Chemical Equilibrium

Transcription

1 Chemical Equilibrium What is equilibrium? Expressions for equilibrium constants, K eq ; Calculating K eq using equilibrium concentrations; Factors that affect equilibrium; Le Chatelier s Principle

2 What is Equilibrium?

3 This is not Equilibrium?

4 Chemical Equilibrium Consider the following reactions: CaCO 3 (s) + CO 2 (aq) + H 2 O(l) Ca 2+ (aq) + 2HCO - 3 (aq)..(1) and Ca 2+ (aq) + 2HCO - 3 (aq) CaCO3 (s) + CO 2 (aq) + H 2 O(l)..(2) Reaction (2) is the reverse of reaction (1). At equilibrium the two opposing reactions occur at the same rate. Concentrations of chemical species do not change once equilibrium is established.

5 Expression for Equilibrium Constant Consider the following equilibrium system: wa + xb yc + zd K eq = [C] y [D] [A] w [B] z x The numerical value of K eq is calculated using the concentrations of reactants and products that exist at equilibrium.

6 Expression and Value of Equilibrium Constant for a Reaction The expression for K depends on the equation The value of K applies to that equation; it does not depend on how the reaction occurs; Concentrations used to calculate the value of K are those measured at equilibrium.

7 Expressions for Equilibrium Constants Examples: N 2 (g) + 3H 2 (g) 2NH 3 (g); K eq = [NH3 ] [N ][H ] 3 PCl 5 (g) PCl 3 (g) + Cl 2 (g); K eq = [PCl3 ][Cl [PCl ] 5 2 ] CH 4 (g) + H 2 (g) CO(g) + 3H 2 (g); K eq = 3 [CO][H 2] [CH ][H O] 4 2

8 Relationships between chemical equations and the expressions of equilibrium constants The expression of equilibrium constant depends on how the equilibrium equation is written. For example, for the following equilibrium: H 2 (g) + I 2 (g) 2 HI(g); 2 [HI] K eq [H ][I 2 2 ] For the reverse reaction: 2HI(g) H 2 (g) + I 2 (g); K [H ][I ' [HI] ] 2 2 eq 1/ 2 K eq

9 Homogeneous & Heterogeneous Equilibria Homogeneous equilibria: CH 4 (g) + H 2 O(g) CO(g) + 3H 2 (g); CO(g) + H 2 O(g) CO 2 (g) + H 2 (g); Heterogeneous equilibria: CaCO 3 (s) CaO(s) + CO 2 (g); HF(aq) + H 2 O(l) H 3 O + (aq) + F - (aq); PbCl 2 (s) Pb 2+ (aq) + 2 Cl - (aq);

10 Examples: Equilibrium Constant Expressions for Heterogeneous System CaCO 3 (s) CaO(s) + CO 2 (g); K = [CO 2 ] or K = P CO2 ; HF(aq) + H 2 O(l) H 3 O + (aq) + F - (aq); K eq [H3O ][F [HF] - ]

11 PbCl 2 (s) Pb 2+ (aq) + 2Cl - (aq); K = [Pb 2+ ][Cl - ] 2

12 Le Chatelier s Principle states that: When a system at equilibrium is stressed, the equilibrium will shift in the direction that will relieve the stress.

13 What are stresses to an equilibrium? change in pressure change in concentration change in temperature

14 Changes in PRESSURE only affect gases RULE: If the pressure on a system increases, the shift will be towards the side of the eqn. with the LOWER # of moles of gas

15 Changes in PRESSURE How do you figure out the number of moles of gas? Add up the coefficients in the balanced eqn.

16 EXAMPLE 1: 3 H 2 (g) + N 2 (g) 2 NH 3 (g) If P increases shift to RIGHT side because there are 4 moles of gas on left side, only 2 moles of gas on right side.

17 EXAMPLE 2: H 2 (g) + I 2 (g) 2 HI (g) If P increases there will be NO SHIFT because there are 2 moles of gas on the left side & 2 moles of gas on the right side.

18 Changes in CONCENTRATION RULE: If the [concentration ] of substance on one side of eqn. increases, equilibrium will shift towards the other side.

19 EXAMPLE 3: 3 H 2 (g) + N 2 (g) 2 NH 3 (g) If [N 2 ] increases shift towards RIGHT side. If [NH 3 ] increases shift towards LEFT side.

20 Changes in CONCENTRATION RULE: If the [concentration] of substance on one side of eqn. decreases, equilibrium will shift towards that side.

21 EXAMPLE 4: 3 H 2 (g) + N 2 (g) 2 NH 3 (g) If [H 2 ] decreases shift towards LEFT side. If [NH 3 ] is removed shift towards RIGHT side.

22 Changes in CONCENTRATION RULE: If the [concentration] of substance on one side of eqn. increases, substances on same side of eqn. will decrease. Substances on other side will increase.

23 In other words Same side of eqn. = opposite direction Opposite side of eqn. = same direction

24 EXAMPLE 5: 4 HCl (g) + O 2 (g) 2 H 2 O (g) + 2 Cl 2 (g) If [O 2 ] decreases shift towards LEFT side. [HCl] increases (same side as O 2, so opposite direction)

25 EXAMPLE 5: 4 HCl (g) + O 2 (g) 2 H 2 O (g) + 2 Cl 2 (g) If [O 2 ] decreases [H 2 O] decreases (opposite side from O 2, so same direction) [Cl 2 ] decreases

26 EXAMPLE 5: still! 4 HCl (g) + O 2 (g) 2 H 2 O (g) + 2 Cl 2 (g) If [H 2 O] increases shift towards LEFT side [HCl] increases

27 EXAMPLE 5: 4 HCl (g) + O 2 (g) 2 H 2 O (g) + 2 Cl 2 (g) If [H 2 O] increases [O 2 ] increases [Cl 2 ] decreases

28 Changes in TEMPERATURE RULE: The word heat or a # of J, kj, or cal should be treated as another reactant or product. Follow same rules as with concentration.

29 If heat is added to start the rxn. & the temp. increases Heat is located on the left side of the eqn. It is an endothermic rxn. The Keq value increases.

30 If heat is given off by the rxn. & the temp. increases Heat is located on the right side of the eqn. It is an exothermic rxn. The Keq value decreases.

31 EXAMPLE 6: 2 H 2 O (g) 2 H 2 (g) + O 2 (g) + 16 kcal (exothermic) If T increases shift towards LEFT side

32 Changes in Keq value RULE: Only changes in temperature affect the Keq value. Changes in pressure and/or concentration do NOT affect the Keq value.