Nonmetallic oxidizing agents such as hydrogen peroxide, sodium hypochlorite, and oxygen gas are discussed. Hydrogen peroxide is an unstable liquid that decomposes into water and oxygen. It is used as a bleaching agent and disinfectant. Sodium hypochlorite is a greenish-yellow solid that decomposes into sodium chloride and chlorine. It is commonly used as a bleach and disinfectant. Oxygen gas makes up 20.8% of the atmosphere and is essential for cellular respiration in living organisms.

1. NONMETALLIC OXIDIZING
AGENTS&OZONOLYSIS
Presented By
P.MOUNIKA
(Ist M.Pharmacy)
Under the Guidence of
N.Vijaya Lakshmi
M.Pharm,(Ph.D.)

2. NON METALLIC OXIDIZING AGENTS
• Non-metals act as oxidizing agent because
they tend to accept electrons i.e. reduction.
The substance which itself gets reduced by
causing oxidation of others is an oxidizing
agent, e.g., nonmetals.

3. Hydrogen peroxide
H2O2
Pure form, it is a pale blue, clear liquid,
slightly more viscous than water.
Hydrogen peroxide is the
simplest peroxide (a compound with an
oxygen–oxygen single bond). It is used
as an oxidizer, bleaching agent
and antiseptic.
Concentrated hydrogen peroxide, or
"high-test peroxide", is a reactive
oxygen species and has been used as
a propellant in rocketry. Its chemistry is
dominated by the nature of its
unstable peroxide bond.

4. • Hydrogen peroxide is unstable and slowly
decomposes in the presence of light. Because
of its instability, hydrogen peroxide is typically
stored with a stabilizer in a weakly acidic
solution. Hydrogen peroxide is found in
biological systems including the human body.
Enzymes that use or decompose hydrogen
peroxide are classified as peroxidases.

5. Structure

6. Structure
• Hydrogen peroxide (H2O2) is a nonplanar molecule as
shown by Paul-Antoine Giguère in 1950 using infrared
spectroscopy, with (twisted) C2 symmetry. Although the
O−O bond is a single bond, the molecule has a relatively
high rotational barrier of 2460 cm−1 (29.45 kJ/mol); for
comparison, the rotational barrier for ethane is 12.5
kJ/mol. The increased barrier is ascribed to repulsion
between the lone pairs of the adjacent oxygen atoms and
results in hydrogen peroxide displaying atropisomerism.
• The molecular structures of gaseous
and crystalline H2O2 are significantly different. This
difference is attributed to the effects of hydrogen bonding,
which is absent in the gaseous state. Crystals
of H2O2 are tetragonal with the space group D4
4P4121.

7. Production
• Previously, hydrogen peroxide was prepared industrially
by hydrolysis of ammonium persulfate, which was itself obtained by
the electrolysis of a solution of ammonium bisulfate (NH4HSO4)
in sulfuric acid.
• Today, hydrogen peroxide is manufactured almost exclusively by
the anthraquinone process, which was formalized in 1936 and
patented in 1939. It begins with the reduction of
an anthraquinone (such as 2-ethylanthraquinone or the 2-amyl
derivative) to the corresponding anthra hydroquinone, typically
by hydrogenation on a palladium catalyst. In the presence
of oxygen, the anthrahydroquinone then undergoes autoxidation: the
labile hydrogen atoms of the hydroxy groups transfer to the oxygen
molecule, to give hydrogen peroxide and regenerating the
anthraquinone. Most commercial processes achieve oxidation by
bubbling compressed air through a solution of the
anthrahydroquinone, with the hydrogen peroxide
then extracted from the solution and the anthraquinone recycled
back for successive cycles of hydrogenation and oxidation.

8. The simplified overall equation for the process is simple.
The economics of the process depend heavily on effective recycling of the
extraction solvents, the hydrogenation catalyst and the expensive quinone.
A process to produce hydrogen peroxide directly from the elements has been
of interest for many years. Direct synthesis is difficult to achieve, as the
reaction of hydrogen with oxygen thermodynamically favours production of
water. Systems for direct synthesis have been developed, most of which are
based around finely dispersed metal catalysts similar to those used
for hydrogenation of organic substrates.None of these has yet reached a point
where they can be used for industrial-scale synthesis.

9. Decomposition
• Hydrogen peroxideis thermodynamically unstable
and decomposes to form water and oxygen with
a ΔHo of −98.2 kJ/mol and a ΔS of
70.5 J/(mol·K).
• The rate of decomposition increases with rising
temperature, concentration and pH, with cool,
dilute, acidic solutions showing the best stability.

10. Organic reactions
• Hydrogen peroxide is frequently used as
an oxidizing agent. Illustrative is oxidation
of thioethers to sulfoxides.
• Alkaline hydrogen peroxide is used
for epoxidation of electron-deficient alkenes such
as acrylic acid derivatives, and for the oxidation
of alkylboranes to alcohols, the second step
of hydroboration-oxidation. It is also the principal
reagent in the Dakin oxidation process.

11. • Hydrogen peroxide is a weak acid,
forming hydroperoxide or peroxide salts with
many metals.
• It also converts metal oxides into the
corresponding peroxides. For example, upon
treatment with hydrogen peroxide, chromic acid
forms an unstable blue peroxide CrO(O2)2.
• This kind of reaction is used industrially to
produce peroxoanions. For example, reaction
with borax leads to sodium perborate, a bleach
used in laundry detergents:

12. Uses
• Bleaching
• Detergents
• Production of organic compounds
• Disinfectant
• Cosmetic applications
• Use in alternative medicine
• Propellant
• Other uses-Glow sticks
Horticulture
Fish aeration

13. Sodium hypochlorite
Sodium hypochlorite is a chemical
compound with the formula NaOCl or
NaClO, comprising a sodium cation (Na+)
and a hypochlorite anion(OCl−or ClO−).
It may also be viewed as the
sodium salt of hypochlorous acid. The
anhydrous compound is unstable and may
decompose explosively. It can be
crystallized as a pentahydrate
NaOCl·5H2O, a pale greenish-yellow
solid which is not explosive and is stable
if kept refrigerated

14. • Sodium hypochlorite is most often encountered as a pale
greenish-yellow dilute solution commonly known
as liquid bleach or simply bleach, a household
chemical widely used (since the 18th century) as
a disinfectant or a bleaching agent. The compound in
solution is unstable and easily decomposes,
liberating chlorine, which is the active principle of such
products. Indeed, sodium hypochlorite is the oldest and
still most important chlorine-based bleach.
• While sodium hypochlorite is non-toxic, its corrosive
properties, common availability, and reaction products
make it a significant safety risk. In particular, mixing
liquid bleach with other cleaning products, such as acids
or ammonia, may produce toxic fumes.

15. Oxidation of organic compounds
• Oxidation of starch by sodium hypochlorite, that
adds carbonyl and carboxyl groups, is relevant to the
production of modified starch products.
• In the presence of a phase-transfer catalyst, alcohols
are oxidized to the corresponding carbonyl compound
(aldehyde or ketone). Sodium hypochlorite can also
oxidize organic sulfides to sulfoxides or sulfones,
disulfides or thiols to sulfonylchlorides or bromides,
imines to oxaziridines.It can also de-aromatize phenols.

16. Oxidation of metals and complexes
Heterogeneous reactions of sodium hypochlorite and metals such
as zinc proceed slowly to give the metal oxide or hydroxide.
NaClO + Zn → ZnO + NaCl
Homogeneous reactions with metal coordination complexes proceed
somewhat faster. This has been exploited in the Jacobsen
epoxidation.
Other reactions
If not properly stored in airtight containers, sodium hypochlorite reacts
with carbon dioxide to form sodium carbonate
2 NaOCl (aq) + CO2 (g) → Na2CO3 (aq) + Cl2 (g)
Sodium hypochlorite reacts with most nitrogen compounds to form
volatile chloramines, dichloramines, and nitrogen trichloride
• NH3 + NaClO → NH2Cl + NaOH
• NH2Cl + NaClO → NHCl2 + NaOH
• NHCl2 + NaClO → NCl3 + NaOH

17. Production
Chlorination of soda
• Potassium hypochlorite was first produced in 1789
by Claude Louis Berthollet in his laboratory on the
Quai de Javel in Paris, France, by
passing chlorine gas through a solution of potash
lye. The resulting liquid, known as "Eau de Javel"
("Javel water"), was a weak solution of potassium
hypochlorite. Antoine Labarraque replaced potash
lye by the cheaper soda lye, thus obtaining sodium
hypochlorite (Eau de Labarraque).
Cl2 (g) + 2 NaOH (aq) → NaCl (aq) + NaClO (aq) + H2O (aq)

18. Electrolysis of brine
• Near the end of the nineteenth century, E. S. Smith
patented the chloralkali process: a method of producing
sodium hypochlorite involving the electrolysis of brine to
produce sodium hydroxide and chlorine gas, which then
mixed to form sodium hypochlorite.The key reactions are:
2 Cl− → Cl2 + 2 e− (at the anode)
2 H2O + 2 e− → H2 + 2 HO− (at the cathode)
From ozone and salt
• Sodium hypochlorite can be easily produced for research
purposes by reacting ozone with salt.
NaCl + O3 → NaClO + O2
• This reaction happens at room temperature and can be
helpful for oxidizing alcohols.

19. Uses
• Bleaching
• Cleaning
• Disinfection
• Deodorizing
• Waste water treatment
• Endodontics
• Nerve agent neutralization
• Reduction of skin damage

20. Oxygen Gas
•Oxygen is the chemical element with the
symbol O and atomic number 8. It is a
member of the chalcogen group on
the periodic table, a
highly reactive nonmetal, and an oxidizing
agent that readily forms oxides with most
elements as well as with other compounds.
•By mass, oxygen is the third-most abundant
element in the universe,
after hydrogen and helium. At standard
temperature and pressure, two atoms of the
element bind to form dioxygen, a colorless
and odorless diatomic gas with the
formula O2. Diatomic oxygen gas constitutes
20.8% of the Earth's atmosphere. As
compounds including oxides, the element
makes up almost half of the Earth's crust.

21. • Dioxygen is used in cellular respiration and many major
classes of organic molecules in living organisms contain
oxygen, such as proteins, nucleic acids, carbohydrates,
and fats, as do the major constituent inorganic compounds of
animal shells, teeth, and bone.
• Most of the mass of living organisms is oxygen as a
component of water, the major constituent of lifeforms.
• Oxygen is continuously replenished in Earth's atmosphere
by photosynthesis, which uses the energy of sunlight to
produce oxygen from water and carbon dioxide.
• Oxygen is too chemically reactive to remain a free element
in air without being continuously replenished by the
photosynthetic action of living organisms.

22. Industrial production
• One hundred million tonnes of O2 are extracted
from air for industrial uses annually by two
primary methods.
• The most common method is fractional
distillation of liquefied air, with N2 distilling as a
vapor while O2 is left as a liquid.
• The other primary method of producing O2 is
passing a stream of clean, dry air through one bed
of a pair of identical zeolite molecular sieves,
which absorbs the nitrogen and delivers a gas
stream that is 90% to 93% O2.

23. • Oxygen gas can also be produced
through electrolysis of water into molecular
oxygen and hydrogen.
• DC electricity must be used: if AC is used, the
gases in each limb consist of hydrogen and
oxygen in the explosive ratio 2:1.
• A similar method is the
electrocatalytic O2 evolution from oxides
and oxoacids.

24. • Another air separation method is forcing air to dissolve
through ceramic membranes based on zirconium dioxide by
either high pressure or an electric current, to produce nearly
pure O2 gas.
•Chemical catalysts can be
used as well, such as
in chemical oxygen
generators or oxygen candles
that are used as part of the
life-support equipment on
submarines, and are still part
of standard equipment on
commercial airliners in case
of depressurization
emergencies.

25. Applications
• Smelting of iron
ore into steel consumes
55% of commercially
produced oxygen.[In this
process, O2 is injected
through a high-pressure
lance into molten iron,
which
removes sulfur impurities
and excess carbon as the
respective
oxides, SO2 and CO2. The
reactions are exothermic,
so the temperature
increases to 1,700 °C.

26. • Another 25% of commercially produced oxygen is used
by the chemical industry. Ethylene is reacted with O2 to
create ethylene oxide, which, in turn, is converted
into ethylene glycol; the primary feeder material used to
manufacture a host of products,
including antifreeze and polyester polymers (the
precursors of many plastics and fabrics).
• Most of the remaining 20% of commercially produced
oxygen is used in medical applications, metal cutting and
welding, as an oxidizer in rocket fuel, and in water
treatment. Oxygen is used in oxyacetylene
welding burning acetylene with O2 to produce a very hot
flame. In this process, metal up to 60 cm (24 in) thick is
first heated with a small oxy-acetylene flame and then
quickly cut by a large stream of O2.

27. OZONOLYSIS
• Ozonolysis is a widely used reaction in organic
synthesis. The reaction was invented by Christian
Friedrich Schoenbein in 1840.
• Alkenes and alkynes are the most common substrates
for the ozonolysis reaction.
• Ozonolysis was an important diagnostic tool for the
determination of the position of unsaturation in
unknown molecules before the invention and
development of spectroscopic techniques for
identification and characterization of organic
molecules.
• The reaction was used for structure elucidation work
because it provided chemists with smaller and more
readily identifiable carbonyl compounds.

28. Ozonolysis of alkenes
• The ozonolysis reaction involves bubbling ozone into a
solution of olefin in an organic solvent.
• The reaction is rapid and produces an intermediate called
ozonide.
• The ozonide is unstable, and hence not isolated, but can be
further reacted with various reagents to give aldehydes,
ketones, carboxylic acids, alcohols etc.
• When the ozonide is treated with mild reducing agents like
phosphines and thio compounds (typically dimethyl sulfide
or thiourea is used) aldehydes and ketones are produced.
• Ozonides can be treated with strong reducing agents like
sodium borohydride to produce alcohols.
• Ozonides when treated with oxidizing agents such as
oxygen or hydrogen peroxide, they produce carboxylic
acids as the products.

29. An example is the ozonolysis of eugenol converting the
terminal alkene to an aldehyde

30. Ozonolysis of alkynes
• Alkynes also undergo ozonolysis but very slowly
compared to alkenes.
• Unlike alkenes, ozonides from alkynes do not
need either an oxidizing agent or reducing agent
to provide end products.
• Ozonides from alkynes upon treatment with water
provide carboxylic acids are products.
• Internal alkynes produce two different carboxylic
acids while terminal alkynes produce carboxylic
acid with one less carbon; the terminal carbon is
converted to carbon dioxide.

31. Ozonolysis of alkanes
Alkanes get oxidized when treated with ozone. The
products formed are alcohols, aldehydes/ketones or
carboxylic acids. The rate of oxidative cleavage of alkanes
is highest for tertiary C-H bond, followed secondary and
primary.

32. Ozonolysis of elastomers
Ozone cracking is a form of stress
corrosion cracking where
active chemical species attack
products of a susceptible
material. Ozone cracking was
once commonly seen in the
sidewalls of tires but is now
rare owing to the use
of antiozonants. Other means
of prevention include replacing
susceptible rubbers with
resistant elastomers such
as polychloroprene, EPDM or
viton.

33. Ozonolysis in industry
• Ozonolysis has been used frequently in major
drug syntheses such as (+)-artemisinin,
indolizidine 251F and D,L-camptothecin, as well
as in fine chemical syntheses such as L-
isoxazolylalanine and prostaglandin
endoperoxides.
• ThalesNano has developed the IceCube reactor to
overcome these disadvantages. When combined
with the ozone module, ozonolysis can be
performed in a safe and highly controlled
manner.

34. Ozonolysis has a number of advantages over
conventional oxidation methods, including:
•Quicker reactions with improved yields
•Cleaner reactions and less side products
•Does not require addition of water