7 7-1 © 2003 thomson learning, inc. all rights reserved bettelheim, brown, and march general,...

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7-1© 2003 Thomson Learning, Inc.All rights reserved

Bettelheim,Bettelheim,Brown, and MarchBrown, and March

General, Organic, and General, Organic, and Biochemistry, 7eBiochemistry, 7e

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Chapter 7Chapter 7

Reaction Rates and Reaction Rates and Chemical EquilibriumChemical Equilibrium

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Chemical KineticsChemical Kinetics• Chemical kinetics:Chemical kinetics: the study of the rates of

chemical reactions• consider the reaction that takes place when

chloromethane and sodium iodide are dissolved in acetone; the net ionic equation for this reaction is

• to determine the rate of this reaction, we measure the concentration of iodomethane at periodic time intervals, say every 10 minutes

CH3-Cl I- CH3-I Cl-+ +Chloro-methane

Iodo-methane

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Chemical KineticsChemical Kinetics• the rate of reaction is the increase in concentration of

iodomethane divided by the time interval• for example, the concentration might increase from 0

to 0.12 mol/L over a 30 minute time period• the reaction rate over this period is

• this unit is read mole per liter per minute

30 min(0.12 mol CH3I/L) - (0 mol CH3I/L)

=0.0040 mol CH3I/L

min

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Reaction RatesReaction Rates• The rates of chemical reactions are affected by

the following factors• molecular collisions• activation energy• nature of the reactants• concentration of the reactants• temperature• presence of a catalyst

• On the following screens, we examine these factors one at a time

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Molecular CollisionsMolecular Collisions• in order for two species, A and B (they may be

molecules or ions), to react, they must first collide• it is possible to calculate how many collisions will take

place between A and B in a given period of time• such calculations indicate that the rate at which A and

B collide is far greater than the rate at which they react• the conclusion is that most collisions do not result in a

reaction• a collision that does result in a reaction is called an

effective collisioneffective collision• there are two main reasons why some collisions are

effective and others are not; activation energy and the orientation of A and B at the time of collision

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Molecular CollisionsMolecular Collisions• Activation energy:Activation energy: the minimum energy required

for a reaction to take place• in most chemical reactions, one or more covalent

bonds must be broken and energy is required for this to happen

• this energy comes from the collision between A and B• if the collision energy is large, there is sufficient

energy to break the necessary bonds, and reaction takes place

• if the collision energy is too small, no reaction occurs

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Molecular CollisionsMolecular Collisions• Orientation at the time of collision

• the colliding particles must be properly oriented for bond breaking and bond making

• for example, to be an effective collision between H2O and HCl, the oxygen of H2O must collide with the H of HCl so that the new O-H bond can form and the H-Cl bond can break

+ +

H2O + HCl H3O+ Cl-+

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Energy DiagramsEnergy Diagrams• Energy diagram for an exothermic reaction

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Energy DiagramsEnergy Diagrams• The reaction of H2 and N2 to form ammonia is

exothermic

• in this reaction, six covalent bonds are broken and six now ones formed

• breaking a bond requires energy, and forming a bond releases energy

• in this reaction, the energy released in making the six new bonds is greater than the energy required to break the six original bonds; the reaction is exothermic

3H2 N2 2NH3+ + energy

H-H H-H H-H N N NH

H

H NH

H

H+ + + +

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Energy DiagramsEnergy Diagrams• Energy diagram for an endothermic reaction

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Energy DiagramsEnergy Diagrams• Transition state:Transition state: a maximum on an energy

diagram• the transition state for the reaction between H2O and

HCl probably looks like this, in which the new O-H bond is partially formed and the H-Cl bond is partially broken

+ +

H2O + HCl H3O+ Cl-+

+ -

transition state

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Factors Affecting RateFactors Affecting Rate• Nature of reactants

• in general, reaction between ions in aqueous solution are very fast (activation energies are very low)

• in general, reaction between covalent compounds, whether in water or another solvent, are slower (their activation energies are higher)

• Concentration• in most cases, reaction rate increases when the

concentration of either or both reactants increases• for many reactions, there is a direct relationship

between concentration and reaction rate; when concentration doubles the rate doubles

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Factors Affecting RateFactors Affecting Rate• Temperature

• in virtually all reactions, rate increases as temperature increases

• an approximate rule for many reactions is that for a 10°C increase in temperature, the reaction rate doubles

• when temperature increases, molecules move faster (have more kinetic energy), which means that they collide more frequently; more frequent collisions mean higher reaction rates

• not only do molecules move faster at higher temperatures, but the fraction of molecules with energy equal to or greater than the activation energy also increases

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Factors Affecting RateFactors Affecting Rate• The distribution of kinetic energies (molecular velocities)

at two temperatures

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Factors Affecting RateFactors Affecting Rate• Catalyst:Catalyst: a substance that increases the rate of a

chemical reaction without itself being used up

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Factors Affecting RateFactors Affecting Rate• Many catalysts provide a surface on which

reactants can meet• the reaction of ethylene with hydrogen is an

exothermic reaction

• if these two reagents are mixed, there is no visible reaction even over long periods of time

• when they are mixed and shaken with a finely divided transition metal catalyst, such as Pd, Pt, or Ni, the reaction takes place readily at room temperature

HC C

H

HH+ H H C C HH

H H

HH+ energy

Ethylene Ethane

Pt

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Reversible ReactionsReversible Reactions• Reversible reaction:Reversible reaction: one that can be made to go

in either direction• if we mix CO and H2O in the gas phase at high

temperature, CO2 and H2 are formed

• we can also make the reaction take place the other way by mixing CO2 and H2

• the reaction is reversible, and we can discuss both a forward reaction and a reverse reaction

CO(g) +H2O(g) CO2(g) +H2(g)

CO(g) +H2O(g)CO2(g) +H2(g)

forward reactionreversereaction

CO(g) + H2O(g) CO2(g) + H2(g)

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Reversible ReactionsReversible Reactions• Equilibrium:Equilibrium: a dynamic state in which the rate of

the forward reaction is equal to the rate of the reverse reaction• at equilibrium there is no change in concentration of

either reactants or products• reaction, however, is still taking place; reactants are

still being converted to products and products to reactants, but the rates of the two reactions are equal

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Equilibrium ConstantsEquilibrium Constants• Equilibrium constant, K:Equilibrium constant, K: the product of the

concentration of products of a chemical equilibrium divided by the concentration of reactants, each raised to the power equal to its coefficient in the balanced chemical equation• for the general reaction

• the equilibrium constant is

aA + bB cC + dD

K =[C]c[D]d

[A]a[B]b

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Equilibrium ConstantsEquilibrium Constants• Problem:Problem: write the equilibrium constant for this

reversible reaction

• solution:solution: for this reaction, K is

• note that no exponents are shown in this equilibrium constant; by convention the exponent “1” is not written

CO(g) +H2O(g) CO2(g) +H2(g)

[CO2][H2]

[CO][H2O]K =

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Equilibrium ConstantsEquilibrium Constants• Problem:Problem: when H2 and I2 react at 427°C, the

following equilibrium is reached

• the equilibrium concentrations are [I2] = 0.42 mol/L, [H2] = 0.025 mol/L, and [HI] = 0.76 mol/L. Using these values, calculate the value of K

• Solution:Solution:

this K has no units because molarities cancel

I2(g) +H2(g) 2HI(g)

[HI]2

[I2][H2]K = = (0.76 M)2

(0.42 M) x (0.025 M)= 55

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Equilibrium and RatesEquilibrium and Rates• There is no relationship between a reaction rate

and the value of K• reaction rate depends on the activation energy of the

forward and reverse reactions; these rates determine how fast equilibrium is reached but not its position

• it is possible to have a large K and a slow rate at which equilibrium is reached

• it is also possible to have a small K and a fast rate at which equilibrium is reached

• it is also possible to have any combination of K and rate in between these two extremes

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LeChatelier’s PrincipleLeChatelier’s Principle• LeChatelier’s Principle:LeChatelier’s Principle: when a stress is applied

to a chemical system at equilibrium, the position of the equilibrium shifts in the direction to relieve the applied stress

• We look at three types of stress that can be applied to a chemical equilibrium• addition of a reaction component• removal of a reaction component• change in temperature

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LeChatelier’s PrincipleLeChatelier’s Principle• Addition of a reaction componentAddition of a reaction component

• suppose this reaction reaches equilibrium

• suppose we now disturb the equilibrium by adding some acetic acid

• the rate of the forward reaction increases and the concentrations of ethyl acetate and water increase

• as this happens, the rate of the reverse reaction also increases

• in time, the two rates will again become equal and a new equilibrium will be established

CH3COH

O

CH3CH2OHH2SO4 CH3COCH2CH3

O

H2O+ +Acetic acid Ethanol Ethyl acetate Water

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LeChatelier’s PrincipleLeChatelier’s Principle• at the new equilibrium, the concentrations of reactants

and products again become constant, but not the same as they were before the addition of acetic acid

• the concentrations of ethyl acetate and water are now higher, and the concentration of ethanol is lower

• the concentration of acetic acid is also higher, but not as high as it was immediately after we added the extra amount

• the system has relieved the stress by increasing the components on the other side of the equilibrium

• we say that the system has shifted to minimize the stress

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LeChatelier’s PrincipleLeChatelier’s Principle• Removal of a reaction componentRemoval of a reaction component

• removal of a component shifts the position of equilibrium to the side that produces more of the component that has been removed

• suppose we remove ethyl acetate from this equilibrium

• if ethyl acetate is removed, the position of equilibrium shifts to the right to produce more ethyl acetate and restore equilibrium

• the effect of removing a component is the opposite of adding one

CH3COH

O

CH3CH2OHH2SO4 CH3COCH2CH3

O

H2O+ +Acetic acid Ethanol Ethyl acetate Water

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LeChatelier’s PrincipleLeChatelier’s Principle• Problem:Problem: when acid rain attacks marble (calcium

carbonate), the following equilibrium can be written

how does the fact that CO2 is a gas influence the equilibrium?

• Solution:Solution: CO2 gas diffuses from the reaction site, and is removed from the equilibrium mixture; the equilibrium shifts to the right and the marble continues to erode

+H2SO4(aq)Sulfuric

acid

CaCO3(s) CaSO4(s) +CO2(g) + H2O(l)Calciumcarbonate

Calciumsulfate

Carbondioxide

Water

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LeChatelier’s PrincipleLeChatelier’s Principle• Change in temperatureChange in temperature

• the effect of a change in temperature on an equilibrium depends on whether the forward reaction is exothermic or endothermic

• consider this exothermic reaction

• we can look on heat as a product of the reaction• adding heat (increasing the temperature) pushes the

equilibrium to the left• removing heat (decreasing the temperature) pushes the

equilibrium to the right

+O2(g)2H2(g) 2H2O(l) + 137 kcal/mol

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LeChatelier’s PrincipleLeChatelier’s Principle• summary of the effects of change of temperature on a

system in equilibrium

Reaction type

exothermic

endothermic

Change intemperature

increase

decrease

increase

decrease

Direction equilibriumis driven

to the left; toward reactants

to the right; toward products

to the right; toward products

to the left; toward reactants

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End End Chapter 7Chapter 7

Chapter 7Chapter 7

7 7-1 © 2003 thomson learning, inc. all rights reserved bettelheim, brown, and march general,...

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