Chemicals GREENHOUSE CHEMICALS by nuhman10

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									                   GREENHOUSE CHEMICALS
                               CARBON DIOXIDE – 83%
          Carbon dioxide, CO2, is one of the gases in our atmosphere, being uniformly
distributed over the earth's surface at a concentration of about 0.033% or 330 ppm.
Commercially, CO2 finds uses as a refrigerant (dry ice is solid CO2), in beverage
carbonation, and in fire extinguishers. In the United States, 10.89 billion pounds of
carbon dioxide were produced by the chemical industry in 1995, ranking it 22nd on
the list of top chemicals produced. Because the concentration of carbon dioxide in
the atmosphere is low, it is not practical to obtain the gas by extracting it from air.
Most commercial carbon dioxide is recovered as a by-product of other processes,
such as the production of ethanol by fermentation and the manufacture of ammonia.
Some CO2 is obtained from the combustion of coke or other carbon-containing fuels.

                         C(coke) + O2(g)               CO2(g)

         Carbon dioxide is released into our atmosphere when carbon-containing
fossil fuels such as oil, natural gas, and coal are burned in air. As a result of the
tremendous world-wide consumption of such fossil fuels, the amount of CO2 in the
atmosphere has increased over the past century, now rising at a rate of about 1 ppm
per year. Major changes in global climate could result from a continued increase in
CO2 concentration.
In addition to being a component of the atmosphere, carbon dioxide also dissolves in
the water of the oceans. At room temperature, the solubility of carbon dioxide is
about 90 cm of CO2 per 100 mL of water. In aqueous solution, carbon dioxide exists
in many forms. First, it simply dissolves.

                             CO2(g)               CO2(aq)

Then, an equilibrium is established between the dissolved CO2 and H2CO3, carbonic

                       CO2(aq) + H2O(l)              H2CO3(aq)

Only about 1% of the dissolved CO2 exists as H2CO3. Carbonic acid is a weak acid
which dissociates in two steps.

                                       +                                      7
                H2CO3                 H + HCO3¯             Ka1 = 4.2 × 10¯
                                       +      2                               11
                HCO3¯                 H + CO3 ¯             Ka2 = 4.8 × 10¯

         As carbon dioxide dissolves in sea water, an equilibrium is established
involving the carbonate ion, CO3 ¯. The carbonate anion interacts with cations in
seawater. According to the solubility rules, "all carbonates are insoluble except those
of ammonium and Group IA elements." Therefore, the carbonate ions cause the
                                               2+        2+
precipitation of certain ions. For example, Ca and Mg ions precipitate from large
bodies of water as carbonates. For CaCO3, the value of Ksp is 5 × 10¯ , and for
MgCO3, Ksp is 2 × 10¯ . Extensive deposits of limestone (CaCO3) and dolomite
(mixed CaCO3 and MgCO3) have been formed in this way. Calcium carbonate is also
the main constituent of marble, chalk, pearls, coral reefs, and clam shells.
       Although "insoluble" in water, calcium carbonate dissolves in acidic solutions.
The carbonate ion behaves as a Brønsted base.

                                +                  2+
                  CaCO3(s) + 2 H (aq)            Ca (aq) + H2CO3(aq)

The aqueous carbonic acid dissociates, producing carbon dioxide gas.

                         H2CO3(aq)             H2O(l) + CO2(g)

        In nature, surface water often becomes acidic because atmospheric CO 2
dissolves in it. This acidic water can dissolve limestone.

         CO2(aq) + H2O(l) + CaCO3(s)                Ca (aq) + 2 HCO3¯(aq)

This reaction occurs in three steps.

                                                             2+       2
                             CaCO3(s)                   Ca (aq) + CO3 ¯(aq)
                       CO2(aq) + H2O(l)                 H2CO3(aq)
              H2CO3(aq) +    CO3 ¯(aq)                  2 HCO3¯(aq)

In the third step, carbonate ions accept hydrogen ions from carbonic acid. This
reaction often occurs underground, when rainwater saturated with CO2 seeps
through a layer of limestone. As the water dissolves calcium carbonate, it forms
openings in the limestone. Caves from which the limestone has been dissolved are
often prevalent in areas where there are large deposits of CaCO3 (e.g., Mammoth
Cave, Carlsbad Caverns, and Cave of the Mounds). If the water containing dissolved
Ca(HCO3)2 reaches the ceiling of a cavern, the water will evaporate. As it evaporates,
carbon dioxide escapes, and calcium carbonate deposits on the ceiling.

          Ca (aq) + 2 HCO3(aq)                H2O(g) + CO2(g) + CaCO3(s)
         A new use for liquid carbon dioxide currently under development
is as a dry-cleaning solvent. Currently, most laundries use chlorinated
hydrocarbons as dry-cleaning solvents. These chlorinated hydrocarbons
are probable human carcinogens, so the search is on for replacements.
Carbon dioxide does not exist in liquid form at atmospheric pressure at
any temperature. The pressure-temperature phase diagram of CO2
shows that liquid carbon dioxide at 20°C requires a pressure of 30
atmospheres. The lowest pressure at which liquid CO2 exists is at the
triple point, namely 5.11 atm at -56.6°C. The high pressures needed for
liquid CO2 require specialized washing machines. Like chlorinated
hydrocarbons, liquid carbon dioxide is an effective solvent for grease and
oils. Liquid CO2 has some advantages over chlorinated hydrocarbons--
items that cannot be dry cleaned with chlorinated hydrocarbons, such as
leather, fur, and some synthetics, can be safely cleaned with liquid
carbon dioxide. More information about alternative dry-cleaning solvents
can be found in the Innovations section of Environmental Health
Perspectives, Volume 104, Number 5.
                                   METHANE – 9%

Methane is a colorless, odorless gas with a wide distribution in nature. It is the
principal component of natural gas, a mixture containing about 75% CH4, 15%
ethane (C2H6), and 5% other hydrocarbons, such as propane (C3H8) and butane
(C4H10). The "firedamp" of coal mines is chiefly methane. Anaerobic bacterial
decomposition of plant and animal matter, such as occurs under water, produces
marsh gas, which is also methane.
At room temperature, methane is a gas less dense than air. It melts at –183°C and
boils at –164°C. It is not very soluble in water. Methane is combustible, and mixtures
of about 5 to 15 percent in air are explosive. Methane is not toxic when inhaled, but it
can produce suffocation by reducing the concentration of oxygen inhaled. A trace
amount of smelly organic sulfur compounds (tertiary-butyl mercaptan, (CH3)3CSH
and dimethyl sulfide, CH3–S–CH3) is added to give commercial natural gas a
detectable odor. This is done to make gas leaks readily detectible. An undetected
gas leak could result in an explosion or asphyxiation. (The attached scratch-and-sniff
sheet from Madison Gas & Electric Company is for your use outside of class.)
Methane is synthesized commercially by the distillation of bituminous coal and by
heating a mixture of carbon and hydrogen. It can be produced in the laboratory by
heating sodium acetate with sodium hydroxide and by the reaction of aluminum
carbide (Al4C3) with water.
In the chemical industry, methane is a raw material for the manufacture of methanol
(CH3OH), formaldehyde (CH2O), nitromethane (CH3NO2), chloroform (CH3Cl), carbon
tetrachloride (CCl4), and some freons (compounds containing carbon and fluorine,
and perhaps chlorine and hydrogen). The reactions of methane with chlorine and
fluorine are triggered by light. When exposed to bright visible light, mixtures of
methane with chlorine or fluorine react explosively.
The principal use of methane is as a fuel. The combustion of methane is highly
   CH4(g) + 2 O2(g)              CO2(g) + 2 H2O(l)                        H = –891 kJ
The energy released by the combustion of methane, in the form of natural gas, is
used directly to heat homes and commercial buildings. It is also used in the
generation of electric power. During the past decade natural gas accounted for about
1/5 of the total energy consumption worldwide, and about 1/3 in the United States.
The cost of natural gas to Wisconsin consumers is regulated by the State Public
Service Commission. Madison Gas Electric Company currently charges its residential
consumers about $0.66 per 100 cubic feet.
Natural gas occurs in reservoirs beneath the surface of the earth. It is often found in
conjunction with petroleum deposits. Before it is distributed, natural gas usually
undergoes some sort of processing. Usually, the heavier hydrocarbons (propane and
butane) are removed and marketed separately. Non-hydrocarbon gases, such as
hydrogen sulfide, must also be removed. The cleaned gas is then distributed
throughout the country through thousands of miles of pipeline. Local utility companies
add an odorant before delivering the gas to their customers.
Some methane is manufactured by the distillation of coal. Coal is a combustible rock
formed from the remains of decayed vegetation. It is the only rock containing
significant amounts of carbon. The elemental composition of coal varies between
60% and 95% carbon. Coal also contains hydrogen and oxygen, with small
concentrations of nitrogen, chlorine, sulfur, and several metals. Coals are classified
by the amount of volatile material they contain, that is, by how much of the mass is
vaporized when the coal is heated to about 900°C in the absence of air. Coal that
contains more than 15% volatile material is called bituminous coal. Substances
released from bituminous coal when it is distilled, in addition to methane, include
water, carbon dioxide, ammonia, benzene, toluene, naphthalene, and anthracene. In
addition, the distillation also yields oils, tars, and sulfur-containing products. The non-
volatile component of coal, which remains after distillation, is coke. Coke is almost
pure carbon and is an excellent fuel. However, it may contain metals, such as arsenic
and lead, that can be serious pollutants if the combustion products are released into
the atmosphere.
         NITROUS OXIDE – 6%

Nitrous oxide, also known as dinitrogen
oxide or dinitrogen monoxide, is a
chemical compound with chemical                                        General
formula N2O. Under room conditions it is
                                              Name                 Dinitrogen oxide
a colourless non-flammable gas, with a
pleasant, slightly sweet odor. It is          Chemical
commonly known as laughing gas due            formula
to the exhilarating effects of inhaling it,   Appearance           Colorless gas
and because it can cause spontaneous
laughter in some people; it's also known                              Physical
as NOS or nitrous in racing and               Formula weight 44.0 u
motorsports, where its usage is               Melting point        182 K (-91 °C)
widespread. It is used in surgery and
dentistry for its anaesthetic and             Boiling point        185 K (-88 °C)
analgesic effects. Nitrous oxide is           Critical
present in the atmosphere where it acts                            309.6 K (36.4 °C)
as a powerful greenhouse gas.
                                                                   7.245 MPa
Chemistry                                                                     3
                                              Density              1.2 g/cm (liquid)
The structure of the nitrous oxide       Solubility                0.112 g in 100g water
molecule is a linear chain of a nitrogen                          Thermochemistry
atom bound to a second nitrogen, which      0
in turn is bound to an oxygen atom. It   ΔfH gas                   82.05 kJ/mol
can be considered a resonance hybrid ΔfH0liquid                    ? kJ/mol
of                                          0
                                         ΔfH solid                 ? kJ/mol
                                              S gas, 100 kPa       219.96 J/(mol·K)
                                              S liquid, 100 kPa    ? J/(mol·K)
                                              S solid              ? J/(mol·K)
                                              n                    See main text. May cause
                                              d                    asphyxiation without warning.
                                                                   Hazardous when cryogenic or
                                                                   Hazardous when cryogenic or
                                                      SI units were used where possible.
Nitrous oxide [[N2O]] should not be
confused with the other nitrogen oxides such as nitric oxide NO and nitrogen dioxide

Note that nitrous oxide is isoelectric with carbon dioxide.

Nitrous oxide can be prepared by heating ammonium nitrate in the laboratory. This is
not, however, advised, since overheated ammonium nitrate can easily explode.
Nitrous oxide can be used to produce nitrites by mixing it with boiling alkali metals,
and to oxidize organic compounds at high temperatures.

The CAS number of nitrous oxide is 10024-97-2 and its UN number is 1070.


Medical grade Nitrous Oxide tanks used in dentistry

Nitrous oxide is a weak general anesthetic, and is generally not used alone in
anaesthesia. However, it has a very low short-term toxicity and is an excellent
analgesic, so a 50/50 mixture of nitrous oxide and oxygen ("gas and air", supplied
under the trade name Entonox) is commonly used during childbirth, for dental
procedures, and in emergency medicine.
In general anesthesia it is often used in an 2:1 ratio with oxygen in addition to more
powerful general anaesthetic agents such as sevoflurane or desflurane. Its lower
solubility in blood means it has a very rapid onset and offset.
It has a MAC of 105% and a blood:gas partition coefficient of 0.46. Less than 0.004%
is metabolised in humans.
Nitrous Oxide is liquid at approximately 760 psi at room temperature, and is usually
stored and shipped as a self-pressurized liquid.

Aerosol propellant

The gas is licensed for use as a food additive, specifically as an aerosol spray
propellant. Its most common uses in this context are in aerosol whipped cream
canisters and as an inert gas used to displace staleness-inducing oxygen when filling
packages of potato chips and other similar snack foods.
The gas is excellently soluble in fatty compounds. In aerosol whipped cream, it is
dissolved in the fatty cream until it leaves the can, when it becomes gaseous and
thus creates foam. One can easily obtain the propellant by slowly turning the canister
upside down (NO SHAKING) and letting all the contents out, leaving you the N2O.
However, if one is using the Nitrous for recreational purposes, using N2O straight
from a whipped cream can is unadvisable due to the fact that it is frequently cut with
certain chemicals that can cause headaches or nausea. There is also usually a
negligible amount of N2O in the cans.

Rocket motors

Nitrous oxide can be used as an oxidizer in a rocket engine. This has the advantages
over other oxidizers that it is non-toxic and, due to its stability at room temperature,
easy to store and relatively safe to carry on a flight.
Nitrous oxide has notably been the oxidizer of choice in several hybrid rocket
designs (using solid fuel with a liquid or gaseous oxidizer). The combination
of nitrous oxide with hydroxy-terminated polybutadiene fuel has been used by
SpaceShipOne and others. It is also notably used in amateur and high power
rocketry with various plastics as the fuel. An episode of MythBusters featured a
hybrid rocket built using paraffin wax mixed with powdered carbon as its
solid fuel and nitrous oxide as its oxidizer.
Internal Combustion Engine

In car racing, nitrous oxide (often just "nitrous" in this context) is sometimes injected
into the intake manifold (or just prior to the intake manifold) to increase power: even
though the gas itself is not flammable, it delivers more oxygen than atmospheric air
by breaking down at elevated temperatures, thus allowing the engine to burn more
fuel and air. Additionally, since nitrous oxide is stored as a liquid, the evaporation of
liquid nitrous oxide in the intake manifold causes a large drop in intake charge
temperature. This results in a smaller, denser charge, and can reduce detonation, as
well as increase power available to the engine.
The same technique was used during by World War II Luftwaffe aircraft with the GM
1 system to boost the power output of aircraft engines. Originally meant to provide
the Luftwaffe standard aircraft with superior high-altitude performance, technological
considerations limited its use to extremely high altitudes. Accordingly, it was only
used by specialized planes like high-altitude reconnaissance aircraft, high-speed
bombers and high-altitude interceptors.
One of the major problems of using nitrous oxide in a reciprocating engine is that it
can produce enough power to destroy the engine. Power increases of 100-300% are
possible, and unless the mechanical structure of the engine is reinforced, most
engines would not survive this kind of operation.
There are several ways of introducing nitrous into a motor. Nitrous kits such as such
as NOS, Nitrous Express, Nitrous Direct brands offer different solutions. You will find
Dry kits, Wet kits & Direct port. See nitrous
It is very important with nitrous oxide augmentation of internal combustion engines to
maintain temperatures and fuel levels so as to prevent preignition, or detonation
(sometimes referred to as knocking, pinging or pinking).

Nitrous oxide in the atmosphere

Nitrogen oxides, nitrous oxide included, are greenhouse gases; per kilogram, nitrous
oxide has 296 times the effect of carbon dioxide for producing global warming [1].
Therefore, nitrogen oxides are a subject of efforts to curb greenhouse gas emissions,
such as the Kyoto Protocol. Behind carbon dioxide and methane, nitrous oxide is the
third most important gas that contribute to global warming.

Nitrous oxide is naturally emitted from soils and oceans. Human activity contributes
to the release of the gas through the cultivation of soil and the production and use of
nitrogen fertilizers, the production of nylon, and the burning of fossil fuels and other
organic matter.

Human activity is thought to account for somewhat less than 2 teragrams (this is
multiplied by appx 300 when calculated as a ratio to Carbon Dioxide) of nitrogen
oxides per year, nature for over 15 teragrams
                                 HFC’s, PCs, SF4s

A haloalkane, also known as alkyl halogenide, halogenalkane, or halogenoalkane,
and alkyl halide is a chemical compound derived from an alkane by substituting one
or more hydrogen atoms with halogen atoms. Substitution with fluorine, chlorine,
bromine and iodine results in fluoroalkanes, chloroalkanes, bromoalkanes and
iodoalkanes, respectively. Mixed compounds are also possible, examples are the
chlorofluorocarbons (CFCs) which are mainly responsible for ozone depletion.
Haloalkanes are used in semiconductor device fabrication, as refrigerants, foam
blowing agents, solvents, aerosol spray propellants, fire extinguishing agents, and
chemical reagents.
Freon is a trade name for a group of chlorofluorocarbons used primarily as a
refrigerant. The word Freon is a registered trademark belonging to DuPont.
Chlorofluoro compounds (CFC, HCFC, HFC)

CFC molecules

Chlorofluorocarbons (CFC) are haloalkanes with both chlorine and fluorine. They
were formerly used widely in industry, for example as refrigerants, propellants, and
cleaning solvents. Their use has been generally prohibited by the Montreal Protocol,
because of effects on the ozone layer (see ozone depletion).
Hydrochlorofluorocarbons (HCFCs) is one of a class of haloalkanes where not all
hydrogen has been replaced by chlorine or fluorine. They are used primarily as
chlorofluorocarbon (CFC) substitutes, as the environmental effects are less than for
CFCs. When the chlorine is reduced to zero, these compounds are known as
hydrofluorocarbons (HFCs), with even less environmental effects.

Bromofluoro compounds (halons)

Halon is the group of haloalkanes with bromine as well as chlorine or fluorine groups.
The two most common ones are bromochlorodifluoromethane (Halon 1211, CF2BrCl)
and bromotrifluoromethane (Halon 1301, CF3Br). Halons are very stable and have
been widely used in fire extinguishers where water and other alternatives would be
ineffective and dangerous (e.g. when dealing with fires involving live electrical
circuits) or cause unacceptable collateral damage (e.g. with electronic equipment).

Polymer haloalkanes

Chlorinated or fluorinated alkenes can be used for polymerization, resulting in
polymer haloalkanes with notable chemical resistance properties. Important
examples include polychloroethene (polyvinyl chloride, PVC), and
polytetrafluoroethene (PTFE, Teflon), but many more halogenated polymers exist.


Original development

Carbon tetrachloride was used in fire extinguishers and glass (anti)-"fire grenades"
from the late nineteenth century until around the end of World War II.
Experimentation with chloroalkanes for fire suppression on military aircraft began at
least as early as the 1920s.
American engineer Thomas Midgley developed Chlorofluorocarbons (CFC) in 1928
as a replacement for ammonia (NH3), chloromethane (CH3Cl), and sulfur dioxide
(SO2), toxic but in common use at the time as refrigerants. The new compound
developed had to have a low boiling point, be non-toxic, and be generally non-
reactive. In a demonstration for the American Chemical Society, Midgley
flamboyantly demonstrated all these properties by inhaling a breath of the gas and
using it to blow out a candle.
Midgley specifically developed CCl2F2. However, one of the attractive features is that
there exists a whole family of the compounds, each having a unique boiling point
which can suit different applications. In addition to their original application as
refrigerants, chlorofluoroalkanes have been used as propellants in aerosol cans,
cleaning solvents for circuit boards, and as blowing agents for making expanded
plastics (such as the expanded polystyrene used in packaging materials and
disposable coffee cups).

Development on alternatives

During World War II, various early chloroalkanes were in standard use in aircraft by
some combatants, but these early halons suffered from excessive toxicity.
Nevertheless after the war they slowly became more common in civil aviation as well.
In the 1960s, fluoroalkanes and bromofluoroalkanes became available, and were
quickly recognised as one of the most effective fire fighting materials discovered.
Much early research with Halon 1301 was conducted under the auspices of the US
Armed Forces, while Halon 1211 was initially mainly developed in the UK. By the late
1960s, they were standard in many applications where water and dry powder
extinguishers posed a threat of damage to the protected property, including computer
rooms, telecommunications switches, laboratories, museums and art collections.
Beginning with warships in the 1970s, bromofluoroalkanes also progressively came
to be associated with rapid knockdown of severe fires in confined spaces with
minimal risk to personnel.
Work on alternatives for chlorofluorocarbons in refrigerants began in the late 1970s
after the first warnings of damage to stratospheric ozone were published in the
journal Nature in 1974 by Molina and Rowland (who shared the 1995 Nobel Prize for
Chemistry for their work). Adding hydrogen and thus creating
hydrochlorofluorocarbons (HCFC), chemists made the compound less stable in the
lower atmosphere enabling them to break down before reaching the ozone layer.
Later alternatives even fully excluded the chlorine, creating hydrofluorocarbons
(HFC) with even shorter lifetimes in the lower atmosphere.
By the early 1980s, bromofluoroalkanes were in common use on aircraft, ships and
large vehicles, as well as in computer facilities and galleries. However, concern was
beginning to be felt about the possible impact of chloroalkanes and bromoalkanes on
the ozone layer. The Vienna Convention on Ozone Layer Protection did not cover
bromofluoroalkanes as it was felt that emergency discharge of systems was too small
in volume to produce a significant impact, and too important to human safety for
restriction. However, by the time of the Montreal Protocol it was realised that
discharges during system tests and maintenance accounted for substantially larger
volumes than emergency discharges, and so halons were brought into the treaty, but
with many exceptions.

Phase out

Use of certain chloroalkanes as solvents for large scale application, such as dry
cleaning, have been phased out, for example by the IPPC directive on greenhouse
gases in 1994 and by the Volatile Organic Compounds (VOC) directive of the EU in
1997. Also chlorofluoroalkanes are minimized to medicinal use only.
At last, bromofluoroalkanes have been generally phased out and the possession of
such equipment is prohibited in some countries like the Netherlands and Belgium
from January 1, 2004, based on the Montreal Protocol and guidelines of the
European Union. Production of new stocks has ceased in most (probably all)
countries as of 1994. However many countries still require aircraft to be fitted with
halon fire suppression systems, as no safe and completely satisfactory alternative
has been discovered for this application. There are also a few other highly
specialised users. These programs recycle halon through "halon banks", coordinated
by the Halon Recycling Corporation, to ensure that discharge to the atmosphere
occurs only in a genuine emergency, and to conserve remaining stocks.


IUPAC nomenclature

The formal naming of haloalkanes should follow IUPAC nomenclature, which put the
halogen as a prefix to the alkane. For example, ethane with bromine becomes
bromoethane, methane with four chlorine groups becomes carbon tetrachloride.
However, many of these compounds have already an established trivial name, which
is endorsed by the IUPAC nomenclature, for example chloroform (trichloromethane)
and methylene chloride (dichloromethane). For unambiguity, this article follows the
systematic naming scheme throughout.

Alternative nomenclature for refrigerants

The refrigerant naming system is mainly used for fluorinated and chlorinated short
alkanes for refrigerant use. The standard is specified in the ANSI/ASHRAE Standard
34-1992 with additional annual supplements [1]. The specified ANSI/ASHRAE
prefixes were FC (fluorocarbon), or R (refrigerant), but today most are prefixed by a
more specific classification:

       CFC - chlorofluorocarbons
         HCFC - hydrochlorofluorocarbons
         HFC - hydrofluorocarbons
         FC - fluorocarbons
         PFC - perfluorocarbons (completely fluorinated)

  The decoding system for CFC-01234a is:

         0 = number of double bonds (omitted if zero)
         1 = Carbon atoms - 1 (omitted if zero)
         2 = Hydrogen atoms + 1
         3 = Fluorine atoms
         4 = eplaced by Bromine ("B" prefix added)
         a = letter added to identify isomers, the "normal" isomer in any number has
          the smallest mass difference on each carbon, and a, b, or c are added as the
          masses diverge from normal.

  Other coding systems are in use as well.

  Overview of named compounds

                                      Overview of haloalkanes
   This table gives an overview of most haloalkanes in general use or commonly known.
          Listing includes bulk commodity products as well as laboratory chemicals.
                                              Common/Trivial                          Chem.
          Systematic name                                                   Code
                                                   name(s)                           formula
Chloromethane                         Methyl chloride                             CH3Cl
Dichloromethane                       Methylene chloride                          CH2Cl2
Trichloromethane                      Chloroform                                  CHCl3
                                      Carbon tetrachloride, tet, Freon
Tetrachloromethane                                                     CFC-14     CCl4
Trichlorofluoromethane                Freon-11, R-11                   CFC-11     CCl3F
Dichlorodifluoromethane               Freon-12, R-12                   CFC-12     CCl2F2
Chlorotrifluoromethane                                                 CFC-13     CClF3
Chlorodifluoromethane                                                  HCFC-22    CHClF2
Trifluoromethane                      Fluoroform                       HFC-23     CHF3
Difluoromethane                                                        HFC-32     CH2F2
Fluoromethane                         Methyl fluoride                  HFC-41     CH3F
Dibromomethane                        Methylene bromide                           CH2Br2
Tribromomethane                       Bromoform                                   CHBr3
Bromochlorodifluoromethane                                             Halon 1211 CBrClF2
Bromotrifluoromethane                                                  Halon 1301 CBrF3
Iodotrifluoromethane                  Trifluoromethyl iodide           Freon 13T1 CF3I
1,1,1-Trichloroethane                 Methyl chloroform, tri                      Cl3C-CH3
Hexachloroethane                                                       CFC-110    C2Cl6
1,1,2-Trichloro-1,2,2-trifluoroethane Trichlorotrifluoroethane         CFC-113    Cl2FC-CClF2
1,1,1-trichloro-2,2,2-trifluoroethane                                  CFC-113a   Cl3C-CF3
                                      Dichlorotetrafluoroethane        CFC-114    ClF2C-CClF2
                                      Chloropentafluoroethane          CFC-115    ClF2C-CF3
2-Chloro-1,1,1,2-tetrafluoroethane                                  HFC-124         CHF2CF3
1,1,2,2,2-pentafluoroethane        Pentafluoroethane                HFC-125         CHF2CF3
1,1,2,2-Tetrafluoroethane                                           HFC-134         F2HC-CHF2
1,1,1,2-Tetrafluoroethane                                                           F3C-CH2F
1,1-Dichloro-1-fluoroethane                                         HCFC-141b       Cl2FC-CH3
1-Chloro-1,1-difluoroethane                                         HCFC-142b       ClF2C-CH3
1,2-Dichloroethane                  Ethylene dichloride             Freon 150       ClH2C-CH2Cl
1,1-Dichloroethane                  Ethylidene dichloride           Freon 150a      Cl2HC-CH3
1,1-Difluoroethane                                                  HFC-152a        F2HC-CH3
Longer haloalkanes, polymers
1,1,1,2,3,3,3-heptafluoropropane                                    FE-227, FM-
                                                                    R610, PFB,      F3C-CF2-
Decafluorobutane                    perfluorobutane
                                                                    CEA-410         CF2-CF3
Polychloroethene                    polyvinyl chloride, PVC
Polytetrafluoroethene                                                               -[CF2-CF2]x-
                                    PTFE, Teflon


   Alkyl halides can be synthesized from alkanes, alkenes, or alcohols.

   From alkanes

   Alkanes react with halogens by free radical halogenation. In this reaction a hydrogen
   atom is removed from the alkane, then replaced by a halogen atom by reaction with a
   diatomic halogen molecule. Thus:

           Step 1: X2 → 2 X· (Initiation Step)

           Step 2: X· + R-H → R· + HX (1st Propagation Step)

           Step 3: R· + X2 → R-X + X· (2nd Propagation Step)

   Steps 2 and 3 keep repeating, each providing the reactive intermediate needed for
   the other step. This is called a radical chain reaction.

   From alkenes

   An alkene reacts with a hydrogen halogenides (HX) like hydrogen chloride (HCl) or
   hydrogen bromide (HBr) to form a haloalkane. The double bond of the alkene is
   replaced by two new bonds, one to the halogen and one to the hydrogen atom of the
   hydrohalic acid. Markovnikov's rule states that in this reaction, the halogen becomes
   attached to the more substituted carbon more likely. Example:

           H3C-CH=CH2 + HBr → H3C-CHBr-CH3 (primary product) + H3C-CH2-
           CH2Br (secondary product).
Alkenes also react with halogens (X2) to form haloalkanes with two neighboring
halogen atoms. This is sometimes known as "decolorizing" the halogen since the
reagent X2 is colored and the product is usually colorless. Example:

        H3C-CH=CH2 + Br2 → H3C-CHBr-CH2Br

From alkanol (alcohol)

Tertiary alkanol reacts with hydrochloric acid directly to produce tertiary chloroalkane,
but if primary or secondary alkanol is used, an activator such as zinc chloride is
needed. Alternatively the conversion may be performed directly using thionyl
chloride. Alkanol may likewise be converted to bromoalkane using hydrobromic acid
or phosphorus tribromide or iodoalkane using red phosphorus and iodine (equivalent
to phosphorus triiodide). Two examples:

        (H3C)3C-OH + HCl.H2O → (H3C)3C-Cl + 2 H2O

        CH3-(CH2)6-OH + SOCl2 → CH3-(CH2)6-Cl + SO2 + HCl

Reactions of haloalkanes

Haloalkanes are reactive towards nucleophiles. They are polar molecules: the carbon
to which the halogen is attached is slightly electropositive where the halogen is
slightly electronegative. This results in an electron deficient (electrophilic) carbon
which, inevitably, attracts nucleophiles.

Substitution reactions

Substitution reactions involve the replacement of the halogen with another molecule -
thus leaving saturated hydrocarbons, as well as the halogen product.
Hydrolysis--a reaction in which water breaks a bond--is a good example of the
nucleophilic nature of halogenoalkanes. The polar bond attracts a hydroxide ion, OH .
(NaOH(aq) being a common source of this ion). This OH is a nucleophile with a clearly
negative charge, as it has excess electrons it donates them to the carbon, which
results in a covalent bond between the two. Thus C-X is broken by heterolytic fission
resulting in a bromide ion, Br . As can be seen, the OH is now attached to the alkyl
group, creating an alcohol. (Hydrolysis of bromoethane, for example, yields ethanol).
One should note that within the halogen series, the C-X bond weakens as one goes
to heavier halogens, and this affects the rate of reaction. Thus, the C-I of an
iodoalkane generally reacts faster than the C-F of a fluoroalkane.
Apart from hydrolysis, there are a few other isolated examples of nucleophilic

       Ammonia (NH3) and bromoethane yields a mixture of ethylamine,
        diethylamine, and triethylamine (as their bromide salts), and
        tetraethylammonium bromide.
        Cyanide (CN ) added to bromoethane will form propionitrile (CH3CH2CN), a
        nitrile, and Br . Nitriles can be further hydrolyzed into carboxylic acids.

Elimination reactions

Rather than creating a molecule with the halogen substituted with something else,
one can completely eliminate both the halogen and a nearby hydrogen, thus forming
an alkene. For example, with bromoethane and NaOH in ethanol, the hydroxide ion
OH attracts a hydrogen atom - thus removing a hydrogen and bromine from
bromoethane. This results in C2H4 (ethylene), H2O and Br .



One major use of CFCs has been as propellants in aerosol inhalers for drugs used to
treat asthma. The conversion of these devices and treatments from CFC to
halocarbons that do not have the same effect on the ozone layer is well under way.
There are some differences between asthma inhalers using CFCs and the newer
propellants, but the conversion has not proven difficult. (By contrast, a significant
amount of development effort has been required to develop non-CFC alternatives to
CFC-based refrigerants, particularly for applications where the refrigeration
mechanism cannot be modified or replaced.)They have now been outlawed in 500
states universally.

Fire extinguishing

At high temperatures, halons decompose to release halogen atoms that combine
readily with active hydrogen atoms, quenching the flame propagation reaction even
when adequate fuel, oxygen and heat remains. The chemical reaction in a flame
proceeds as a free radical chain reaction; by sequestering the radicals which
propagate the reaction, halons are able to "poison" the fire at much lower
concentrations than are required by fire suppressants using the more traditional
methods of cooling, oxygen deprivation, or fuel dilution.
For example, Halon 1301 total flooding systems are typically used at concentrations
no higher than 7% v/v in air, and can suppress many fires at 2.9% v/v. By contrast,
carbon dioxide fire suppression flood systems are operated from 34% concentration
by volume (surface-only combustion of liquid fuels) up to 75% (dust traps). Carbon
dioxide can cause severe distress at concentrations of 3 to 6%, and has caused
death by respiratory paralysis in a few minutes at 10% concentration. Halon 1301
causes only slight giddiness at its effective concentration of 5%, and even at 15%
persons remain conscious but impaired and suffer no long term effects.
(Experimental animals have also been exposed to 2% concentrations of Halon 1301
for 30 hours per week for 4 months, with no discernible health effects at all.) Halon
1211 also has low toxicity, although it is more toxic than Halon 1301, and thus
considered unsuitable for flooding systems.
However, Halon 1301 fire suppression is not completely non-toxic; very high
temperature flame, or contact with red-hot metal, can cause decomposition of Halon
1301 to toxic byproducts. The presence of such byproducts is readily detected
because they include hydrobromic acid and hydrofluoric acid, which are intensely
irritating. Halons are very effective on Class A (organic solids), B (flammable liquids
and gases) and C (electrical) fires, but they are totally unsuitable for Class D (metal)
fires, as they will not only produce toxic gas and fail to halt the fire, but in some cases
pose a risk of explosion. Halons can be used on Class K (kitchen oils and greases)
fires, but offer no advantages over specialised foams.
Halon 1211 is typically used in hand-held extinguishers, in which a stream of liquid
halon is directed at a smaller fire by a user. The stream evaporates under reduced
pressure, producing strong local cooling, as well as a high concentration of halon in
the immediate vicinity of the fire. In this mode, extinguishment is achieved by cooling
and oxygen deprivation at the core of the fire, as well as radical quenching over a
larger area. After fire suppression, the halon moves away with the surrounding air,
leaving no residue.
Halon 1301 is more usually employed in total flooding systems. In these systems,
banks of halon cylinders are kept pressurised to about 4 MPa (600 PSI) with
compressed nitrogen, and a fixed piping network leads to the protected enclosure.
On triggering, the entire measured contents of one or more cylinders are discharged
into the enclosure in a few seconds, through nozzles designed to ensure uniform
mixing throughout the room. The quantity dumped is pre-calculated to achieve the
desired concentration, typically 3-7% v/v. This level is maintained for some time,
typically with a minimum of ten (10) minutes and sometimes up to a twenty (20)
minute 'soak' time, to ensure all items have cooled so reignition is unlikely to occur,
then the air in the enclosure is purged, generally via a fixed purge system that is
activated by the proper authorities. During this time the enclosure may be entered by
persons wearing SCBA. (There exists a common myth that this is because halon is
highly toxic; in fact it is because it can cause giddiness and mildly impaired
perception, and also due to the risk of combustion byproducts.)
Flooding systems may be manually operated or automatically triggered by a VESDA
or other automatic detection system. In the latter case, a warning siren and strobe
lamp will first be activated for a few seconds to warn personnel to evacuate the area.
The rapid discharge of halon and consequent rapid cooling fills the air with fog, and is
accompanied by a loud, disorienting noise.
Due to environmental concerns, alternatives are being deployed. [2]
Halon 1301 is also used in the F-16 fighters to prevent the fuel vapors in the fuel
tanks from becoming explosive; when the aircraft enters area with the possibility of
unfriendly fire, Halon 1301 is injected into the fuel tanks for one-time use. Due to
environmental concerns, trifluoromethyl iodide (CF3I) is being considered as an
alternative. [3]

Environmental issues
Ozone-depleting gas trends

Since the late 1970s the use of CFCs has been heavily regulated because of its
destructive effects on the ozone layer. This damage was discovered by Sherry
Rowland and Mario Molina, who first published a paper suggesting the connection in
1974. It turns out that one of CFCs' most attractive features—their unreactivity—has
been instrumental in making them one of the most significant pollutants. CFCs' lack
of reactivity gives them a lifespan which can exceed 100 years in some cases. This
gives them time to diffuse into the upper stratosphere. Here, the sun's ultraviolet
radiation is strong enough to break off the chlorine atom, which on its own is a highly
reactive free radical. This catalyses the break up of ozone into oxygen by means of a
variety of mechanisms, of which the simplest is:

        Cl + O3 → ClO + O2
        ClO + O → Cl + O2

Since the chlorine is regenerated at the end of these reactions, a single Cl atom can
destroy many thousands of ozone molecules. Reaction schemes similar to this one
(but more complicated) are believed to be the cause of the ozone hole observed over
the poles and upper latitudes of the Earth. Decreases in stratospheric ozone may
lead to increases in skin cancer.
In 1975, the US state of Oregon enacted the world's first ban of CFCs (legislation
introduced by Walter F. Brown). The United States and several European countries
banned the use of CFC's in aerosol spray cans in 1978, but continued to use them in
refrigeration, foam blowing, and as solvents for cleaning electronic equipment. By
1985, scientists observed a dramatic seasonal depletion of the ozone layer over
Antarctica. International attention to CFCs resulted in a meeting of world diplomats in
Montreal in 1987. They forged a treaty, the Montreal Protocol, which called for drastic
reductions in the production of CFCs. On March 2, 1989, 12 European Community
nations agreed to ban the production of all CFCs by the end of the century. In 1990,
diplomats met in London and voted to significantly strengthen the Montreal Protocol
by calling for a complete elimination of CFCs by the year 2000. By the year 2010
CFCs should be completely eliminated from developing countries as well.
Because the only available CFC gases in countries adhering to the treaty is from
recycling, their prices have gone up considerably. A worldwide end to production
should also terminate the smuggling of this material, such as from Mexico to the
United States.
A number of substitutes for CFC's have been introduced. Hydrochlorofluorocarbons
(HCFCs) are much more reactive than CFC's, so a large fraction of the HCFCs
emitted break down in the troposphere, and hence are removed before they have a
chance to affect the ozone layer. Nevertheless, a significant fraction of the HCFCs do
break down in the stratosphere and they have contributed to more chlorine buildup
there than originally predicted. Development of non-chlorine based chemical
compounds as a substitute for CFCs and HCFCs continues. One such class are the
hydrofluorocarbons (HFCs), which contain only hydrogen and fluorine. One of these
compounds, HFC-134a, is now used in place of CFC-12 in automobile air
There is concern that halons are being broken down in the atmosphere to bromine,
which reacts with ozone, leading to depletion of the ozone layer (this is similar to the
case of chlorofluorocarbons such as freon). These issues are complicated: the kinds
of fires that require halon extinguishers to be put out will typically cause more
damage to the ozone layer than the halon itself, not to mention human and property
damage. However, fire extinguisher systems must be tested regularly, and these
tests may lead to damage. As a result, some regulatory measures have been taken,
and halons are being phased out in most of the world.
In the United States purchase and use of freon™ gases is regulated by the
Environmental Protection Agency, and substantial fines have been levied for their
careless venting. Also, licenses, good for life, are required to buy or use these
chemicals. The EPA website discusses these rules in great detail, and also lists
numerous private companies that are approved to give examinations for these
There are two kinds of licenses. Obtaining a "Section 609" license to use CFCs to
recharge old (pre-1993 model year) car air conditioners is fairly easy and requires
only an online multiple choice test offered by several companies. Companies that use
unlicensed technicians for CFC recharge operations are subject to a US$15,000 fine
per technician by the EPA.
The "Section 608" license, needed to recharge CFC-using stationary and non-
automobile mobile units, is also multiple choice but more difficult. A general
knowledge test is required, plus separate exams for small size (such as home
refrigerator) units, and for high and low pressure systems. These are respectively
called Parts I, II, and III. A person who takes and passes all tests receives a
"Universal" license; otherwise, one that is endorsed only for the respectively passed
Parts. While the general knowledge and Part I exams can be taken online, taking
them before a proctor (which has to be done for Parts II and III) lets the applicant
pass these tests with lower scores.

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