1 / 22

Chemical Equilibrium: Quantitatively

Chemical Equilibrium: Quantitatively. Chapter 12. Equilibrium Constant (Kc). Demonstrates the relationship between concentration of products and reactants at equilibrium. Concentrations are raised to the power of the stoichiometric coefficients. We use only gaseous or aqueous concentrations.

shelby-hensley
Download Presentation

Chemical Equilibrium: Quantitatively

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Chemical Equilibrium: Quantitatively Chapter 12

  2. Equilibrium Constant (Kc) • Demonstrates the relationship between concentration of products and reactants at equilibrium. Concentrations are raised to the power of the stoichiometric coefficients. • We use only gaseous or aqueous concentrations

  3. We omit the concentration of pure substance (either solid or liquid) as they remain constant throughout a reaction • Concentrations are expressed in mol/L

  4. Example: • In the reaction: 2A + 4B ↔ 3C + D The equilibrium constant is calculated using the following expression: Kc = [C]3 [D] . [A]2 [B]4

  5. Example: • At 365.0 oC a mixture at equilibrium in a 6.00 L flask contains 1.62˙moles H2, 0.546 moles N2 and 0.0378 moles NH3. N2 (g) + 3H2 (g)↔ 2NH3 (g)

  6. 1. Calculation of concentration at Equ.

  7. 2. Calculation of the equ. constant of the synthesis of ammonia

  8. 3. Calculation of the equ. constant of the decomposition of ammonia

  9. Example: • At 125.0 oC, a system at equilibrium in a 3.0 L flask contains 0.056 moles I2 and 0.12 moles Br2. What is the concentration of IBr if Kc = 167.28? I2 (g) + Br2 (g)↔ 2IBr (g)

  10. 1. Calc. of the conc. of iodine & bromine at equ

  11. 2. Calc. of the conc of iodine bromine

  12. The value of the constant • Because the constant is a ratio of products over reactants: • If K is GREATER that 1, there are more products at equilibrium • If K is LESS than 1, there are more reactants at equilibrium

  13. Calculating Equilibrium Concentration – ICE Tables • Using Initial concentrations, Change in concentrations and Equilibrium concentrations we can find missing values • We can use the ratio of coefficients in the equation to determine change values in the table

  14. We express concentrations in mol/L • We use equilibrium values to calculate Kc

  15. Example #1: • At a given temperature 8.0 moles of nitrogen oxide and 6.0 moles of oxygen are placed in a 4.0 L container. At equilibrium 3.0 moles of oxygen remain. Calculate Kc. 2NO (g) + O2 (g)↔ 2 NO2 (g)

  16. 1. Calculation of conc. at equilibrium

  17. 2. Recording of Data and use of the ICE Table

  18. Step A: Calculate the change in O2

  19. Step B: Use the stoichiometric ratios to determine the changes of the other substances

  20. Step C: Calculate the concentrations at Equ.

  21. 3. Calculation of the equilibrium constant

  22. Example #2: (use quadratic equation, example pg 317) • At a given temperature, the equilibrium constant for the following reaction is 34. Initially, 3.0 mol of hydrogen and 4.5 moles of iodine were placed in a 2.0 L container. What is the concentration of EACH substance at equilibrium? H2 (g) + I2 (g)↔ 2 HI (g)

More Related