CHM 120 CHAPTER 16 KINETICS: Rates and Mechanisms of Chemical Reactions

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CHM 120CHAPTER 16KINETICS Rates and
Mechanisms of Chemical Reactions

• Dr. Floyd Beckford
• Lyon College

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CHEMICAL KINETICS

• Two factors control the outcome of chemical
• reactions
• Chemical Thermodynamics
• Chemical Kinetics
• Chemical Kinetics study of rates of chemical
• reactions and mechanisms by which they occur
• A reaction may be spontaneous but does not
• occur at measurable rates

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REACTION RATES

• Rate of reaction describes how fast
• reactants are used up and products are formed
• There are 4 basic factors that affect
• reaction rates
• Concentration
• Physical state
• Temperature
• Catalysts

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• For every reaction the particles must come into
• intimate contact with each other
• High concentrations by definition implies that
• particles are closer together (than dilute
• solutions)
• So rate increases with concentration
• The degree of intimacy of particles obviously
• depends on the physical nature of the particles
• Particles in the liquid state are closer than in
• the solid state

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• Likewise, particles in a finely divided solid
  will
• be closer than in a chunk of the solid
• In both situations, there is a larger surface
• area available for the reaction to take place
• This leads to an increase in rate
• Temperature affects rate by affecting the
• number and energy of collisions
• So an increase in temperature will have the
• effect of increasing reaction rate

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• Rate of reaction is typically measured as the
• change in concentration with time
• This change may be a decrease or an increase
• Likewise the concentration change may be of
• reactants or products

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• Rate has units of moles per liter per unit time
• - M s-1, M h-1
• Consider the hypothetical reaction
• aA bB ? cC dD
• We can write

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• Note the use of the negative sign
• - rate is defined as a positive quantity
• - rate of disappearance of a reactant is
• negative
• 2N2O5(g) ? 4NO2(g) O2(g)

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• Rate may be expressed in three main ways
• Average reaction rate a measure of the
• change in concentration with time
• 2. Instantaneous rate rate of change of
• concentration at any particular instant during
  the
• reaction
• 3. Initial rate instantaneous rate at t 0
• - that is, when the reactants are first
• mixed

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RATE LAW

• Consider the following reaction
• aA bB ? products
• Rate of reaction changes as concentration of
• reactants change at constant temperature
• RATE LAW equation describing the relationship
• between concentration of a reactant and the rate
• Rate kAmBn
• where k is called the rate constant

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• m, n are called reaction orders
• - they indicate the sensitivity of the rate to
• concentration changes of each reactant
• NOTE the orders have nothing to do with the
• stoichiometric coefficients in the balanced
• overall equation
• An exponent of 0 means the reaction is zero
• order in that reactant - rate does not depend
• on the concentration of that reactant

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• An exponent of 1 ? rate is directly proportional
• to the concentration of that reactant
• - if concentration is doubled, rate doubles
• - reaction is first order in that reactant
• An exponent of 2 ? rate is quadrupled if the
• concentration of that reactant is doubled
• - reaction is second order in that reactant
• The overall reaction order is the sum of all the
  
• orders

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• Rate kAB0 m 1 and n 0
• - reaction is first order in A and zero order
• in B
• - overall order 1 0 1
• - usually written Rate kA
• Remember the values of the reaction orders
• must be determined from experiment they
• cannot be found by looking at the equation

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DETERMINATION OF THE RATE LAW

• The method of initial rates may be used
• - involves measuring the initial rates as a
• function of the initial concentrations
• - avoids problems of reversible reactions
• - initially there are no products so they
• cannot affect the measured rate
• In this method the experiments are chosen so
• as to check the effect of a single reactant on
• the rate

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THE RATE CONSTANT
1. The units of k depends on the overall order of
reaction 2. The value of k is independent of
concentration and time 3. The value refers to a
specific temperature and changes if we change
temperature 4. Its value is for a specific
reaction
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THE INTEGRATED RATE EQUATION

• This is the equation that relates concentration
• and time
• Consider a first-order reaction
• aA ? products
• Rate kA

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• The equation may be written in the form for a
• linear plot

• A plot of log At vs. t is linear plot with
  slope
• -k/2.303
• Note that this plot gives a straight line ONLY
• if the reaction is first-order

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• Half-life
• The half-life, t1/2, is defined as the time it
• takes for the reactant concentration to drop
• to half its initial value

• Note the half-life for a first order reaction
• does not depend on the initial concentration
• The value of the half-life is constant

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• Second order reactions
• Consider a reaction that is 2nd order in
  reactant
• A and 2nd overall
• aA ? products and Rate kA2

• A plot of 1/At vs. t gives a straight line
  with
• slope k

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RATES AND TEMPERATURE

• Recall that temperature is the only factor that
• affects the rate constant
• In general rates increase with temperature

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• This is ARRHENIUS EQUATION
• Can be arranged in the form of a straight line
• ln k (-Ea/R)(1/T) ln A
• Plot ln k vs. 1/T ? slope -Ea/R

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• Another form of Arrhenius equation

• COLLISION THEORY a reaction results when
• reactant molecules, which are properly oriented
• and have the appropriate energy, collide
• The necessary energy is the activation energy,
• Ea

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• Not all collisions leads to a reaction
• For effective collisions proper orientation of
• the molecules must be possible

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TRANSITION STATE THEORY

• During a chemical reaction, reactants do not
• suddenly convert to products
• The formation of products is a continuous
• process of bonding breaking and forming
• At some point, a transitional species is formed
• containing partial bonds
• This species is called the transition state or
• activated complex

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• The transition state is the configuration of
• atoms at the maximum of the reaction energy
• diagram
• The activation energy is therefore the energy
• needed to reach the transition state
• Note also that the transition state can go on
• to form products or break apart to reform the
• reactants

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REACTION MECHANISMS

• MECHANISM the step-by-step pathway by
• which a reaction occurs
• Each step is called an elementary step
• NO2(g) CO(g) ? NO(g) CO2(g)
• Mechanism
• NO2(g) NO2(g) ? NO(g) NO3(g)
• NO3(g) CO(g) ? NO2(g) CO2(g)
• NO3 is a reaction intermediate

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• Elementary reactions are classified by the
• molecularity
• A ? B C unimolecular
• A B ? C D bimolecular
• A 2B ? E termolecular
• Termolecular reactions are very unlikely
• For ANY SINGLE ELEMENTARY
• REACTION reaction orders are equal to the
• coefficients for that step

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• Rate kelemAB
• The slow step is called the rate-determining
• step (RDS)
• A reaction can never occur faster than its
• slowest step
• Overall reaction sum of all elementary steps
• 2. The mechanism proposed must be consistent
• with the rate law

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• There may be more than one plausible
• mechanism
• The experimentally determined reaction orders
• indicate the number of molecules of the reactants
• - in the RDS (if it occurs first)
• - the RDS and any fast steps before it

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CATALYSIS

• Reaction rates are also affected by catalysts
• Catalyst a substance that increases the rate of
• a reaction without being consumed in the reaction
• Catalysts work by providing alternative pathways
• that have lower activation energies
• A catalyst may be homogeneous or heterogeneous
• Homogeneous catalyst and reactants are in the
• same phase

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• Heterogeneous catalyst in a different phase
• Typically a solid in a liquid
• An important example catalytic converters in
• automobile
• - convert pollutants to CO2 H2O, O2, N2
• - usually Pt, Pd, V2O5, Cr2O3, CuO
• Cars must use unleaded fuels lead poisons the
• catalytic bed