Chemical Equilibrium

1. Chemical Equilibrium • Collision theory • Rates of reactions • Catalysts • Reversible reactions • Chemical equilibrium • Le Chatelier’s Principle • Concentration • Temperature • Volume • Catalysts

2. A. Collision Theory • Reaction rate depends on the collisions between reacting particles. • Successful collisions occur if the particles... • collide with each other • have the correct orientation • have enough kinetic energy to break bonds

3. A. Collision Theory • Particle Orientation Required Orientation Unsuccessful Collisions Successful Collision

4. Exothermic Endothermic Energy Energy Time Time Activation energy: minimum energy required for a reaction to occur Activation energy Energy of reaction

5. Ea A. Collision Theory • Activation Energy • depends on reactants • low Ea = fast rxn rate

6. 16.2: Rates of Reactions Chemical kinetics: the study of the rate (the speed) of a reaction Rate of a chemical reaction depends on: 1. SURFACE AREA 2. CONCENTRATION of reactants 3. TEMPERATURE (T) of reactants 4. Presence/absence of a CATALYST

7. SURFACE AREA • Surface Area • high SA = fast rxn rate • more opportunities for collisions • Increase surface area by… • using smaller particles • dissolving in water

8. Effect of Concentration on Rate Concentration: • increasing concentration of reactants results in more collisions. • More collisions = increased rate of reaction

9. Effect of Temperature on Rate Temperature: • Increasing T increases particle speed. • Faster reactants means more collisions have the activation energy, which increases the rate of the reaction.

10. Temperature Analogy: 2-car collision 5 mph “fender bender” 50 mph “high-speed crash”

11. Energy Time Effect of Catalysts on Rate A catalyst: • A chemical that influences a reaction, but is not consumed in the reaction. (It can be recovered unchanged at the end of the reaction.) • Lowers the activation energy of the reaction. Activation energy Activation energy with catalyst

12. Catalysts • Enzyme Catalysis

13. 16.1: Reversible Reactions * Thus far, we have considered only one-way reactions: A + B → C + D Some reactions are reversible: • They go forward (“to the right”) : A + B → C + D and backwards (“to the left”) : A + B ← C + D • Written with a two-way arrow: A + B ↔ C + D Examples: • Boiling & condensing • Freezing & melting

16. Reversible Reactions • At chemical equilibrium there is no net change in the actual amounts of the components of the system. • And although the rates of the forward & reverse rxns are equal at chemical equilibrium, the concentrations of the components on both sides of the chem-ical eqn are not necessarily the same. • *In fact they can be dramatically different.

17. Consider a set of escalators as being like the double arrows in a dynamic equilibrium. • The # of people using the up escalator must be the same as the # of people using the down escalator for the # of people on each floor to remain at equilibrium • However, the # of people upstairs do not have to equal the # of people downstairs • Just the transfer between floors must be consistent

18. Examples of irreversible reactions: • Striking a match / burning paper • Dropping an egg • Cooking (destroys proteins)

19. C C C D D D + + + + + + A A A B B B 16.3: Chemical Equilibrium For a reversible reaction, when the forward rateequals the backward rate, a chemical equilibrium has been established. • Both the forward and backward reactions continue, but there is a balance of products “un-reacting” and reactants reacting. A + B ↔ C + D

20. Chemist’s generally express the position of equilibrium in terms of numerical values These values relate the amounts of reactants to products at equilibrium Consider this hypothetical rxn… wA + xB yC + zD Equilibrium Expression • Where “w” mols of reactant A and “x” mols of reactant B react to give “y” mols of product C and “z” mols of productD at equil.

21. We can write a mathematical expression to show the ratio of product concs to reactant concs called anequilibrium expression [C]y [D]z [A]w [B]x Equilibrium Expression • The concentration of each substance is raised to a power equal to the # of mols of that substance in the balanced rxn eqn. • The square brackets indicate concentration in Molarity (mol/L)

22. The resulting ratio of the equilibrium is called theequilibrium constant or Keq The Keq is dependent on the temp If the temp changes so does the Keq Keq= [C]y [D]z [A]w [B]x Equilibrium Expression NOTE: this is only for gases!!!

23. Equilibrium constants provide valuable chemical information They show whether products or reactantsare favored in a rxn always written as a ratio of products over reactants a value of Keq > 1 means that products are favored Keq < 1 than reactants are favored Equilibrium Constant

24. Keq > 1 products favored at equil Keq < 1 reactants favored at equil

25. Sample Problem 1 Dinitrogen tetroxide (N2O4), a colorless gas, and nitrogen dioxide (NO2), a brown gas, exist in equilibrium with each other according to the following eqn: N2O4(g) 2NO2(g) A liter of gas mixture at 10C at equilibrium contains 0.0045mol N2O4 & 0.030 mol NO2. Write the Keq expression and calculate Keq for the reaction.

26. Known: [N2O4] =.0045mol/L [NO2] =.030mol/L Unknown: Keq expression = ? Keq = ? Analyze: list what we know • At equilibrium, there is no net change in the amount of N2O4 or NO2 at any given instant

27. The only product of the rxn is NO2, which has a coefficient of 2 in the balanced eqn The only reactant N2O4 has a coefficient of 1 in the balanced eqn The equilibrium expression is: [N2O4]1 [.030M]2 Keq= [.0045M]1 Calculate: solve for unknowns [NO2]2 Keq= • Keq is equal to: Keq= 0.20 • Keq < 1, therefore rxn doesn’t favor products

28. A mixture at equilibrium at 827°C contains 0.552 M CO2, 0.552 M H2, 0.448 M CO, and 0.448 M H2O. CO2(g)+ H2(g)<==> CO(g) + H2O(g) Write the equilibrium expression for the above rxn. Calculate Keq at this temp? More CO2 is added to the system, which direction will the reaction shift? Are the reactants or products favored in this reaction? Classwork:

29. * Le Chatelier’s Principle is about reducing stress – a stress applied to a chemical equilibrium Relax! Reduce stress brought on by chemical equilibrium with me, Henri Le Chatelier! (1850 – 1936)

30. 16.4: Le Chatelier’s Principle Le Chatelier’s Principle: • When a stress is applied to a system (i.e. reactants and products) at equilibrium, the system responds to relieve the stress. • The system shifts in the direction of the reaction that is favored by the stress. • A stress is a change in: • Concentration • Temperature • Volume

31. 16.5: Stress: Change Concentration Ex: Co(H2O)62+ + 4 Cl1- ↔ CoCl42- + 6 H2O (pink) (blue) StressResult Add Cl1- Forward rxn favored Shifts forward to reduce extra Cl1- More CoCl42- will form Add H2OBackward rxn favored Shifts backward to reduce extra H2O More Co(H2O)62+ will form

32. 16.7: Stress: Change Temperature Ex: heat + Co(H2O)62+ + 4 Cl1- ↔ CoCl42- + 6 H2O (pink) (blue) This reaction is endothermic. For Le Chatelier’s principle, consider “heat” as a chemical. StressResult Increase T Forward rxn favored; shifts forward to reduce extra heat More CoCl42- will form Decrease T Backward rxn favored; shifts backward to replace “lost” heat More Co(H2O)62+ will form

33. 16.6: Stress: Change Volume Ex: 1 N2 (g) + 3 H2(g) ↔ 2 NH3(g) (1 + 3 = 4 moles of gas) ↔ (2 moles of gas) StressResult Decrease V Forward rxn favored; shifts forward to side with fewer moles of gas (reduces # of molecules packed into this smaller volume) Increase V Backward rxn favored; shifts backward to side with more moles of gas (to fill the larger volume with more molecules)

34. 16.7: Catalysts & Equilibrium MnO2 Ex: 2 H2O2 (aq) ↔ 2 H2O (l) + O2 (g) • Since a catalyst increases the forward and backward rates equally, it will not shift the equilibrium.