Chem342 Environmental Chemistry - science.kln.ac.lkscience.kln.ac.lk/Chemistry/Teaching_Resources/Documents/CHEM 1… · CHEM 11132 Basic Physical Chemistry I Course content: Thermodynamics

Chemical Kinetics

CHEM 11132

Dr. M. P. Deeyamulla

([email protected])

CHEM 11132 Basic Physical Chemistry I

Course content:

Thermodynamics (10 L)

Electrochemistry (10 L)

Chemical Kinetics (10 L) :

Basic concepts; rates of reactions, elementary reactions, rate expressions, order and

the rate constant of a reaction, molecularity. Experimental determination of rate

laws; fitting data to rate laws, obtaining data for different timescales. Introduction to

theories about reaction rates; collision theory and activated complex theory.

Complex reactions and reaction mechanisms; rate determining steps, pre-

equilibrium hypothesis, steady-state approximation and their applications.

Temperature dependence of reaction rates: Arrhenius rate law and deviation. Chain

reactions, fast reactions and catalysis.

CHEM 11111 Calculations in Chemistry

Chemistry…….....…?

CHEMISTRY…

IS THIS WHAT YOU THINK?

Linus Pauling

‘Chemistry is wonderful.

I feel sorry for people who don’t know

anything about chemistry. They are missing

an important source of happiness.

Most people don’t feel that.’

Chemical Reactions

• Definition: A Process which produces chemical change.

• Review- Chemical vs. Physical Change

• Bonds are broken

• Reactants and Products

Reactants Products

3 Major Groupings of Chemical

Reactions

1. Precipitation

Reactions

2. Oxidation-Reduction

Reactions

3. Acid-Base

Neutralization

Reactions

SOLID - SOLID ?

Elephant’s Toothpaste

Mix: Saturated Solution of KI with 30%

solution of Hydrogen Peroxide and

detergent.

)g(OOH2)aq(OH222

KI

22

Chemistry

Biology

Plant Sciences

Geology

Environmental Science

Health and MedicineNuclear Chemistry

Physics

Astronomy

Biochemistry

Biology

“The Central Science”

Chemistry in the Home

Everything in your home is Chemistry

Cooking is Chemistry

Chemical changes are responsible for changes in flavour and texture

Chemistry keeps food fresh

Water-Wine-

Milk-Beer

CHEM 11132 Basic Physical Chemistry I

Course content:

Thermodynamics (10 L)

Electrochemistry (10 L)

Chemical Kinetics (10 L) :

Basic concepts; rates of reactions, elementary reactions, rate expressions, order and

the rate constant of a reaction, molecularity. Experimental determination of rate

laws; fitting data to rate laws, obtaining data for different timescales. Introduction to

theories about reaction rates; collision theory and activated complex theory.

Complex reactions and reaction mechanisms; rate determining steps, pre-

equilibrium hypothesis, steady-state approximation and their applications.

Temperature dependence of reaction rates: Arrhenius rate law and deviation. Chain

reactions, fast reactions and catalysis.

CHEM 11111 Calculations in Chemistry

Textbooks

P. W. AtkinsThe Elements of Physical Chemistry (Third Ed., Ch. 10)

P. W. AtkinsPhysical Chemistry

Michael J. Pilling & Paul W. Seakins

Reaction Kinetics

Any Physical Chemistry Text Book; Levine, Daniel &

Alberty, Barrow.

Chemistry…….....…?

Chemical reaction……?

slide 1.ppt

Combustion Reactions with Hydrogen

2H2 + O2 → 2H2O

Chemical reaction……• bonds break - this requires energy

• bonds form - this releases energy

• overall for the reaction– exothermic reaction - releases energy

– endothermic reaction - absorbs energy

slide 1.ppt

"Everyone has Problems -

but Chemists have Solutions"

Will the reaction occur?

How far will the reaction proceed?

How fast will the reaction proceed?

How?

Chemical kinetics

Kinetics …

• studies the rates at which chemical

reactions occur.

• gives information about how the

reaction occur, that is, the reaction

mechanism

CHEMICAL KINETICS

• Area of chemistry concerned with rates of reactions

– How rapidly food spoils

– Rate of fuel burning in automobiles

– How quickly medicines work

– Development of catalysts

© Boardworks Ltd 200740 of 39

What does rate of reaction mean?

The speed of different chemical reactions varies hugely.

Some reactions are very fast and others are very slow.

What is the rate of these reactions?

The speed of a reaction is called the rate of the reaction.

rusting baking explosion

slow fast very fast

Airbag Reaction

INTERNATIONAL FOOD POLICY RESEARCH INSTITUTE

65g of sodium azide ≈ N2 (35 litres)

fully inflated within about 30 milliseconds

Airbag reaction


Requirements for reaction

• Molecules must meet (collisions)

• Molecules must transfer enough energy to

overcome the activation barrier

• They must meet in the right orientation

Activation Energy, Ea

Energy barrier (hump) that must be

overcome for a

chemical reaction to proceed

A2(g) + B2(g) 2AB(g)

How is the reaction going to occur?

Activated complex-hypothetical

an effectivecollision

comesapart

Animations

Chemistry 30 Chemical Kinetics - The Collision Theory.swf

collis11.swf


Factors That Affect Reaction Rates

1. Physical state of the reactants

Important when reactants are of different phases (higher SA faster reaction)

2. Concentration of Reactants

For most reactions higher [reactant] faster reaction

3. Temperature

Increasing T will also increase the reaction rate

4. Presence of Catalyst


Effect of concentration on rate of reaction

The higher the concentration of a dissolved reactant, the

faster the rate of a reaction.

Why does increased concentration increase the rate of

reaction?

At a higher concentration, there are more particles in the

same amount of space. This means that the particles are

more likely to collide and therefore more likely to react.

higher concentrationlower concentration

The ratio of successful collisions to unsuccessful collisions

will stay the same, but there will be more successful

collisions.


Concentration and particle collisions


Animations

Effect of temperature and catalyst on the reaction rate.swf


Rate of reaction

reaction rate: change in concentration of a

product or a reactant per unit time.

Rate – change in some variable per unit time

timerate

1

Δt

t in c, ion,concentrat in change

Rate = ______________ = ______________change in timechange in time

in [products] in [reactants]

reactant

Rates of Chemical Reactions

Rate of a chemical reaction refers to the change in

concentration of a substance per unit of time

Let’s consider the rate at which you give me 3 rupees….

your 3 rupees my 3 rupees

Let’s say that it took you 5 seconds to give it to me.

your 3 rupees my 3 rupees

reactants products

What is the rate of the reaction with respect to me [products]?

Rate = change in concentration of money

change in time

Remember, change () is always [ final – initial]

Rate = + [3-0 rupees]

[5-0 secs]

Positive because I am the product

which gains the money

Rate = 0.6 rupees/sec

What is the rate with respect to you?

Rate of reaction = [0-3 rupees]

[5-0 secs]

Negative because you are the

reactant and you are losing money

Rate of reaction = 0.6 rupees/sec

***Therefore, you can determine the rate of reaction either

by using the reactants or the products. It will give you the

same rate of reaction****



rate = + [product] = - [reactant]

t t

reactant

Let’s consider the rates for chemical reaction

NO(g) + ½ O2 (g) NO2(g)

Rate of the disappearance of NO:

Rate = -[NO]

t Rate of the disappearance of O2:

Rate = -[O2]

t

Rate of the appearance of NO2:

Rate = +[NO2]

t

67

© 2009 Brooks/Cole - Cengage

Reaction rate = change in concentration of a reactant or product with time.

–initial rate

–average rate

–instantaneous rate

Reaction Rates


What equipment is needed to investigate the rate of

hydrogen production?

gas syringe

rubber bung

rubber connecterglass tube

conical

flask

magnesium

hydrochloric

acid


..\Animations\rxnRate01.swf


hydro

gen

pro

duce

d (

cm3)

time (seconds)10 20 30 40 50

10

20

30

40

50

60

70

0

0

x

y

Calculating rate of reaction from graphs

rate of reaction =

x

y

rate of reaction =

20s

45cm3 rate of reaction = 2.25cm3/s

The gradient of the graph is equal to the initial rate of reaction

at that time

How can the rate of reaction be calculated from a graph?

c

Δt

Chemical

Kinetics

Reaction Rates

The average rate of

the reaction over

each interval is the

change in

concentration divided

by the change in time:

C4H9Cl(aq) + H2O(l) C4H9OH(aq) + HCl(aq)

Average Rate, M/s

Chemical

Kinetics

Reaction Rates

• A plot of concentration

vs. time for this reaction

yields a curve like this.

• The slope of a line

tangent to the curve at

any point is the

instantaneous rate at

that time.


Chemical

Kinetics

Reaction Rates

• The reaction slows

down with time because

the concentration of the

reactants decreases.


Br2 (aq) + HCOOH (aq) 2Br- (aq) + 2H+ (aq) + CO2 (g)

average rate = -[Br2]

t= -

[Br2]final – [Br2]initial

tfinal - tinitial

slope of

tangentslope of

tangentslope of

tangent

instantaneous rate = rate for specific instance in time

From last two lectures……..

• Reaction Rates

initial rate - instantaneous rate at t = 0

average rate - Δ[A] over a specific time interval

instantaneous rate - rate at a specific time

The rate is the instantaneous slope, and this varies

with time

Chemical

Kinetics

Reaction Rates

• Note that the average

rate decreases as the

reaction proceeds.

• This is because as the

reaction goes forward,

there are fewer

collisions between

reactant molecules.




rate = + d[product] = - d[reactant]

dt dt

reactant

i.e.

[(amount of material)(volume)-1] [time]-1

common units: mol dm-3 s-1

Units

ratio of concentration upon time,

Exercise 1Symbolically, [A], [Br]

rate = + [ ]

t

Reaction Rates and Stoichiometry

• The molar ratios between reactants and products

correspond to the rates of reaction.

• Relative rates – relationship between rates of

reactant disappearance and product appearance at a

given time.

O2 + 2H2 →2 H2O

Or, equivalently

dt

Ad1r

A

Where vA is the stoichiometric coefficient of

species A

example: express the rate of reaction for each

reactant and product in the reaction:

4 NH3(g) + 5 O2(g) 4 NO(g) + 6H2O (g)

d[NH3]

dt- 1

4NH3:

d[O2]

dt- 1

5O2:

d[NO]

dt+ 1

4NO:

d[H2O]

dt+

16

H2O:

Reaction Rates and Reaction

Stoichiometry

Look at the reaction

O3(g) + NO(g) NO2(g) + O2(g)

dt

]Od[+ =

dt

]NOd[+ =

dt

d[NO]- =

dtOd

- = rate 223

Another Example

2 NOCl (g) 2 NO + 1 Cl2 (g)

dt

d[Cl+ =

dt

d[NO]

2

1 =

dt

NOCld

2

1- = rate 2 ]

WHY? 2 moles of NOCl disappear for every 1 mole Cl2 formed.

The General Case

a A + b B c C + d D

rate = -1 d[A] = -1 d[B] = +1 d[C] = +1 d[D]

a dt b dt c dt d dt

Why do we define our rate in this way?

Obtain a single rate for the entire equation, not

just for the change in a single reactant or product.

Chemical

Kinetics

Concentration and Rate

Each reaction has its own equation that

gives its rate as a function of reactant

concentrations.

this is called its Rate Law

To determine the rate law we measure the rate

at different starting concentrations.

Chemical

Kinetics


Compare Experiments 1 and 2:

when [NH4+] doubles, the initial rate doubles.

Chemical

Kinetics


Likewise, compare Experiments 5 and 6:

when [NO2-] doubles, the initial rate doubles.

Chemical

Kinetics


This equation is called the rate law, and

k is the rate constant.

RATE LAW

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 LAW

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 = k[A]m[B]n

where k is called the rate constant and is

independent of concentration

k increases with T

The rate law expresses the relationship of the rate of a

reaction to the rate constant and the concentrations of

the reactants raised to some powers.

The reaction orders are empirically determined.

For the general reaction:

the rate equation

aA + bB + cC … mM + nN ….

dt= k [A]x [B]y [C]z …

-d[A]

• the rate law can only be determined by experiment, not from the stoichiometric equation

• x is the order of the reaction with respect to A,y is the order of the reaction with respect to B…

• the overall order of the reaction is given by x + y + z …

• there is no relationship between a and x, b and y ….

slide 2.ppt

slide 2.ppt

slide 2.ppt

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

example

H2 (g) + 2 ICl (g) 2 HCl (g) + I2 (s)

dt= k[H2][ICl]rate =

-d[H2]from experiment:

the reaction is:

• first order with respect to H2

• first order with respect to ICl

• second order overall

questionfor the reaction:

2 NO (g) + O2 (g) 2 NO2 (g)

rate = k[NO]2[O2]

what is the order of reaction with respect to the

reactants and the overall order of reaction?

• second order with respect to NO

• first order with respect to O2

• third order overall

Units

Reaction order enables us to understand how the reaction depends on reactant concentrations.

aA + bB bC + dD

This reaction is zero order in A, first order in B, and first order overall. The exponent zero tells us that the rate of this reaction is independent of the concentration of A.

www.deakin.edu.au<[email protected]>

107

Example : A + B → C

Changing [B] ⇒ no effect ⇒ rate ∝ [B]0

Double [A] ⇒ rate ∝ 22 ⇒ rate ∝ [A]2

Rate = k [A]2 [B]0 = k [A]2

Quick check

Experiment [A] / M [B] / M Init. rate / M s-1

1 0.10 4.0 × 10-5

2 4.0 × 10-5

0.10

0.10 0.20

3 16 × 10-50.20 0.10

Method of Initial Rates

a. Determine the rate law of the reaction

b. Calculate the rate constant

c. Calculate the rate when [NO]= 0.050 M and [H2]

= 0.150

Temperature and Rate

• Generally, as

temperature

increases, so does

reaction rate.

• This is because k is

temperature

dependent.

Temperature and Rate

• In a chemical reaction, bonds are broken and new

bonds are formed. In order for molecules to react,

they must collide.

• Collisions are either effective or ineffective due to

orientation of molecules.

• Collisions must have enough energy to overcome the

barrier to reaction, the activation energy.

• Temperature affects the number of collisions.

slide 3.ppt

• Not all collisions leads to a reaction

• For effective collisions proper orientation of

the molecules must be possible

Chemical

Kinetics

Maxwell–Boltzmann Distributions

• Temperature is

defined as a

measure of the

average kinetic

energy of the

molecules in a

sample.

• At any temperature there is a wide distribution of

kinetic energies.

Chemical

Kinetics


• As the temperature

increases, the curve

flattens and

broadens.

• Thus at higher

temperatures, a

larger population of

molecules has

higher energy.

Chemical

Kinetics


• 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.

Chemical

Kinetics


This fraction of molecules can be found through the expression:

where R is the gas constant and T is the temperature in Kelvin .

Arrhenius Equation

Arrhenius developed an equation for the mathematical

relationship between k and Ea.

RT

EexpAk a

Ea = activation energy (kJ mol-1), and is the

minimum kinetic energy required to allow reaction to occur

A = the frequency factor or pre-exponential factor (same units as k),

is the fraction of sufficiently energetic collisions that actually

lead to reaction.

T = Kelvin temperature

R = ideal gas constant (8.314 J mol-1 K-1)

k is the rate constant

Ea/RTdecreases

-Ea/RTincreases

e-Ea/RT

increases kincreases

REACTIONSPEEDS UP

If Tincreases

RT

EexpAk a

Arrhenius Equation

AlnT

1

R

Ekln

RT

EAlnkln

Aek

a

a

RT

Ea

y = m x + c

RT

EexpAk a



Last week……

Rate law

Reaction order

Rate constant

Units


Arrhenius Equation

RT

EexpAk a

Types of Rate Laws

1. Differential rate law or rate law

Shows how the reaction rate changes with concentration

2. Integrated rate law

Shows how concentration changes with time

Graphical determination of the order

TIME OUT FOR CALCULUS

Deriving the Integrated Rate

Expressions

• First-order reactions –

A B, then the rate of disappearance of A is:k

][][

Akdt

AdR

Rearranging gives:

kdtA

Ad

][

][

At time t = 0, [A] = [A]0And when t = t, [A] = [A]

Integrating:

.constkt]Aln[ xdx

x

that

call

ln1

Re

.const]Aln[

]A[]A[,0tat

0

0

ln[A] = ln[A]0 - kt

Integrated form of the

1st order rate expression

y = c + mx

0]Aln[kt]Aln[

dtk]A[

]A[d

ln[A]

t / s

slope = -k

Intercept = ln[A]0

Recall ln[A] = ln[A]o - kt

Antilog gives:

[A] = [A]0 e-kt

Intercept = [A]0

Example 01The concentration of N2O5 in liquid bromine varied with time as follows:

t/s 0 200 400 600 1000

[N2O5] /(mol L-1) 0.110 0.073 0.048 0.032 0.014

Confirm that the reaction is first order in N2O5 and determine the rate constant.

[Answer: 2.1 x 10-3 s-1]

ln[A]

t / s

slope = -k

Intercept = ln[A]0

• Second-order reactions –

Two possible cases:

Case I : A Products

Case II : A + B Products

2]A[kdt

]A[dr

Rearranging gives:kdt

]A[

]A[d2

At time t = 0, [A] = [A]0And when t = t, [A] = [A]

Integrating:

dtk]A[d]A[

12

xx

xdxxdx

x

1

12

1 112

2

2

.constkt]A[

1

0]A[

1kt

]A[

1

Integrated form of the

2nd order rate expression

.const]A[

1

]A[]A[,0tat

0

0

(1/[A]) / dm3 mol-1

t / s

slope = k

Intercept = 1/[A]0

kt]A[

1

]A[

1

0

y = c + mx

A + B Products

]B][A[kdt

]A[dr

2]A[kdt

]A[dr

If t=0, [A] = [B]

Case II :

• Zero-order reactions –

A Products

0]A[kdt

]A[dr k

dt

]A[d

constkt]A[

.const]A[

]A[]A[,0tat

0

0

0]A[kt]A[

Plotting [A] versus t will give a straight line with slope -k.

dtk]A[d1

slide 4.ppt

slide 4.ppt

slide 4.ppt

slide 4.ppt

slide 4.ppt

Order

zero 1st 2nd

Rate law rate = k rate = k[A] rate = k[A]2

Integrated

rate law[A]=−kt+[A]0 ln[A]=−kt+ln[A]0 1/[A]=kt+1/[A]0

Straight-

line plot[A] vs. t ln[A] vs. t 1/[A] vs. t

Slope −k −k k

Half-life

(t1/2)

[A]o/2k 0.693/k 1/k[A]0

Experimental determination of the

rate laws and the rate constants

1. Integral Methods

• Test the data against an appropriate integral

rate law.

• e.g. ln[A] vs t or 1/[A] vs t.

Rate laws have to be determined experimentally.

Zero order:

1st order:

2nd order:

2]A[kdt

]A[d

0]A[

1kt

]A[

1

Example: 2 H2O2 2 H2O + O2

Time(min) [H2O2](mol/L)

0 0.0200

200 0.0160

400 0.0131

600 0.0106

800 0.0086

1000 0.0069

This technique simplifies the rate law by making all the

reactants except one, in large excess.

Therefore,

The dependence of the rate on each reactant can be found

by isolating each reactant in turn and keeping all other

substances (reactants) in large excess.

Using as example: r = k[A] [B]2

Make B in excess, so [B]>>[A].

Hence, by the end of the reaction [B] would not have

changed that much, although all of A has been used up

And we can say, [B] [B]0

2. Isolation Method

r = k’[A] , where k’ = k[B]02

Since A is the reactant that changes, then the rate

becomes dependent on A, and we can say

Created a ‘false’ first-order (imitating first-order)

PSEUDO-FIRST-ORDER,

where k’ is the pseudo-first-order rate constant

Keff , effective rate constant

slide 5.ppt

r = k’’[B]2 , where k’’ = k[A]0

PSEUDO-SECOND-ORDER,

r = k’[A] , where k’ = k[B]02

Plot of k’ vs [B]02

Isolation method

- all reactants in large excess except one

» This means concentration of all reactants except one are

constant

The other values would be lumped into the rate constant

determined

Order thus determined is called psuedo- nth order

Example say rate is v = k [A][B]2

If B is in large excess, v = k’[A] pseudo-first order

k’ = k[B0]2

If A is in large excess, v= k’[B]2 psuedo 2nd order

3. Differential Method:

n]A[kr

]Aln[nklnrln

4. Initial Rate Method:

- often used in conjunction with the isolation method,

Recall

A + B P,

Taking ‘logs’

log Rate0 = log k + a log [A]0 + b log[B]0

y m xc

** Keep [A]0 constant for varying values of [B]0 to find b

Log Ro

log[B]0

slope = b

Intercept = log k + a log[A]0

b

0

a

00 ]B[]A[krate

** Keep [B]0 constant for varying values of [A]0 to find a from the slope

of the graph, log R0 vs log [A]0

1st order

y = mx + c

Plot: ln[A] vs. t

slope = − k

][][

Akdt

Adrate

ot AktA ]ln[]ln[

integrate

Last week……

2nd order

y = mx + c

Plot: 1/[A] vs. t

slope = k

2][][

Akdt

Adrate

ot Akt

A ][

1

][

1

integrate

zero order

y = mx + c

Plot: [A] vs. t

slope = − k

kAkdt

Adrate 0][

][

ot AktA ][][

integrate

Order

zero 1st 2nd

Rate law rate = k rate = k[A] rate = k[A]2

Integrated

rate law[A]=−kt+[A]0 ln[A]=−kt+ln[A]0 1/[A]=kt+1/[A]0

Straight-

line plot[A] vs. t ln[A] vs. t 1/[A] vs. t

Slope −k −k k

Half-life

(t1/2)

[A]o/2k 0.693/k 1/k[A]0

4. Half life method:

The half-life, t1/2, is defined as the time it takes for

the reactant concentration to drop to half its initial

valueIt is a useful indication of the rate of a chemical reaction.

157


• Reaction is 1st order decomposition of H2O2.

Half-Life

158


Half-Life

• Reaction after 1 half-life.

• 1/2 of the reactant has been consumed and 1/2 remains.

159


Half-Life

• After 2 half-lives 1/4 of the reactant remains.

160


Half-Life

• A 3 half-lives 1/8 of the reactant remains.

2/1

0

0 @2

1

][

][

2

][][ t

A

AAA t

t

• First-order reactions –

Remember that for a 1st order reaction: ln[A]t = ln[A]0 - kt

At time t = 0, [A] = [A]0Then at time t = t½ (half-life), [A]t½ = [A]0/2

Substituting into above equation,

ln([A]0/2) = ln[A]o – kt½ln([A]0/2) – ln[A]0 = -kt½

2/1

0

0

][

2/][ln kt

A

A

2/12

1ln kt

ln 1 – ln 2 = -kt½, where ln 1 = 0

Therefore, ln 2 = kt ½

Hence,

kt

2ln2/1 or

kt

693.02/1

What is/are the main point(s) to note from this expression??

For a 1st order reaction, the half-life is independent of reactant

concentration but dependent on k.

The half-life is constant for a 1st order reaction

time

concentration

[A]0

[A]0/2

[A]0/4

[A]0/8

Recall: [A]t = [A]0e-kt

t1/2

t1/2t1/2

Note: Radioactive decay

follows 1st order kinetics.

slide 6.ppt

• Zeroth-order reactions –

Dependent on initial concentration

k2

]A[t 0

2/1

slide 6.ppt

• Second-order reactions –

kt]A[

1

]A[

1

0t

At time t = 0, [A] = [A]0And when t = t½, [A]t½ = [A]0/2

2/1

00

kt]A[

1

2

]A[

1

2/1

00

kt]A[

1

]A[

2

2/1

0

kt]A[

1

0

2/1]A[k

1t

So t1/2 for 2nd order reactions

depends on initial concentration

slide 6.ppt

Therefore, larger initial concentrations imply shorter half-lives

(so faster the reaction).

concentration

[A]0

[A]0/2

[A]0/4

[A]0/8

time

t1/2

t1/2

t1/2

0

2/1]A[k

1t

Obtaining kinetic data

Variation of reactant concentration with time

• Classical methods

Example: acid catalyzed hydrolysis of methyl acetate at 30C

COOHCHOHCHCOOCHCH 33

OH/HCl

332

n33COOCHCHkrate

If n = 1, we can use the integrated equation

slide 6.ppt

0]Aln[kt]Aln[

0t33tt33 COOCHCHktCOOCHCHln

t/s

_

_

_

_

_

_

_

_

tt33COOCHCHln

tt30t33tt33 COOHCHCOOCHCHCOOCHCH

slide 5.ppt

Electrochemical method

Example: oxidation of formic acid by bromine in aqueous solution

At the platinum electrode the redox process is

Conductivity measurements

Conductance, which is the reciprocal of resistance, is

directly proportional to the concentration of the ions

Eg. The hydrolysis of acetic anhydride to acetic acid can be studied by measuring the conductivity.

-Gas pressure

- Volume change


Using spectroscopy [Light absorbance]



Studying fast reactions

Method of measuring concentration has to be

fast enough to make measurements over the

time scale of the reaction


Continuous flow method

Stopped flow method

Flash photolysis