1. UNIT 10 COLLISION THEORY, RATE OF REACTION, LE CHATELIER PRINCIPLE

2. 1. COLLISION THEORY In order to react, molecules must collide. If the two molecules A and B are to react, they must come into contact with sufficient force so that chemical bonds break. We call such an encounter a collision. The number of effective collisions depend on the activation energy and the orientation of the reactants

3. COLLISION THEORY  states that the rate of a chemical reaction is proportional to the number of collisions between reactant molecules.  The more often reactant molecules collide, the more often they react with one another, and the faster the reaction rate.  In reality, only a small fraction of the collisions are effective collisions. Effective collisions are those that result in a chemical reaction.

4. FACTORS THAT INCREASE THE FREQUENCY OF COLLISIONS: Favorable geometry Increasing temperature Increasing concentration Increasing the surface area Adding a catalyst

5. products no products Favourable Geometry Poor Geometry 1. Favorable geometry –molecules have to line up in a specific way in order for the reactants to form products

6. 2. Increasing temperature Increased temperature will increase the kinetic energy of molecules and increase the number of collisions

7. 3. Increasing concentration Higher concentration increases the probability of successful collisions.

8. 4. Increasing the surface area

9. 5. Adding a catalyst Activation energy is the minimum amount of energy required for a successful collision. Adding a catalyst will decrease the activation energy.

10. EXPLAIN EACH SCENARIO BY USING THE COLLISION THEORY Example 1: 1. Balloon is filled with a mixture of H 2 and O 2 2. H 2 and O 2 gases in a balloon do not react at room temperature 3. Activation energy is too high for the room temperature collisions 4. A small spark ignites causes an explosion

11. Example 2: 1.A candle does not burn at room temperature by itself 2.A match causes the candle to burn 3.The candle continues to burin

12. Example 3: 1. H 2 O 2 decomposes very slowly at room temperature. 2H 2 O 2(aq) → O 2(g) + 2H 2 O (l) 2. Catalyst KI is added and the rate of decomposition increases rapidly.

13. 2. REVERSIBLE REACTIONS Some reactions do not go to completion as we have assumed They may be reversible – a reaction in which the conversion of reactants to products and the conversion of products to reactants occur simultaneously Forward: 2SO 2(g) + O 2(g) → 2SO 3(g) Reverse: 2SO 2(g) + O 2(g) ← 2SO 3(g)

14. REVERSIBLE REACTIONS The two equations can be combined into one, by using a double arrow, which tells us that it is a reversible reaction: 2SO 2(g) + O 2(g) ↔ 2SO 3(g) A chemical equilibrium occurs, and no net change occurs in the actual amounts of the components of the system.

15. REVERSIBLE REACTIONS Molecules of SO 2 and O 2 react to give SO 3. Molecules of SO 3 decompose to give SO 2 and O 2.

16. Even though the rates of the forward and reverse are equal, the concentrations of components on both sides may not be equal An equilibrium position may be shown: A B or A B 1% 99% 99% 1% Note the emphasis of the arrows direction It depends on which side is favored; almost all reactions are reversible to some extent

17. REVERSIBLE REACTIONS 3. Chemical Equilibrium When the rates of the forward and reverse reactions are equal, the reaction has reached a state of balance called chemical equilibrium. At chemical equilibrium, both the forward and reverse reactions continue, but because their rates are equal, no net change occurs in the concentrations of the reaction components.

18. Notice that after a certain time, the concentrations remain constant. This graph shows the progress of a reaction that starts with concentrations of SO 2 and O 2, but with zero SO 3. This graph shows the progress of the reaction that begins with an initial concentration of SO 3, and zero concentrations for SO 2 and O 2.

19. REVERSIBLE REACTIONS Chemical equilibrium is a dynamic state. When the store opens, only the forward reaction occurs as shoppers head to the second floor. Equilibrium is reached when the rate at which shoppers move from the first floor to the second is equal to the rate at which shoppers move from the second floor to the first.

20. 4. LE CHATELIER’S PRINCIPLE The French chemist Henri Le Chatelier (1850- 1936) studied how the equilibrium position shifts as a result of changing conditions Le Chatelier’s principle: If stress is applied to a system in equilibrium, the system changes in a way that relieves the stress Stress: Concentration Temperature Pressure

21. FACTORS AFFECTING EQUILIBRIUM: LE CHÂTELIER’S PRINCIPLE a. Concentration Changing the amount, or concentration, of any reactant or product in a system at equilibrium disturbs the equilibrium. Adding more reactant produces more product, and removing the product as it forms will produce more product

22. Consider the decomposition of carbonic acid (H 2 CO 3 ) in aqueous solution. The system has reached equilibrium. The amount of carbonic acid is less than 1%. H 2 CO 3 (aq) CO 2 (aq) + H 2 O(l) < 1%> 99%

23. Suppose carbon dioxide is added to the system. This increase in the concentration of CO 2 causes the rate of the reverse reaction to increase. Adding a product pushes the reaction in the direction of the reactants. H 2 CO 3 (aq) CO 2 (aq) + H 2 O(l) Add CO 2 Direction of shift

24. Suppose carbon dioxide is removed. This decrease in the concentration of CO 2 causes the rate of the reverse reaction to decrease. Removing a product always pulls a reversible reaction in the direction of the products. H 2 CO 3 (aq) CO 2 (aq) + H 2 O(l) Add CO 2 Direction of shift Remove CO 2 Direction of shift

25. During exercise, the concentration of CO 2 in the blood increases. This shifts the equilibrium in the direction of carbonic acid. The increase in the level of CO 2 also triggers an increase in the rate of breathing. With more breaths per minute, more CO 2 is removed through the lungs. The removal of CO 2 causes the equilibrium to shift toward the products, which reduces the amount of H 2 CO 3. FACTORS AFFECTING EQUILIBRIUM: LE CHÂTELIER’S PRINCIPLE An equilibrium between carbonic acid, carbon dioxide, and water exists in your blood.

26. b. Temperature increasing the temperature causes the equilibrium position to shift in the direction that absorbs heat If heat is one of the products (just like a chemical), it is part of the equilibrium cooling an exothermic reaction will produce more product, and heating it would shift the reaction to the reactant side of the equilibrium: C + O 2(g) → CO 2(g) + 393.5 kJ

27. Heat can be considered to be a product, just like NH 3. Heating the reaction mixture at equilibrium pushes the equilibrium position to the left, which favors the reactants. Cooling, or removing heat, pulls the equilibrium position to the right, and the product yield increases. N 2 (g) + 3H 2 (g) 2NH 3 (g) + heat Add heat Direction of shift Remove heat (cool) Direction of shift

28. c. Pressure changes in pressure only effect gases Increasing the pressure will usually favor the direction that has fewer molecules N 2(g) + 3H 2(g) ↔ 2NH 3(g) For every two molecules of ammonia made, four molecules of reactant are used up – this equilibrium shifts to the right with an increase in pressure A shift will occur only if there are an unequal number of moles of gas on each side of the equation

29. Pressure Initial equilibriumEquilibrium is disturbed by an increase in pressure. When the plunger is pushed down, the volume decreases and the pressure increases. A new equilibrium position is established with fewer molecules.

30. When two molecules of ammonia form, four molecules of reactants are used up. A shift toward ammonia (the product) will reduce the number of molecules. You can predict which way the equilibrium position will shift by comparing the number of molecules of reactants and products. N 2 (g) + 3H 2 (g) 2NH 3 (g) Add pressure Direction of shift Reduce pressure Direction of shift

31. Fritz Haber and Karl Bosch figured out how to increase the yield of ammonia when nitrogen and hydrogen react. Their success came from controlling the temperature and pressure. In which direction did they adjust each factor and why? CHEMISTRY & YOU

32. Fritz Haber and Karl Bosch figured out how to increase the yield of ammonia when nitrogen and hydrogen react. Their success came from controlling the temperature and pressure. In which direction did they adjust each factor and why? An increase in pressure and a decrease in temperature would increase the yield of ammonia by shifting the equilibrium toward the production of ammonia. CHEMISTRY & YOU

33. Catalysts and Equilibrium Catalysts decrease the time it takes to establish equilibrium. However, they do not affect the amounts of reactants and products present at equilibrium.

34. 5. EQUILIBRIUM CONSTANTS: K EQ Chemists generally express the position of equilibrium in terms of numerical values, not just percent These values relate to the amounts (Molarity) of reactants and products at equilibrium This is called the equilibrium constant, and abbreviated K eq

35. consider this reaction (the capital letters are the chemical, and the lower case letters are the balancing coefficient): aA + bB  cC + dD The equilibrium constant (K eq ) is the ratio of product concentration to the reactant concentration at equilibrium, with each concentration raised to a power (which is the balancing coefficient).

36. consider this reaction: aA + bB  cC + dD Thus, the “equilibrium constant expression” has this general form: [C] c x [D] d [A] a x [B] b (brackets: [ ] = molarity concentration) K eq = Note that K eq has no units on the answer; it is only a number because it is a ratio

37. the equilibrium constants provide valuable information, such as whether products or reactants are favored: if K eq > 1, products favored at equilibrium if K eq < 1, reactants favored at equilibrium if K eq = 1, the amounts of product and reactants will be approximately equal