SCH4U Unit #2: EQUILIBRIUM

1. SCH4UUnit #2: EQUILIBRIUM Ms. Cornacchione Mon Mar 3rd2014

2. Unit #2: EquilibriumTOPICS • Equilibrium (EQM) Introduction • EQM Law & EQM Constant (Kc) • Qualitative Changes (Le Chatelier’s Principle) • Quantitative Changes & RXN Quotient (Q) • Solubility Equilibria • Learn To Calc: • Find an R or P EQM concentration given another EQM concentration and stoichiometry

3. When is the rate of reaction the fastest? Why? When is the rate of reaction the slowest? Why? 2 NO2 (g) 2 NO (g) + O2 (g) RateFORWARD = k [NO2(g)]x RateREVERSE = k [NO(g)]y[O2(g)]z

4. time time time RateFWD >> RateREV RateFWD > RateREV RateFWD =RateREV A B RateFWD =RateREV RateFWD =RateREV

5. Equilibrium Introduction • Most reactions DO NOT go to completion • Reactions that do not go to completion are REVERSIBLE • Reversible reactions proceed in both forward and reverse directions, existing in a state of EQUILIBRIUM • Chemical equilibrium is the state of a reaction in which all reactants and products have reached constant concentrations in a closed system • Equilibrium is reached when the rate of the forward reaction equals the rate of the reverse reaction

6. + A B Chemical Equilibrium 1. 4. 2. 3. + + + C A D B C + A D B • Products formed • Collisions b/w reactants decrease • Rate of forward reaction DECREASES • Reverse reaction begins • Reaction begins • No products yet formed • High rate of collisions between A & B. • Rate of forward reaction HIGH • Rate of forward reactionEQUAL torate of reverse reaction • Dynamic equilibriumexists • Concentrations constant

7. Chemical Equilibrium + + + + C A A D B B C + A D B Reactants Concentration DYNAMIC CHEMICAL EQUILIBRIUM Products time Forward Reaction Rate Reverse time

8. RateFWD = RateREV Dynamic Equilibrium

9. Dynamic Equilibrium

10. Dynamic Equilibrium In order for a chemical system (such as a reaction or phase change) to be at equilibrium: • It must be a reversible process • It must be taken place in a closed system Example: H2O (l)  H2O (g) in an open system H2O (l) H2O (g) in a closed system

11. Forward and Reverse Reaction For a closed EQM system, the same EQM concentrations are reached regardless of the direction by which EQM was reached

12. Stoichiometry & EQM You can predict the changes in concentration of reactants and products at equilibrium from the coefficients of a balanced chemical equation

13. ICE Table EXAMPLE(Initial, Change, Equilibrium)

14. ICE Table EXAMPLE

15. Equilibrium Systems

16. Equilibrium SystemsPractice Makes Perfect • Study Sample Problem 1 & 2 (Pg 425 & Pg 426) • Do Practice Problems (Page 428)

17. Unit #2: EquilibriumTOPICS • Equilibrium (EQM) Introduction • EQM Law & EQM Constant (Kc) • Qualitative Changes (Le Chatelier’s Principle) • Quantitative Changes & RXN Quotient (Q) • Solubility Equilibria • Learn To Calc: • Write an Expression for KC • Calculate KC • Find an R or P EQM concentration given KC

18. Equilibrium Law & Equilibrium Constant • The Equilibrium Law mathematically describes a chemical system at equilibrium [ ] C • Equilibrium Constant(KC)defines the equilibrium law for a given system. K is a constant for a reaction at a specific temperature. IT IS UNITLESS!!  [ ] concentrations at equilibrium A, B, C, D are reactants and products in (g) or (aq) state ONLY a, b, c, d are the coefficients in the balanced chemical equation

19. Equilibrium Constant EXAMPLES [CO2]3[H2O]4 [C3H8] [O2 ]5 Write expressions for Kc for each of the following reactions: • C3H8(g) + 5O2(g)  3CO2(g) + 4H2O(g) • Ca(s) + 2H2O(ℓ)  Ca(OH)2(aq) + H2(g) • AgNO3(aq) + NaCℓ(s) AgCℓ(s)+ NaNO3(aq) D. Na2CO3(s) + 2HCℓ (aq) 2NaCℓ(aq) + H2O(ℓ) + CO2(g) Kc = [Ca(OH)2][H2] Kc = [NaNO3] [AgNO3] Kc = [NaCl]2[CO2] [HCl]2 Kc =

20. Equilibrium Law & Equilibrium Constant The value of K is a constant for a given reaction at a certain Temperature, regardless of initial concentration

21. Equilibrium Law & Equilibrium Constant The value of KC is a constant for a given reaction at a certain Temperature, regardless of initial concentration

22. KC & EQM Position The equilibrium constant, K, will tell you if the equilibrium position is far to the left (reactants) or far to the right (products)

23. Calculate KC – EXAMPLE 1

24. Calculate KC – EXAMPLE 2

25. Heterogeneous vs. Homogeneous EQM Homogeneous – all R & P in the same state of matter Heterogeneous – R & P present in at least two different states If pure solids (s) or liquids (l) are involved in an EQM system, their concentrations are NOT included in the equilibrium law equation

26. Heterogeneous vs. Homogeneous EQM Write the Equilibrium Law equation for the following reaction: Why not include (s)?

27. Calculating EQM Concentration Example Methane gas reacts with water vapour to produce carbon monoxide gas and hydrogen gas according to this equation: CH4(g) + H2O(g) ↔ CO(g) + 3H2(g) At equilibrium: [CO] = 0.300 M, [H2] = 0.800 M, [CH4] = 0.400 M. If K is 5.67, calculate the concentration of water vapour

28. Equilibrium Law and Equilibrium Constant

29. EQM Constant - Practice Makes Perfect • Study Sample Problems (Pg 430 & 434) • Do Practice Problems (Page 436) and check your answers! • “EQM Law & EQM Constant” Worksheet