reactants

  7. Equilibrium Constant, KC • For a Specific Temperature, the Concentration Ratio of Products to Reactants is a Constant. Where: • Shows the Relationship Between the Concentrations • Illustrates How EachProduct orReactantConcentration Can be Manipulated to Affect the Others • (KC Has No Units)

  8. KC

  9. Catalyst • Speeds up the Rate of Reaction • Does NOT alter the equilibrium constant • The equilibrium does not change, the system just reaches equilibrium quicker

  10. Gases • The “Concentration” of a Gas in a Mixture is Measurable by its Partial Pressure • Therefore, While: • The Equilibrium Constant:

  11. Converting kp to kc • Where:

  12. Size of kc or kp • Remember That Equilibrium Constants are Related to the Position of the Equilibrium • Small Numerical Value: • More Reactants • Fewer Products • Large Numerical Value: • Fewer Reactants • More Products

  13. manipulating equilibrium equations • Flipping: • Multiplying by a Coefficient: • Adding Equations Together:

  14. Reaction quotients • The Reaction Quotient is Calculated Using the Same Equation as forthe KC • Is frequently compared to the equilibrium constant to determine if the reaction is at equilibrium • If not, then they may be compared to determine which direction the reaction is going at that certain point • Q = K • The reaction has reached equilibrium • Q < K • Reaction will proceed forward • More products will form • Q > K • Reaction will continue backwards • More reactants will form

  15. Consider this problem • A mixture of 0.0100 moles each of carbon monoxide and water were added to a 100-mL flask at 900 K and the reaction allowed to come to equilibrium. The experimental value of K = 1.56 at this temperature. • Calculate the concentrations of all substances at equilibrium.

  16. Calculating equilibrium concentrations • ICE Method (Initial Change Equilibrium) • Write out equilibrium equation: • Write expression for equilibrium constant: • Make an ICE chart:

  17. Calculating Equilibrium Concentrations (continued) • Substitute the equilibrium values into the equilibrium constant expression: • Solve:

  18. Plug and check!

  19. Systems with small equilibrium constants • Consider: • Now, Calculate all equilibrium concentrations when (1) 5.00 moles of COCl2 and (2) 0.100 moles of COCl2 decompose in a 10.0-L flask.

  20. Part 1 – 5.00 moles of COCl2 • K expression: • ICE chart: • Since K is very small, what can we say about x?

  21. Part 1 (continued) • If K is so small, then the reaction must not produce a lot of its products, therefore we can assume x is relatively small and won’t make a large impact on the concentration of the COCl2 • Let’s rewrite the K expression, substituting our equilibrium values in and making the appropriate assumption:

  22. Part 2 – 0.100 moles of COCl2 • Not Fair to Assume that x is Unimportant in This Case • That Assumption May Only be Made if: • K is very small • ~OR~ • The initial concentration is very high compared to the value of K • As a General Rule: • If 100×K < [A]0, then you can approximate x to be unimportant to the overall value of [A]0

  23. Another Trick • Similar to the Method Used in the Previous Example, if K is Very Large, then You Can Assume (When Calculating Equilibrium Concentrations) that the Reaction Goes to Completion • Example: • If given the number of moles of the reactants, just add them to calculate the moles of the products and assume there will be no reactants remaining at equilibrium (keep in mind that there may be a limiting reactant!)

  24. The method of successive approximations • If, After Plugging in Equilibrium Values from an ICE Chart, You Get Something that Looks Like this: • Where K is Small and Therefore We Can Assume x is Small: • Now Plug this Back in For x in the Original Equation: • Repeat: • Keep Repeating Until x ‘Stabilizes at a Fixed Value (at Step 4 this Case stabilizes at 3.4 × 10‐4) • This Method Works as Long as K < 4C Where:

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