Week 14/Th: Lecture Units 34 & 35

Transcription

1 Week 14/Th: Lecture Units 34 & 35 Unit 33: Colligative Properties Unit 34: Introduction to Equilibria -- Rate of reaction -- Reaction pathway -- Forward / Reverse Rxns. Unit 35: Equilibrium Constants -- Equilibrium Constants & Concentration -- Reaction Quotients & Change Unit 36: Equilibria in Solution Issues: Homework Set 10 due on 08:00AM Levitated Mass LACMA, Los Angeles

2 Week 14/Tu: Meaning of Chemical Equilibrium Chemical Thermodynamics tells us about the direction of spontaneous chemical change based on the energy and entropy changes associated with the process. The Gibbs Free Energy: ΔG o = ΔH o TΔS o ΔG o is negative, rxn. is spontaneous in forward direction ΔG o = 0, reaction stops at equilibrium ΔG o is positive, rxn. is spontaneous in reverse direction If we look at the microscopic level, at equilibrium we would see that chemical reactions don t stop, but that the rates of the forward and reverse reactions are equal. Unfortunately, rates of change (for example, the change of a concentration with time) are not given by Thermodynamics.

3 Week 14/Tu: Kinetics, Path of Reaction Chemical Kinetics is used to describe the rates of reactions and the speed depends on the path between the initial & final states. Recall that we discussed the chemical energy and we found that the NET change does not depend on the path. In general there is not a straight-line path from reactants to products. Height Above Sea level Chemical Energy Reactants Activated Complex at peak & Activation Energy Products time Extent of Reaction or Reaction Coordinate

4 Week 14/Tu: Activated Complex The path between the reactants & products can depend on many things including the velocities & orientation of the molecules when they collide. For example, in the gas phase: N 2 + C à N + C 2 ΔG o = -170 kj Chemical Energy C N BTW: which molecule moves faster, on average? Extent of Reaction or Reaction Coordinate

5 Week 14/Tu: Activated Complex The path between the reactants & products can depend on many things including the shapes & orientation of the molecules when they collide. For example, in the gas phase: N 2 + C à N + C 2 ΔG o = -170 kj BTW: which molecule is the slowest, on average? Chemical Energy N Extent of Reaction or Reaction Coordinate

6 Week 14/Th: Demo Catalysis Decomposition of hydrogen peroxide, ΔG o = -234 kj as written 2 H 2 2 (aq) à 2 H 2 (l) + 2 (g) VSEPR Arrangement tetrahedral tetrahedral linear at xygen 3D structure bent bent linear Bond rder 1 & H H H H Adding a Catalyst provides an alternate pathway with a lower activation energy and allows the reaction to proceed faster. H 2 2 (aq) + I - (aq) I - (aq) + H 2 (l) H 2 2 (aq) + I - (aq) I - (aq) + H 2 (l) + 2 (g)

7 Week 14/Th: Forward & Reverse Reactions Up to this point we have used the idea that reaction mixtures will contain only reactants and will proceed forward if ΔG o is negative. nce we have a mixture of reactants and products then the reverse reaction becomes possible. An example: 2 N N 2 ΔG o = kj Forward Reaction Rate = k f [N] 2 [ 2 ] 1 Reverse Reaction Rate = k r [N 2 ] 2 At equilibrium the reaction rates become equal. k f [N] 2 [ 2 ] = k r [N 2 ] 2 k f k r = [N 2 ]2 [N] 2 [ 2 ] = K EQ

8 Week 14/Th: Equilibrium Constant There is one value of the equilibrium constant for a given chemical reaction we can calculate the number if we know ΔG o but we won t do that here. K EQ = [Product]C [Product] D [Reactant] A [Reactant] B = a number It is also clear that we can mix together reactants and products in an infinite number of combinations that will not equal to the ratio of the equilibrium constant. We describe the (real, lab) mixture by a Reaction Quotient, Q. Q = [Actual Product]C [Actual Product] D [Actual Reactant] A [Actual Reactant] B

9 Week 14/Th: K EQ vs. Q The equilibrium constant describes the thermodynamic balanced system whereas Q describes a system that we put together in the lab. They are both numbers, so there are three familiar possibilities: Q > K EQ Too much product, system reverses Q = K EQ = [Product]C [Product] D [Reactant] A [Reactant] B Q < K EQ Too much reactant, system goes forward [Henry-Louis] Le Chatelier s Principle: an equilibrium (balanced) chemical system will react to change in a way to restore equilibrium. Nothing special here we could just evaluate Q for the change and compare the new value to K EQ

10 Week 14/Th: K EQ vs. Q, Examples a) PCl 5 PCl 3 + Cl 2 (all gases, ΔG o = -64 kj) [Equil] Mole/L Add Cl à out of balance New Equil 3+x 2-x 2-x K EQ = [PCl 3][Cl 2 ] [PCl 5 ] K EQ = [PCl 3 ][Cl 2 ] [PCl 5 ] = 2 *1 3 = 2 3 Q = 2 * 2 3 = 4 3 ( = 2 x )* 2 x 3+ x ( ) ( ) = 2 3

11 Week 14/Th: K EQ vs. Q, Examples b) 2 S S 3 (all gases) [Equil] Mole/L K EQ = [S 3 ] 2 [S 2 ] 2 [ 2 ] = * 5 = 5 Add S à out of balance New Equil 2-2x 5-x 5+2x K EQ = [S 3 ] 2 [S 2 ] 2 [ 2 ] = Q = * 5 = 5 4 ( 5+ 2x) 2 ( 2 2x) 2 ( 5 x) = 5

12 Week 14/Th: K EQ vs. Q, Examples 1) C (g) + 2H 2 (g) CH 3 H (g) ΔH o = -18 kj a) increase pressure of container à Forward b) increase temperature of container à Reverse c) add more C (g) à Forward 2) 2 N (g) + Cl 2 (g) 2 NCl (g) ΔH o = -77 kj a) remove N à Reverse b) increase temperature of container à Reverse 3) 2C (s) + 2 (g) 2C (g) ΔH o = -220 kj a) add more C (s) à No Effect, WHY? b) increase pressure of container à Reverse