chapt. 16 – reaction rates 16.1a model for reaction rates 16.2 factors affecting reaction rates...
TRANSCRIPT
Chapt. 16 – Reaction Rates
16.1 A Model for Reaction Rates16.2 Factors Affecting Reaction Rates16.3 Reaction Rate Laws16.4 Instantaneous Reaction Rates and
Reaction Mechanisms
Section 16.1 A Model for Reaction Rates
• Calculate average rates of chemical reactions from experimental data, including data in both tabulated and graphical form.
• Use stoichiometry to calculate rates of other reactants or products given the rate for a specific reactant or product.
• Convert a reaction rate expressed in terms of a given unit of time to a rate in a different unit of time.
Collision theory is the key to understanding why some reactions are faster than others.
Section 16.1 A Model for Reaction Rates
• Describe the three key aspects of collision theory and their relationship to reaction rates.
• Explain the meaning of the terms activation energy and activated complex (transition state) and how they are related to collision theory.
• Use an energy versus reaction progress diagram to determine if a reaction is endo- or exothermic and to identify the forward and reverse activation energies and the activated complex (transition state).
• State the relationship between reaction rate and the spontaneity of a reaction.
Key Concepts
• The rate of a chemical reaction is expressed as the rate at which a reactant is consumed or the rate at which a product is formed.
• Reaction rates are generally calculated and expressed in moles per liter per second (mol/(L ● s)).
• In order to react, the particles in a chemical reaction must collide.
• The rate of a chemical reaction is unrelated to the spontaneity of the reaction.
Section 16.1 A Model for Reaction Rates
Expressing Reaction Rates
Very wide range of reaction time scales
Slow - Geologic time (> 104 years)• C(graphite) C(diamond)• Formation of oil underground
Very fast - psec (10-12 s) or faster• Initial electron transfer process in
photosynthesis
Expressing Reaction RatesA + B C
Average rate – change in some quantity related to reactants or products per some specific time interval
Avg Rate = quantity t
Expressing Reaction RatesA + B C
Avg Rate = quantity t
quantity of interest usually molar concentration (M) of reactant (A, B) or product (C)
• Units of rate: mol/(L•s)• Square brackets used for molarity• [A] = concentration of A in mol/L• May be gas phase or in solution
Expressing Reaction Rates
A + B CFor most reaction rate problems, think of reaction equation with some initial concentration of reactants and no products
t = 0 s [A] = 2.0 [B] = 2.0 [C] = 0.0
t = 10 s [A] = 1.8 [B] = 1.8 [C] = 0.2
Expressing Reaction Rates
A B
t= 0 min t= 20 min t= 40 min
Expressing Reaction Rates
Time (min)
Num
ber
of m
oles
A
B
Expressing Reaction Rates
CO(g) + NO2(g) CO2(g) + NO(g)
Avg Rate = quantity t
Have data of [NO] at times t1 and t2
Avg Rate = [NO]2 – [NO]1 = [NO] t2 - t1 t
t1= 0.0 s t2= 2.0 s [NO]1=0 [NO]2= 0.010M
Rate =(0.010M/2.0 s) = 5.0x10-3 mol/(L•s)
Expressing Reaction Rates
CO(g) + NO2(g) CO2(g) + NO(g)
t1= 0.0 s t2= 2.0 s [NO]1=0 [NO]2= 0.010M
Rate =(0.010M/2.0 s) = 5.0x10-3 mol/(L•s)
Above quantity is average rate of production of NO(g)
Could also have expressed rate in terms of consumption of CO(g)
• Get negative number but use + value
Expressing Reaction Rates
2H2(g) + O2(g) 2H2O(l)
Stoichiometry applies to rates• Rate of consumption of H2(g) =
twice rate of consumption of O2(g)
Time unit is not always seconds
4 mol/(L•s) 60 s/min =
240 mol/(L•min)
Expressing Reaction Rates
Example Problem 16.1, page 562
Reaction: butyl chloride (C4H9Cl) & H2O
t1= 0.0 s t2= 4.00 s
[C4H9Cl]1= 0.220M [C4H9Cl]2= 0.100M
Avg Rate = – (0.100 mol/L – 0.220 mol/L) (4.00 s – 0.00 s)(minus sign needed for consumption rate to get positive value for rate)
Avg Rate = 3.00x10-2 mol / (Ls)
Practice
Calculating reaction rate (and converting to different units of time)
Problems 1- 3 page 563
Problem 12 page 567
Problems 47 – 49 page 586
Problems 1 page 987
Collision Theory
A2 + B2 2AB
At molecular level, A2 and B2 must collide with each other for reaction to occur
Large numbers of collisions occur per unit time, but only small fraction of collisions result in reaction
Collision Theory
A2 + B2 2AB
A2 B2 AB AB
Thanks to Sam Mantel
Collision Theory - Orientation
Molecules must have proper orientation relative to each other so that new bond can form
Collisions with correct orientation yield short-lived activated complex (AC), also known as transition state
AC may form products or break apart and re-form reactants
Collision Theory - Orientation
Collision Theory - Orientation
Short-Lived Activated Complex
(Transition State)
+
Nature of the transition state in the reaction between CH3Br and OH-
Activation Energy
Ea = Activation energy = minimum amount of energy needed by reacting particles to form activated complex
High Ea – few collisions have required energy to produce products
• Low rate of reaction relative to collision rate
Low Ea – most collisions have energy required to produce products
Activation Energy
Collision Theory – Activation Energy
Activation Energy – Exothermic Reaction
Activation Energy – Endothermic Reaction
Collision Theory Summary
For reactants to produce products:1. Reacting particles must collide
2. Collision must have proper orientation
3. Collision energy must be activation energy (Ea)
Activation Energies
Energy level diagram allows you to determine Ea for both forward and reverse reactions
A + B C + D forward
C + D A + B reverse
Ea (forward)Ea (reverse)
Ene
rgy
Spontaneity & Rate
Reaction rate not related to spontaneity GHighly spontaneous reactions (large and negative G) may be very slow
Slightly spontaneous reactions may be very fast
Practice
Concepts related to modeling reaction rates
Problems 4 – 11 page 567
Problem 40 – 46 page 586
Chapt. 16 – Reaction Rates
16.1 A Model for Reaction Rates16.2 Factors Affecting Reaction Rates16.3 Reaction Rate Laws16.4 Instantaneous Reaction Rates and
Reaction Mechanisms
Section 16.2 Factors Affecting Reaction Rates
• Identify the factors that affect the rates of chemical reactions and the direction in which they act (speed up or slow down).
• Give specific examples of substances which are highly reactive and which are relatively inert.
• Explain the role of a catalyst.• Explain the role of an inhibitor.
Factors such as reactivity, concentration, temperature, surface area, and catalysts affect the rate of a chemical reaction.
Section 16.2 Factors Affecting Reaction Rates
• Explain the difference between homogeneous and heterogeneous catalysts and provide specific examples of each.
• Use a rule-of-thumb to estimate the effect of a given temperature change on the rate of reaction.
• Explain the effect of a temperature change on reaction rate in terms of the change in the distribution of reactant energies with temperature and also on its effect on collision frequencies.
Key Concepts
• Key factors that influence the rate of chemical reactions include reactivity, concentration, surface area, temperature, and catalysts.
• Raising the temperature of a reaction generally increases the rate of the reaction by increasing the collision frequency and the number of collisions that form an activated complex.
• Catalysts increase the rates of chemical reactions by lowering activation energies.
Section 16.2 Factors Affecting Reaction Rates
Rates and Nature of Reactants
Similar types of reactions occur at different ratesReactions of metal with water:Fast
2 Na(s) + 2H2O(l) H2(g) + 2 NaOH(aq)Slower
Ca(s) + 2H2O(l) H2(g) + Ca(OH)2(aq)Slow
Fe(s) + 2H2O(l) H2(g) + Fe(OH)2(aq)
Rates and Nature of Reactants
What are some highly reactive materials?
?
Highly combustible or explosive• TNT, nitroglycerin, gasoline,methane
Unreactive materials?
?
Teflon, water, titanium, nitrogen, helium
Surface Area and Rate
Increasing area provides more opportunity for collisions – rates increase
• Steel wool burns, nails don’t• Pulverizing (powdering) a solid makes it
dissolve faster in liquid than it would if left as big chunk
Temperature and Rate
Generally, increasing T will increase rate (see next slide)
Rule-of-thumb: increasing T by 10 degrees (C or K) will double reaction rate
Temperature and Rate
Temperature (K)
Rel
ativ
e R
ate
Increasing T by 10 K can
~ double rate
Temperature and Rate
Increasing T will increase rate
Effect due to:• Increased collision frequency• Increased collision energies
Greater number of collisions have energy activation energy Ea
Particle Energy and Temperature
T2 >T1Fraction of particles having energy > Ea larger at T2 than at T1
higher rate of reaction at T2
Activation Energy
Ea
Catalyst and Rate
Catalyst – substance that increases rate of chemical reaction without itself being consumed in reaction
Catalysts do not change amount of product produced, just rate at which it is produced
Catalysts function by lowering activation energy
Catalyst and Rate
Reaction Progress
Ene
rgy
Reactants
Products
Ea no catalyst
Ea with catalyst
Catalyzed reaction pathway
Uncatalyzed reaction pathway
Metal-catalyzed hydrogenation of ethyleneH2C=CH2(g) + H2(g) H3C-CH3
H2 & C2H4 approach & absorb to metal surface
Rate-limiting step: H—H bond breakage
Metal-catalyzed hydrogenation of ethyleneH2C=CH2(g) + H2(g) H3C-CH3
One H atom bonds to adsorbed C2H4
Another C—H bond forms and C2H6 leaves surface
Catalytic Converter in Engines
Catalysts and Oil Refining
Fluid Catalytic Cracking Unit (FCCU) is the heart of a petroleum refinery
Role is to upgrade gas oil to high octane (profitable) gasoline
http://www.unb.ca/che/che5134/fcc.html
FCCU – Simplified Schematic
Hot petroleum stream injected into catalyst
stream
Reaction generates vapor – velocities ~ 20 m/s; residence
time ~ 5 sec
Separation of solids and vapor
Fluid Catalytic Cracking UnitCatalyst: fine zeolite – has pore size that allows reactant molecules to fit in
http://www.unb.ca/che/che5134/fcc.html
Zeolite catalyst structurehttp://www.unb.ca/che/che5134/fcc.html
EnzymesCatalysts of biochemical worldAlmost all are proteinsEnzymes identified by –ase ending in name; root part of name indicates function of that enzyme
• Oxidase/Reductase – catalyze oxidation or reduction of something
• Transferase - catalyze transfer of chemical groups
• Hydrolase – catalyze hydrolysis
Enzymes
Activities determined by their 3-dimensional structure
Most much larger than substrates they act on; only small portion of enzyme (~ 3–4 amino acids) directly involved in catalysis
Active site: region that contains these catalytic residues, binds substrate, and then carries out reaction
Enzymes
Usually very specific as to which reactions they catalyze and substrates involved in these reactions
Complementary shape, charge and hydrophilic/hydrophobic characteristics of enzymes and substrates are responsible for this specificity
Enzyme Catalysis: Lock & Key Model
Substrate
Enzyme Enzyme-Substrate Complex
Enzyme
Products
Shapes of enzyme & substrate have “lock and key” relationship; this particular enzyme’s function is to split substrate apart
Enzyme Catalysis: Induced Fit ModelActive site continually reshaped by interactions with substrateSubstrate does not simply bind to rigid active site; amino acid side chains which make up active site are molded into precise positions that enable enzyme to perform its catalytic function
Inhibitors and Rate
Inhibitor: substance that slows down, or inhibits reaction rates
In food industry, inhibitors called preservatives
Inhibited No inhibitor
Types of Catalysts
Heterogeneous catalyst – exists in physical state different than that of reaction it catalyzes
• Metal (solid) catalytic convertor catalyzes reactions of gases in exhaust
Homogeneous catalyst – exists in same physical state as reaction it catalyzes
• Enzymes (aqueous solution)
Practice
Concepts related to factors affecting reaction rates
Problems 13 – 18 page 573
Problems 50 – 59 page 586
Chapt. 16 – Reaction Rates16.1 A Model for Reaction Rates16.2 Factors Affecting Reaction Rates16.3 Reaction Rate Laws16.4 Instantaneous Reaction Rates and
Reaction Mechanisms
Section 16.3 Reaction Rate Laws
• Express the relationship between reaction rate and concentration.
• Determine the proper units for the rate constant for a given rate law.
• Describe several experimental methods that can be used to collect concentration versus time data.
• Determine reaction orders and rate constants using the method of initial rates.
The reaction rate law is an experimentally determined mathematical relationship that relates the speed of a reaction to the concentrations of the reactants.
Key Concepts
• The mathematical relationship between the rate of a chemical reaction at a given temperature and the concentrations of reactants is called the rate law.
rate = k[A]rate = k[A]m[B]n
• The rate law for a chemical reaction is determined experimentally using the method of initial rates.
Section 16.3 Reaction Rate Laws
Varying Reaction Rates
Avg Rate = concentration t
If pick constant t (e.g., 10 min) and monitor reaction for an hour, generally will find that average rate decreases with time
See following slide for reaction
A B
Expressing Reaction Rates
Time (min)
Num
ber
of m
oles
A
B
Varying Reaction Rates
Reaction rates generally decrease because as reactants are depleted to form products:• reactant concentration drops • collision frequency between reactants
drop
Use rate law to quantify how instantaneous rate depends on instantaneous concentration of reactants
Reaction Rate Laws
A B
Rate law for this reaction: Rate = k [A]
k = specific rate constant
Units for k depend on form of law• Rates always have units of mol /(L•s)• For rate law above, k has units of s-1
k unique for each reaction• k depends on T, P, other conditions
Units for Rate Constant
2NO(g) + 2H2(g) N2(g) + 2H2O(g)
Rate law
Rate = k[NO]2[H2]
What units does k have in this rate law?
mol = ? mol2 mol(L•s) L2 L
? = mol•L3 = L2 _
mol3•L•s mol2•s
Reaction Rate LawsA B Rate = k [A]
Double [A], rate doubles• k large, fast rate• k small, slow rate• for the particular rate law shown, k is
analogous to 1/(RC) in an R-C circuit
Rate constants determined by fitting experimental data to an assumed rate law
Measuring Reaction Rates
There exist a variety of experimental methods to determine rates
Most appropriate method depends on specific nature of reaction (gas phase, solution phase, etc.) as well as on the speed of the reaction
Spectrophotometric MonitoringReactions Producing Change in
Light-Absorbing Species
Spectrophotometric Monitoring Student Version
Conductometric MonitoringMonitor Conductivity – Reactions That
Produce/Remove Ions
Manometric MonitoringMonitor Pressure for Reactions Involving Gases
Reaction Order
A B Rate = k [A]
Reaction order for reactant defines how rate affected by concentration of that reactant
Rate law above = k [A]1
Reaction order: power that concentration raised to in rate law
Above rate law first order in [A]
Reaction Order
A B Rate = k [A]
Reaction is first order in reactant A• Double [A], double rate
Nuclear decay follows first order rate process
• All first order processes have an associated ½ half – time necessary to decrease reactant concentration by 1/2
Determining Reaction Order
Like k, reaction orders for reactants must be experimentally determined
Examine how changing concentration of a particular reactant changes rate
Reaction Order
aA +bB products
General rate law Rate = k[A]m[B]n
Reaction said to be:• mth order in A• nth order in B
m and n:• Do not have to be integers• Can be zero
Reaction Order
aA +bB products
General rate law Rate = k [A]m [B]n
Overall reaction order: sum of orders for individual reactants in rate law
For above, overall order = m + n
Under special circumstances (reaction occurs in single step & with single activated complex), then: m = a, n = b
Reaction Order
2NO(g) + 2H2(g) N2(g) + 2H2O(g)
Experimentally determined rate law
Rate = k[NO]2[H2]
Reaction second order in NO• [NO] 2 [NO], rate 4 x rate
Reaction first order in H2
• [H2] 2 [H2], rate 2 x rate
Overall reaction: third order (2+1)
Determining Reaction Order
One approach: method of initial rates
Compare initial rates of reactions carried out with varying reactant concentrations
Method of Initial RatesRate of decomposition of H2O2 (produces O2(g))
Measure initial rate vs concentration – get data shown in graph
Rate = k[H2O2]
First order reaction: initial rate proportional to initial concentration of H2O2.
Method of Initial Rates: Find Rate LawaA +bB Products (see table 16.2)
Trial Initial [A] Initial [B] Initial Rate (mol/(L•s)
1 0.100 0.100 2.00 10-3
2 0.200 0.100 4.00 10-3
3 0.200 0.200 16.00 10-3
Trial 1 vs 2 m = 1Overall
reaction order = 2+1 = 3
x2 x2
x4x2
Trial 2 vs 3 n = 2
Rate = k[A]m[B]n
Method of Initial Rates: Find Rate Law
Trial Initial [A] Initial [B] Initial Rate (mol/(L•s)
1 0.100 0.100 2.00 10-3
2 0.200 0.100 4.00 10-3
3 0.200 0.200 16.00 10-3
Rate Law: Rate = k[A][B]2 Value of k?Using trial 1: Rate = 2.00 10-3 mol/(L•s) = k(0.100 mol/L)(0.100 mol/L)2 = k 1.00x10-3 mol3/L3 k = 2.00 L2/(mol2•s)
Practice
Rate laws and using the initial rates method to determine reaction order
Problems 18 – 19, p 577 (rate laws)
Problems 21 - 22, p 577 (initial rates)
Problems 64 - 68 p 587
Problems 2 - 4 page 987
Chapt. 16 – Reaction Rates
16.1 A Model for Reaction Rates16.2 Factors Affecting Reaction Rates16.3 Reaction Rate Laws16.4 Instantaneous Reaction Rates and
Reaction Mechanisms
Section 16.4 Instantaneous Reaction Rates and Reaction Mechanisms
• Identify the difference between average and instantaneous rates of chemical reactions.
• Calculate instantaneous rates of chemical reactions from a rate law and both average and instantaneous rates from graphical data.
• Identify intermediates and catalysts in complex reaction mechanisms.
The slowest step in a sequence of steps determines the rate of the overall chemical reaction.
Section 16.4 Instantaneous Reaction Rates and Reaction Mechanisms
• Identify intermediates and activated complexes (transition states) for a complex reaction on an energy versus reaction progress diagram.
• Identify the rate determining step of a complex reaction.
Key Concepts
• The reaction mechanism of a chemical reaction must be determined experimentally.
• For a complex reaction, the rate-determining step limits the instantaneous rate of the overall reaction.
Section 16.4 Instantaneous Reaction Rates and Reaction Mechanisms
Instantaneous Rate of Reaction
Avg Rate = concentration t
Average rate applies to relatively long time interval (one long enough so that measurable change in concentration has occurred)
Really want to know rate at specific instant of time = instantaneous rate
Instantaneous Reaction Rate
Average reaction rate depends on size of time intervalAs Dt approaches zero, average reaction rate approaches value equal to slope of line drawn tangent to curve at time tValue called instantaneous rate of reaction
Instantane-ous Rate
Given by slope
of tangent to curve at time
of interest
Instantaneous = [H2O2] Rate Dt
[H
2O2]
Relative Time
Dt
D[H
2O2]
C4H9Cl(aq) + H2O(l) →
C4H9OH(aq) + HCl(aq)
Instantaneous Rate of Reaction
=
Average Rate
H2 + 2ICl I2 + 2HCl
0
0.20.40.60.81.01.2
0 1 2 3 4 5 6 7 8 9
Time (s)
Mol
es/L
H2
s
mol/L
0.10.8
65.020.0
t
H
2rate Avg = = 6.4x10-2 mol/(Ls)
Average rate over this interval?
=
Instantaneous Rate
H2 + 2ICl I2 + 2HCl
0
0.20.40.60.81.01.2
0 1 2 3 4 5 6 7 8 9
Time (s)
Mol
es/L
H2
s
mol/L
1.18.3
30.060.0
= 0.11 mol/(Ls)
Instantaneous rate at 2.0 sec?
Rate = slope of tangent line =
Instantaneous Rate of Reaction
Besides determining from graph of concentration vs time, can get instantaneous rate directly from rate law for known concentration
2N2O5 4NO2(g) + O2(g)
Rate = k[N2O5] k = 1.0x10-5 s-1
For [N2O5] = 0.350 MRate = 1.0x10-5 s-1 0.350 M = 3.5x10-6 mol/(Ls)
Practice
Instantaneous rate of reaction
Problems 31- 33 page 579
Problem 39 page 582
Problems 73 - 74 page 587
Problem 5(a-b) page 988
Reaction Mechanisms
Most reactions are multi-step (complex) – consist of sequence of two or more simpler reactions, called elementary steps
• Generally involve one, two, or three (rare) molecules as reactants
• Statistically unlikely to have more than 3 molecules simultaneously collide – even 3 is relatively uncommon
Reaction MechanismsKnowing reaction’s mechanism means knowing complete sequence of elementary steps
Elementary steps contain intermediates – substances produced in one elementary step and consumed in another
Catalysts may be involved – speed rate without being consumed in reaction
Mechanism – Ozone Decomposition
2O3(g) 3O2(g) Complex overall reaction
ES = Elementary step
ES1 Cl + O3 O2 + ClO
ES2 O3 O2 + O
ES3 ClO + O Cl + O2
ES1 + ES2 + ES3 = overall reaction
Catalyst (Cl) regenerated
Cl catalyst
Intermediates produced and consumed
ClO intermediate
O intermediate
Rate Determining Step
Complex (multi-step) reactions can proceed no faster than their slowest elementary step – rate-determining step
Rate Determining Step
2NO(g) + 2H2(g) N2(g) + 2H2O(g)
ES1 2NO N2O2 (fast)
ES2 N2O2 + H2 N2O + H2O (slow)
ES3 N2O + H2 N2 + H2O (fast)
ES2 is rate-determining step – limits overall reaction rate
Intermediates in this reaction?N2O2 N2O
Reaction with Two Intermediates
Reaction Progress
Activated Complex #1
Products
Activated Complex #3
Activated Complex #2
Intermediates #1
Intermediates #2
Reactants
Ene
rgy
Reaction with Two Intermediates
Reaction Progress
Products
Intermediates #1
Intermediates #2
Reactants
Ene
rgy
Highest Ea for forward reaction
– rate determining step
Ea
Practice
Reaction mechanism, rate determining step
Problems 36 – 38 page 582
Problems 69 – 72 page 587
Problem 79, page 588
End of Chapter