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

There are many questions about chemical reactions 2 Au (s) + 3 H2O (l)  Au2O3 (s) + 3 H2 (g) This reaction is balanced but does not occur. Your gold jewelry is safe in the shower. Why do reactions occur? If it is balanced does that mean it will occur? Prentice Hall © 2007 Chapter Seven

Chapter 7 Goals 1. What Energy changes take place during a reaction? Explain the 2 factors that influence energy in a chemical reaction 2. What is “Free Energy,” and what is the criterion for spontaneity? Define enthalpy, entropy, and free energy change and explain how these values affect a chemical reaction 3. What determines the rate of a chemical reaction? Explain activation energy, and the other factor that determines the rate. Prentice Hall © 2007 Chapter Seven

Chapter 7 Goals 4. What is a chemical equilibrium? What occurs in a reaction at equilibrium and write the equilibrium constant expression. 5. What is Le Châtelier’s Principle? State Le Châtelier’s Principle, use it to predict the effect of temperature, pressure, and concentration on a reaction. Prentice Hall © 2007 Chapter Seven

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 Prentice Hall © 2007 Chapter Seven

Chapter 7 2 factors determine if a reaction goes to products Total Energy Is the reaction fast or slow Can we influence the 2 factors above? Prentice Hall © 2007 Chapter Seven

7.1 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. Prentice Hall © 2007 Chapter Seven

1. What Energy changes take place during a reaction 1. What Energy changes take place during a reaction? Explain the 2 factors that influence energy in a chemical reaction The first factor is Heat. The term in chemistry is called enthalpy, and it’s symbol is “H”. Mathematically, it is always treated as ΔH (Delta H, mean change in heat). We always compare the amount of heat at the beginning of a reaction, we call it bond disassociation, and compare it to the amount of heat at the end, we call it bond forming, of a chemical reaction; (the second factor is “S”, slide 19). Prentice Hall © 2007 Chapter Seven

7.2 Heat Changes During Chemical Reactions The first step in any reaction is 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 second step will be Bond Formation energy. The triple bond in N2 has a bond dissociation energy 226 kcal/mole, while the single bond in Cl2 has a bond dissociation energy 58 kcal/mole. Prentice Hall © 2007 Chapter Seven

Measure of Heat in a reaction We always want to know if the total heat is: Endothermic: A process or reaction that absorbs heat and has a positive DH. Exothermic: A process or reaction that releases heat and has a negative DH. Examples of each follow Prentice Hall © 2007 Chapter Seven

7.3 Exothermic and Endothermic Reactions The seconds step is bond formation. 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. Prentice Hall © 2007 Chapter Seven

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. Prentice Hall © 2007 Chapter Seven

A thermite reaction is: Fe2O3 + Al  Fe + Al2O3 ΔH = -851.3 KJ/mol Is this endothermic or exothermic? Reactions that give off heat are _________________? Figure: 07-00UN05 Title: Exothermic reactions Caption: The so-called “thermite” reaction of aluminum with iron(III) oxide is so strongly exothermic that it melts iron. Notes: An exothermic reaction is one that releases heat. Keywords: exothermic, aluminum, iron(III) oxide exothermic

When potassium is added to water contained in a beaker, the reaction shown below occurs, and the beaker feels hot to the touch. K(s) + 2 H2O(l)  2 KOH(aq) + H2(g) This reaction is endothermic and H = – endothermic and H = + exothermic and H = – exothermic and H = +

When potassium is added to water contained in a beaker, the reaction shown below occurs, and the beaker feels hot to the touch. K(s) + 2 H2O(l)  2 KOH(aq) + H2(g) This reaction is endothermic and H = – endothermic and H = + exothermic and H = – exothermic and H = +

This desert has 550 kcal of potential energy, heat releasing calories. Food is “potential energy.” Does our body use food in an endothermic or exothermic manner? Figure: 07-00UN08 Title: Calories in food Caption: Eating this dessert gives your body 550 Calories. Burning the dessert in a calorimeter releases 550 kcal as heat. Notes: The caloric value in food is given in calories where 1 Cal= 1000cal = 1 kcal. Keywords: calorimeter, calories

The second influence on energy is called Spontaneity. The symbol is S. We’ve seen how heat is a factor in a chemical reaction (bond breaking, bond forming). Heat releasing (exothermic) reactions are often “favorable” to us. However, there is one more factor in a chemical reaction. 1. What Energy changes take place during a reaction? Explain the 2 factors that influence energy in a chemical reaction The second influence on energy is called Spontaneity. The symbol is S. Prentice Hall © 2007 Chapter Seven

Once at the top will water flow down Figure: 07-00UN07 Title: Spontaneous reactions Caption: Events that lead to lower energy tend to occur spontaneously. Thus, water always flows down a waterfall, not up. Notes: Events that lead to lower energy tend to occur spontaneously. Keywords: spontaneous reactions

7.4 Why Do Chemical Reactions Occur? Free Energy Spontaneous process: A process that, once started, proceeds without any external influence. Spontaneity does not care about how long it takes, only once it starts, it will go. For instance, it takes a long time for a car to rust, though it is spontaneous. Once gasoline is ignited, it will continue to burn on it’s own. We measure spontaneity by measuring the “disorder” of a system. Once at the top will water flow down? Prentice Hall © 2007 Chapter Seven

7.4 Why Do Chemical Reactions Occur? Free Energy 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. An example: solid CO2 has a lower entropy value then gaseous CO2…why? Prentice Hall © 2007 Chapter Seven

The beaker on the right has a more “+” entropy value If I shake the beaker on the right, will they ever line up like the beaker on the left? The beaker on the right has a more “+” entropy value Prentice Hall © 2007 Chapter Seven

7.4 Why Do Chemical Reactions Occur? Free Energy 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. An example: solid CO2 has a lower entropy value then gaseous CO2…why? Gas molecules are more disordered then solids. Prentice Hall © 2007 Chapter Seven

Entropy In 7.4 your book talks about melting ice. Taking ice from the freezer and letting it melt on the table is an endothermic process—so in your head, you might think that this is “Not Favorable,” in the sense that burning a match is heat “favorable.” However, the ice does melt. It melts because liquid molecules are more disordered then solid molecules. At room temperature, the ice cube will not only melt but Evaporate into a gas…why? Going from order to disorder is a powerful driving factor in a chemical reaction. Prentice Hall © 2007 Chapter Seven

Entropy Is this an entropy favorable, S = +, reaction? Dynamite, better known as TNT 2 C7H5N3O6 (s)  3 N2 (g) + 5 H2O (g) + 7 CO (g) + 7 C (g) It goes from 2 moles of a soild to 22 moles of a gas. Prespective: it starts out as a cup of solid, and ends up as 2083 cups of gas Prentice Hall © 2007 Chapter Seven

When potassium is added to water contained in a beaker, the reaction shown below occurs, and the beaker feels hot to the touch. K(s) + 2 H2O(l)  2 KOH(aq) + H2(g) During this reaction entropy decreases and S = – entropy decreases and S = + entropy increases and S = – entropy increases and S = +

When potassium is added to water contained in a beaker, the reaction shown below occurs, and the beaker feels hot to the touch. K(s) + 2 H2O(l)  2 KOH(aq) + H2(g) During this reaction entropy decreases and S = – entropy decreases and S = + entropy increases and S = – entropy increases and S = +

Both Heat and Entropy affect the Total Energy of a chemical reaction. 2. What is “Free Energy,” and what is the criterion for spontaneity? Define enthalpy, entropy, and free energy change and explain how these values affect a chemical reaction Both Heat and Entropy affect the Total Energy of a chemical reaction. Prentice Hall © 2007 Chapter Seven

Free energy change (DG): Free energy change is used to describe spontaneity of a process. It takes both DH and DS 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. Prentice Hall © 2007 Chapter Seven

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

When potassium is added to water contained in a beaker, the reaction shown below occurs, and the beaker feels hot to the touch. K(s) + 2 H2O(l)  2 KOH(aq) + H2(g) This reaction is nonspontaneous and G = – nonspontaneous and G = + spontaneous and G = – spontaneous and G = +

When potassium is added to water contained in a beaker, the reaction shown below occurs, and the beaker feels hot to the touch. K(s) + 2 H2O(l)  2 KOH(aq) + H2(g) This reaction is nonspontaneous and G = – nonspontaneous and G = + spontaneous and G = – form heat a gas and ions spontaneous and G = +

When is it Spontanious? 7.40. For the reaction: H2 + Br2  2HBr ∆H = -17.4 kcal/Kmol, ∆S = 27.2 cal/Kmol Is this reaction spontaneous at all temperatures? No At what temperature is the reaction spontaneous? When ∆G = 0, anything above 0 is spontaneous, so solve for T; ∆G = ∆H – T ∆S, 0 = -17.4 kcal/Kmol– T(.0272kcal/Kmol) = 640. K or 367ºC Prentice Hall © 2007 Chapter Seven

3. What determines the rate of a chemical reaction 3. What determines the rate of a chemical reaction? Explain activation energy, and the other factor that determines the rate. Prentice Hall © 2007 Chapter Seven

7.5 How Do Chemical Reactions Occur? Reaction Rates The value of DG 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. Prentice Hall © 2007 Chapter Seven

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. Prentice Hall © 2007 Chapter Seven

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. Prentice Hall © 2007 Chapter Seven

Activation energy (Ea): 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. 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. Prentice Hall © 2007 Chapter Seven

Shown is an energy diagram for a reaction with a small activation energy and products having less energy than reactants. This reaction is nonspontaneous and fast. nonspontaneous and slow. spontaneous and fast. spontaneous and slow.

Shown is an energy diagram for a reaction with a small activation energy and products having less energy than reactants. This reaction is nonspontaneous and fast. nonspontaneous and slow. spontaneous and fast. spontaneous and slow.

Draw a reaction diagram for: A reaction with a ∆G of -50 kcal/Kmol and Ea of 25kcal/mol. Label axis, and if the reaction is endergonic or exergonic. -50 kcal/mol 25 kcal/mol

7.6 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. Prentice Hall © 2007 Chapter Seven

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. Prentice Hall © 2007 Chapter Seven

A catalyzed reaction has a lower activation energy. 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 catalyzed reaction has a lower activation energy. Prentice Hall © 2007 Chapter Seven

Figure: 07-05 Title: Catalytic converters Caption: A catalytic converter. Notes: Exhaust gases from an automobile pass through a two-stage catalytic converter. In one stage, carbon monoxide and unburned hydrocarbons are converted to carbon dioxide and water. In the second stage, NO is converted to nitrogen gas and oxygen gas. Keywords: catalyst, reaction rate, activation energy, catalytic converter

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. Prentice Hall © 2007 Chapter Seven

When a catalyst is added to a reaction to increase its rate of reaction the activation energy is lowered and G becomes more negative. activation energy is lowered and G remains unchanged. activation energy is raised and G becomes more positive. activation energy is raised and G remains unchanged.

When a catalyst is added to a reaction to increase its rate of reaction the activation energy is lowered and G becomes more negative. activation energy is lowered and G remains unchanged. activation energy is raised and G becomes more positive. activation energy is raised and G remains unchanged.

Figure: 07-06UN18 Title: Key concepts problem 7.20 Caption: Two curves are shown on the energy diagram. Which represents the faster reaction and which the slower? Which represents a spontaneous reaction? Notes: The faster reaction is the one with the smaller activation energy which is the reaction indicated by the blue line. The spontaneous reaction should have a negative delta G which would be the reaction indicated by the red line. Keywords: energy diagram, activation energy, endergonic, exergonic, spontaneity, spontaneous process Two curves are shown on the energy diagram. Which represents the faster reaction and which the slower? Which represents a spontaneous reaction?

Figure: 07-06UN18 Title: Key concepts problem 7.20 Caption: Two curves are shown on the energy diagram. Which represents the faster reaction and which the slower? Which represents a spontaneous reaction? Notes: The faster reaction is the one with the smaller activation energy which is the reaction indicated by the blue line. The spontaneous reaction should have a negative delta G which would be the reaction indicated by the red line. Keywords: energy diagram, activation energy, endergonic, exergonic, spontaneity, spontaneous process The faster reaction is the one with the smaller activation energy which is the reaction indicated by the blue line. The spontaneous reaction should have a negative delta G which would be the reaction indicated by the red line.

Figure: 07-06UN19 Title: Key concept problem 7.21 Caption: Two curves are shown in the following energy diagram. Which represents the catalyzed reaction? Notes: Catalysts lower the activation energy, so the curve with the lower hump is the catalyzed reaction. Keywords: energy diagram, catalyst, activation energy Two curves are shown in the following energy diagram. Which represents the catalyzed reaction?

Figure: 07-06UN19 Title: Key concept problem 7.21 Caption: Two curves are shown in the following energy diagram. Which represents the catalyzed reaction? Notes: Catalysts lower the activation energy, so the curve with the lower hump is the catalyzed reaction. Keywords: energy diagram, catalyst, activation energy Catalysts lower the activation energy, so the curve with the lower hump is the catalyzed reaction.

You’ve just digested a lot about chemical reactions. Free Energy, ΔG, tells us if we make products or not. The Eact gives us the rate, or how fast. To complete this chapter, we’re going to look at what commonly occurs in many chemical reactions, that is, often there isn’t a large ΔG, and the stability of the products and reactants is about the same, AND there is usually a very low Eact. So what happens is, it is about as easy to make products as reactants Prentice Hall © 2007 Chapter Seven

Be able to answer the following question in the next section 4. What is a chemical equilibrium? What occurs in a reaction at equilibrium and write the equilibrium constant expression. Prentice Hall © 2007 Chapter Seven

What occurs in a reaction at equilibrium. When the number of people moving up is the same as the number of people moving down, the number of people on each floor remains constant, and the two populations are in equilibrium. Equilibrium occurs when the forward and reverse reactions are the same. Prentice Hall © 2007 Chapter Seven

7.7 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; like people on the escalator. Prentice Hall © 2007 Chapter Seven

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. Prentice Hall © 2007 Chapter Seven

7.8 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. Prentice Hall © 2007 Chapter Seven

The value of K varies with temperature. 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 each substance raised to a power equal to its coefficient in the balanced equation. The value of K varies with temperature. Prentice Hall © 2007 Chapter Seven

Figure: 07-06UN05 Title: Key concept problem 7.14 Caption: The pictures shown represent two similar reactions that have achieved equilibrium. Which has the larger equilibrium constant, and which has the smaller equilibrium constant? Notes: In the first box very few product has formed and there are a lot of reactants left over, therefore the K is small. In the second box a lot of product has formed and there are very few reactants left, therefore the K is large. Keywords: equilibrium, equilibrium constant, equilibrium equation The pictures shown represent two similar reactions that have achieved equilibrium. Which has the larger equilibrium constant, and which has the smaller equilibrium constant?

The pictures below represent four similar reactions that have achieved equilibrium. A atoms are unshaded. B, C, D, and E atoms are shaded. Which reaction has the largest equilibrium constant? A2 + B2  2 AB A2 + C2  2 AC A2 + D2  2 AD A2 + E2  2 AE

The pictures below represent four similar reactions that have achieved equilibrium. A atoms are unshaded. B, C, D, and E atoms are shaded. Which reaction has the largest equilibrium constant? A2 + B2  2 AB A2 + C2  2 AC A2 + D2  2 AD A2 + E2  2 AE

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. Prentice Hall © 2007 Chapter Seven

7.9 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. Prentice Hall © 2007 Chapter Seven

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: Prentice Hall © 2007 Chapter Seven

The forward and reverse reaction rates adjust until they are again equal and equilibrium is reestablished. At this new equilibrium state, the value of [H2] will be lower, because more has reacted with the added CO, and the value of [CH3OH] will be higher. The changes offset each other, however, so the value of the equilibrium constant K remains constant. Prentice Hall © 2007 Chapter Seven

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. Prentice Hall © 2007 Chapter Seven

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. Prentice Hall © 2007 Chapter Seven

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

Practice 7.83: For the reaction: 2Al + 3Cl2  2AlCl3 + 336.6 kcal/mol How much heat is released when 5.00 grams of Al reacts with excess chlorine? Prentice Hall © 2007 Chapter Seven

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. Prentice Hall © 2007 Chapter Seven

Chapter Summary Contd. Spontaneous reactions are favored by a release of heat, H <0, and an increase in entropy, S >0. The free-energy change, G = H - T S, takes both factors into account. G<0 indicates spontaneity, 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. Prentice Hall © 2007 Chapter Seven

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. Prentice Hall © 2007 Chapter Seven

End of Chapter 7 Prentice Hall © 2007 Chapter Seven