1 / 26

Standard 9: Chemical Equilibrium chapter 18

Standard 9: Chemical Equilibrium chapter 18. Vocabulary : Equilibrium position Equilibrium constant Reversible reaction Rate Concentration Le Chatelier’s Principle. Chemistry. Ms. Siddall. Most reactions are ‘reversible’ Forward reaction : reactants make products

jerod
Télécharger la présentation

Standard 9: Chemical Equilibrium chapter 18

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. Standard 9: Chemical Equilibriumchapter 18 Vocabulary: Equilibrium position Equilibrium constant Reversible reaction Rate Concentration Le Chatelier’s Principle Chemistry. Ms. Siddall

  2. Most reactions are ‘reversible’ Forward reaction: reactants make products e.x. 3O2(g)  2O3(g) Reverse reaction: products make reactants e.x. 2O3(g)  3O2(g) In a reversible reaction the forward and reverse reactions occur at the same time e.x. 3O2(g) 2 O3(g) Reversible Reactions Standard 9b: equilibrium conditions

  3. Summary 1 • What is a reversible reaction?

  4. Reversible reactions reach equilibrium: a balance between reactants and products Conditions of Equilibrium: • rate of forward reaction = rate of reverse reaction • Concentration of reactants and products is constant (does not change) NOTE: Rate = speed Concentration = number of particles or moles example: [HCl] = concentration of HCl 6M HCl = 6mole/L HCl = 6 moles of HCl per liter of solution

  5. Summary 2 • What is equilibrium?

  6. Reversible reaction: X  Y [X] X  Y concentration rate [Y] Y  X time time Concentrations are constant Reaction rates are equal equilibrium

  7. Summary 3 • Describe the part of each graph that illustrates equilibrium conditions.

  8. 9a: Le Chatelier’s Principle Le Chatelier’s Principle • A system in equilibrium will react to relieve stress (change) and re-establish equilibrium • Stress: • Adding reactants or products • Removing reactants or products • Changing temperature • Changing pressure (for gases only)

  9. Summary 4 • According to Le Chatelier’s Principle, what will happen to a system at equilibrium if more reactants or products are added?

  10. Example: N2(g) + 3H2(g) NH3(g) + heat • Stress: Add N2 • Stress relief: • Forward reaction (→) to get rid of N2 • H2 is used up (↓) • NH3 and Heat are produced (↑)

  11. Summary 5: N2(g) + 3H2(g) NH3(g) + heat • Stress: remove N2 • Stress relief: • Which way does equilibrium shift? • What happens to [H2]? • What happens to [NH3]? • What happens to heat?

  12. Stress relief. • Adding products or reactants • Equilibrium shifts to remove addition • Removing products or reactants • Equilibrium shifts to replace what has been removed • Gasses • Equilibrium shifts to produce: • more gas at low pressure • Less gas at high pressure

  13. Summary 6 • Why would a gas equilibrium system produce more gas at low pressure and less gas at high pressure?

  14. Haber Process: N2(g) + 3H2(g) NH3(g) + heat

  15. Haber Process: N2(g) + 3H2(g) 2NH3(g) + heat

  16. Summary 7 • According to Le Chatelier’s Principle: • Increasing reactant concentration will cause: • other reactants to __________? • products to __________? • Decreasing reactant concentration will cause: • other reactants to __________? • products to __________?

  17. A(g) + B(g) AB(g) + heat

  18. Summary 8 A(g) + B(g) AB(g) + heat • Complete the table of equilibrium changes

  19. Equilibrium Constant: Keq HONORS Standard 9c: equilibrium constant • At equilibrium concentrations are constant • Keq represents concentrations of reactants and products at equilibrium • Example: aA + bB  cC + dD • Keq = [C]c[D]d [A]a[B]b

  20. Summary 9 • write Keq expression for the Haber-Bosch Process: N2(g) + 3H2(g) 2NH3(g)

  21. Concentrations calculated in mol/L (M) • Only solutions(aq) & gases(g) are considered • No solids (s) • No liquids (l) Example: 2H2O(l) 2H2(g) + O2(g) Keq = [H2]2[O2]

  22. Summary 10 • Fe(OH)2(aq) + 2HSO3(aq)Fe(SO3)2(aq) + 2H2O(l) • Find Keq

  23. What Keq tells us • If Keq ≤ 1 There are more reactants than products at equilibrium • If Keq ≤ 1/100 There are mostly reactants at equilibrium • If Keq ≥ 1 There are more products than reactants at equilibrium • If Keq ≥ 100 There are mostly products at equilibrium

  24. Summary 11 CO(g) + 2H2(g) CH3OH(g) Keq=290 at 430°C • Write the expression for Keq • Reaction is… (mostly products or reactants?)

  25. Solubility • Ksp is the equilibrium constant for solubility • Example: AgCl(s) Ag+(aq) + Cl-(aq) • Ksp AgCl = 1.77 x 10-10 • Does not really dissolve, mostly solid • Example: AgNO3(s) Ag+(aq) + NO3-(aq) • Ksp AgNO3~ 1 x 1010 • Very soluble

  26. Summary 13 • Write the balanced equation for the dissolving of sodium sulfate. • Write a Ksp expression for the reaction.

More Related