1. Chemical Reactions: Energy, Rates, and Equilibrium
Chapter Seven
Chemical Reactions: Energy, Rates, and Equilibrium

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
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3. Goals
1. What energy changes take place during reactions? Be able to explain the factors that influence energy changes in chemical reactions.
2. What is “free energy,” and what is the criterion for spontaneity in chemistry? Be able to define enthalpy, entropy, and free-energy changes, and explain how the values of these quantities affect chemical reactions.
3. What determines the rate of a chemical reaction? Be able to explain activation energy and other factors that determine reaction rate.
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4. Goals Contd.
4. What is chemical equilibrium? Be able to describe what occurs in a reaction at equilibrium and write the equilibrium equation for a given reaction.
5. What is Le Châtelier’s principle? Be able to state Le Châtelier’s principle and use it to predict the effect of changes in temperature, pressure, and concentration on reactions.
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5. 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.
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6. 7.2 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 N2 has a bond dissociation energy 226 kcal/mole, while the single bond in Cl2 has a bond dissociation energy 58 kcal/mole.
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7. 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.
Law of conservation of energy: Energy can be neither created nor destroyed in any physical or chemical change.
Heat of reaction: Represented by DH, is the difference between the energy absorbed in breaking bonds and that released in forming bonds. DH is also known as enthalpy change.
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8. 7.3 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.
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9. 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.
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10. 7.4 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.
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11. (omit)
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.
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12. (omit) DG = DH - TDS H S G (-) favorable (+) favorable
(-) spontaneous always
(+) unfavorable
(-) unfavorable
(+) nonspontaneous always
(-) Low T
(+) High T
(+) Low T
(-) High T
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13. 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.
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14. 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.
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15. 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.
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16. 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.
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17. 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.
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18. 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.
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19. 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.
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20. 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.
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21. 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.
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22. 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.
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23. 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.
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24. 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.
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25. Very large K: Reaction goes essentially to completion.
Very small K: More reactants than products are present at equilibrium.
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26. 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.
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27. 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:
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28. 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.
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29. 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.
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30. 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.
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31. The effects of changing reaction conditions on equilibria are summarized below.
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32. 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.
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33. 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.
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34. 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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35. Key Words Activation energy Bond dissociation energy Catalyst
Chemical equilibrium
Endergonic
Endothermic
Enthalpy
Enthalpy change
Entropy (S)
Equilibrium constant (K)
Exergonic
Exothermic
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36. Key Words Contd. Free-energy change Heat Heat of reaction
Kinetic energy
Law of conservation of energy
Le Châtelier’s principle
Potential energy
Reaction rate
Reversible reaction
Spontaneous process
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