Transcript http://college.hmco.com/chemistry/shared/media/animations/oscillatingreaction.html Chemical Kinetics Concerns reaction rates: Speed with which reactants are converted to products 11/6/2015

http://college.hmco.com/chemistry/shared/media/animations/oscillatingreaction.html
Chemical Kinetics
Concerns reaction rates:
Speed with which reactants are converted
to products
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Fireworks explode/give bright colors due to very
fast chemical reactions
The rusting of a bridge is also a chemical
reaction, but it is very slow.
The eroding of a mountain takes even longer
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Rate: the speed at which
something happens
Average rate or speed = ∆quantity
∆t
Ratechem reaction = ∆ [reactant] or [product] = x mol
unit time
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L⋅s
Always positive
[ ] = molar concentration
Forward Rate: reactants  products
Reverse Rate: products recombine to become reactants
Net Rate: forward rate - reverse rate
Average Rate: speed of entire reaction from start to finish
Instantaneous Rate: speed of reaction at one moment in time
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The Collision Theory
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Atoms, ions, and
molecules (reacting
substances) must collide
in order to react:
Speed of
molecules
Do they touch?
Too slowly
Just right
No
Yes, with correct
orientation
Too fast
Yes, but bounce
off each other
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Chemical
reaction
occurs?
No
New bonds
formed – Yes
No
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Activated complex: temporary,
unstable arrangement of atoms that
may form products or may break
apart to reform reactants
Reacting
substances
must
collide with
sufficient
energy to
form the
activated
complex
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http://college.hmco.com/chemistry/shared/media/animations/uctransitionstates.html
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When reactant molecules collide
w/enough energy, orbitals overlap
• Transition state
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Reactants broken down into particles
Not reactants or products
As likely to form reactants as
products
Reactant molecules undergo
further change
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Rearrangement of molecular
structure
Breaking of existing bonds
Making new ones
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Enough kinetic energy for
reactants to leap over
reaction barrier?
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Exothermic reactions release energy and
form products at lower energy level.
Endothermic reactions absorb energy and
form products at lower energy level.
∆H = (-)
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Covers period of time between
mixing of reactants and point at
which chemical reaction stops or
reaches equilibrium
How fast chemicals react and factors that
influence rate allows chemists to exercise
precise control over chemical reactions
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Chemical Kinetics
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Understand steps by which reaction takes place
Called reaction mechanism
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• Rate of chemical reaction dictates whether reaction
can occur
• Knowing rate law (expression relating rate to
concentrations of reactants) helps chemist adjust
reaction conditions to get more suitable rate
• If 2 reactions competing for single reagent, knowing
rate law lets you favor exclusive formation of single
product
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Measurement of reaction rate
based on rate of appearance
of product or disappearance
of reactant
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Determined by measuring concentration
of one or more chemicals at different
times during course of reaction
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Use kinetic curve (concentration vs.
time curve)
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If [reactant] measured, decrease in
[ ] w/time is recorded
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Slope of curve at point in time from
kinetic curve determines rate of
product production (or reactants’
disappearance)
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Change can be
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Disappearance of reactants
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Decrease
Rate has negative sign
Appearance of product
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Increase
Rate has positive sign
Always define rate as positive quantity
[ ] indicate concentration in mol/L
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AB
• Rateaverage = -Δ[A] = +Δ[B]
Δt
Δt
• RateA disappearing = RateB appearing
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Rate of appearance (or disappearance) of
substance is divided by its stoichiometric coefficient
aA + bB  cC + dD
• Rate = -ΔA = -ΔB = +ΔC = +ΔD
aΔt bΔt cΔt dΔt
• Get same reaction rate no matter which is
measured
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Rate is change in [ ] of some species
w/respect to time, then measure [ ] of all
species at multiple times to determine rate.
Initial [ ]
H2/I2 equal
at all times
and initial [ ]
product is
zero
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• Rate of formation of HI is twice rate of
disappearance of H2 or I2 at any given time
• Rate slows in time due to decreasing [ ] of
reactants
• Stated mathematically, relationship between
formation of products and disappearance of
reactants for reaction is
• H2 + I2 ↔ 2HI
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Instantaneous rate
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Rate not constant
Decreases w/time ([ ] change with time)
Rate measured for given point in time
Rates in tables: average rate over specific
time period
Determined by finding slope of line tangent to
point representing particular time
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Rate = -(slope of tangent line)
Slope of tangent line = change in y/change in x
= d[A] /dt (derivative): rates measured over time
intervals that approach 0 second
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Decomposition of NO2
2NO2(g)  2NO(g) + O2(g)
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2NO2(g)  2NO(g) + O2(g)
Rate of consumption of NO2 =
Rate of production of NO =
2(Rate of production of O2) because rate of
production of NO is twice that of O2
–Δ[NO2] = Δ[NO] = (Δ[O2] )
2Δt
2Δt
Δt
(multiply all by 2)
–Δ[NO2] = Δ[NO] = 2(Δ[O2] )
Δt
Δt
Δt
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http://www2.wwnorton.com/college/chemist
ry/gilbert/tutorials/interface.swf?chapter=ch
apter_14&folder=reaction_rate
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Factors affecting
rate of reaction
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Nature of reactants
• Some naturally occur faster than other
• Number of reacting species
• State of reacting species (particles forming solids
move slower than gases/in solution)
• Complexity of reaction
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Change in concentration
of reactants
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↑ [ ] reactants increases rate
For reaction to occur, particles must collide
• If [ ] higher, chance of collision greater
Cases where this doesn't affect rate of reaction
• Where catalyst is already working as fast as it can
• Steps in certain multi-step reactions likely to have
widely different rates (some fast/some slow)
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• Overall rate governed by how fast A splits up to
make X and Y (rate determining step)
• Increasing [A] increases chances of step
happening
• Increasing [B] speeds up 2nd step which is
“waiting around” for 1st to happen, so hardly
makes difference to overall rate
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Temperature
• Particles only react when they collide
• Most particles don’t have energy needed to react
• Increasing that # increases rate of collisions/probability
of favorable collisions
• Increasing temperature always increases reaction rates
• Decreasing temperature always decreases reaction
• For many reactions occurring at around room T, rate of
reaction doubles for every 10°C (9/11°C rise in temperature
• # degrees needed to double rate changes gradually as
T
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Entropy & temperature
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1.
2.
3.
Place 3 drops of blue food coloring in 3 flasks.
Place 3 drops of yellow food coloring in 3 more.
Keep 1 of each at room temperature, and warm 1, and cool
the other.
Using note cards, invert the blue flask over the yellow and
remove the card. Secure the flask with a clamp.
Allow the flasks to stand and record the time necessary for
both flasks to become the same shade of green.
Is the entropy greatest when the colors of the flasks are
different or the same? Explain.
Explain the relationship between diffusion and entropy.
What influence does an increase/decrease in temperature
have on the diffusion observed?
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•Particles present have very wide range of energies
•For gases, shown on Maxwell-Boltzmann Distribution
•Plot # particles having each particular energy
•Area under curve = total # particles present
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•For reaction to occur, particles must collide with
energies ≥ activation energy
•Many don't have enough energy to react when they
collide
•Either change shape of curve, or move activation
energy to left
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Pressure
• Increasing pressure
increases rate of
reaction
• Particles are closer
together
• Have increased
probability of
favorable collisions
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Use a
Catalyst
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Substance which speeds up reaction
Chemically unchanged at end of reaction
Usually increases rate , but some can
decrease it
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Provide alternative pathway w/lower
activation energy
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http://college.hmco.com/chemistry/shared/media/animations/homogeneouscatalysis.html
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Homework:
Read 17.1-.2, pp. 529-541, 17.4, pp.
546-547
Q pp. 554-555, #34, 37, 48-49, 62, 66,
71
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Rate Laws
The rate of a chemical reaction dictates whether
a reaction can occur
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Rate Laws:
An introduction
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Chemical reactions are reversible
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With time, enough products accumulate
Reverse reaction becomes important
Concentration of reactants depends on difference in
rates of forward and reverse reactions
Focus on reaction soon after reactants are mixed,
before products build up to significant levels
Reaction rate depends only on [ ] of reactants
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Rate = k[A]n
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Rate law depends on [ ] of reactants
Must be determined by experiment
− Proportionality constant "k"-rate constant
– Relates [ ] and orders to rate of reaction
– Constant value at fixed T for given reaction
– Units of rate constant depend on order of
reaction
− "n"-order of reactant
– Can be integer (including zero) or fraction
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Important Rate law points
1.
[products] do not appear
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2.
Value of exponent "n" must be determined by
experiment
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Reaction rate studied under conditions where reverse
reaction does not contribute to overall rate
Cannot be written from balanced equation
Differential Rate Law (Rate law) - expresses how
rate depends on concentration
Integrated rate law - expresses how concentration
depends on time
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Possible to write rate law that relates reaction
rate to concentration of reactants
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Rate = k[A]x[B]y[C]z
• x/y/z: reaction orders (reactant with
which it is associated)
− May be zero, positive, negative, integer,
or fraction
− Must be determined by laboratory
experiments
− No relationship to stoichiometric
coefficients of balanced chemical
equation
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concentration-time graph
If reaction is
zero order with
respect to
reactant
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Rate does not
depend on
concentration
of that reactant
Doubling
concentration
will not
increase or
decrease rate
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rate-concentration graph
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Rate directly
proportional to
concentration of
that reactant
Doubling
concentration
doubles rate
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concentration-time graph
If first order with
respect to reactant
rate-concentration graph
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concentration-time graph
If second order
with respect to a
reactant
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Rate is directly
proportional to
square of
concentration of
that reactant
Doubling
concentration
quadruples rate
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rate-concentration graph
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Overall order of a reaction: sum of
individual reaction orders (above
reaction-overall order is x + y + z).
Concentration of
Reactant against Time
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Rate of Reaction against
Concentration of Reactant
Easier to use
rateconcentration
graph rather
than
concentrationtime graph to
decide if
reaction is
first or
second-order
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Units used:
• Rate measured in change in concentration
per second - units are mol· dm-3· s -1
• Units for concentrations of A/B are mol·
dm-3, but if orders are more or less, they
change
• If order is 0, just ignore them
• If order is 1, it's mol· dm-3
• If order is 2, units are [mol· dm-3]2 =
mol2· dm-6
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http://www2.wwnorton.com/college/chemistry
/gilbert/tutorials/interface.swf?chapter=chapt
er_14&folder=reaction_order
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NH4+ + NO2-  N2 + 2 H2O
Experiment
[NH4+]o
[NO2-]o
initial rate
[N2]/s
1
0.100
0.005
1.35 x 10-7
2
0.100
0.010
2.70 x 10-7
3
0.200
0.010
5.40 x 10-7
rate = k[NH4+]1[NO2-]1 or simply k[NH4+][NO2-]
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1.35 x 10-7 M/s = k(0.100 M)(0.005 M)
solving, k = 2.7 x 10-4/Ms or 2.7 x 10-4 L/mol·s
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NO2 + CO  NO + CO2
initial rate
[CO2]/s
Experiment
[NO2]o
[CO]o
1
0.10
0.0010
0.18
2
0.10
0.0020
0.18
3
0.20
0.0020
0.73
rate = k[NO2]2[CO]0 or simply k[NO2]2
0.18 M/s = k(0.10 M)2(0.0010 M)0
k = 18/Ms or 18 L/mol·s
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2 NO2 + 2 H2  N2 + 2 H2O
[NO2]o
[H2]o
1
initial rate
[N2]/s
0.0050 0.0020 1.3 x 10-5
2
3
0.010
0.010
0.0020
0.0040
Experiment
5.1 x 10-5
1.0 x 10-4
rate = k[NO2]2[H2]1 or simply k[NO2]2[H2]
1.3 x 10-5 M/s = k(0.0050 M)2(0.0020 M)1
k = 2.6 x 102/M2s
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Example
•
The initial rate of the reaction BrO3-(aq) + 5 Br-(aq) +
8 H+(aq)  3 Br2(l) + H2O(l) has been measured at
the reactant concentrations shown (in mol/L):
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Experiment [BrO3-] [Br-] [H+] Initial rate (mol/Ls)
1
0.10 0.10 0.10 8.0 x 10-4
2
0.20 0.10 0.10 1.6 x 10-3
3
0.10 0.20 0.10 1.6 x 10-3
4
0.10 0.10 0.20 3.2 x 10-3
Rate = k[A]x[B]y so Rate = k[BrO3-]1[Br-]1[H+]2
Rate = 1 + 1 + 2 = 4
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Example
•
The reaction of iodide ion with hypochlorite ion, OCl- (which is
found in liquid bleach), follows the equation
OCl- + I-  OI- + Cl•
It is a rapid reaction that gives the following rate data:
Initial Concentrations
Rate of Formation (mol/Ls)
[OCl-]
[I-]
[Cl-]
1.7 X 10-3 1.7 X 10-3
1.75 X 104
3.4 X 10-3 1.7 X 10-3
3.50 X 104
1.7 X 10-3 3.4 X 10-3
3.50 X 104
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Determine value of rate constant. Rate = k[A]x[B]y
• 1.75 x 104 = k(1.7 x 10-3)1 (1.7 x 10-3)1
• rate = 6.06 x 109 L/mol sec [OCl-] [I-]
•
Rate = 1 + 1 = 2
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Homework:
• Read 17.3, pp. 542-545
• Q pp. 554-556, #56, 69-70, 74
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Reaction Mechanisms-
how atoms or molecules interact
to make products
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Complex reaction-two or more elementary steps
• Mechanism may never be proven
• Cannot see chemical reaction
• Based on scientific evidence available for reaction
• Guess at identity of many intermediates
Intermediates usually so reactive, can’t be isolated
• Produced in one elementary step and consumed in another
• Does not appear in overall reaction
Mechanism for decomposition of ozone, O3
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•Represents single
collision or vibration that
leads to chemical change
•Combine to produce path
molecules take from reactant(s)
to product(s) in chemical
reaction-sum is overall reaction
•Determines overall rate of
reaction
•Reaction can only go as fast
as its slowest step-highest
activation energy
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• Use rate laws to determine mechanism for
chemical reaction
• A+BY+Z
• Not sure how process actually occurs, but overall
reaction rate expressed by Rate = k[B]
• Predict if following mechanism is valid or invalid
• Step 1: 2A  Y + D
Slow process
• Step 2: B + D  A + Z
Fast process
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• Reaction mechanisms must conform to rules
1. Overall equation for reaction needs to equal sum of
elementary steps
2A  Y + D
B+DA+Z
2A + B + D  Y + D + A + Z
A + B  Y + Z (So far, looks good)
2. Rate law for rate-determining step should match rate
law for overall reaction
• 2A  Y + D
• Rate = k[A]2 (doesn’t match rate = k[B])-invalid
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Rate determining step
• Overall reaction
• 2 NO2 + F2 --> 2 NO2F
• Mechanism (two elementary steps)
• NO2 + F2  NO2F + F (slow)
• F + NO2  NO2F (fast)
• Rate determining step (RDS) determines overall
reaction rate law
• Rate overall  rate for rds step = k1 [NO2] [F2]
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Homework:
• Read 17.4, pp. 546-549
• Q pp. 554-556-#57, 76
• Do test, pg. 557
• Use link for quiz and submit as before.
http://www.glencoe.com/qe/science.php?qi=1001
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