Transcript Topic 6 - uaschemistry

Topic 6
REACTION RATES
IB Core Objective
6.1.1 Define the term rate of
reaction.
6.1.1 Define the term rate of reaction.
 Different chemical reactions occur at different
rates (ie. speeds)
 Rapid reactions, eg. The neutralisation of a strong
acid by a strong base in aqueous solution
 Slow reactions, eg. Rusting of iron
 Rate of a chemical reaction is a measure of the
speed at which products are formed, measured
as the change in concentration divided by the
change in time (units = mol dm-3s-1)
IB Core Objective
6.1.2 Describe suitable experimental
procedures for measuring rates of
reactions.
6.1.2 Describe suitable experimental procedures
for measuring rates of reactions.
 Different changes can be measured vs. time to
give a rate:





Rate = Change/time
Rate = Δ[Concentration]/ Δ[time]
Rate = Δ[Pressure]/ Δ[time]
Rate = Δ[Absorbance/ transmittance]/ Δ[time]
Rate = Δ[pH]/ Δ[time]
 Reaction rate describes how fast a reaction will
take place.
6.1.2 Describe suitable experimental procedures
for measuring rates of reactions.
 What factors will determine rate?




Surface area (usually refers to solid)
Concentration or pressure (if gas)
Temperature
Catalyst
 How is it found?
 Through experimentation only
 Rate of decreasing reactants
 −Δ[Concentration]/ Δ[time]
 Rate of increasing products
 Δ[Concentration]/ Δ[time]
Notice the
negative sign
6.1.2 Describe suitable experimental procedures
for measuring rates of reactions.
 pH changes (pH meter or acid/base titration)
 Volume, pressure or mass changes
 Conductivity (aq) (Conductivity meter or titration)
 Spectrometer or colorimeter for colour change
6.1.2 Describe suitable experimental procedures
for measuring rates of reactions.
 Data loggers and probes
 Conductivity
 Acid/base
 Temperature
 Titrations
 Acid/ base
 Concentration
6.1.2 Describe suitable experimental
procedures for measuring rates of reactions.
TITRATION
 Removal of small samples from
the reaction mixture at different
times and then titrating the
sample to determine the
concentration of either one of
the reactants or one of the
products at this time.
 Results can then be used
directly to generate a graph of
concentration against time.
 Best for quite slow reactions
6.1.2 Describe suitable experimental procedures for
.
measuring rates of reactions
 Spectrometer
 Absorbance and
transmittance
 Measures concentration
Collecting gas
6.1.2 Describe suitable experimental
procedures for measuring rates of reactions.
COLLECTION OF EVOLVED GAS
 Gas produced in the reaction is collected either in a
gas syringe or in a graduated vessel over water.
 The volume of gas collected at different times can
be recorded.
 This technique is limited to reactions that produce a
gas (obviously!) PLUS if the gas is to be collected
over water, the gas must not be water soluble.
 An alternative technique is to carry out the reaction
in a vessel of fixed volume and monitor the increase
in the gas pressure
6.1.2 Describe suitable experimental
procedures for measuring rates of
reactions.
COLLECTION OF
EVOLVED GAS
 Example: Measuring the
rate of reaction between a
moderately reactive metal
(such as zinc) and an acid
(such as hydrochloric acid).
Zn (s) + 2 H+ (aq) →
Zn2+ (aq) + H2 (g)
6.1.2 Describe suitable experimental
procedures for measuring rates of reactions.
MEASUREMENT OF THE MASS OF REACTION MIXTURE
 The total mass of the reaction mixture will only vary
if a gas is evolved.
 The gas should have a high molar mass (ie. not hydrogen)
so there is a significant change in mass
 Example: measuring the rate of reaction between a
metal carbonate (such as calcium carbonate, marble
chips) and an acid (such as hydrochloric acid) by
measuring the rate of mass loss resulting from the
evolution of carbon dioxide:
CaCO3 (s) + 2H+ (aq) → Ca2+ (aq) + H2O (l) + CO2 (g)
IB Core Objective
 6.1.3 Analyse data from rate experiments.
 Students should be familiar with graphs of
changes in concentration, volume, and mass
against time.
6.1.3 Analyse data from rate experiments.
 Numerical value will vary according to amount
of the substance involved in the stoichiometric
equation
MnO4- (aq) + 8 H+ (aq) + 5 Fe2+ (aq) → Mn2+ (aq) 4 H2O (l) + 5 Fe3+ (aq)
 The rate of appearance of Fe3+ is five times as
great as the rate at which MnO4- is consumed
Rate = - ∆ [MnO4-] = 1 ∆ [Fe3+]
∆t
5 ∆t
Units: mol dm-3 s-1
6.1.3 Analyse data from rate experiments.
 Or more simply for a reaction: a A → b B,
then
Rate = - 1 ∆ [A] = 1 ∆ [B]
a ∆t
b ∆t
General Formula
 Any property that differs between the
reactants and products can be used to
measure the rate of the reaction
6.1.3 Analyse data from rate experiments.
 To find rate at an instant
in time (instantaneous
rate) we calculate the
slope of a tangent on the
experimentally obtained
graph.
 Here, Δx is the change in
time and Δy is the change
in concentration.
IB Core Objective
How does kinetic
 6.2.1 Describe the kinetic theory
terms
of the
energyinand
contact
movement of particles whose differ
average
energy is
for solids
proportional to temperaturewhen
in kelvins.
compared to
liquids and gases?
 As heat is supplied to a substance, the velocity
(kinetic energy) of the particles will increase.
 When the velocity increases, so does the
temperature.
 Therefore, the absoluteSolids
temperature
are fixed in kelvin is
proportional to the average
kinetic
energy of all
and only
vibrate,
the particles.
so kinetic energy
is limited
IB Core Objective
 6.2.2 Define the term activation energy, Ea.
 The minimum amount of energy required
for reaction is the activation energy, Ea.
6.2.2 Define the term activation energy, Ea.
• Just as a ball cannot get over a hill if it does not
roll up the hill with enough energy, a reaction
cannot occur unless the molecules possess
sufficient energy to get over the activation
energy barrier.
IB Core Objective
 6.2.3 Describe the collision theory
 Just because chemicals collide /interact does
not mean they will react!
 Reaction rate depends on:
 Collision frequency
 Number of particles with E ≥ Ea.
 Appropriate collision geometry or orientation.
IB Core Objective
 6.2.4 Predict and explain, using collision
theory, the qualitative effects of particle size,
temperature, concentration and pressure on
the rate of a reaction.
6.2.4 Surface Area
 By increasing the surface area we increase
the contact area.
 Collision rate will increase.
Marble
Marble chip
6.2.4 Concentration & Pressure
 Increasing the concentration or pressure
will increase opportunity for molecules to
react
 Collision rate will increase
 Pressure only affects gases
6.2.1 and 6.2.4 Temperature and Rate
• Temperature is directly
related to kinetic energy or
particle speed.
• Faster particles increases
probability of molecule
interaction
• Will increase # of particles
with sufficient energy to over
come activation energy.
Orientation
IB Core Objective
 6.2.6 Describe the effect of a catalyst on a
chemical reaction.
 What do you know about catalysts? How
would you define one in your own terms?
 Have you heard the expression “catalyst for
change”?
Catalysts 6.2.6
 Catalysts provide
are just an
a facilitator
alternate reaction pathway.
Theythat
are takes
not used
 One
lessup
energy. Similar to a tunnel
a mountain
side, whether
you go up
 through
They lower
the activation
energy allowing
forand
more
over,
or go
still end
up atrequirement
the same place.
particles
tothrough
have theyou
correct
energy
Analogy
 Catalysts are like a dating service
 They bring compatible people together
 They are not involved in what happens after the pairs
are together
 They can be reused again and again.
 This is not to say that people
will not get together on their
own, however it lowers the
energy required to find a
match
6.2.6 Describe the effect of a catalyst on a chemical reaction.
One way a catalyst can speed up a reaction is by
holding the reactants together and helping bonds
to break.
Draw an Energy Diagram for a
catalyzed reaction
IB Core Objectives
 6.2.5 Sketch and explain qualitatively the
Maxwell-Boltzmann energy distribution curve
for a fixed amount of gas at different
temperatures and its consequences for
changes in reaction rate.
 Students should be able to explain why the
area under the curve is constant and does not
change with temperature.
Maxwell–Boltzmann
Distributions 6.2.5
Activation Energy
Maxwell–Boltzmann Distributions
6.2.5
 As the temperature
increases, the curve
flattens and broadens.
 At higher
This system has a
fixed number of
particles
temperatures, a
larger number of
molecules has higher
energy.
Maxwell–Boltzmann Distributions
 If the dotted line represents the activation energy,
as the temperature increases, so does the fraction
of molecules that can overcome the activation
energy barrier.
• As a result, the
reaction rate
increases.
Be sure to draw the higher
temp. Curve shorter and
wider than the original.
IB Core Objective
 6.2.7 Sketch and explain Maxwell-Boltzmann
curves for reactions with and without
catalysts.
Catalysts and Activation Energy 6.2.7
Is this a
MaxwellBoltzmann
distribution
curve?
Why or why
not?
Sketch what a
MaxwellBoltzmann
curve would
look like.
0. 693
1 [lnA12]
1
t1 / 2   k t0 
[[A
A]] t k2k
[kA] 0 [ kA] 0
HL
Rate Law
Return
Order in [A]
Rate
Law
Integrated Form,
y = mx + b
Straight
Line Plot
Half-Life
t1/2
zeroth
order
(n = 0)
rate = k [A]o= k
[A]t = - k t +[A]o
[A]t vs. t
(slope = - k)
t½ = [A]0/2k
ln[A]t vs. t
t½ = ln 2/k
first
order
(n = 1)
second
order
(n = 2)
rate = k
[A]1
rate = k
[A]2
ln[A]t = - k t + ln[A]o
1/[A]t = kt + 1/[A]o
(slope = - k)
1/[A]t vs. t
(slope = k)
t½ = 1/(k[A]o)
HL Objective
 16.1.1 Distinguish between the terms rate
constant, overall order of reaction and order of
reaction with respect to a particular reactant.
16.1.1 Distinguish between the terms rate
constant, overall order of reaction and order of
reaction with respect to a particular reactant.
 Rate Rxn = Δ[A]/ Δt




‘k’ represents a
constant and does
not change EXCEPT
•Temperature!!
•Particle size (solids)
Rate = k[A]m [B]n
k = rate constant
A and B are reactants.
m and n represent the order of reaction for
each reactant.
m + n = overall order of reaction
HL Objective
 16.1.2 Deduce the rate expression for a
reaction from experimental data.
16.1.2 Deduce the rate expression for a
reaction from experimental data.
 1) If concentration of [A] is doubled yet no
change in rate, than order = zero. [A]0
 2) If doubling conc. of [A] = a doubling of rate,
than order = one. [A]1
 3) If doubling conc. of [A] = a quadrupling of
rate, than order = 2. [A]2
16.1.2 Deduce the rate expression for a
reaction from experimental data.
For the reaction A + 2B → C
Expt. Number
1.
2.
3.
4.
[A] mol dm-3
[B] mol dm-3
Rate of formation
of C mol dm-3 s-1
1
0.10
0.05
0.02 x 102
2
0.10
0.10
0.04 x 102
3
0.05
0.10
0.01 x 102
4
0.10
0.20
0.08 x 102
Deduce the order of reaction in respect to reactant A
Deduce the order of reaction in respect to reactant B
State what the overall order of the reaction is.
Deduce the rate expression for this reaction
A: 1. 2nd order. 2. 1st order 3. 2+1=3 3rd order 4. rate = k[A]2[B]
IB HL Objective
 16.1.3 Solve problems involving the rate
expression.
16.1.3 Solve problems involving the rate
expression.
 From the previous problem, rate = k[A]2[B].
 To find the rate constant, plug in numbers
from the results:
Expt. Number
[A] mol dm-3
[B] mol dm-3
Rate of formation
of C mol dm-3 s-1
1
0.10
0.05
0.02 x 102
2
0.10
0.10
0.04 x 102
3
0.05
0.10
0.01 x 102
4
0.10
0.20
0.08 x 102
A: k= 4 x 103 mol-2 dm6 s-1
Notice the
units!
16.1.3 Solve problems involving the rate
expression.
Units for rate constant
Zero order overall: mol dm-3 s-1
First order overall: s-1
Second order overall: mol-1 dm3 s-1
Third order overall: mol-2 dm6 s-1
IB HL Objective
 16.1.4 Sketch, identify and analyse graphical
representation for zero-, first- and secondorder reactions.
 Note: Students should be familiar with both
concentration-time and rate-concentration
graphs.
16.1.4 Sketch, identify and analyse graphical
representation for zero-, first- and second-order
reactions.
Concentration
Order
of
Reaction
 Zero order reaction
doubled. The rate:
Does not change
Triples (x3)
x9
Rate
Concentration
st order reaction

1
Does not change
Zero
 2nd order reaction
Doubles (x2)
1st
Quadruples (x4)
2nd
Concentration
tripled. The rate:
Time
Time
16.1.4 Sketch, identify and analyse graphical
representation for zero-, first- and second-order
reactions.
 Zero order = linear. Decrease in
concentration does not affect the rate of
reaction.
 First order = exponential decay (same as
radioactive decay). (t1/2)
 Second order = parabolic, because it depends
on the square of the concentration.
Other Useful Information
 FOR 1ST ORDER REACTIONS ONLY
 Half life: The time it takes for half a
substance to decay/ disappear
 t1/2 = ln(2)/k
Concentration / Pressure
Calculate the half life for
the previous question
Half life
IB HL Objective
 16.2.1 Explain that reactions can occur by
more than one step and that the slowest step
determines the rate of reaction (ratedetermining step).
16.2.1 Explain that reactions can occur by more than one step and
that the slowest step determines the rate of reaction (ratedetermining step).
 Several reactions may have simple equations,
but sometimes there are several intermediate
steps which occur to get to the final
product(s).
 These various intermediate steps can occur at
different rates.
 The slowest step is the rate-determining step.
16.2.1 Explain that reactions can occur by more than one step and
that the slowest step determines the rate of reaction (ratedetermining step).
 Unimolecular: One species breaks up or rearranges to
form products.
 Bimolecular: Two species collide and interact to form
the product. (Doubling the concentration of either
will double the collision rate. What would be the
overall order for this?)
Molecularity Reaction
Rate Law
Uni
A Product
Rate = k[A]
Bi
A +B  Product
Rate = k[A][B]
Bi
A + A  Product
Rate = k[A]2
IB HL Objective
 16.2.2 Describe the relationship between
reaction mechanism, order of reaction and
rate-determining step.
Mechanism
 Must accurately represent the original
stoichiometric equation
2XY2  2XY + Y2
Possible Mechanism
XY + XY  X Y
X Y  X + 2Y
X + Y  2XY
2
2
2
2
4
2
2
2
4
2
Determining the Slowest Step
 1) Determine the overall reaction equation
 2) Check to see if adding the reactions = the
overall stoichiometric equation
 3) Check for consistency with regards to the
rate equation and the rate determining step
 Does the proposed rate equation = the overall
reaction equation?
Question
 Through experimentation it was found that
overall rate equation for:
NO2(g) + CO(g)  NO(g) + CO2(g)
Rate = k[NO2]2
Found to be
zero order
 The proposed reaction mechanism is:
(Slow) NO2 + NO2  N2O4
(Fast) N2O4 + CO  NO + CO2 + NO2
NO2(g) + CO(g)  NO(g) + CO2(g)
Does the sum match the original stoichiometric
equation?
Question Continued
 The proposed reaction mechanism is:
(Slow) NO2 + NO2  N2O4
(Fast) N2O4 + CO  NO + CO2 + NO2
NO2(g) + CO(g)  NO(g) + CO2(g)
 Does the RDS rate eqn = the overall rate eqn?
 RDS rate = k[NO2][NO2]  k[NO2]2
Appears to be consistent with the original rate equation
Question 2
2NO(g) + 2H2(g) → N2(g) + 2H2O(g)
rate = k[H2][NO]2
 State and explain whether this mechanism
agrees with the experimental rate expression in
H2 + NO  X
X + NO Y + H2O
Y + H2  N2 + H2O
fast step
slow step
fast step
NO is part of two steps, therefore must be considered in the rate
expression.
Question 3
2H2O2 Br- cat.
2H2O + O2
 Which of the following mechanisms are correct?
Mech. 1
H2O2  2OH
2OH + Br-  BrO- + 2H2O2
2H2O2 BrO-  2H2O + O2 + BrMech. 2
H2O2 + Br-  H2O + BrOBrO- + H2O2  H2O + O2 + BrA: Mech 2
IB HL Objective
 16.3.1 Describe qualitatively the relationship
between the rate constant (k) and
temperature (T).
 16.3.2 Determine activation energy (Ea)
values from the Arrhenius equation by a
graphical method.
16.3.1 Describe qualitatively the relationship between the
rate constant (k) and temperature (T).
 The Arrhenius Equation is k = Ae(-Ea/RT). This
equation can be found in your data booklet.
 A is the Arrhenius constant. It is dependent
on collision rate and steric factors (geometry
of the colliding particles).
 This expression indicates that the rate
constant k depends exponentially on
temperature. Temperature has a large effect
on the reaction rate.
Arrhenius Equation (16.3)
 The relationship between rate constant
and temperature.
 The equation of a straight line
k = Ae(-Ea/RT)
can be determined by taking
the natural log.
Activation Energy can be
found by varying the
temperature of an experiment
lnk = lnA –(Ea/R)1/T
Arrhenius Equation (16.3)
 Graphing ln k VS 1/T will give a straight line
 Slope will = -Ea/R
ln k
Extrapolate to find:
ln [A]
1/T