Chapter 1
Introduction
Introduction
Nature is prevalent chemist with its vast chemistry and is always inspirational for
those working in different areas, including from creating intelligent to discovering new
building blocks. In this respect, organic chemistry is no exception, although chemists
have yet to match the efficiency by which nature synthesizes complex and chiral
molecules. Organic chemistry began as an attempt to understand chemistry of life. It
creates itself as it grows. Organic chemistry studies life by creating new molecules that
gives the information which is not already available. Organic chemistry is a vibrant and
growing scientific discipline that covers vast numbers of scientific areas.
Most organic reactions are belong to one of these five main classes are substitution,
addition, elimination, rearrangement and redox reactions which are used in the
construction or synthesis of new organic molecules. Oxidation reactions belong to one
such class of reactions (i. e. redox reactions) that plays a vital role in organic synthesis.
Synthesis of many biologically active compounds, pharmaceuticals, fine chemicals, dyes,
plastics, food additives and fabrics have been involve through oxidation reactions.[1]
Therefore, it is important to know the basic information about such important strategy of
oxidation reactions.
1.1 Oxidation
All living organisms have chemically reactive systems and depends on the
continuance of oxidation and reduction reactions. For example plants utilize solar energy
to reduce carbon dioxide into sugar molecules and releases the oxygen while in another
reaction glucose and other compounds are oxidized to produce carbon dioxide, water, and
releases chemical energy which is utilized by animals for living.[2] So it is task for
organic chemists to understand and provide clarity by which these redox reactions take
place in simple molecules, so that biologists can have strong basis to understand the
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behavior of complex organisms. Though, the term oxidation equally well known in both
organic and inorganic chemistry, the present chapter has major concern of oxidation of
various organic substrates.
1.1.1 The meaning of oxidation
The term oxidation has been a part of common chemical usage since the time of
Lavoisier a French chemist.[3] In an inorganic chemistry loss of electrons from atom or
molecule refers to oxidation. In an organic chemistry, the oxidation reactions are different
from inorganic oxidations because many reactions carry the name oxidation but do not
actually involve electron transfer as it occurs in electrochemistry. Also, for most of
organic chemistry where molecules have many atoms, keeping track on oxidation states
of atoms in the electrochemical sense is cumbersome and not very useful. Therefore the
practice in organic chemistry has been to set up series of functional groups in order of
increasing oxidation states and then to define the oxidation as the conversion of a
functional group in a molecule from one category to a higher one (one oxidation state to
another as shown in figure 1.1). [5]
Fig.1.1: Different oxidation states of carbon
1.1.2 Classification of oxidation processes
The organic compounds are covalent in nature and have their valency electrons
associated in pairs. Moreover they are mostly composed of carbon skeleton surrounded
by hydrogens or other atoms replacing hydrogens and consequently have few superficial
electrons accessible for attack by colliding reagents. Covalent bond fission is an
important feature in organic chemistry and it can be affected by two pathways, viz.
homolytic reactions in which electron pairs are symmetrically disrupted and heterolytic
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reactions in which shared electrons are completely transfer from one atom to other one.
Electrons removal by these two pathways has clearly distinguishable characteristics.
1.1.2.1 Homolytic oxidation
In homolytic oxidations, electrons are removed singly from organic molecules by
active atoms, such as chlorine (Eq. 1), or by any active free radicals. Generally molecules
containing unshared electrons are oxidized by this way, but organic molecules involve
removal of one electron along with removal of hydrogen nucleus.
The initial radical organic product has an unpaired electron (Eq. 1) and undergoes
similar radical reaction (Eq. 2) or undergoes dimerization (Eq. 3) to form the stable
products. All such homolytic oxidation reactions require less activation energy, hence
once started, they proceeds very rapidly. The traces of radicals required for initiation of
homolytic oxidation may be formed by thermal dissociation or by photolysis.
1.1.2.2 Heterolytic oxidation
Heterolytic oxidations involve the attack of organic compounds on electrophilic
reagents which can, by a single process, gain control of a further electron pair. For
example, heterolytic oxidants attack the exposed π electrons of olefins (Eq. 4) or
unshared electron pair of atom such as oxygen (Eq. 5), nitrogen, or sulphur. These
processes require higher activation energies than homolytic reactions and giving stable
molecular or ionic products either in one or two stage. Therefore, heterolytic oxidation
processes seldom leads to a chain reactions.
Stereochemical considerations are of much greater significance in heterolytic
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oxidations than the homolytic oxidation because in the heterolytic oxidation tetrahedral
symmetry of the product is much related to starting substrate while in homolytic process
radical intermediate has planar distribution of the three covalences and lead to mixture of
stereo isomers.
1.1.3 Autoxidation
The term autoxidation is applied to process of slow oxidations which can be affected
by free oxygen at moderate temperatures. It can be considered to be a slow, flameless
combustion of materials by reaction with oxygen. Autoxidation are promoted by light,
small quantities of catalysts notably oxides and peroxidic substances. Again these
processes are retarded by traces of oxidizable organic substances such as phenols and
amines. It is well known that oxides and peroxides are characteristic products of
autoxidation. Generally all types of organic materials can undergo areal oxidation over a
period of time. Certain types of organic substrates are particularly prone to undergo
autoxidation, including unsaturated compounds that have allylic or benzylic hydrogens;
these substrates are converted to hydroperoxides by autoxidation. The study of
autoxidation is important because it is a useful reaction for converting compounds to
oxygenated derivatives, and also because it occurs in situations where it is not desired.
1.2 Oxyhalogenation
Halogenated organic compounds play a very important role in organic chemistry.
They are important as starting substrates as well as versatile intermediate in the organic
synthesis.[6] Laboratory halogenations involves hazardous, toxic and corrosive moleculer
halogens and furthermore they are performed in the chlorinated solvents.[7] Therefore
direct use of halonium species in the reactions is no more recommendable and alternative
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strategy is required for their use. Electrophilic halogenation in nature mainly takes place
through oxidative halogenation with the exception of fluorination (It is too difficult to
oxidize fluoride). The better approach could be a in-situ generation of halonium species
via the oxidation of halides using a suitable oxidants under moderate reaction conditions
to overcome the current drawbacks. An increasing environmental concerns and advanced
studies on oxidative halogenations, it is desirable for synthetic chemists to study and
understand in-depth the oxidative halogenations. Hence efforts are being made for
development of sustainable oxyhalogenation during present thesis work using greener
reagents. The halogens and other halogenating reagents are employed for the purpose of
oxyhalogenation.
1.2.1 Halogens
The Swedish chemist Baron Jons Jakob Berzelius coined the term "halogen" for an
element that produces a salt when it forms a compound with a metal.[8] The halogens or
halogen elements are a series of nonmetal elements from Group 17 [IUPAC Style
(formerly: VII, VIIA)][9] of the periodic table, comprising fluorine (F), chlorine (Cl),
bromine (Br), iodine (I), and astatine (At). The artificially created element 117,
provisionally referred to by the systematic name ununseptium[10] may also be a halogen.
The halogens have seven electrons in their outer shells. They have s2p5 electronic
configuration and are just one P-electron less than that of the next noble gas. Thus,
halogens complete their octet either by gain of one electron or sharing of electrons with
another atom. Compounds of halogen with metals are typically ionic, while those with
non-metal are covalent.
1.2.1.1 Abundance [11]
Due to the high reactivity of halogens they are not found in their free state, instead of
they are present in compounds or as halide salts. Halide ions and oxyanions such as
iodate (IO3−) can be found in many minerals and in seawater. Halogenated organic
compounds can also be found as natural products and in living organisms. In their
elemental forms, the halogens exist as diatomic molecules, but these only have a fleeting
existence in nature and are much more common in the laboratory and in industry. At
room temperature and pressure, fluorine and chlorine are gases, bromine is a liquid and
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iodine and astatine are solids; therefore Group 17 is only periodic table group exhibiting
all three states of matter at room temperature.
1.2.1.2 Properties
The halogens show a series of trends when moving down the group in the periodic
table, [11] for instance, decreasing electronegativity and reactivity, increases the melting
and boiling points.
Table 1.1: Periodic properties of halogens
Halogen
Standard At.Wt.
(g/mol)
Melting Point
Boiling Point
(K)
(K)
Electronegativity
(Pauling Scale)
Fluorine
18.998
53.53
85.03
3.98
Chlorine
35.453
171.60
239.11
3.16
Bromine
79.904
265.8
332
2.96
Iodine
126.904
386.85
457.40
2.66
Astatine
(210)
575
610
2.20
1.2.1.3 Reactivity
Generally, halogens are highly reactive and as such can be harmful or lethal to
biological organisms in sufficient quantities. This high reactivity is due to the atoms
being highly electronegative due to their high effective nuclear charge. Fluorine is one of
the most reactive elements in existence and can attack the inert materials such as glass as
well forms compounds with the heavier noble gases. It is a corrosive and highly toxic
gas. The reactivity of fluorine is such that if used or stored in laboratory glassware, it can
react with glass in the presence of small amounts of water to form silicon tetra fluoride
(SiF4). Thus fluorine must be handled with substances such as Teflon (which is itself an
organofluorine compound), extremely dry glass, and metals such as copper or steel which
form a protective layer of fluoride on their surface.
Chlorine and bromine react with most of the elements less vigoursly than does
fluorine. Both chlorine and bromine are used as disinfectants for drinking water,
swimming pools, fresh wounds, spas, dishes, and surfaces. They kill bacteria and other
potentially harmful microorganisms through a process known as sterilization. Their
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reactivity is also utilized in the bleaching. Sodium hypochlorite, which is produced from
chlorine, is the active ingredient of most fabric bleaches and also chlorine-derived
bleaches are used in the production of some paper products. Chlorine also reacts with
sodium to create sodium chloride, which is another name for table salt. Iodine is less
reactive and requires activation and does not combine with elements such as sulphur and
selenium. In summary, some halogens have less reactivity to get the desired products
while some have more activity that they are potentially dangerous in many reactions
hence proper care should be taken while handling.
1.2.1.4 Oxidizing power
Halogens act as oxidizing agents. The strength of an oxidizing agent depends on
several energy factors. Among the halogens, fluorine is the strongest oxidizing agent and
it will replace Cl- ions both in solution and in dry conditions. Similarly, chlorine will
replace the bromide in solution. In general, any halogens of low atomic number will
oxidize halogens of higher atomic number.
1.2.1.5 Reactivity with water
Fluorine reacts vigorously with water to produce oxygen and hydrogen fluoride (Eq.
6).
Chlorine has minimum solubility of 0.7g per kg of water at ambient temperature
(21oC).[12] Dissolved chlorine in water reacts to form hydrochloric acid and hypochlorous
acid (Eq. 7), a solution that can be used as a disinfectant or bleach. Bromine has a
solubility of 3.41g per 100g of water, but it slowly reacts to form hydrogen bromide and
hypobromous acid (Eq. 8).[13]
Iodine, however, has less solubility in water [0.03g per 100g water (20 °C)] and does not
react with it.[14] However; iodine will form an aqueous solution in the presence of other
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iodide ions, such as by addition of potassium iodide (KI) which convert iodine to
triiodide ion (I3-).
At present, the electrochemical data for all the elements as well for large number of
reactions has been well established. This electrochemical data provides idea about the
redox chemistry of various elements. Some fundamental principles of both organic and
inorganic chemistry are often illustrated by describing properties of halogens and halide
compounds. The present thesis is mainly focused on oxybromination strategy; therefore,
the introduction of the subject matter of the thesis is reviewed with general history and
properties of the bromine.
1.3 Bromine
Bromine is a chemical element with the symbol Br, an atomic number of 35 and an
atomic mass of 79.904. It belongs to halogen family in the periodic table.
1.3.1 Characteristics
1.3.1.1 Physical characteristics
Elemental bromine exists as a diatomic molecule, Br2. It is a dense, mobile, slightly
transparent reddish-brown liquid, which evaporates easily at standard temperature and
pressures to give a red vapor. It has a strongly disagreeable odor similar that of chlorine.
Bromine is only non-metal which is a liquid at room temperature. Bromine with mercury
are only two elements on the periodic table that are liquids at room temperature. At a
pressure of 55 GPa bromine converts to a metal, at 75 GPa it converts to a face centered
orthorhombic structure and at 100 GPa it converts to a body centered orthorhombic
monoatomic form.[15]
1.3.1.2 Chemical characteristics
Bromine is less reactive than chlorine and more reactive than iodine. Moreover,
bromine reacts vigorously with metals, especially in the presence of water, to give
bromide salts. It is also reactive toward most organic compounds, particularly in
photochemical conditions which favor the dissociation of the diatomic molecule into
bromine radicals. It bonds easily with many elements and has a strong bleaching action.
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Bromine is slightly soluble in water, but it is highly soluble in organic solvents such as
carbon disulfide, aliphatic alcohols, and acetic acid.
1.3.2 History of Bromine
Bromine was discovered independently by two chemists Antoine Balard[16] and Carl
Jacob Lowig [17,21, 32] in 1826 and 1825 respectively. The first attempt for the discovery of
bromine was done by Carl Jacob Lowig
[18, 21, and 32]
in 1825 as he treated his reaction
mixture of mineral origin with gaseous chlorine and as obtained that red solution, having
very unpleasant smell. The ether extraction of that mother liquor gave a new pink colored
liquid which was responsible for the red color of liquor in Carl Jacob Lowig’s bottle.
That extracted pink colored liquid was bromine.
Balard found bromide chemicals in the ash of seaweed from the salt marshes of
Montpellier in 1826. The seaweed was used to produce iodine. Balard distilled the
solution of seaweed ash saturated with chlorine. The properties of the resulting substance
resembled that of an intermediate of chlorine and iodine; with those results he tried to
prove that the substance was iodine monochloride (ICl), but after failing to do so he was
sure that he had found a new element and named it muride, derived from the Latin word
muria for brine.[16] The experiments of Balard were approved by the French chemists
Louis Nicolas Vauquelin, Louis Jacques Thénard and Joseph-Louis Gay-Lussac. The
results were presented at a lecture of the Académie des Sciences and published in Annales
de Chimie et Physique. In his publication, Balard stated that he changed the name from
muride (name of bromine) to brome on the proposal of M. Anglada which derives from
the Greek.[16] Some reports claim that the French chemist and physicist Joseph-Louis
Gay-Lussac suggested the name brome for the bromine due to its characteristic smell of
the vapors.[16] The first commercial use of bromine was in the photographic process
(except some minor medical applications). In 1840 it was discovered that bromine had
some advantages over the previously used iodine vapor to create the light sensitive silver
halide layer used for photographic process.[19]
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1.3.3 Properties of Bromine[11]
Table 1.2: Properties of Bromine
General properties
Heat of
vaporization
(Br2) 29.96 kJmol−1
Name
Bromine
Heat of fusion
(Br2) 10.571 kJmol−1
Symbol
Br
Specific heat
capacity
(25 °C) (Br2) 75.69
Jmol−1K−1
Atomic Number
35
Chemical Series
Group, period, block
Atomic properties
Crystal structure
Orthorhombic
Halogen
Oxidation states
7, 6, 5, 4, 3, 1, -1
17, 4, p
Electronegativity
2.96 (Pauling scale)
Ionization energies
1st: 1139.9.0 kJmol−1
Appearance
2nd: 2103.0 kJmol−1
Reddish
liquid
Standard
weight
atomic
3rd: 3470.0 kJmol−1
brown
79.904 g/mol
Atomic radius
115 pm
Electron
configuration
[Ar] 4s2 3d10 4p5
Atomic radius
(calc.)
94 pm
Electrons per shell
2, 8, 18, 7
Covalent radius
114 pm
Vander Waals
Radius
185 pm
Physical Properties
Phase
Miscellaneous Properties
Liquid
3
Density
3.1028 g/cm
Magnetic ordering
Non-magnetic
Melting point
265.8 K (-7.2 °C,
9 °F)
Thermal
conductivity
(300 K) 0.122 Wm−1K−1
Boiling point
332 K (58.8 °C,
137.8 °F)
CAS registry
number
7726-95-6
Critical point
588 K, 10.34 MPa
Electrical
resistivity
(20 °C) 7.8×1010Ω· m
1.3.4 Occurrence and abundance
Due to high reactivity, free form of bromine does not occur in nature, but occurs as
soluble crystalline halide salts, analogous to table salt in the crustal rock. However, the
high solubility of bromide ion has caused its accumulation in the oceans. Commercially
the bromine is easily extracted from brine pools, mostly in the United States, Israel and
China. China's bromine reserves are located in the Shandong Province and Israel's
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bromine reserves are contained in the waters of the Dead Sea. The largest bromine
reserve in the United States is located in Columbia and Union County, Arkansas.[20]
About 556,000 metric tons of bromine was produced worldwide in 2007,[18] an amount
similar to the far more abundant element magnesium. The ocean water contain about 67
mg per liter,[21] and varies in parallel with the concentration of other salts. Saline is the
another source of bromine, one such saline is found at Kharaghoda, near to the river
Indus in India which is also called the Rann of Kutch spread over an area of about 18,000
Km2.
Bromine can be economically recovered from bromide-rich brine wells and from the
Dead Sea waters (up to 50000 ppm).[22] It exists in the Earth's crust at 0.4 ppm. The
bromine concentration in soils varies normally between 5-40 ppm, but some volcanic
soils can contain up to 500 ppm. The concentration of bromine in the atmosphere is
extremely low, at only a few ppt. A large number of organobromine compounds are
found in nature but in small amounts.
1.3.5 Isotopes
There are 23 isotopes are known to exist.[23] Many of the bromine isotopes are
fission products of radioactive series. Several heavier isotopes of bromine are obtained
from fission and are delayed neutron emitters. All radioactive bromine isotopes have less
half-life period. A number of the bromine isotopes exhibit metastable isomers. Stable
79
Br exhibits a radioactive isomer, with a half-life of 4.86 seconds.
1.3.6 Production
In the bromine production, bromine reach brines are treated with chlorine, flushing
through with air. In this treatment, the bromide anions are oxidized to bromine by
chlorine (Eq. 9).
It can be affected by chemical or electrochemical method. In chemical methods,
oxidizing agents such as elemental chlorine or oxo-compounds like MnO2 or salts like
BrO3− and ClO3− are useful. A. Frank,[24a] for the first time, reported the extraction
process by the oxidation of bromide using MnO2 in sulfuric acid medium at 60°C.
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Electrochemical methods are based on the principle that bromide undergoes oxidation by
one-electron per atom at suitable anode materials. This is possible even in the presence of
large concentrations of chloride due to the fact that the decomposition of bromide is
lower by 0.29 V (the standard oxidation potentials of chloride and bromide are +1.356 V
and +1.065 V respectively). Electrochemical methods for the oxidation of bromide
solution are low intensive in energy and thus more economic than using chlorine by
chemical method. However, the electrolytic method of oxidation of bromide in the
bromine recovery has not yet been proved superior to the chemical methods. This could
be due to low concentration of bromine in the source, deposition of magnesium
hydroxide and the presence of sulfates. The two electrochemical methods that were tried
in Germany are Wunsche process[24a] and the Kossuth process.[24c] There are some
important methods involved in the bromine productions and it is prepared from
laboratory scale to industrial scale.
1.3.6.1 Laboratory methods of production
Generally bromine is not prepared in the laboratory, because of its commercial
availability and long shelf-life. However sometimes it can be generated through the
reaction of solid sodium bromide with concentrated sulfuric acid (H2SO4). The first stage
is formation of hydrogen bromide (Eq. 10), but under the reaction conditions some of the
HBr is oxidized further by the sulfuric acid to form bromine (Br2) and sulfur dioxide (Eq.
11). Reactions involving a strong oxidizing agent, such as potassium permanganate, on
bromide ions in the presence of an acid also give bromine. An acidic solution of bromate
and bromide ions wills also comproportionate slowly to give bromine.
1.3.6.2 Industrial scale production
A) Streaming out process and Dow’s blowing out process : Bromine could be
produced by two processes[25] namely the continuous process[26] and the periodic
process.[27] The continuous process depends upon the decomposition of the bromide by
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chlorine, which is generated in special stills. A regular current of chlorine was purged
from the bottom of a tall tower filled with broken bricks, where it reacts with stream of
hot bittern. The chlorine oxidizes bromide to bromine; the liberated bromine was swept
out of the tower together with some chlorine by the current of steam, and then condensed
in a worm. Any uncondensed bromine vapors were absorbed by moist iron borings, and
the resulted iron bromide was used for the manufacture of potassium bromide.
Alternatively, the periodic process involves reaction between manganese dioxide,
sulphuric acid and a bromide salt, and the operation were carried out in sandstone stills
were heated to 60° C and the product being condensed as in the continuous process. The
crude bromine thus obtained was purified by the repeating shaking with potassium,
sodium or ferrous bromide and subsequent re-distillation. Commercial bromine is not
obtained in pure form; it contains mainly chlorine, hydrobromic acid and bromoform as
impurities.[28] To obtained pure bromine E. Gessner[29] removed chlorine by repeated
shaking with water, followed by distillation over sulfuric acid. Hydrobromic acid is
removed by distillation with pure manganese dioxide or mercuric oxide and the product
was dried over sulfuric acid. Potassium bromide is also used by J. S. Stas,[27] to prepare
chemically pure bromine. During his research, he produced bromine by converting
bromide into the bromate which was purified by repeated crystallization. Bromate was
partially converted into bromide by heating and the resulting mixture was distilled with
sulfuric acid. The distillate was further purified by digestion with milk of lime and
precipitation with water which was further digested with calcium bromide and barium
oxide and was finally redistilled. Before Dow got into the bromine business, bromine-rich
brine was evaporated by heating with wood scraps and then crystallized sodium chloride
was removed. The next step was addition of an oxidizing agent, which form the bromine
in the solution. Finally, then bromine was distilled. These methods were are very
complicated and not suitable for industrial production of bromine.[30] The Dow process
was the electrolytic method of bromine extraction from brine invented by Herbert Henry
Dow, founder member of Dow’s chemicals. It was revolutionary process for generating
bromine commercially.[31] The first two steps, of the Dow’s process are similar to that of
streaming out process where acidification and oxidation of bittern liberates the aqueous
bromine.[32] The air blown crude bromine was absorbed in alkali to form the mixture of
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bromide and bromate salts, whereas the acidification of aqueous bromide and bromate
liberates the bromine fumes, which is condensed to obtained pure liquid bromine (Eq. 9,
12 and 13). T. Jakagi et. al.[33] reported a process of manufacture of bromine by ion
exchange electrolysis. The catalytic process for production of bromine reported by P.
Schubert et al.,[34] bromide salts were acidified to generate a gaseous hydrobromic acid
which was further oxidised to produces bromine. R. C. Williams et. al.[35] described an
automatically operated electrolytic cell assembly and a method of efficiently providing
brominated water for swimming pools and for other similar applications. In R. Khamizov
et. al.[36]process for bromine extraction from sea water, the sea water was concentrated to
obtain about ≥ 5 g/l of bromide ion content. This processes requires large amount of
stronger anionic exchanger for bromide ion absorption and desorption occurs at 50 − 95
°C. M. Jean-Charles et. al.[37] reported a combustion and high temperature-high pressure
process for the recovery of bromine from sea bittern/liquid effluents. The combustion
gases were cooled and subjected to hetero-azeotropic distillation to obtain the gaseous
aqueous bromine mixture. The aqueous mixture was further subjected to distillation to
recover pure bromine with purity of 99.9%. M. Yamada et. al.[38] processed the
photographic wastewaters for the recovery of bromine. More recently, the continuous
manufacture of bromine and its steam distillation was described.[39] The aqueous
solutions of bromides like HBr and NaBr were oxidized with chlorine. The systematic
flow diagram for the industrial preparation of liquid bromine is shown in figure 1.2.[32]
1.3.7 National status of bromine [32]
The bromine production during 1995-2000 in India was found to be around 20002200 tons in addition to the 500-750 tons of imported material. CSMCRI is pioneer in
this research field and carried out developmental studies on the recovery of bromine from
Indian sea brines. The institute has developed a streaming out process for the recovery of
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bromine from end liquor (bittern) of salt industry in early seventies and has released
know how to few entrepreneurs in the country. Before 1970, bromine production of the
Fig 1.2: Flow diagram for the industrial preparation of liquid bromine
country was less than 400 tons. This has increased to about 2,200 tons annually mainly
because of increase in the number of bromine producing plants installed based on
CSMCRI know how.
1.3.8 Bromine compounds and their chemistry
1.3.8.1 Bromine in organic chemistry
Organic compounds are brominated by either addition or substitution reactions.
Bromine undergoes addition to the unsaturated hydrocarbons (alkenes and alkynes) via a
cyclic bromonium intermediate.[40] In non-aqueous solvents such as carbon disulfide, it
gives di-bromo products. For example, reaction of bromine with ethylene will produce
1,2-dibromoethane as product. When bromine is used in presence of water, a small
amount of the corresponding bromohydrin will form along with desired dibromo
compounds.[41] Bromine also gives electrophilic nuclear bromination of phenols and
anilines.[42] Due to this properties, bromine water was employed as a qualitative reagent
to detect the presence of alkenes, phenols and anilines in a particular system. Like the
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other halogens, bromine also participates in free radical reactions. In presence of catalytic
amount of phosphorus, bromine brominates carboxylic acids at the α-carbon (HellVolhard-Zelinsky reaction).[43] Though bromine has many application in chemistry as a
reagent, it has some disadvantages also whenever disposed to environment. Some
bromine-related compounds have been evaluated to have an ozone depletion potential or
bio accumulate in living organisms.[44] As a results, many industrial bromine compounds
are no longer manufactured and are being banned.
1.3.8.2 Bromine in inorganic chemistry
Bromine forms a different bromine compounds which adopts a variety of oxidation
states from -1 to +7 (Table 1.3).[45] Bromine is an oxidizer and it will oxidize iodide ions
Table 1.3: Different oxidation states of bromine[11]
Compound HBr
Br2
BrCl
BrF3
BrO2
BrF5
LiBrO4
BrO-4
Oxidation
State
0
+1
+3
+4
+5
+6
+7
-1
to iodine (Eq. 14), metals and metalloids to the corresponding bromides. Anhydrous
bromine is less reactive toward many metals than hydrated bromine, however, dry
bromine reacts vigorously with aluminium, titanium, mercury as well as alkaline earths
and alkali metals. Dissolving bromine in alkaline solution gives a mixture of bromide and
hypobromite (Eq. 15). Bromine reacts violently and explosively with aluminium metal,
forming aluminium bromide (Eq. 16). Bromine reacts with hydrogen in gaseous form and
gives hydrogen bromide (Eq. 17). Bromine reacts with alkali metal iodides in a
displacement reaction. This reaction forms alkali metal bromides and produces elemental
iodine.
1.3.9 Applications
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Bromine and its various derivatives (Organic and inorganic bromo compounds) have
wide utility in academia and industries and hence for society. Many industries worldwide
are involved in the bromine manufacture and supply because of its large demand.
Following are some premium suppliers of liquid-bromine in India.[46]
a) Emichem Private Limited
b) Mody Chemi-Pharma
c) Ekta International
d) Northern Alliance
e) Meru Chem Pvt. Ltd.
f) Forbes Pharmaceuticals
g) Ang-Froid Chemicals Pvt. Ltd.
h) Sang-froid Chemicals Pvt. Ltd.
i) Bombay Lubricants Oil Co.
j) Sri Venkateswara Exports
k) Taj Pharmaceuticals Ltd.
l) Laurel Research Lab.
m) Vir Chemtech
n) Gopsi Pharma Pvt. Ltd.
o) Altret Performance Chemicals Gujarat Pvt. Ltd.
Some of premium applications of bromine are as follow.
1.3.9.1 As a reagent
Elemental bromine is employed in number of reactions, both in organic and
inorganic chemistry. It is greatly used as reagent for analytical and synthetic purposes.
Bromine is used extensively in organic chemistry for the substitution and addition
reactions to give an important halogenated derivative as well as an oxidizing agent. For
example, bromine is used to prepare N-bromosuccinimide which is commonly used as a
mild and selective reagent.[47]
1.3.9.2 Flame retardant
At high temperatures, organobromine compounds were easily converted to free
bromide atoms called radicals, and further it produces hydrobromic acid which interferes
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in the radical chain reaction of the oxidation reaction of the fire. In this mechanism, the
highly reactive hydrogen, oxygen and hydroxyl radicals react with hydrobromic acid and
form less reactive bromine radicals (free bromine atoms).[48] This makes such compounds
useful fire retardants and this is bromine's primary industrial use, consuming more than
half of world production of the element. For example, TBBPA (Tetrabromobisphenol A)
decabromodiphenyl ether, and vinyl bromide[32] are used as reactive intermediates,
additives and flame retardants.
1.3.9.3 Gasoline additive
Ethylene bromide was an additive in gasoline containing lead (Pb) anti-engine
knocking agents. It scavenges lead by forming volatile lead bromide, which is exhausted
from the engine.
1.3.9.4 Pesticide
Poisonous methyl bromide was widely used as pesticide to fumigate soil and to
fumigate housing, by the tenting method. In the year 1991, an estimated 35,000 metric
tons of the chemicals were used to control nematodes, fungi, weeds and other soil-borne
diseases.[ 49]
1.3.9.5 Other uses
Bromides in the form of simple salts are used as anticonvulsants in both veterinary
and human medicine. Ethidium bromide is a popular fluorescent gel stain for DNA
detection through intercalation.[50] Pure cesium bromide obtained has been used in the
manufacturing of optical prisms which are highly transparent to infrared radiations.
Potassium bromide is used in some photographic developers to inhibit the formation of
fog (undesired reduction of silver). Bromine is also used to reduce mercury pollution
from coal-fired power plants. This can be achieved either by treating activated carbon
with bromine or by injecting bromine compounds onto the coal prior to combustion. Soft
drinks containing brominated vegetable oils are sold in the US (2011).[51a] Various
bromine containing compounds are used in various pharmaceutical applications such as
brompheniramine,[51b] bromocriptine (parkinsons disease),[51c] citalopram hydrobromide
(antidepressant),[52] homatropine methyl bromide (anticholinergic),[53] propantheline
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Table 1.4: Applications of bromine and its some organobromine compounds
Bromocompounds
Bromine
Applications
Manufacture of pharmaceuticals, organic preparations, dyes, pesticides, fire
retardants, germicides, photographic chemicals, perfumes, heavy brines and
other chemicals
Bromoxynil
Extensively used as herbicides in the control of broad-leafed weeds
Allyl bromide
Organic synthesis
Methyl bromide
As fumigant for soil fumigation and for space fumigation to control insects
in stored products
Ethylene dibromide
As fumigant and an important constituent of ethyl petrol used as a motor
fuel. Ethylene dibromide finds use as an agricultural fumigant but this use in
the United States is now prohibited because its presense found in ground
water
Methyl Bromide
As fumigant for soil fumigation and for space fumigation to control insects
in stored products
Ethyl bromide
Organic synthesis, used in the manufacture of pharmaceutical like Vitamin
A etc., in flame-retardants, refrigerant
N-Butyl bromide
Intermediate for pharmaceuticals, insecticides, quaternary ammonium
compounds and pigments.
Isobutyl bromide
Organic synthesis, pharmaceutical intermediates
Bromobenzene
Solvent, organic synthesis, agricultural intermediates, pharmaceutical
intermediates
Cetylbromide
Organic synthesis, surfactant manufacturing, imaging chemicals
Bronopol/2-bromo-2-
In cosmetics as a bactericide, antiseptic and preservative. Preservatives for
nitro-1,3-propane diol
coatings, slurries and to control microbial fouling in paper mills and process
water systems
2,4,6-Tribromophenol
Intermediate for high molecular weight flame retardant. Effective fungicide
and wood preservative
2,4,4,6-Tetrabromo-2,
As selective brominating reagent and as many functional group
5- cyclohexadienone
transformation in organic synthesis
Tetrabromobisphenol-
A Solaris FR10 is highly effective when reacted into epoxy resin system,
A
due to its structural compatibility, high bromine content and thermal stability
Its high purity allows its use as a reactive flame retardant for polycarbonates
and as an additive for styrenic thermoplastics such as ABS and HIPS
4-Bromoanisole
Organic Synthesis, pharmaceutical Intermediates
1-Bromo-2-naphthol
In organic synthesis
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2,4,6-Tribromoaniline
Industrial intermediate
4-BromoN, N-
As dye intermediate
dimethyl aniline
4-Bromoacetanilide
Drug intermediate in pharmaceutical industry
N-Bromosuccinimide
As a secondary brominating reagent and also as functional group
transformation
N-Bromosaccharine
As a secondary brominating reagent
4-Nitro benzyl
Intermediate in organic synthesis
bromide
bromide (antimuscarinic),[54a] and pyridostigmine bromide (cholinesterase inhibitor).[54]
Applications of bromine and some of its organobromine compounds in industry/academia
are summarized in Table 1.4.[32]
1.3.10 Biological role
Direct applications of bromine has no role in human or mammalian health, but it’s
inorganic and organobromine compounds are occurs naturally, and some of them may be
of useful to higher organisms in dealing with parasites. For example, in the presence of
H2O2 formed by the eosinophil, and either chloride or bromide ions, eosinophil
peroxidase provides a potent mechanism by which eosinophils kill multicellular parasites
and also certain bacteria (such as tuberculosis bacteria).[55] Most of the organobromine
compounds in nature arise from the sea, via the action of a unique algal enzyme,
vanadium bromoperoxidase.[56] Though this enzyme has the most prolific creator of
organic bromides by living organisms, other bromoperoxidases exist in nature that do not
use vanadium. Tyrian purple is the famous example of a bromine-containing organic
compound that has been used by humans since ancient times for dying the fabrics.[57]
Bromine can also be substituted for the methyl substituent in the nitrogenous base
thymine of DNA, creating the base analog 5-bromouracil. When this base is incorporated
into DNA its different hydrogen bonding properties may cause mutation at the site of that
base pair.[58]
1.3.11 Safety
The liquid bromine rapidly attacks the skin and other tissues to produce irritation and
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necrosis.[21] Bromine is not a carcinogen but liquid and gaseous bromine however pose
serious health hazard. Exposure to bromine vapors can cause painful irritation to the eyes
and inflammation of the respiratory tract. Less than 1 ppm is the maximum concentration
of bromine which is considered safe for an eight-hour exposure. Inhalation of bromine is
dangerous to health, and even small amounts (more than 10 ppm) of bromine may cause
great discomfort. Inhalation of high concentrations may cause inflammatory scratches of
the mucous membranes of the upper respiratory tract, fatal chemical burns and even
respiratory failure. Bromine may be deposited in the tissues as bromides and accumulate
to cause central nervous system disorders.[32] Intake of excessive bromide can induce a
condition termed ‘bromism’. The pathology of animals exposed to 300 ppm for 3 hours
showed pulmonary edema, pseudo membranous deposit on the trachea and bronchi, and
hemorrhages of the gastric mucosa. Prolonged or repeated exposure of liquid bromine
causes headache, pain in the region of the heart, increasing irritability, loss of appetite,
joint pains, dyspepsia loss of corneal reflexes, pharyngitis, vegetative disorders, thyroid
hyperplasia accompanied by thyroid dysfunction and bone marrow depression.
Cardiovascular disorders may occur in the form of myocardial degeneration and
hypotension. Functional and secretory disorders of the digestive tract may also occur.
Hematologic effects may include inhibition of leucopoiesis, leukocytosis, moderate
hypoglycemia or altered blood sugar curves, hypercholesterolemia, reduction of total
bilirubin, decreased hemoglobin concentration and increased erythrocyte sedimentation
rates. Elemental bromine is toxic and causes burns. As an oxidizing agent, it is
incompatible with most organic and inorganic compounds. Care needs to be taken when
transporting bromine; it is commonly carried in steel tanks lined with lead, supported by
strong metal frames. When certain ionic compounds containing bromine are mixed with
potassium permanganate (KMnO4) and acidic substance, they will form a pale brown
cloud of bromine gas. This gas smells like bleach and is very irritating to the mucous
membranes. Upon exposure, one should move to fresh air immediately. If symptoms of
bromine poisoning arise, medical attention is needed.
FIRST AID[32]. Remove immediately from exposure area to fresh air if inhaled or
contact. If breathing has stopped, perform artificial respiration. Keep the person warm
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and at rest. Treat symptomatically and supportively. Get medical attention immediately.
Remove contaminated clothing and shoes immediately. Wash affected area with soap or
mild detergent and large amounts of water until no evidence of chemical. In case of
chemical burns, cover area with sterile, dry dressing. Bandage securely, but not too
tightly. Wash eyes immediately with large amounts of water. Continue irrigating with
normal saline until the pH has returned to normal. Cover with sterile bandages. Do not
use emesis. Dilute chemical immediately by drinking large amounts of water or milk. If
vomiting persists, administer fluids repeatedly. Do not give an unconscious person
anything to drink. Get medical attention.
1.4 Bromide/Bromate couple
Recently an eco-friendly brominating reagent (comprising of bromide/bromate
couple in various mole ratious) has been developed at CSMCRI, Bhavnagar for diverse
applications such as addition, substitution and oxidation reactions.[59] The reagent is
bromide bromate couple in different stoichiometry which in presence of acid (H+)
generate the reactive species bromine and BrOH. At present the three different
brominating reagents are prepared and used as a) 2:1 bromide:bromate known as ‘BR-S’
for oxidation, oxybromination and substitution reactions b) 1:8 bromide:bromate known
as ‘BR-O’ for oxidation. c) 5:1 bromide:bromate known as ‘BR-A’ for addition reactions.
The details about synthesis, reactivity and applications of Green Brominating Reagent
will be described in details in chapter 2 titled “Review of reagents and catalysis”.
1.5 Catalytic approach
A catalyst is a substance that changes the rate of reaction without being consumed in
the reaction and this phenomenon of reactions is called catalysis. Catalysts are employed
in a number of industrial processes. Nature is the master designer and user of catalysts.
For example, the simplest bacterium produces thousands of biological catalysts called
enzymes to speed up its cellular reactions.[60] Every living organism depends upon the
enzymes to sustain life. Catalysts that speeds the reactions are called positive catalysts
while those retards the reactions are called negative catalysts. Catalytic reactions have a
lower rate-limiting free energy of activation than the corresponding uncatalyzed reaction,
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resulting in higher reaction rate at the same temperature. However, the mechanistic
explanation of catalysis is complex. Catalysts may affect the reaction environment
favorably, or bind to the reagents to polarize bonds. Catalysis is high on approaches to
minimize green chemistry ‘E-Factor’.
1.5.1 Metal catalysis
Some of the most exciting reactions in organic chemistry are based on transition
metal catalysts. For example Heck reaction,[61] which allows nucleophilic addition to an
unactivated alkene. Another example is Pauson-Khand[62] reaction, is special method
making of five membered rings from three components an alkene, an alkyne and carbon
monoxide (CO). Reagents and complexes containing transition metals are important in
modern synthetic organic chemistry because they allow apparently impossible reactions
to occurs easily. This chemistry complements traditional functional group based
chemistry and significantly broaden the scope of organic chemistry. Many industries and
academia now routinely uses transition metals catalyzed reactions because of their
selectivity towards desired products, cost effectiveness and environmental safety. In the
present thesis, such vital approaches are successfully attempted particularly copper metal
and copper salts catalyzed reactions were developed for the important functional group
transformations. Hence brief introduction about relevant copper chemistry is described.
Copper: Copper is the base metal of group 11 at the top of transition series along with
other base metals silver and gold. It is a ductile metal with very high thermal and
electrical conductivity. Pure copper is soft and malleable; an exposed surface has a
reddish-orange tarnish. It is used as a conductor of heat and electricity, a building
material, and a constituent of various metal alloys. Its various general, atomic, physical
and miscellaneous properties are summarized in the following Table 1.5.[11,63] Although
these metals are mostly thought to be used for ornaments and other domestic purpose
from long ago, the arrival of synthetic modern chemistry using these metals has a far
different lineage. Among the copper, silver and gold, copper has rich history in
organometallic chemistry, starting with Gilman in 1950s.[64] Certaintly a copper catalyzed
Grignard or Zinc reagents, or stoichiometric lithiocuprates are important, but these
methods in organocopper chemistry are now textbook chemistry. Now a day’s field has
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Table 1.5: Properties of copper
300.4 kJ·mol−1
Heat of
General properties
vaporization
Copper
Heat of fusion
13.26 kJ.mol−1
24.440 J.mol−1·K−1
Symbol
Cu
Specific heat
capacity
Atomic
Number
29
Name
Chemical
Series
Group,
period, block
Atomic properties
Crystal structure
Orthorhombic
Oxidation states
+1,
+2,
+3,
(mildly basic oxide)
Electronegativity
1.90 (Pauling scale)
Ionization
energies
1st: 1st: 745.5 kJ·mol−1
2nd: 1957.9 kJ·mol−1
−1
3rd: 3555 kJ·mol
Transition metal
11, 4, d
Appearance
Native copper
Standard
atomic weight
63.546 g/mol
Atomic radius
128 pm
Electron
configuration
[Ar] 3d10 4s1
-
-
Electrons per
shell
2, 8, 18, 1
Covalent radius
132 pm
Van der
radius
140 pm
Physical Properties
Phase
waals
+4
Miscellaneous Properties
Solid
3
Density
8.94 g/cm
Magnetic ordering
Dimagnetic
Melting point
1357.77 K (1084.62 °C,
1984.32 °F)
Thermal
conductivity
401 Wm−1K−1
Boiling point
2835 K (2562 °C, 4643 °F)
CAS registry
number
7440-50-8
changed, catalysis has become a driving force for organometallic chemistry. Copper is
now well known for variety of bond constructions, the advances in stereo control and
exclusive chemoselectivity.[65] In particular asymmetric catalysis mediated by copper
reagents has large and significant contribution to the synthetic organic chemistry.[66]
Many reviews and monographs over the decades have nicely served the needs of
synthetic community devoted to this vibrant and continuously emerging field.
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1.5.2 Non-metal catalysis
Usually in many catalytic processes, the catalyst means transition metals and their
complexes were employed which activates the reaction. The dependency of organic
chemists only on metal catalysis is certaintly not healthy situation for the future
advancement and developments. Now-a-days a major research has been going for the
development of sustainable methodologies and to find better alternative for traditional
methods. Non-metal catalysts have gained much attention in recent years. The reason for
the non-metallic catalytic system is that: a number of such catalysts have become readily
accessible; on the other hand, such catalysts are quite resistant toward self-oxidation and
compatible under aerobic and aqueous reaction conditions unlike many other metal
catalysts. One of the major advantages of metal-free organic catalysts is their better
environmental acceptance compared to transition-metal catalysts, because most of the
metals are toxic/hazardous to the environment.
In present thesis work, we have focused on non-metal catalytic systems for the
hydroarylation of styrenes. Various inorganic acids are attempted to get effective
conversion and selectivity for desired products. Potassium hydrogen sulphate KHSO4 a
Bronsted acid was found to be a highly effective catalyst for hydroarylation. The detailed
studies will be explained in the preceding chapters separately. Such non-metallic catalysis
has attained prominence in the synthesis field in view of their efficiency and their
potential for future developments. Under mild reaction conditions, non-metal catalytic
process holds much promise for future practical applications too.
Objectives
The specific objectives of the present work are as follows:
1. Synthesis of α-bromoketones from olefins using Green Brominating Reagent.
2. Direct synthesis of imines from amines using copper catalyst.
3. Arylation of styrenes using KHSO4 as a catalyst.
4. Oxidation of methyl arenas to carboxylic acids and direct synthesis of
benzaldehyde from benzyl bromides.
5. Regioselective synthesis of bromohydrins and α-bromoketones from olefins using
HBr/H2O2 system.
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6. Direct synthesis of bromohydrins and α-bromoketones from vic-dibromo using
H2O2 system.
7. Direct synthesis of α,α-dibromoketones from alkynes using HBr/H2O2 system.
8. Process development and technology transfer for important brominated
compounds.
References & Notes
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[4]. R. Stewart, W. A. Benzamin, Oxidation Mechanism Applications to Organic
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[5]. J. March, Advanced Organic Chemistry reactions, mechanisms and structure (3rd
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[6]. M. J. Dagani, H. J. Barda, T.J. Benya, D. C. Sanders, Ullmann’s Encyclopedia of
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magnesium sulfate. It is a water solution of bromides, magnesium, and calcium salts
remaining after sodium chloride is crystallized out from seawater.
[25]. http://encyclopedia.jrank.org/BRI_BUN/BROMINE_symbol_Br_atomic_weight.ht
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