introduction 1

Decomposition of Hydrogen Peroxide

INTRODUCTION

The branch of physical chemistry which deals with the rate of

reactions in called chemical kinetics.

The study of chemical kinetics includes:

The rate of the reactions and rate laws

The factors are temperature, pressure concentration and catalyst that

influence the rate of a reaction.

The mechanism or the sequence of steps by which a reaction occurs.

The knowledge of the rate of reactions in very valuable to understand

the chemical of reactions.

It is also great importance in selecting optimum conditions for an

industrial process so that it proceeds at a rate to give maximum yield.

The rate of reactions is defined as change in concentration of any of

reactant or products per unit time.

Eg: Consider a simple reaction

A B

The concentration of the reactant a decrease and that of B increases as

time passes.

rate=−d [ A ]

dt

rate =+ d [ B ]dt

Where [ ] represents the concentration in mole per liter

Sahyadri Science College (Autonomous), Shimoga 1


dt = Infinitesimally small change in concentration

The units of reactions rates expressed in moles per liter

Factors influencing the reaction rate

The rate of chemical reaction is not a fixed quantity and varies with

the experimental conditions.

Rate of reactions influences by the following factors.

i. Nature of reactants

ii. Concentration of the reactant

iii. Temperature

iv. Presence of catalyst

v. Surface area of the catalyst or reactant

vi. Radiations or presence of light

Rate Law and Rate Constant

The rate of reaction is directly proportional to the reactant

concentration each concentration being raised to some power.

Thus, Rate α [A]n

Rate = K [A]n

For reaction

2A+B Products

The rate reaction expressed

Rate = K [A]n [B]n

An expression which shows how the reaction rate is related to

concentration is need the rate law are rate equations.



Reaction

2N2O5 4NO2+O2

Rate = K [N2O5]2

H2+I2 2HI

Rate = K [H2] [I2]

Order of reaction

The order of reaction is defined as the of the sum of the power of

concentration in the rate law

Rate = K[A]m [B]n

Order of reaction in ( m + n)

Rate = K [N2O5]2

Rate = K [H2][I2]

Rate = K [NO2]2

m+n = 1, it is first order reaction

m+n = 2, it is second order reaction

m+n = 3, it is third order reaction

First Oder reaction

A reaction is said to be first order if the rate is given by the expression

of the type.

R = K1CA



Best example for the first order reaction in decomposition of hydrogen

peroxide.

Hydrogen peroxide always campore exothermically into water and

oxygen gas spontaneously.

2H2O2 2H2O + O2 ↑

The liberation of oxygen and energy in decomposition has dangerous

side effect splitting high concentration of hydrogen peroxide on a flammable

substance can cause immediate.

The rate of decomposition in dependent on the temperate and

concentration of the hydrogen peroxide as well as the pH, presence of

impurities and stabilities.

The decomposition occurs more rapidly in alkali so acid is of then

added as stabilizer.



ABOUT HYDROGEN PEROXIDE

Hydrogen peroxide

IUPAC name Hydrogen peroxide

Other names µ-1KO, 2KO’ – Dioxidodihydrogen

Dihydrogen dioxide

Hydrogen dioxide

Dioxidane

Properties

Molecular formula H2O2

Molar mass 34.0147 g/mol

Appearance Very light blue color; colorless in solution

Density 1.463 g/cm3

Melting point -0.43 °C, 273 K, 31 °F

Boiling point 150.2 °C, 423 K, 302 °F

Solubility in water Miscible

Acidity (p Ka) 11.62

Viscosity 1.245 cP (20 °C)

Dipole moment 2.26 D


http://en.wikipedia.org/wiki/File:Wasserstoffperoxid.svg

http://en.wikipedia.org/wiki/File:Hydrogen-peroxide-3D-balls.png


Hydrogen peroxide (H2O2) is a very pale liquid which appears

colourless in a dilute solution, slightly more viscous than water. It is a weak

acid. It has strong oxidizing properties and is therefore a powerful bleaching

agent that is mostly used for bleaching paper, but has also found use as a

disinfectant, as an oxidizer, as an antiseptic, and in rocketry ( particularly in

high concentrations as high –test peroxide or HTP) as a monopropellant, and

in bipropellant systems. The oxidizing capacity of hydrogen peroxide is so

strong that the chemical is considered a highly reactive oxygen species.

Hydrogen peroxide is naturally produced as a by-product of oxygen

metabolism, and virtually all organisms possess enzymes known as

peroxidases, which harmlessly and catalytically decompose low

concentrations of hydrogen peroxide to water and oxygen.

History

Hydrogen peroxide was first isolated in 1818 by Louis Jacques

Thénard by reacting barium peroxide with nitric acid. An improved version

of this process used hydrochloric acid, followed by sulfuric acid to

precipitate the barium sulfate byproduct. Thénard's process was used from

the end of the 19th century until the middle of the 20th century. Modern

production methods are discussed below.

For a long time, pure hydrogen peroxide was believed to be unstable,

because attempts to separate the hydrogen peroxide from the water, which is

present during synthesis, failed. This was because traces of solids and heavy

metal ions led to a catalytic decomposition or explosions of the hydrogen

peroxide. 100 % pure hydrogen peroxide was first orbited through vacuum

distillation by Richard Wolffenstein in 1894. At the end of 19th century,



Petre Melikishvili and his pupil L. Pizarjevski showed that of the many

proposed formulas of hydrogen peroxide, the correct one was H-O-O-H.

Storage

Regulations vary, but low concentrations, such as 3%, are widely

available and legal to buy for medical use. Higher concentrations may be

considered hazardous and are typically accompanied by a material safety

data sheet(MSDS). In high concentrations, hydrogen peroxide is an

aggressive oxidizer and will corrode many materials, including human skin.

In the presence of a reducing agent, high concentrations of H2O2 will react

violently.

Hydrogen peroxide should be in a cool, dry, well –ventilated area and

away from any flammable or combustible substances. It should be stored in a

container composed of non reactive materials such as stainless steel or glass

( other materials including some plastics and aluminum alloys may also be

suitable). Because it breaks down quickly when exposed to light, it should be

stored in an opaque container, and pharmaceutical formulations typically

come in brown bottles that filer out light.

Physical Properties

H2O2 adopts a nonplanar structure of C2 symmetry. Although chiral,

the molecule undergoes rapid racemization. The flat shape of the anti


http://en.wikipedia.org/wiki/File:H2O2_structure.png


conformer would minimize steric repulsions, the 90° torsion angle of the syn

conformer would optimize mixing between the filled p-type orbital of the

oxygen (one of the lone pairs) and the LUMO of the vicinal O-H bond.[3] The

observed anticlinal "skewed" shape is a compromise between the two

conformers.

While the anti conformer would minimize steric repulsions, a 90°

torsion angle would optimize mixing between the filled p-type orbital of the

oxygen (one of the lone pairs) and the LUMO of the vicinal O-H bond.[8]

Reflecting a compromise between the two interactions, gaseous and liquid

hydrogen peroxide adopts an anticlinal "skewed" shape. This rotational

conformation is a compromise between the anti conformer, which would

minimize steric repulsion, and between the syn conformer that associates

0¬H bonds with lone pairs on the oxygen atoms. Despite the fact that the 0-0

bond is a single bond, the molecule has a remarkably high barrier to

complete rotation of 29.45 kllmol (compared with12.5 kllmol for the

rotational barrier of ethane). The increased barrier is also attributed to

repulsion between one lone pair and other lone pairs. The bond angles are

affected by hydrogen bonding, which is relevant to the structural difference

between gaseous and crystalline forms; indeed a wide range of values is seen

in crystals containing molecular H2O2.

Chemical properties

H2O2is one of the most powerful oxidizers known stronger than

chlorine, chlorine dioxide and potassium permanganate. Also, through

catalysis, H2O2 can be converted into hydroxyl radicals (.OH) with reactivity

second only to fluorine.



Oxidant Oxidation potential, V

Fluorine 3.0

Hydroxyl radical 2.8

Ozone 2.1

Hydrogen peroxide 1.8

Potassium permanganate 1.7

Chlorine dioxide 1.5

Chlorine 1.4

Hydrogen peroxide can decompose spontaneously into water and

oxygen. It usually acts as an oxidizing agent, but there are many reactions

where it acts as a reducing agent, releasing oxygen as a by-product. It also

readily forms both inorganic and organic peroxides.

Decomposition

Hydrogen peroxide always decomposes (disproportionates)

exothermically into water and oxygen gas spontaneously:

2H2O2 2H2O + O2

This process is very favorable thermodynamically. It has a H0 of

98.2 kJ mol-1 and a G0 -119.2 kJ mol-1 and a S of 70.5 J mol -1 per K-1

The rate of decomposition is dependent on the temperature and

concentration of the peroxide, as well as the pH and the presence of

impurities and stabilizers. Hydrogen peroxide is incompatible with many

substances that catalyse its decomposition, including most of the Transition

metals and their compounds. Common catalysts include manganese dioxide,

and silver. The same reaction is catalysed by the enzyme catalase, found in

the liver, whose main function in the body is the removal of toxic byproducts



of metabolism and the reduction of oxidative stress. The decomposition

occurs more rapidly in alkali, so acid is often added as a stabilizer.

The liberation of oxygen and energy in the decomposition has

dangerous side effects. Spilling high concentrations of hydrogen peroxide on

a flammable substance can cause an immediate fire, which is further fueled

by the oxygen released by the decomposing hydrogen peroxide. High-

strength peroxide (also called high-test peroxide, or HTP) must be stored in a

suitable vented container to prevent the buildup of oxygen gas, which would

otherwise lead to the eventual rupture of the container. In the presence of

certain catalysts, such as Fe2+ or Ti3+, the decomposition may take a different

path, with free radicals such as HO' (hydroxyl) and HOO' being formed. A

combination of H2O2 and Fe2+ is known as Fenton's reagent.

A common concentration for hydrogen peroxide is "20 volume",

which means that when 1 volume of hydrogen peroxide is decomposed, it

produces 20 volumes of oxygen. A 20 "volume" concentration of hydrogen

peroxide is equivalent to 1.67 mol/dm3 (Molar solution) or about

6%.Hydrogen peroxide available at drug stores is three percent solution. In

such small concentrations, it is less stable, and decomposes faster. It is

usually stabilized with acetanilide, a substance which has toxic side effects

in significant amounts.

Redox reactions

In aqueous solution, hydrogen peroxide can oxidize or reduce a

variety of inorganic ions. When it acts as a reducing agent, oxygen gas is

also produced. In acidic solutions Fe2+ is oxidized to Fe3+,

2Fe2+(aq) +H2O2 + 2H+(aq) 2Fe+3 (aq) +2H2O (l)



and sulphite (SO32-) is oxidized to sulphate (SO4

2-). However, potassium

permanganate is reduced to Mn2+ by acidic H2O2. Under alkaline conditions,

however, some of these reactions reverse; for example, Mn2+ is oxidized to

Mn4+ (as MnO2).

Another example of hydrogen peroxide acting as a reducing agent is

the reaction with sodium hypochlorite, which is a convenient method for

preparing oxygen in the laboratory.

NaOCl + H2O2 O2 + NaCl + H2O

Hydrogen peroxide is frequently used as an oxidizing agent in organic

chemistry. One application is for the oxidation of thioethers to sulfoxides.

For example, methyl phenyl sulfide was oxidized to methyl phenyl sulfoxide

in 99% yield in methanol in 18 hours (or 20 minutes using a TiCh catalyst).

Ph-S-CH3 + H2O2 Ph-S(O)-CH3 + H2O

Preparation of Oxalic acid solution

Weigh accurately 0.315g of oxalic acid into a clean 100ml volumetric

flask and dissolve the crystals. Dilute the crystals up to the mark with

distilled water. Mix the solution for uniform concentration.

Weight of oxalic acid = 0.315g

0.315 x 10 Normality of oxalic acid = --------------------- = 0.05 N

63 Preparation of potassium permanganate solution

Weigh accurately 0.790g of potassium permanganate crystals into

clean 500ml beaker and dissolve the crystals. Dilute the solution up to the

mark with distilled water. Mix the solution for uniform concentration.



Standardization of potassium permanganate

Pipette out 10 ml of standard 0.05 N oxalic acid and 1 test tube 2 N

sulphuric acid solutions into a clean 250ml conical flask. Heat the solution

and titrate it with KMnO4 till the color changes from colorless to pale pink.

Note down the titrate value and continue the titration for agreeing values.

Tabular column:

Burette - KMnO4 solution

Conical flask - 10ml oxalic acid + 1 test tube 2N H2SO4

Indicator - Self indicator

End point - Colorless to pale pink color

Burette readings in ml Trial I Trial 2

Final burette reading 10.2 10.2

Initial burette reading 0.0 0.0

Volume of KMnO4 consumed 10.2 10.2

V = 10.2 ml

(NV)KMnO4 = (NV) oxalic acid

(N x V) Oxalic acid (V)KMnO4 = ----------------------------

(V) KMnO4

0.05 x 10 = ----------------------------

10.2

= 0.049 N



Preparation of hydrogen peroxide solution

Take 5 ml of 30 % H2O2 in a clean measuring jar and pour it into a

beaker containing 500 ml of distilled water. Mix thoroughly for uniform

concentration.

Standardization of Hydrogen Peroxide

Pipette out 10 ml of prepared H2O2 solution into a clean 250 ml

conical flask and add 1 test tube 2 N sulphuric acid solution and titrate it

with standard KMnO4 till the color changes from colorless to pale pink. Note

down the titrate value and continue the titration for agreeing values.

Tabular column:

Burette - Standard KMnO4 solution

Conical flask- 10ml H2O2 + 1 test tube 2N H2SO4

Indicator - self indicator

End point - colorless to pale pink color

Burette readings in ml Trial 1 Trial 2Final burette reading 30.1 30.1

Initial burette reading 0.0 0.0

Vol. of KMnO4 consumed 30.1 30.1

V=30.1 ml

(NV) H2O2 = (NV) KMnO4

(NV) KMnO4

(NxV) H2O2 = --------------------------

(V) H2O2

0.049 x 30.1= ------------------------- 10= 0.147 N

STUDIES OF KINETICS



Take 50 ml of 0.147 N standard H2O2 in two separate reagent bottles.

Maintain the temperature of one reagent bottle at room temperature (28oC)

and another at 35°C. Add 3 ml of 3% catalyst to both the reagent bottles.

Start the stop watch, when half of the volume of catalyst solution has been

poured into the reagent bottles. This is taken as zero time. Mix the contents

in each reagent bottles thoroughly and immediately pipette out 10ml of

reaction mixture into a conical flask containing ice cold water and one test

tube 2N H2SO4. The solution is titrated against O.05N standard KMnO4

Solution. Note down the end point by color change. Let the titer value at zero

time be Vo Repeat the same procedure at the interval of 5 minutes. Let the

titer value at any time be Vt. The Vt values decreases as the concentration of

H2O2 in the reaction mixture gradually falls. The calculated value of K can

be obtained by using the formula,

K=2 .303×log

V o

V t

t /s

Plot the graph log Vo/Vt versus time, a straight line is obtained which

pass through the origin. Slope of the line gives values of rate constant (K).

Theoretical value of k is compared with graphical value of K.

We know that the rate constant k is related to the activation energy

(Ea).Thus, by measuring the rate constant (k) at two different temperatures,

we can evaluate activation energy (Ea).

Ea= logK2

K1×2.303×R×( T1×T 2

T 2 −T1) KJ /mol

Similarly the above procedure is carried out by adding 5 ml catalyst to

95 ml H2O2 solution at two different temperatures.



The above procedure is followed for these catalysts,

a) Ferric chloride

b) Cupric chloride

c) Stannous chloride

d) Cobalt chloride

e) Zinc Chloride

f) Nickel Chloride

g) Potassium Iodide

h) Mixture of mohr salt of Fe2+ and Fe3+

i) Ferric ammonium sulphate

Catalytic Variation

Catalysis is the change in rate of a chemical reaction due to the

participation of a substance called a catalyst. Unlike other reagents that

participate in the chemical reaction, a catalyst is not consumed by the

reaction itself. A catalyst may participate in multiple chemical

transformations. Catalysts that speed the reaction are called positive

catalysts. Substances that interact with catalysts to slow the reaction are

called inhibitors (or negative catalysts). Substances that increase the activity

of catalysts are called promoters, and substances that deactivate catalysts are

called catalytic poisons.

Kinetically, catalytic reactions are typical chemical reactions, i.e. the

reaction rate depends on the frequency of contact of the reactants in the rate-



determining step. Usually, the catalyst participates in this slow step, and

rates are limited by amount of catalyst and its "activity". In heterogeneous

catalysis, the diffusion of reagents to the surface and diffusion of products

from the surface can be rate determining. Analogous events associated with

substrate binding and product dissociation apply to homogeneous catalysts.



A) POTASSIUM IODIDE

1. Reaction mixture = 1 ml (0.09g) Potassium Iodide + 50 H2O2 + 49 g

H2O correspondingly for 2 and 3 ml the volume of water changed. i.e. 48 and

47 ml respectively and those are tabulated below.

Table 1:

Catalyst: Potassium Iodide (KI)

Temperature: 28oC

Concentration: 1 ml

Vo : 19.5ml

Time in

minutes

Vt in

mllog (V o

V t) K=2 .303

tlog ( V o

V t)

0 0 0 0

5 18.4 0.0252 0.0116

10 18.3 0.0275 0.0063

15 18.2 0.0299 0.0043

20 18.0 0.0348 0.0040

25 17.4 0.0519 0.0047

30 17.0 0.0595 0.0045

35 16.8 0.0647 0.0042

40 16.5 0.0725 0.00417

45 16.0 0.0859 0.0043



Table 2:

Catalyst : Potassium Iodine (KI)

Temperature: 28oC

Concentration: 2 ml

Vo: 19.5ml

Time in minutes

Vt in ml log (V o

V t) K=2 .303

tlog ( V o

V t)

0 19.5 0 05 19.0 0.0112 0.005110 18.1 0.032 0.007415 17.4 0.049 0.007520 16.7 0.067 0.007525 16.3 0.077 0.007130 15.5 0.199 0.007635 14.8 0.11 0.007840 14.4 0.13 0.007545 13.9 0.014 0.0075

Table 3:

Catalyst: Potassium Iodine (KI)

Temperature: 28oC

Concentration: 3 ml

Vo : 19.3ml

Time in minutes

Vt in ml log (V o

V t) K=2 .303

tlog ( V o

V t)

0 Vo=19.3 0 05 18.1 0.027 0.012810 17.9 0.032 0.007515 17.3 0.047 0.007220 16.3 0.073 0.008425 15.5 0.095 0.008730 15.0 0.10 0.008435 14.5 0.12 0.008140 13.6 0.15 0.008545 13.3 0.16 0.0082



B) NICKEL CHLORIDE

1. Reaction mixture = 1 ml (0.09g) Nickel chloride + 50 H2O2 + 49 g



Table 4 :

Catalyst : NiCl2

Temperature: 28oC

Concentration : 1 ml

Vo : 18.5 ml

Time in minutes

Vt in ml log (V o

V t) K=2 .303

tlog ( V o

V t)

0 0 0 05 18.8 0 010 18.5 0.0069 0.001615 18.3 0.0117 0.001720 18.0 0.0188 0.002125 18.2 0.014 0.001230 18.1 0.016 0.001235 17.9 0.021 0.001440 17.8 0.023 0.001345 17.5 0.031 0.0015



Table 5 :

Catalyst : NiCl2

Temperature: 28oC


Vo : 18.5 ml

Time in minutes

Vt in ml log (V o

V t) K=2 .303

tlog ( V o

V t)

0 0 0 05 18.3 0.0047 0.002110 18.2 0.071 0.001615 17.8 0.016 0.002520 18.3 0.0047 0.005425 17.7 0.019 0.001730 17.5 0.024 0.001835 17.4 0.026 0.001740 17.4 0.026 0.001545 17.4 0.026 0.0013

Table 6 :

Catalyst : NiCl2

Temperature: 26oC


Vo : 18.5 ml

Time in minutes

Vt in ml log (V o

V t) K=2 .303

tlog( V o

V t)

0 0 0 05 17.9 0.014 0.0065910 17.4 0.026 0.0061315 17.2 0.031 0.004820 17.1 0.034 0.003925 16.9 0.039 0.003630 16.8 0.041 0.003235 16.5 0.049 0.003240 16.1 0.060 0.003445 15.8 0.068 0.0035



C) FERRIC CHLORIDE

1. Reaction mixture = 1 ml (0.09g) ferric chloride + 50 H2O2 + 49 g



Table 6 :

Catalyst : FeCl3

Temperature: 27oC


Vo : 19 ml

Time in minutes

Vt in ml log (V o

V t) K=2 .303

tlog ( V o

V t)

Unit 0 0 0 05 15.6 0.085 0.0310 15.0 0.102 0.02315 14.3 0.123 0.01820 13.1 0.161 0.01825 12.7 0.161 0.014830 11.8 0.174 0.013435 11.4 0.20 0.013640 11.0 0.23 0.013645 9.5 0.30 0.0154



Table 7 :

Catalyst : FeCl3

Temperature: 27oC


Vo : 16.8ml

Time in minutes

Vt in ml log (V o

V t) K=2 .303

tlog ( V o

V t)

0 0 0 05 15.5 0.03 0.016110 14.1 0.076 0.017515 12.7 0.12 0.018620 11.6 0.16 0.018525 10.3 0.212 0.019530 9.6 0.243 0.018635 8.6 0.290 0.019140 7.9 0.327 0.018845 7.5 0.3502 0.0179

Table 8 :

Catalyst : FeCl3

Temperature: 27oC


Vo : 16.5 ml

Time in minutes

Vt in ml log (V o

V t) K=2 .303

tlog ( V o

V t)

0 0 0 05 15.4 0.029 0.0138010 13.6 0.083 0.019315 11.5 0.196 0.030020 9.0 0.26 0.030325 7.5 0.342 0.031530 6.4 0.411 0.031535 5.5 0.477 0.031340 4.7 0.545 0.031445 4.2 0.594 0.03042



D) COPPER CHLORIDE

1. Reaction mixture = 1 ml (0.09g) copper chloride + 50 H2O2 + 49 g



Table 9 :

Catalyst : CuCl2

Temperature: 27oC

Concentration : 0.3 ml

Vo : 17.5 ml

Time in minutes

Vt in ml log (V o

V t) K=2 .303

tlog( V o

V t)

0 0 0 05 17.2 0.075 0.003410 16.8 0.017 0.004015 16.2 0.033 0.005720 15.9 0.044 0.005125 15.3 0.058 0.005330 14.8 0.072 0.005535 14.5 0.081 0.005140 14.0 0.096 0.005545 13.9 0.10 0.0051



Table 10 :

Catalyst : CuCl2

Temperature: 27oC


Vo : 17.2 ml

Time in minutes

Vt in ml log (V o

V t) K=2 .303

tlog ( V o

V t)

0 0 0 05 16.2 0.026 0.011910 15.0 0.059 0.01315 14.8 0.065 0.01020 14.7 0.068 0.007825 14.5 0.075 0.006830 14.1 0.086 0.006635 13.7 0.098 0.006540 13.5 0.105 0.006045 13.0 0.121 0.0062

Table 11 :

Catalyst : CuCl2

Temperature: 27oC


Vo : 17.0 ml

Time in minutes

Vt in ml log (V o

V t) K=2 .303

tlog ( V o

V t)

0 0 0 05 16.0 0.026 0.01210 15.5 0.040 0.009215 15.0 0.054 0.0083420 14.8 0.060 0.006925 14.3 0.075 0.006930 13.8 0.090 0.006935 13.6 0.096 0.006340 12.9 0.11 0.006945 12.6 0.13 0.0066



e) MIXTURE Of Fe+2and Fe+3 AMMONIUM SULPHATE

1. Reaction mixture = 1 ml (0.09g) Mixture of Fe+2and Fe+3

Ammonium sulphate + 50 H2O2 + 49 g H2O correspondingly for 2 and 3 ml

the volume of water changed. i.e. 48 and 47 ml respectively and those are

tabulated below.

Table 13 :

Catalyst : Mixture of Fe+2and Fe+3 Ammonium sulphate

Temperature: 27oC


Vo : 16.4 ml

Time in minutes

Vt in ml log (V o

V t) K=2 .303

tlog ( V o

V t)

0 Vo=19.3 0 05 16.3 0.0026 0.0012210 15.0 0.038 0.008915 14.4 0.056 0.008620 13.9 0.075 0.008225 13.2 0.094 0.0086830 12.8 0.107 0.0082635 12.3 0.124 0.0082240 11.8 0.142 0.0082345



Table 14 :


Temperature: 27oC


Vo : 16.1 ml

Time in minutes

Vt in ml log (V o

V t) K=2 .303

tlog ( V o

V t)

0 0 0 05 14.6 0.042 0.019510 12.9 0.096 0.02215 11.5 0.146 0.02220 10.1 0.20 0.02325 9.1 0.24 0.02230 8.5 0.27 0.021235 7.0 0.36 0.02340 6.2 0.414 0.02345 5.3 0.482 0.022

Table 15 :


Temperature: 27oC


Vo : 16.0 ml

Time in minutes

Vt in ml log (V o

V t) K=2 .303

tlog ( V o

V t)

0 0 0 05 14.6 0.039 0.018210 13.4 0.077 0.017715 11.1 0.158 0.02420 9.6 0.22 0.02525 8.1 0.295 0.02730 7.1 0.35 0.02735 6.2 0.41 0.02740 5.4 0.47 0.02745



TEMPERATURE VARIATION

Temperature is a physical property of matter that quantitatively

expresses the common notions of hot and cold.

Temperature relates to the thermal energy held by an object or a

sample of matter, which is the kinetic energy of the random motion of the

particle constituents of matter. While the thermal energy of an object is

proportional to the amount of matter it contains, temperature measures

thermal energy in a manner that is independent of size; it is an intensive

property, while thermal energy is an extensive property.

Temperature is one of the principal properties studied in the field of

thermodynamics. The empirical definition of temperature arises from the

conditions of thermodynamic equilibrium,

On the molecular level, temperature is the result of the motion of the

particles that constitute the material. Moving particles carry kinetic energy.

Temperature increases as this motion and the kinetic energy increase. The

motion may be the translational motion of particles, or the energy of the

particle due to molecular vibration or the excitation of an electron energy

level.

Similarly the above procedure is carried out by adding 2 ml catalyst

to correspondingly varying water with 50 ml of H2O2 solution at different

temperature



Table 16 : KI

Temperature: 32oC


Vo : 32 ml

Time in minutes

Vt in ml log (V o

V t) K=2 .303

tlog( V o

V t)

0 Vo= 32 0 05 29.3 0.0175 0.038210 28.0 0.0133 0.057915 27.4 0.0103 0.067320 25.5 0.0113 0.098625 25.0 0.0098 0.107230 24.0 0.0095 0.124935 22.8 0.0096 0.147240 21.6 0.0098 0.170645 21.0 0.0093 0.1829

Table 17 : NiCl2

Temperature: 32oC


Vo : 25.7 ml

Time in minutes

Vt in ml log (V o

V t) K=2 .303

tlog ( V o

V t)

0 25.7 0 05 15.2 0.02280 0.105010 6.5 0.05970 0.137415 3.0 0.093281 0.143220 1.6 1.2058 0.138825 1.2 1.330 0.122530 1.0 1.4099 0.108235 0.9 1.3944 0.0914045



Table 18 : FeCl3

Temperature: 32oC


Vo : 28.8 ml

Time in minutes

Vt in ml log (V o

V t) K=2 .303

tlog( V o

V t)

0 28.8 0 05 25.8 0.021 0.047710 23.8 0.0019 0.082815 21.7 0.018 0.122920 20.8 0.016 0.141325 19.0 0.016 0.180630 18.0 0.015 0.20435 16.8 0.015 0.23440 15.3 0.014 0.27445

Table 20 : Fe2+ & Fe3+ Mixture

Temperature: 32oC


Vo : 8.0 ml

Time in minutes

Vt in ml log (V o

V t) K=2 .303

tlog( V o

V t)

0 8 0 05 7.6 0.0222 0.10210 6.6 0.0835 0.019215 6.3 0.1037 0.015920 5.8 0.1396 0.016025 5.4 0.1706 0.015730 4.8 0.221 0.016935 4.7 0.230 0.015140 4.5 0.249 0.0143



Table 21 : CuCl2

Temperature: 32oC


Vo : 8 ml

Time in minutes

Vt in ml log (V o

V t) K=2 .303

tlog( V o

V t)

0 38.8 0 05 37.7 0.0124 0.0057110 37 0.020 0.004615 36.2 0.030 0.004120 35.7 0.036 0.004125 35.3 0.041 0.003730 34.7 0.048 0.003635 33.7 0.061 0.004040 33 0.070 0.040



RESULTS AND DISCUSSION

Catalyst Conc. in ml

T1 T2 K1 K2 Ea

FeCl3 2 ml /2ml 27oC 32oC 4.606 3.45 43.77x103

KI 2 ml /2ml 28oC 32oC 3.822 2.118 112.5x103

CuCl2 0.3 ml /0.3 ml

26oC 32oC 3.454 1.888 76.88x103

NiCl2 2 ml /2ml 28oC 32oC 2.39512 3.4545 69.88x103

Mixture of Fe2+ & Fe3+

3 ml /3ml 27oC 32oC 3.4545 3.9381 19.93x103



CONCLUSION

1. Rate of decomposition of H2O2 is directly proportional concentration

of the catalyst.

R Concentration of the Catalyst

i.e. the rate constant increases as the concentration of catalyst

(FeCl3, CuCl2, KI, mixture of Fe2+ and Fe3+, NiCl2) increases as shown

in the above tables 1 – 15.

2. Rate of decomposition of H2O2 is directly proportional concentration

of the temperature of the reaction media.

R temperature of the Catalyst

i.e. the rate constant increases as the temperature of catalyst

(FeCl3, CuCl2, KI, mixture of Fe2+ and Fe3+, NiCl2) increases as shown

in the above tables 16 – 21.

3. Rate of decomposition of H2O2 is independent on the concentration of

following catalyst.

( SnCl2, ZnCl2, CoCl2, Ferric ammonium sulphate)



BIBLIOGRAPHY

1. C.W. Jones, J.H. Clark, Application of Hydrogen peroxide and

Derivatives. Royal Society of Chemistry, 1999.

2. Hydrogen Peroxide MSDS

3. Ozone Lab Peroxide Compatibility

4. Principles of Physical Chemistry by Puri, Sharma, Kalia

5. www.wikipedia.org


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