Reaction Rates and Chemical Equilibrium

Reaction Rates and Chemical Equilibrium 12-1 12.1 Reaction Rates a measure of how fast a reaction occurs. Some reactions are inherently fast and some are slow 12-2 12.2 Collision Theory In order for a reaction to occur, reactant molecules must collide with proper orientation with enough energy to break existing bonds Only a small fraction of the collisions that do occur meet these requirements. Consider the following reaction that occurs in smog O 3(g) + NO (g) O 2(g) + NO 2(g) Which of the following collisions has a proper orientation? 12-3 1

12.2 Collision Theory Energy Diagrams Figure 12.6 12-4 12.2 Collision Theory Energy Requirements The minimum energy required for a reaction to be initiated. Collisions that have the proper orientation & have at least the minimum E a can convert to products. Reactions with large activation energies tend to be slow - Relatively small fraction of reactants have sufficient energy for an effective collision Reactions with small activation energies tend to be fast - A large fraction of reactants have sufficient energy for an effective collision Short-lived, unstable, high-energy chemical species that must be achieved before products can form. Formed from reactant molecules that collide with the proper orientation and sufficient energy Actual structure is unknown Each reaction has its own reaction diagram, which shows the energy changes as reactants convert to products. 12-5 12.2 Collision Theory Energy Diagrams Figure 12.7 12-6 2

Energy Diagrams The following reaction is an endothermic reaction 2NO 2(g) 2NO (g) + O 2(g) Draw an energy diagram that shows the relative energies of the reactants, products, and the activated complex. Label the diagram with molecular representations of reactants, products, and a possible structure for the activated complex. 12-7 12.3 Conditions That Affect Rxn Rates Increases cause an increase in the number of effective collisions, resulting in increased reaction rate. Steel wool (right) reacts faster than the iron nail (left) More iron atoms are exposed to the oxygen in the atmosphere. 12-8 12.3 Conditions That Affect Rxn Rates Increasing the of one or more reactants increases the number of effective collisions by increasing the total number of collisions but fraction of collisions that are effective remains the same. 12-9 3

12.3 Conditions That Affect Rxn Rates increased T increases rate of the reaction by increasing the number of effective collisions by increasing the total number of collisions (moving faster) increasing the fraction of collisions that are effective (because the average kinetic energy is higher). 12-10 Consider Collision Rates A + B AB Collision rate between (25 C) = 10,000 collisions/second Effective collisions (25 C) = 100 collisions/second How will each of the following affect the total number of collisions and the fraction of effective collisions a) The temperature decreases b) The concentration of reactant B decreases 12-11 12.3 Conditions That Affect Rxn Rates a substance that increases the rate of reaction by changing the mechanism to a reaction with a lower energy of activation. A catalyst is not a reactant or product. It interacts with the reactants, but is not permanently changed during the reaction. Since catalysts are recycled, small amounts are needed and last a long time. Adding an appropriate catalyst increases the number of effective collisions by lowering the activation energy. This also increases the fraction of collisions that are effective. 12-12 4

12.3 Conditions That Affect Rxn Rates Catalytic converters consist of metal catalysts that convert toxic gases (like CO) to harmless ones (like CO 2 ) quickly. Figure 12.11 12-13 12.3 Conditions That Affect Rxn Rates The thousands of enzymes in our bodies act to catalyze specific biological processes. The enzyme sucrase catalyzes the decomposition of sucrose by making bond-breaking easier 12-14 12.3 Conditions That Affect Rxn Rates Cl atoms from CF 2 Cl 2 catalyze the decomposition of ozone in the stratosphere by the following proposed 2-step mechanism O 3(g) + Cl (g) ClO (g) + O 2(g) ClO (g) + O 3(g) Cl (g) + 2 O 2(g) 2 O 3(g) + ClO (g) + Cl (g) ClO (g) + Cl (g) + 3 O 2(g) Net Reaction 2 O 3(g) 3 O 2(g) used early on, and produced later. produced early on and consumed later. 12-15 5

Identifying Intermediates & Catalysts Ethene (H 2 C=CH 2 ) can be converted to ethanol (CH 3 CH 2 OH) by a three-step process. Identify any catalysts and/or intermediates. H 2 C=CH 2 + H 3 O + H 3 C-CH 2 + + H 2 O H 3 C-CH 2 + + H 2 O CH 3 CH 2 OH 2 + CH 3 CH 2 OH 2 + + H 2 O CH 3 CH 2 OH + H 3 O + Net H 2 C=CH 2 + H 2 O CH 3 CH 2 OH 12-16 Catalysis The decomposition of HI is an exothermic reaction 2 HI (g) H 2(g) + I 2(g) Draw an energy diagram for the uncatalyzed reaction. Label reactants and products. Sketch a possible activated complex. Use a dotted line to show the energy changes when platinum metal a catalyst that increases the rate of the reaction, is added to the system. 12-17 12.4 Chemical Equilibrium a state reached by a chemical reaction where there is no net change in the concentrations of reactants and products. Equilibrium is established when reactants are converted to products, and those products are converted back to reactants at an equal rate. Rate of forward reaction occurs at the same rate as the reverse reaction. Both reactants and products exist simultaneously when a system is at equilibrium The concentrations of reactants and products is no longer changing N 2 O 4(g) 2 NO 2(g) 12-18 6

Chemical Equilibrium At what point is equilibrium first reached? N 2 O 4(g) 2 NO 2(g) 12-19 12.4 Chemical Equilibrium A description of equilibrium in terms of favoring products or reactants. When a system is in equilibrium, if there are Reactant favored more reactants than products. Product favored more products than reactants. When there are similar amounts of reactants and products, the position of equilibrium is about in the middle. A value that expresses whether a system at equilibrium favors the products or the reactants. K eq is calculated by using an equilibrium constant expression. For the generic reaction aa + bb cc + dd The equilibrium expression is # $% = ' ( ) * Where [ ] is concentration in molarity +, -. 12-20 N 2(g) + 3 H 2(g) 2 NH 3(g) What is the equilibrium constant expression? What is the value of the equilibrium constant? Are reactants or products favored? 12-21 7

12.5 Equilibrium Constants aa + bb cc + dd Equilibrium constant expression K 01 = 2 3 4 5 6 7 8 9 Table 12.2 12-22 Equilibrium Constants At 25 C a pure sample of N 2 O 4 is placed into a reaction container and allowed to reach equilibrium 2 NO 2(g) N 2 O 4(g) The equilibrium concentrations are determined to be [NO 2 ] = 0.0750 M & [N 2 O 4 ] = 1.25 M 1. Write the equilibrium constant expression for this reaction. 2. Calculate the value of the equilibrium constant at 25 C. 2. Describe the position of equilibrium. 12-23 12.5 Equilibrium Constants Is this mixture of reactants and products at equilibrium? If not which direction will the reaction proceed? 12-24 8

Predicting the Direction of a Reaction S 2 Cl 2(g) + Cl 2(g) 2 SCl 2(g) K eq = 4 1. Write the equilibrium constant expression for this reaction. 2. Given the following molecular picture, is the reaction at equilibrium? If not, which way will it shift? 12-25 12.5 Equilibrium Constants same physical state reactants and products are in the reactants and products are not all in the same physical state Consider the evaporation of bromine in a closed container Br 2(l) Br 2(g) The concentration of bromine vapor, [Br 2 ], at equilibrium is a constant, and is independent of the amount of bromine liquid. Concentrations of liquids and solids are constant, therefore they are left out of the equilibrium constant expression. Only (g) and (aq) phase substances are included in the equilibrium expression. 12-26 Heterogeneous Equilibrium 1. What is the equilibrium constant expression for the decomposition of calcium carbonate? CaCO 3(s) CaO (s) + CO 2(g 2. Write the equilibrium constant expression for the following reactions. a) Mg (s) + CO 2(g) MgO (s) + CO (g) b) PbCl 2(s) Pb 2+ (aq) + 2Cl - (aq) 12-27 9

12.6 Le Chatelier s Principle a system at equilibrium will react to a disruption in equilibrium in such a way to counteract that disruption and reestablish equilibrium. Ways to disrupt a chemical equilibrium Adding or removing a reactant or product Changing the volume of the reaction container - gases only Changing the temperature - changes K eq value 12-28 12.6 Le Chatelier s Principle Fe 3+ (aq) + NCS - (aq) FeNCS 2+ (aq) # $% = ;<=>?@ ; A@ <=> B Reactants Equilibrium mixture If you add either Fe 3+ or NCS - New Equilibrium mixture 12-29 12.6 Le Chatelier s Principle Increased concentrations of reactants or products will shift the equilibrium away from it to consume some of the added substance. Decreased concentrations of reactants or products will shift the equilibrium toward what was removed, to produce more of it. 12-30 10

Changes in Concentration Consider the following system at equilibrium AgI (s) Ag + (aq) + I - (aq) Predict the effect of the following changes when the system is initially in a state of equilibrium. 1. Ag + is removed from the system by adding NaOH. 2. Solid AgNO 3 is added to the system. (AgNO 3 is water soluble) 3. Solid NaI is added to the system. (NaI is water soluble) 12-31 12.6 Le Chatelier s Principle In which direction does the reaction proceed? CaCO 3(s) CaO (s) + CO 2(g) K eq = [CO 2 ] Reducing the volume of the container causes the concentration of ALL gases to increase. The system shifts to reestablish equilibrium concentrations. 12-32 12.6 Le Chatelier s Principle N 2 O 4(g) 2 NO 2(g) Colorless Brown # $% = CD E E In which direction does the reaction proceed? C E D F 12-33 11

12.6 Le Chatelier s Principle 12-34 Changes in Volume How will a decrease in container volume affect each of the following reactions at equilibrium? Explain. 1) 2 NOBr (g) 2 NO (g) + Br 2(g) 2) CuO (s) + H 2(g) Cu (s) + H 2 O (g) 12-35 12.6 Le Chatelier s Principle N 2 O 4(g) 2 NO 2(g) Colorless Brown In which direction does the equilibrium shift when the temperature is increased? decreased? 12-36 12

12.6 Le Chatelier s Principle To predict the effect of temperature on the position of equilibrium, we must know whether a reaction is endothermic or exothermic. We can think of heat as a reactant or a product Endothermic heat + N 2 O 4(g) 2 NO 2(g) Exothermic 2SO 2(g) + O 2(g) 2 SO 2(g) + heat 12-37 Temperature Changes Consider the following equilibrium Fe 3+ (aq) + NCS - (aq) FeNCS 2+ (aq) When the temperature is increased, the solution turns darker, indicating a higher concentration of the FeNCS 2+ product. Is this reaction endothermic or exothermic? 12-38 12.6 Le Chatelier s Principle A catalyst does not affect the position of equilibrium. - It speeds up both the forward and the reverse reaction. A catalyst only increases the rate at which equilibrium is reached. Just because a reactions equilibrium constant suggests the reaction is unfavorable (Keq < 1), does not mean you cannot do things to push the formation of products. N 2(g) + 3 H 2(g) 2 NH 3(g) (exothermic) Under which conditions of temperature and volume can the yield of NH 3 be maximized? High or low temperatures? Large or small volumes? 12-39 13

Le Chatelier s Principle Consider the following reaction that has reached a state of equilibrium at 500 C PCl 5(g) PCl 3(g) + Cl 2(g) endothermic Determine whether each of the following will increase the equilibrium concentration of Cl 2 product. Explain your reasoning 1. Add more PCl 3 gas 2. Remove PCl 3 gas 3. Increase the temperature 4. Reduce the volume 5. Add a catalyst that increases the reaction rate 12-40 Le Chatelier s Principle CO (g) + H 2 O (g) CO 2(g) + H 2(g) exothermic Predict the direction the equilibrium will shift after each stress is applied 1. Add CO (constant V) 2. Remove H 2 O (constant V) 3. Increase volume 4. Increase temperature 5. Add a catalyst 12-41 12.6 Le Chatelier s Principle K eq is temperature dependent How temperature affects K eq depends on reaction -Endothermic reactions increase K eq with increased temperature -Exothermic reactions decrease K eq with increased temperature Is this reaction endothermic or exothermic? N 2(g) K eq = 1.0 10 6 at 1500 K K eq = 6.2 10 4 at 2000 K + O 2(g) 2 NO (g) 12-42 14