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Rate of Reactions{Chemical Kinetics}
Chemical Kinetics
• Chemical kinetics is the branch of chemistry concerned with the rates of chemical reactions
• A rate is a change in a measurable quantity over time
• Different chemical reactions proceed from reactants to products at different rates (e.g. combustion of propane and oxidation of silver)
Reaction Rate
• The reaction rate of a chemical reaction is the change in concentration of a reactant or product per unit time
• Δc is measured in mol/L
• Δt is measured in s
• r is measured in mol/(Ls)
Fast or Slow Reactions
• Extremely slow reactions– Iron rusting– Limestone weathering
• Extremely fast reactions– Explosion
Measuring Rate of Reactions
• Some rate of reactions have detectable change with respect to time
• Changes that are observable like– When a volume of gas is given off– When this is a change in mass during the
reaction– When there are temperature changes– When there are colour changes– When a precipitate forms– When there are pH changes
Speed= Speed=
=
= 300km/h
=
= 10km/h
According to Collision Theory what factors determine whether a reaction happens or not?
According to Thermochemistry, what two factors determine whether a reaction is favoured?
Collision theory
• Theory that explains how factors affect rate of reaction
• Before reaction can occur, particles must first collide
• Particles collide so that bonds are broken and new bonds can form
Collision theory
• Collision of particles must be effective collisions to produce result
• Effective collisions are collisions which – produce enough energy to overcome
energy of activation– correct orientation
Orientation
Collision theory
• Activation energy is the minimum energy the colliding particles must overcome so that reaction occur
• In order for particles to overcome the activation energy, several factors are involved
• So lets say collision theory is determining that a reaction can happen….what factors determine how fast it will happen?
• Mr McCormack is cooking a roast dinner but his over-sized potatoes are not boiling fast enough….what measures can he take to increase the rate of reaction???
How factors affect rate of reaction according to collision theory
How factors affect rate of reaction according to collision theory
• Size of reactant for solid reactant• Temperature of reactant mixture• Concentration of solution reactant• Presence of catalyst• Pressure on the reaction
Concentration
• Rate of reaction increases with increasing concentration
• Higher concentration, more reacting particles are present
• Greater probability of an effective collision
• Faster rate of reaction
Concentration – Same no. of molesAmt of product formed
Time/s
Higher concentration
Lower concentration
Concentration – Different no. of molesAmt of product formed
Time/s
Higher concentration
Lower concentration
Temperature
• Rate of reaction increases with increasing temperature
• High temperature, particles have greater heat energy
• Particles move faster with greater kinetic energy
• Leading to more collisions between particles
• Increased probability of effective collision• Reactions take place faster
• Speed of reaction doubles when the temperature
rises by 10 C
TemperatureAmt of product formed
Time
Higher temperature
Lower temperature
Catalyst
• Presence of catalyst increases rate of reaction
• (Presence of inhibitors decreases rate of reaction)
• Catalysts lower activation energy of reactants
• Aids the formation of unstable intermediate products
• Increases probability of formation of products
• Faster rate of reaction
POTENTIAL ENERGY DIAGRAM
REACTION COORDINATE
POTETIAL
ENERGY
REACTANTS
PRODUCTS
A + B
AB
ACTIVATIONENERGYWITHOUTCATALYST
LOWER ACTIVATONENERGY PATH
WITH CATALYST
CatalystEnergy
Time/s
Use of catalyst
Absence of catalyst
Ea
Ea
CatalystAmt of product formed
Time/s
Use of catalyst
Absence of catalyst
Catalyst
Definition: A substance which increases the rate of a chemical reaction by providing an alternative pathway with a lower activation energy but remains unchanged at the end of the reaction
Pressure
• Rate of reaction increases with increasing pressure
• Higher pressures, reacting particles are closer together
• Increasing concentration per unit volume
• Greater probability of an effective collision
• Faster rate of reaction
PressureAmt of product formed
Time/s
Higher pressure
Lower pressure
Particle Size
• Rate of reaction increases when particle size decreases
• Smaller particles has greater surface area than larger particles of the same mass
• Greater surface area for collision by another reacting particle
• Greater probability of an effective collision• Faster rate of reaction
HOW PARTICLE SIZE AFFECTS CHEMICAL REACTION RATE!
Particle sizeAmt of product formed
Time/s
Smaller particle size
Larger particle size
A reaction is fast , the time taken for the reaction is short .
A reaction is slow, the time taken for the reaction is long .
The rate of reaction depends to the speed of reaction .
If a reaction is fast, its rate of reaction is high .
If a reaction is slow, its rate of reaction is low .
The rate of reaction is inversely proportional with time .
Rate ά 1
time taken for reaction
Rate of reaction = change in quantity of product / reactant
time taken
For gas product , rate of reaction = volume of gas
time
From a graph , the average rate of reaction = the gradient
HOW TO DETERMINE RATE OF REACTION
FROM GRAPH
AVERAGE RATE OF
REACTION
INSTANTANEOUS RATE OF REATION
AVERAGE RATE OF REACTION
For the whole exp
From X
min to Y min
On the X min (2nd) ( from 2nd to 1st)
For first x
(3) min( from 0 to
3rd )
INSTANTANEOUS RATE OF REACTION
( rate of reaction at that time) Draw tangent to the
graph
Y
X
Rate = Y/X
b) Example from the graph, determine:i) The rate of reaction at 120 s
Instantaneous rate of reaction= Draw tangent to the graph
= 56 – 20 = 0.176 cm3 s-1
222-18
Changes to the graph
Changes to the CURVE part of graph
II
IIII
Volume of gas
Time
• Use positive catalyst• Increase temperature •Increase total surface area
Use negative catalyst
Decrease temperature
Decrease TSA
Changes to the graph
d) Curve I represents the result of the experiment using excess zinc powder and 50cm3 of 1.0 moldm-3 dilute hydrochloric acid
II
IIII
Volume of gas/ cm3
Time/s
• Use positive catalyst• Increase temperature of reactant
Lower concentration of hydrochloric acid
(h)Time (seconds) 0 30 60 90 120 150 180 210 240Burette reading (cm3) 49.5 33.5 23.5 16.0 10.5 5.0 2.0 2.0 2.0Volume CO2 (cm3) 0 16 26 33.5 39 44.5 47.5 47.5 47.5
Time(second) 0 30 60 90 120
Burette reading(cm3)
x
49.5
y
33.5
z
23.5
Total volume of gas(cm3)
x-x
0.00
x-y
16.00
x-z
26.00
Volume of CO2, cm3
Time , s
Connect the point without using ruler!Not all the point is connected
Volume of CO2 cm3
Time s
Cannot make this graphStraight lineIt’s must be
smooth graph
Average Rate Of reaction
The average rate of reaction in the first 90 seconds.= The total volume of gas released in the first 90 seconds
Time taken
(i)
Time (seconds) 0 30 60 90 120 150 180 210 240Burette reading (cm3) 49.5 33.5 23.5 16.0 10.5 5.0 2.0 2.0 2.0Volume CO2 (cm3) 0 16 26 33.5 39 44.5 47.5 47.5 47.5
=
33.5÷90=0.372
cm3s-1unit
i(ii)Time (seconds) 0 30 60 90 120 150 180 210 240Burette reading (cm3) 49.5 33.5 23.5 16.0 10.5 5.0 2.0 2.0 2.0Volume CO2 (cm3) 0 16 26 33.5 39 44.5 47.5 47.5 47.5
47.5÷180= 0.264 cm3s-1
The average rate of reaction in the whole experiment.= The total volume of gas released in the whole experiment
Time taken
=
Total volume of Hydrogen gas/cm3
Time (second)
Analysis of Data
t
p
q
Rate of reaction at t second = gradient AB
= p/q cm3 s-1
A
B
Tangent
Cannot take directly at x
Tangent is a line that touch just 1 point of graph in order to calculate gradient
Tangent
Only touch 1 point of curve
Cannot touch more than 2 points because each of point has a different gradient
tangentα
Total Volume of CO2(cm3)
Time (second)
Analysis of data
A
B
C
D
E
F
t1 t3t2
Rate of reaction at t1 = gradient ABRate of reaction at t2 = gradient CDRate of reaction at t3 = gradient EFEach of point has a
different gradient!
Two methods to calculate tangent:
Total volume of Hydrogen gas/cm3
Time (second)
A
B
Tangent
number of small
boxes × value of 1 small unit box
Y
X
Total volume of Hydrogen gas/cm3
Time (second)
A
B
Tangent
Gradient of graph:
m = ΔY
ΔX
x1 x2
y2
y1
m = Y2-y1
X2-x1
First Method
Total volume of Hydrogen gas/cm3
Time (second)
Analysis of Data
t
p
q
Rate of reaction at t second = gradient AB
= p/q cm3 s-1
A
B
Tangent
Total Volume of CO2(cm3)
Time (second)
Analysis of data
A
B
C
D
E
F
t1 t3t2
Rate of reaction at t1 = gradient ABRate of reaction at t2 = gradient CDRate of reaction at t3 = gradient EF
The Rate Law: Reactant Concentration and Rate
The rate of a chemical reaction depends on several
factors ….
Relating Reactant Concentrations and Rate
Consider the general reaction below.
aA + bB cC + dD
This reaction occurs at constant temperature
The reactant formulas are represented by A and B
The stoichiometric coefficients are represented by
a and b
In this section we will look at reaction rates that are not
affected by concentrations of products
In general, the rate of a reaction increases when the
concentrations of reactants increases.
The dependence of of the rate of a reaction on the con-
centration is given by:
OR Rate=k[A]m
This relationship can be expressed in a general
equation called the rate law equation
For any reaction, the rate law equation expresses the
relationship between the concentrations of the
reactants and the rate of the equation.
The letter k represents the proportionality constant
called the rate constant
There is a different rate constant for each reaction at
any given temperature
The exponents m and n must be determined by exp-
eriment.
The do not necessarily correspond to the coefficients
of their reactants
They are usually 1 or 2, but values of 0, 3 even fractions
can occur
HOW CONCENTRATION EFFECTS REACTION RATES (REACTION ORDER)
• REACTIONS WITH RATE EQUATIONS HAVING n = 0 ARE ZERO ORDER REACTIONS. THOSE WITH n = 1 ARE FIRST ORDER AND THOSE WITH n = 2 ARE SECOND ORDER.
• IN ZERO ORDER REACTIONS, CHANGING THE CONCENTRATION OF THE REACTANT HAS NO EFFECT ON THE RATE.
• IN FIRST ORDER REACTIONS, RATE CHANGES ONE FOR ONE WITH CONCENTRATION CHANGE. FOR EXAMPLE, DOUBLING CONCENTRATION DOUBLES THE RATE.
HOW CONCENTRATION EFFECTS REACTION RATES (REACTON ORDER)
• IN SECOND ORDER REACTIONS, RATE CHANGES RELATIVE TO THE SQUARE OF THE CONCENTRATION CHANGE. FOR EXAMPLE, DOUBLING THE CONCENTRATION OF THE REACTANT RESULTS IN THE RATE INCREASING 4 TIMES.
• THIS KNOWLEDGE OF HOW RATE CHANGES WITH CONCENTRATION DEPENDING ON THE ORDER LETS US FIND REACTION ORDERS BY AN EXPERIMENTAL PROCESS CALLED “METHOD OF INITIAL RATES”
• REACTION ORDERS MUST BE DETERMINED EXPERIMENTALLY. THEY CAN NEVER BE DETERMINED FROM THE CHEMICAL EQUATION.
RXN RATE?
RXN RATE ?FORWARD OR
REVERSE RXN?
RXN RATE?FORWARD OR REVERSE?
RATE = 0
RXN RATE?FORWARD OR REVERSE?
REACTION RATES
TIME
MOLES
A
TIME
MOLES
A
TIME
MOLES
A
TIME
MOLES
A
A B + C
FORWARD RXNRATE = CONSTANT
FORWARD RXNRATE = VARIABLE REVERSE RXN
RATE = VARIABLE
SLOPE OF A TANGENT LINE TO AN AMOUNT VS. TIME GRAPH = RATE
GRAPH 1 GRAPH 2
GRAPH 3 GRAPH 4
HOW CONCENTRATION EFFECTS REACTION RATES (INITIAL RATES)
• USING THE METHOD OF INITIAL RATES REQUIRES THAT A REACTION BE RUN AT SERIES OF DIFFERENT STARTING CONCENTRATIONS AND THE RATE BE DETERMINED FOR EACH.
• GIVEN THE FOLLOWING DATA FOR THE REACTION A B + C(TABLE 1)• EXPT [A] RATE (M/SEC) 1 1 x 10 -3 4 x 10 -1
2 2 x 10 -3 8 x 10 -1
3 4 x 10 -3 16 x 10 -1
• AS CONCENTRATION OF A DOUBLES, RATE DOUBLES. THE REACTION IS FIRST ORDER IN REACTANT A
• RATE = k[A]1 OR RATE = k[A]
HOW CONCENTRATION EFFECTS REACTION RATES (INITIAL RATES)
• FOR THE REACTION: A + B C + D• (TABLE 2) 1 1 x 10 -3 1 x 10 –3 4 x 10 -1
2 2 x 10 -3 1 x 10 -3 8 x 10 -1
3 1 x 10 -3 2 x 10 -3 16 x 10 –1
• USING EXPT 1 AND 2, [A] DOUBLES AND [B] IS CONSTANT. THE DOUBLING OF THE RATE IS THEREFORE CAUSED BY REACTANT A AND THE ORDER WITH RESPECT TO A IS FIRST.
• USING EXPT 1 AND 3, [A] IS CONSTANT AND [B] IS DOUBLED. THE FOUR TIMES RATE INCREASE IS THEREFORE CAUSED BY REACTANT B AND THE ORDER WITH RESPECT TO B IS SECOND.
• RATE = k[A]1[B]2 OR RATE = k[A][B]2
HOW CONCENTRATION EFFECTS REACTION RATES (RATE CONSTANTS)
• FROM INITIAL RATES DATA TABLE 1,
• RATE = k [A] , k CAN BE CALCULATED BY SUBSTITUTING ANY DATA SERIES INTO THE RATE EQUATION, FOR EXAMPLE, FROM EXPT 1 ON TABLE 1
• [A] = 1 x 10 –3 M , RATE = 4 x 10 –1 M/SEC• 4 x 10 –1 M/SEC = k (1 x 10 –3 M ) • k = 1 x 102 SEC-1 OR 1 x 102 / SEC
HOW CONCENTRATION EFFECTS REACTION RATES (RATE CONSTANTS)
• FROM INITIAL RATES DATA TABLE 2,
• RATE = k [A] x [B]2, k CAN BE CALCULATED BY SUBSTITUTING ANY DATA SERIES INTO THE RATE EQUATION, FOR EXAMPLE, FROM EXPT 1 ON TABLE 2
• [A] = 1 x 10 –3 M , [B] = 1 x 10 –3 M
• RATE = 4 x 10 –1 M/SEC
• 4 x 10 –1 M/SEC = k (1 x 10 –3 M ) x (1 x 10 –3 M )2
• k = 4 x 105 M-1 SEC-1 OR 4 x 105 / M x SEC
Let us see how the initial rates method works
2N2O3(g) 2NO2(g) + O2(g)
The general rate law equation for this reaction is:
Rate = k [N2O3]m
To determine the value of m, a chemist performs three
experiments. A different initial concentration of [N2O35]0
Is used for each experiment. The 0 represents t=0
Experiment Initial [N2O3]0 (mol/L) Initial rate (mol/(L.s))
1 0.010 4.8 x 10 -6
2 0.020 9.6 x 10 -6
3 0.030 1.5 x 10-5
Value of m can be determined with at least two different methods
by inspection rate law equation
1. By inspection
When the [N2O5] is doubled expts 1 and 2 doubles
the rate also doubles
when [N2O5] is tripled, (expts 1 and 3) the rate triples
this indicates a first-order relationship as follows
Rate = k [N2O5]1
2. Compare rate law equation using ratiosthis method very useful when relation between conc and rate are not
immediately obvious from data
Write the rate expressions for expts 1 and 2 as follows:
Rate1 = k [0.010]m
= 4.8 x 10-6 mol/(L.s)
Rate2 = k [0.020]m
= 9.6 x 10-6 mol/(L.s)
Create a ratio to compare the two rates
Rate1 = k(0.010 mol/L)m = 4.8 x 10-6 mol/L.s
Rate 2 k(0.020 mol/L)m 9.6 x 10-6 mol/L.s
Since k is a constant at constant temp you can cancel out
k(0.010 mol/L)m = 4.8 x 10-6 mol/L.s k(0.020 mol/L)m 9.6 x 10-6 mol/L.s
(0.5)m = 0.5 m = 1 (by inspection)
Determining the Rate Constant
Once you know the rate law equation for a reaction,
you can calculate the rate constant using results from
any of the experiments
Rate = k [N2O5]1
You can use data from any of the three experiments to
calculate k
4.8 x 10-6 mol/(L.s) = k(0.010 mol/L)
k = 4.8 x 10-6 mol/ (L.s) 0.010 mol/L
= 4.8 x 10-4 s-1
TEMPERATURE & REACTION RATE• AT ANY TEMPERATURE THE MOLECULES IN A SYSTEM HAVE A
DISTRIBUTION OF KINETIC ENERGIES (SIMILAR TO THE DISTRIBUTION OF THE SPEEDS OF CARS ON A HIGHWAY).
• AS THE TEMPERATURE INCREASES, THE AVERAGE KINETIC ENERGY OF THE MOLECULES INCREASE (THEY MOVE FASTER AT HIGHER TEMPERATURES) AND THEREFORE COLLIDE WITH EACHOTHER MORE FREQUENTLY AND HIT HARDER WHEN THEY DO COLLIDE.
• AS A RESULT, REACTION RATE INCREASES WITH TEMPERATURE.
KINETIC ENERGY DISTRIBUTION CURVE
TEMPERATURE = T1
KINETIC ENERGY
NUMBER
OF
MOLES
LOW ENERGYMOLECULES
HIGH ENERGYMOLECULES
AVERGE ENERGYMOLECULES
KINETIC ENERGY DISTRIBUTION CURVE
TEMPERATURE = T1
TEMPERATURE = T2
KINETIC ENERGY
NUMBER
OF
MOLES
TEMP 2 > TEMP 1AT HIGHER
TEMPERATURES, MOLECULES HAVE HIGHER ENERGIES
ON AVERAGE
TEMPERATURE & REACTION RATE
• IN ORDER TO REACT, MOLECULES MUST COLLIDE WITH SUFFICIENT ENERGY. THIS MININIUM ENERGY FOR REACTION IS CALLED ACTIVATION ENERGY.
• AS THE TEMPERATURE OF A SYSTEM IS INCREASED, THE NUMBER OF MOLECULES WITH THE NECESSARY ENERGY FOR REACTION (THE ACTIVATION ENERGY) INCREASES.
KINETIC ENERGY DISTRIBUTION CURVE
TEMPERATURE = T1
TEMPERATURE = T2
KINETIC ENERGY
NUMBER
OF
MOLES
ACTIVATION ENERGY
MOLECULES WITHSUFFICIENT ENERGY
TO REACT AT T1
AS TEMPERATURE , REACTION RATE
MOLECULES WITHSUFFICIENT ENERGY
TO REACT AT T2
TEMPERATURE & REACTION RATE• THE ENERGY CHARACTERISTICS OF A CHEMICAL
REACTION CAN BE SHOWN ON A POTENTIAL ENERGY DIAGRAM.
• THIS GRAPH SHOWS THE ENERGY STATE OF THE SYSTEM AS REACTANTS PROCEED THROUGH THE ACTIVATED COMPLEX TO FORM THE PRODUCTS. THE ACTIVATED COMPLEX IS THE INTERMEDIATE STATE (MOLECULAR FORM) WHICH REACTANTS GO THROUGH AS THEY CONVERT INTO THE PRODUCTS.
• THE ACTIVATION ENERGY IS THE ENERGY REQUIRED TO FORM THE INTERMEDIATE ACTIVATED COMPLEX MOLECULE.
POTENTIAL ENERGY DIAGRAM
REACTION COORDINATE
POTETIAL
ENERGY
REACTANTS
PRODUCTS
ACTIVATEDCOMPLEX
A + B
AB
H
ACTIVATIONENERGY
EXOTHERMIC
H = (-)ENERGY ISRELEASED
The Effect of Surface Area on Reaction Rate
Which will cook faster?
One 60g whole potato…
…or 60g of small potato
pieces?
The small potato pieces…
Why do they cook faster?
Because if you compare a 60g potato with 60g of small potato pieces, the small potato pieces
have a LARGER SURFACE AREA
So what?
• The larger the surface area, the more collisions that can happen between the boiling water and the potato at the same time
• This means the potato cooks faster
SURFACE AREA & REACTION RATE
REACTANT THAT IS INTERIOR CANNOT BE ATTACKED UNTIL THE EXTERIOR REACTANT IS CONSUMED. THE REACTION RATE IS SLOW!
SURFACE AREA & REACTION RATE
WHEN SURFACE AREA IS INCREASED MORE REACTANTS ARE EXPOSEDTO EACHOTHER SIMULTANEOUSLY AND THE REACTION IS RAPID.
REACTION MECHANISMS• ALL REACTIONS, NO MATTER HOW SIMPLE FOLLOW A PARTICULAR
REACTION PATHWAY CALLED A MECHANISM. IT IS A SERIES OF STEPS WHICH LEAD TO THE FORMATION OF PRODUCTS.
• DURING THE PROCESS OFTEN MOLECULES CALLED INTERMEDIATES ARE FORMED AND SUBSEQUENTLY CONSUMED.
• THE SLOWEST STEP IN THE SERIES OF STEPS THAT MAKE UP THE MECHANISM IS CALLED THE RATE DETERMINING STEP.
• THE SPEED OF THE RATE DETERMINING STEP DEPENDS ON THE COMPLEXITY OF THE STEP (NUMBER OF MOLECULES INVOLVED CALLED THE MOLECULARITY) AND THE BOND STRENGTHS OF THE REACTING COMPONENTS
NO
NONO2
NO2
O2
REACTION MECHANISM FOR;2 NO + O2 2 NO2
STEP I 2 NO N2O2 (fast)STEP II N2O2 + O2 2 NO2 (slow)
N2O2 INTERMEDIATE
REACTION MECHANISMS• UNLIKE OVERALL REACTIONS, THE REACTON ORDERS FOR
THE STEPS IN A MECHANISM ARE BASED ON THE NUMBER OF MOLECULES REACTING IN THAT STEP.
IN THE REACTION MECHANISM FOR:2 NO + O2 2 NO2
STEP I 2 NO N2O2 (fast)STEP II N2O2 + O2 2 NO2 (slow)
SINCE THERE IS ONLY ONE N2O2 AND ONE O2 IN THE RATE DETERMING STEP (THE SLOW STEP), THE RATE EQUATION FOR THE REACTION IS:
RATE = k [N2O2] x [O2]
THE REACTION IS FIRST ORDER IN BOTH REACTANTS
