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Chapter Seven Chemical Reactions: Energy, Rates, and Equilibrium Fundamentals of General, Organic and Biological Chemistry 5th Edition James E. Mayhugh.

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Presentation on theme: "Chapter Seven Chemical Reactions: Energy, Rates, and Equilibrium Fundamentals of General, Organic and Biological Chemistry 5th Edition James E. Mayhugh."— Presentation transcript:

1 Chapter Seven Chemical Reactions: Energy, Rates, and Equilibrium Fundamentals of General, Organic and Biological Chemistry 5th Edition James E. Mayhugh Oklahoma City University  2007 Prentice Hall, Inc.

2 Prentice Hall © 2007 Chapter Seven 2 Outline ►7.1 Energy and Chemical Bonds ►7.2 Heat Changes during Chemical Reactions ►7.3 Exothermic and Endothermic Reactions ►7.4 Why Do Chemical Reactions Occur? Free Energy ►7.5 How Do Chemical Reactions Occur? Reaction Rates ►7.6 Effects of Temperature, Concentration, and Catalysts on Reaction Rates ►7.7 Reversible Reactions and Chemical Equilibrium ►7.8 Equilibrium Equations and Equilibrium Constants ►7.9 Le Châtelier’s Principle: The Effect of Changing Conditions on Equilibria

3 Prentice Hall © 2007 Chapter Seven Energy and Chemical Bonds ►There are two fundamental kinds of energy. ►Potential energy is stored energy. The water in a reservoir behind a dam, an automobile poised to coast downhill, and a coiled spring have potential energy waiting to be released. ►Kinetic energy is the energy of motion. When the water falls over the dam and turns a turbine, when the car rolls downhill, or when the spring uncoils and makes the hands on a clock move, the potential energy in each is converted to kinetic energy.

4 Prentice Hall © 2007 Chapter Seven Heat Changes During Chemical Reactions ►Bond dissociation energy: The amount of energy that must be supplied to break a bond and separate the atoms in an isolated gaseous molecule. ►The triple bond in N 2 has a bond dissociation energy 226 kcal/mole, while the single bond in Cl 2 has a bond dissociation energy 58 kcal/mole.

5 Enthalpy of Reaction  H  H = Energy of bonds formed – Energy of bonds broken Prentice Hall © 2007 Chapter Seven 5

6 Prentice Hall © 2007 Chapter Seven 6 ►Endothermic: A process or reaction that absorbs heat and has a positive  H. ►Exothermic: A process or reaction that releases heat and has a negative  H. ►Law of conservation of energy: Energy can be neither created nor destroyed in any physical or chemical change. ►Heat of reaction: Represented by  H, is the difference between the energy absorbed in breaking bonds and that released in forming bonds.  H is also known as enthalpy change.

7 Prentice Hall © 2007 Chapter Seven Exothermic and Endothermic Reactions When the total strength of the bonds formed in the products is greater than the total strength of the bonds broken in the reactants, energy is released and a reaction is exothermic.

8 Prentice Hall © 2007 Chapter Seven 8 When the total energy of the bonds formed in the products is less than the total energy of the bonds broken in the reactants, energy is absorbed and the reaction is endothermic.

9 Energy from food kcal = Calorie ►Fat 9 kcal/g ►Protein 4 kcal/g ►Carbohydrate 4 kcal/g ►Ethanol 7.1 kcal/g ►Hamburger patty 245 kcal ►Pizza slice 290 kcal ►Cup of ice cream 270 kcal Prentice Hall © 2007 Chapter Seven 9

10 Nutrition label Prentice Hall ©

11 Prentice Hall © 2007 Chapter Seven Why Do Chemical Reactions Occur? Free Energy ►Spontaneous process: A process that, once started, proceeds without any external influence. ►Entropy: The symbol S is used for entropy and it has the unit of cal/mole·K. The physical state of a substance and the number of particles have a large impact on the value of S.

12 Entropy – a measure of disorder Prentice Hall © 2007 Chapter Seven 12

13 Prentice Hall © 2007 Chapter Seven 13 ►Free energy change (  G): Free energy change is used to describe spontaneity of a process. It takes both  H and  S into account. ►Exergonic: A spontaneous reaction or process that releases free energy and has a negative  G. ►Endergonic: A nonspontaneous reaction or process that absorbs free energy and has a positive  G.

14 Prentice Hall © 2007 Chapter Seven 14 HH SS GG (-) favorable(+) favorable(-) spontaneous always (+) unfavorable(-) unfavorable(+) nonspontaneous always (-) favorable(-) unfavorable(-) Low T (+) High T (+) unfavorable(+) favorable(+) Low T (-) High T  G =  H - T  S

15 Prentice Hall © 2007 Chapter Seven How Do Chemical Reactions Occur? Reaction Rates ►The value of  G indicates whether a reaction will occur but it does not say anything about how fast the reaction will occur or about the details of the molecular changes that takes place. ►For a chemical reaction to occur, reactant particles must collide, some chemical bonds have to break, and new bonds have to form. Not all collisions lead to products, however.

16 Prentice Hall © 2007 Chapter Seven 16 One requirement for a productive collision is that the colliding molecules must approach with the correct orientation so that the atoms about to form new bonds can connect.

17 Prentice Hall © 2007 Chapter Seven 17 Another requirement for a reaction to occur is that the collision must take place with enough energy to break the appropriate bonds in the reactant. If the reactant particles are moving slowly the particles will simply bounce apart. Activation energy (E a ): The amount of energy the colliding particles must have for productive collisions to occur.. The size of the activation energy determines the reaction rate, or how fast the reaction occurs.

18 Prentice Hall © 2007 Chapter Seven 18 ►The lower the activation energy, the greater the number of productive collisions in a given amount of time, and faster the reaction. ►The higher the activation energy, the lower the number of productive collisions, and slower the reaction.

19 Problem 7.6 ►Melting of solid ice to liquid water has:  H = 1.44 kcal/mol  S = cal/mol. K What is  G at -10 o C, 0 o C and 10 o C?  G =  H - T  S Prentice Hall © 2007 Chapter Seven 19

20 At -10 o C  G =  H – T  S  H = 1.44 kcal/mol T = = 263K  S = 5.26 cal/mol. K x 1 kcal/1000 cal = kcal/mol. K  G = 1.44kcal/mol – [(263 K)( kcal/mol. K)] = kcal/mol Not spontaneous Prentice Hall © 2007 Chapter Seven 20

21 At 0 o C  G =  H – T  S  G = 1.44kcal/mol – [(273 K)( kcal/mol. K)] = kcal/mol (or 0.00 kcal/mol to hundredths place) Prentice Hall © 2007 Chapter Seven 21

22 Prentice Hall © 2007 Chapter Seven Effects of Temperature, Concentration, and Catalysts on Reaction Rates Reaction rates increase with temperature. With more energy the reactants move faster. The frequency of collisions and the force with which collisions occur both increase. As a rule of thumb, a 10°C rise in temperature causes a reaction rate to double.

23 Prentice Hall © 2007 Chapter Seven 23 ►A second way to speed up a reaction is to increase the concentrations of the reactants. ►With reactants crowded together, collisions become more frequent and reactions more likely. Flammable materials burn more rapidly in pure oxygen than in air because the concentration of molecules is higher (air is approximately 21% oxygen). ►Hospitals must therefore take extraordinary precautions to ensure that no flames are used near patients receiving oxygen.

24 Prentice Hall © 2007 Chapter Seven 24 ►A third way to speed up a reaction is to add a catalyst—a substance that accelerates a chemical reaction but is itself unchanged in the process. ►A catalyst lowers the activation energy.

25 Prentice Hall © 2007 Chapter Seven 25 The thousands of biochemical reactions continually taking place in our bodies are catalyzed by large protein molecules called enzymes, which promote reaction by controlling the orientation of the reacting molecules. Since almost every reaction is catalyzed by its own specific enzyme, the study of enzyme structure, activity, and control is a central part of biochemistry.

26 Prentice Hall © 2007 Chapter Seven Reversible Reactions and Chemical Equilibrium Imagine the situation if you mix acetic acid and ethyl alcohol. The two begin to form ethyl acetate and water. But as soon as ethyl acetate and water form, they begin to go back to acetic acid and ethyl alcohol. Such a reaction, which easily goes in either direction, is said to be reversible and is indicated by a double arrow in equations.

27 Prentice Hall © 2007 Chapter Seven 27 Both reactions occur until the concentrations of reactants and products reach constant values. The reaction vessel contains both reactants and products and is said to be in a state of chemical equilibrium. A state in which the rates of forward and reverse reactions are the same.

28 Prentice Hall © 2007 Chapter Seven Equilibrium Equations and Equilibrium Constants ►Consider the following general equilibrium reaction: aA + bB + …  mM + nN + … ►Where A, B, … are the reactants; M, N, …. Are the products; a, b, ….m, n, …. are coefficients in the balanced equation. At equilibrium, the composition of the reaction mixture obeys an equilibrium equation.

29 Prentice Hall © 2007 Chapter Seven 29 ►The equilibrium constant K is the number obtained by multiplying the equilibrium concentrations of the products and dividing by the equilibrium concentrations of the reactants, with the concentration of each substance raised to a power equal to its coefficient in the balanced equation. ►The value of K varies with temperature.

30 Prentice Hall © 2007 Chapter Seven 30 ►K larger than 1000: Reaction goes essentially to completion. ►K between 1 and 1000: More products than reactants are present at equilibrium. ►K between 1 and 0.001: More reactants than products are present at equilibrium. ►K smaller than 0.001: Essentially no reaction occurs.

31 Equilibrium Equation K eq Prentice Hall © 2007 Chapter Seven 31

32 Write the equilibrium equation Prentice Hall © 2007 Chapter Seven 32

33 Problem 7.57 Prentice Hall © 2007 Chapter Seven 33

34 Problem 7.59 Having solved for Keq, Determine the [O 2 ] at equilibrium when ►[OF 2 ] = 0.18 mol/L ►[F 2 ] = mol/L Prentice Hall © 2007 Chapter Seven 34

35 Solve for [O 2 ] Prentice Hall © 2007 Chapter Seven 35

36 Problem Solved! Prentice Hall © 2007 Chapter Seven 36

37 Prentice Hall © 2007 Chapter Seven Le Châtelier's Principle: The Effect of Changing Conditions on Equilibria ►Le Châtelier's Principle: When a stress is applied to a system at equilibrium, the equilibrium shifts to relieve the stress. ►The stress can be any change in concentration, pressure, volume, or temperature that disturbs original equilibrium.

38 Prentice Hall © 2007 Chapter Seven 38 ►What happens if the concentration of CO is increased? ►To relieve the “stress” of added CO, according to Le Châtelier’s principle, the extra CO must be used up. In other words, the rate of the forward reaction must increase to consume CO. ►Think of the CO added on the left as “pushing” the equilibrium to the right:

39 Prentice Hall © 2007 Chapter Seven 39 ►The forward and reverse reaction rates adjust until they are again equal and equilibrium is reestablished. ► At this new equilibrium state, the value of [H 2 ] will be lower, because more has reacted with the added CO, and the value of [CH 3 OH] will be higher. ►The changes offset each other, however, so the value of the equilibrium constant K remains constant.

40 Prentice Hall © 2007 Chapter Seven 40 ►Le Châtelier’s principle predicts that an increase in temperature will cause an equilibrium to shift in favor of the endothermic reaction so the additional heat is absorbed. ►You can think of heat as a reactant or product whose increase or decrease stresses an equilibrium just as a change in reactant or product concentration does.

41 Prentice Hall © 2007 Chapter Seven 41 ►Pressure influences an equilibrium only if one or more of the substances involved is a gas. As predicted by Le Châtelier’s principle, increasing the pressure shifts the equilibrium in the direction that decreases the number of molecules in the gas phase and thus decreases the pressure. ►For the ammonia synthesis, increasing the pressure favors the forward reaction because 4 moles of gas is converted to 2 moles of gas.

42 Prentice Hall © 2007 Chapter Seven 42 The effects of changing reaction conditions on equilibria are summarized below.

43 Problem 7.40 The reaction of gaseous H 2 and liquid Br 2 to give gaseous HBr has:  H = kcal/mol  S = 27.2 cal/mol. K a) Write a balanced reaction for this reaction. b) Does entropy increase or decrease in this process? c) It this process spontaneous at all temperatures? d) What is the value of  G at 300K? Prentice Hall © 2007 Chapter Seven 43

44 Problem 7.62 Oxygen can be converted into ozone by the action of lightning or electrical sparks (  H = + 68 kcal/mol): Explain the effect on the equilibrium of: a) increasing the pressure b) increasing the concentration of O 2 c) increasing the concentration of O 3 d) adding a catalyst e) increasing the temperature Prentice Hall © 2007 Chapter Seven 44

45 Prentice Hall © 2007 Chapter Seven 45 Chapter Summary ►The strength of a covalent bond is measured by its bond dissociation energy. ►If heat is released,  H is negative and the reaction is said to be exothermic. If heat is absorbed,  H is positive and the reaction is said to be endothermic. ►Spontaneous reactions are those that, once started, continue without external influence; nonspontaneous reactions require a continuous external influence. ►Spontaneity depends on two factors, the amount of heat absorbed or released in a reaction and the entropy change.

46 Prentice Hall © 2007 Chapter Seven 46 Chapter Summary Contd. ►Spontaneous reactions are favored by a release of heat,  H 0. ►The free-energy change,  G =  H - T  S, takes both factors into account. ►  G 0 indicates nonspontaneity. ►Chemical reactions occur when reactant particles collide with proper orientation and energy. The exact amount of collision energy necessary is the activation energy. ►Reaction rates can be increased by raising the temperature, by raising the concentrations of reactants, or by adding a catalyst.

47 Prentice Hall © 2007 Chapter Seven 47 Chapter Summary Contd. ►At equilibrium, the forward and reverse reactions occur at the same rate, and the concentrations of reactants and products are constant. Every reversible reaction has an equilibrium constant, K. The forward reaction is favored if K>1; the reverse reaction is favored if K<1. ►Le Châtelier’s principle states that when a stress is applied to a system in equilibrium, the equilibrium shifts so that the stress is relieved. ►Applying this principle allows prediction of the effects of changes in temperature, pressure, and concentration.


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