Chemical For Consumer Form 4

7/31/2019 Chemical For Consumer Form 4

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SULPHURIC ACID

Sulfuric acid (alternative spelling sulphuric acid) is a strong mineral acid with the molecular

formula H2SO4. Its historical name is vitriol. The salts of sulfuric acid are called sulfates. Sulfuric acid is

soluble in water at all concentrations.

Sulfuric acid has many applications, and is a central substance in the chemical industry. Principal uses

include lead-acid batteries for cars and other vehicles, ore processing, fertilizermanufacturing, oil

refining, wastewater processing, and chemical synthesis

Uses Of Sulphuric Acid

Sulfuric acid is one of the most important industrial chemicals. More of it is made each year

than is made of any other manufactured chemical; more than 40 million tons of it wereproduced in the United States in 1990. It has widely varied uses and plays some part in the

production of nearly all manufactured goods. The major use of sulfuric acid is in the

production of fertilizers, e.g., superphosphate of lime and ammonium sulfate. It is widely

used in the manufacture of chemicals, e.g., in making hydrochloric acid, nitric acid, sulfate

salts, synthetic detergents, dyes and pigments, explosives, and drugs. It is used in

petroleum refining to wash impurities out of gasoline and other refinery products. Sulfuric

acid is used in processing metals, e.g., in pickling (cleaning) iron and steel before plating

them with tin or zinc. Rayon is made with sulfuric acid. It serves as the electrolyte in the

lead-acid storage battery commonly used in motor vehicles (acid for this use, containing

about 33% H2SO4and with specific gravity about 1.25, is often called battery acid).

Sulphuric acid is used to produce chemical fertilizer such as ammonium sulphate and

potassium sulphate, which is highly soluble in water and can be easily absorbed by

plants.

Car batteries contain sulphuric acid which is used as the electrolyte.

Sulphuric acid also used in the making of artificial silk-like fibers and rayon.

Chemical like paints, dyes and drugs use sulphuric acid as one of their component

materials.

Sulphuric acid then reacts with sodium hydroxide to form sodium alkyl sulphonate,

which is detergent.

It is used by the manufacturers of iron and steel sheets to make iron and steel sheets free

from rust and oxidation before selling it to automobile industries.
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Uses Of

Sulphuric Acid

Sulphuric acid is used to produce

chemical fertilizer such as ammonium

sulphate and potassium sulphate, which

is highly soluble in water and can be

easily absorbed by plants.

Car batteries contain

sulphuric acid which is used

as the electrolyte

Sulphuric acid also used in

the making of artificial silk-

like fibers and rayon

Chemical like paints, dyes and

drugs use sulphuric acid as one

of their component materials.

Sulphuric acid then reacts with

sodium hydroxide to form sodium

alkyl sulphonate, which is detergent.

It is used by the manufacturers of iron and

steel sheets to make iron and steel sheets

free from rust and oxidation before selling it

to automobile industries.


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Manufacture Of Sulphuric Acid

Sulphuric acid is manufactured in industry though contact process. The process contain three

stage

The Contact Process:

makes sulphur dioxide;

convers the sulphur dioxide into sulphur trioxide (the reversible reaction at the

heart of the process);

converts the sulphur trioxide into concentrated sulphuric acid.

Making the sulphur dioxide

This can either be made by burning sulphur in an excess of air:

. . . or by heating sulphide ores like pyrite in an excess of air:

In either case, an excess of air is used so that the sulphur dioxide produced is alreadymixed with oxygen for the next stage

STAGE1: Production Of SulphurDioxide From Sulphur

i.Combustion of sulphur or sulphide ores in the air produce sulphur dioxide SO2.

S(s)+O2(g) SO2(g)

sulphur

ii.sulphur dioxide is dried and purified.

STAGE2: Production Of Sulphur Trioxide From Sulphur Dioxide

I .The purified sulphur dioxide SO2 and excess air are passed over vanadium(V)

oxide V2O5 at controlled optimum condition optimum condition to produce sulphur

trioxide SO3.

2SO2(g)+O2(g) 2SO3(g)

II .The optimum used are

a) Temperature:450-500C

b)Pressure: 2-3 atmospheres

c)Catalyst: Vanadium(V) oxide


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iii.Under controlled optimum conditions, 98% conversion is possible. Sulphur dioxideand

oxygen that have not reacted are allowed to flow back again over the catalyst inthe

converter.

This is a reversible reaction, and the formation of the sulphur trioxide isexothermic.

A flow scheme for this part of the process looks like this:


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STAGE3: Conversion of trioxide to sulphuric acid

This can't be done by simply adding water to the sulphur trioxide - the reaction is souncontrollable that it creates a fog of sulphuric acid. Instead, the sulphur trioxide is firstdissolved in concentrated sulphuric acid:

The product is known as fuming sulphuric acidor oleum.

This can then be reacted safely with water to produce concentrated sulphuric acid -twice as much as you originally used to make the fuming sulphuric acid.

iii.The addition of sulphur trioxide directly into is not carried out because the reaction is vary

vigorous; a lot of heat is given off. As a result, alarge cloud of sulphuric acidfumes is produced,

which is corrosive and causes severe air pollution.

The Contact Process Of Sulphuric Acid


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The Contac Process Of Sulphuric Acid


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Main Sources Of Sulphur Dioxide and The Environment Pollution

Sulphur dioxide is a yellowish, acrid gas. It is highly water-soluble and has a typical odour of

burnt matches. Sulphur dioxide is one of the main components of acid rain. The main

source of sulphur dioxide is probably coal-fired power generation, since this often uses low

grade coal, which may have a high level of sulphur in it. Sulphur dioxide is also producedduring the combustion of hydrogen sulphide, which may be present in biogas from animal

sources. Sulphur dioxide levels are regulated in most countries due to its aggressive nature

and effect on plant and animal life. Many of the sources of sulphur dioxide are not easily

avoided, so the main remedy is filtration to remove sulphur dioxide and other toxic

components.

Main Sources Of Sulphur Dioxide


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The Effect Of Sulphur Dioxide Towards Human Being

Sulfur dioxide (SO2) is a moderate to strong irritant. Most inhaled SO2 only penetrates as far

as the nose and throat with minimal amounts reaching the lungs unless the person isbreathing heavily, breathing only through the mouth or the concentration of SO2 is high.

Sensitivity varies among people, however, short exposure (1-6 hours) to concentrations as

low as 1 ppm may produce a reversible decrease in lung function. A 10 to 30 minuteexposure to concentrations as low as 5 ppm has produced constriction of the bronchiole

tubes. Only one of eleven volunteers showed any effects at 1 ppm. A 20-minute exposure to

8 ppm has produced reddening of the throat and mild nose and throat irritation. About 20ppm is objectionably irritating, although people have been reported to work inconcentrations exceeding 20 ppm. 500 ppm is so objectionable that a person cannot inhale

a single deep breath.

In severe cases where very high concentrations of SO2

have been produced in closed

spaces, SO2 has caused severe airways obstruction, hypoxemia (insufficient oxygenation ofthe blood), pulmonary edema (a life threatening accumulation of fluid in the lungs), anddeath in minutes. The effects of pulmonary edema include coughing and shortness of breath

which can be delayed until hours or days after the exposure. These symptoms areaggravated by physical exertion. As a result of severe exposures, permanent lung injurymay occur.

The Effects On Human

SKINS

The gas will react with moisture on the skin and cause irritation. Liquid SO2 may cause

burns due to freezing. Symptoms of mild frostbite include numbness, prickling and

itching in the affected area. The skin may become white or yellow. Blistering, necrosis

(dead skin) and gangrene may develop in severe cases.

EYES

Volunteers exposed to 5.4 ppm SO2 experienced mild irritation, while 9.1 ppm cause

moderate to severe irritation. At 8-12 ppm, smarting of the eyes and lachrymation

(tears) began. There is strong irritation at 50 ppm. In severe cases, (very high

concentrations in confined spaces), SO2 has caused temporary corneal burns. Liquid

SO2 can burn the eye and permanently affect vision. Injury from contact with liquid

SO2 may not be immediately noticed by the victim because SO2 damages the nerves of

the eye. Any eye contact should be treated as very serious.


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RESPIRATORY EFFECTS

Several human studies have shown that repeated exposure to low levels of SO2 (below5 ppm) has caused permanent pulmonary impairment. This effect is probably due to

repeated episodes of bronchoconstriction. One study has found a decrease in lung

function in smelter workers exposed for over 1 year to 1-2.5 ppm SO2. No effect was

seen in the same study in workers exposed to less than 1 ppm. In another study, a highincidence of respiratory symptoms was reported in workers exposed to 20-30 ppm for

an average of 4 years. Workers exposed to daily average values of 5 ppm SO 2 (withoccasional peaks of 53 ppm) had a much higher incidence of chronic bronchitis thancontrols.

CAUSES OF CANCER

Several epidemiological studies have examined the possibility that sulfur dioxide maycause cancers such as lung cancer, stomach cancer or brain tumours. In all of the

studies, there were uncontrolled confounding factors, such as concurrent exposure toother chemicals. The International Agency for Cancer (IARC) has reviewed these studiesand concluded there is inadequate evidence for carcinogenicity in humans. However,there is limited evidence of carcinogenicity in animals. Their overall evaluation is that

sulfur dioxide is not classifiable as to its carcinogenicity to humans (Group 3).

ACCULUMATION IN HUMAN BODY

SO2 may enter the body by the respiratory tract or following dilution in saliva. Most

studies in both man and animals have indicated that 40-90% or more of inhaled SO2is

absorbed in the moist upper respiratory tract. SO2 is quickly converted to sulfurous acidupon contact with moist mucous membranes. Inhaled SO2 is only slowly removed from

the respiratory tract. After absorption in the blood stream, the sulfurous acid is widely

distributed throughout the body, quickly converted to sulfite and bisulfite, which in turnis oxidized to sulfate and excreted in the urine.

REPRODUCTIVE EFFECTS

A number of epidemiological studies have suggested that exposure to SO2 may be

related to adverse reproductive effects. However, it is not clear that SO2 caused the

effects observed in any of these studies. There are no relevant results from animal

studies.In animal studies, no teratogenic effects were observed. However, slightfetotoxicity such as reduced birth weight and functional deficits have been reported at

doses which were probably toxic to the mother.


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ACID RAIN

There are many different kinds of acids. The proteins in our food, and in our bodies, are made upof amino acids. Motor cars start because of the sulphuric acid in their batteries. Swimming pools

need hydrochloric acid, commonly known as `pool acid'.

Some acids are weak, e.g acetic acid (vinegar) and lemon juice. They are not harmful and are

used in preparing our food. Othershowever, such as sulphuric acid (battery acid) are strong andcan burn holes in our clothes.

Formation Of Acid Rain

Acid rain is caused by the release of the gases SO2 (sulphur dioxide) and NOX (nitrous oxides).

The main sources of SO2 in South Africa are coal-fired power stations and metal workingindustries. The main sources of NOX emissions are vehicles and fuel combustion.

Sulphur dioxide reacts with water vapour and sunlight to form sulphuric acid. Likewise NOXform nitric acid in the air. These reactions takes hours, or even days, during which polluted air

may move hundreds of kilometres. Thus acid rain can fall far from the source of pollution.

When mist or fog droplets condense they will remove pollutants from the air and can become

more strongly acid than acid rain. Even snow can be acid. Gases and particles, not dissolved in

water, with a low pH can also be deposited directly onto soil, grass and leaves. It is possible that

even more acidity is deposited in this way than by rain! Not much is known about this process,

and it is particularly difficult to study.

The Effects Of Acid Rain

I. Acid rain can increase the acidity of lakes, dams and streams and cause the death ofaquatic life.

II. Acid rain can increase the acidity of soil, water and shallow groundwater.III. Acid rain has been linked with the death of trees in Europe and North America. In

spite of a great deal of research, no one yet knows exactly how acid rain harms

forests. Most of the forests of Europe consist of huge areas of one tree species. Thisencourages the spread of plant pests and diseases. It seems likely that acid rain

weakens the trees, perhaps helped by other pollutants such as ozone, and then leavesthe trees open to attack by disease. Acid rain also disrupts the availability of soilnutrients. The final death of a tree may result from a combination of stresses such as

heat, cold, drought, nutrient disruption and disease. It seems that the slow-growing,

longer lived forests of the North may be more susceptible than the faster growing,

shorter lived forests of South Africa.IV. Acid rain erodes buildings and monuments. Acid particles in the air are suspected of

contributing to respiratory problems in people.


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Acid rain causes a lot of damages.


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AMMONIAAmmonia is acompoundofnitrogenandhydrogen(and so is covalently bonded) with theformulaNH3. It

is a colourlessgaswith a characteristicpungentodour. Ammonia contributes significantly to

thenutritionalneeds of terrestrial organisms by serving as a precursor tofoodandfertilizers. Ammonia,either directly or indirectly, is also a building block for the synthesis of manypharmaceuticals. Although in

wide use, ammonia is bothcausticandhazardous. In 2006, worldwide production was estimated at 146.5

million tonnes.[5]

It is used in commercial cleaning products.

Ammonia, as used commercially, is often called anhydrous ammonia. This term emphasizes the absence

of water in the material. Because NH3 boils at -33.34 C (-28.012 F), the liquid must be stored under high

pressure or at low temperature. Itsheat of vapourizationis, however, sufficiently high so that NH3 can be

readily handled in ordinarybeakers, in afume hood(i.e., if it is already a liquid it will not boil readily).

"Household ammonia" or "ammonium hydroxide" is a solution of NH3 in water. The strength of such

solutions is measured in units ofbaume(density), with 26 degrees baume (about 30% w/w ammonia at

15.5 C) being the typical high concentration commercial product.[6]

Household ammonia ranges in

concentration from 5 to 10 weight percent ammonia
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Uses of Ammonia

Agricultural industriesare the major users of ammonia, representing nearly 80% of all

ammonia produced in the United States. Ammonia is a very valuable source of nitrogen that

is essential for plant growth. Depending on the particular crop being grown, up to 200 pounds

of ammonia per acre may be applied for each growing season.

Ammonia is used in the production of liquid fertilizer solutions which consist of ammonia,

ammonium nitrate, urea and aqua ammonia. It is also used by thefertilizer industryto

produce ammonium and nitrate salts.

Ammonia and urea are used as a source of protein in livestock feeds for ruminating animals

such as cattle, sheep and goats. Ammonia can also be used as a pre-harvest cotton defoliant,

an anti-fungal agent on certain fruits and as preservative for the storage of high-moisture

corn.

Dissociated ammonia is used in suchmetal treating operationsas nitriding, carbonitriding,

bright annealing, furnace brazing, sintering, sodium hydride descaling, atomic hydrogen

welding and other applications where protective atmospheres are required.

Ammonia is used in themanufactureof nitric acid; certain alkalies such as soda ash; dyes;

pharmaceuticals such as sulfa drugs, vitamins and cosmetics; synthetic textile fibers such as

nylon, rayon and acrylics; and for the manufacture of certain plastics such as phenolics and

polyurethanes.

The petroleum industry utilizes ammonia inneutralizing the acidconstituents of crude oil

and for protection of equipment from corrosion. Ammonia is used in the mining industry forextraction of metals such as copper, nickel and molybdenum from their ores.

Ammonia is used in several areas ofwater and wastewater treatment, such as pH control, in

solution form to regenerate weak anion exchange resins, in conjunction with chlorine to

produce potable water and as an oxygen scavenger in boiler water treatment.

Ammonia is used instack emission control systemsto neutralize sulfur oxides from

combustion of sulfur-containing fuels, as a method of NOx control in both catalytic and non-

catalytic applications and to enhance the efficiency of electrostatic precipitators for particulate

control.

Ammonia is used as the developing agent in photochemical processes such as white printing,

blue printing and in the diazo duplication process.

Ammonia is a widely used refrigerant inindustrial refrigeration systemsfound in the food,

beverage, petro-chemical and cold storage industries.
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Ammonia is used in therubber industryfor the stabilization of natural and synthetic latex to

prevent premature coagulation.

Thepulp and paper industryuses ammonia for pulping wood and as a casein dispersant in

the coating of paper.

Thefood and beverage industryuses ammonia as a source of nitrogen needed for yeast and

microorganisms.

The decomposition of ammonia serves as a source of hydrogen for some fuel cell and other

applications.

Ammonia is used by theleather industryas a curing agent, as a slime and mold preventative

in tanning liquors and as a protective agent for leathers and furs in storage.

Weak ammonia solutions are also widely used as commercial and household cleaners and

detergents.

The uses of ammonia.
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USES OF

AMMONIA

Agricultural

industries

Fertilizer

industry

Metal

treatingoperation

Manufacture

of nictric acid

Nuetralizing

of acid

Wateer and

wastewater

treatmentStack

emissioncontrol

system

Industrial

refrigeration

systems

Rubber

industry

Pulp and

paper

industry

Leather

industry

Householdcleaners and

detergents

Food and

beverage

industry


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Properties Of Ammonia

Ammonia is a colourless gas with a characteristic pungent smell. It is lighter than air, its density being

0.589 times that of air. It is easily liquefied due to the strong hydrogen bonding between molecules;

the liquid boils at 33.3 C, and freezes at 77.7 C to white crystals.[12]

The crystal symmetry is

cubic, Pearson symbol cP16, space group P213 No.198, lattice constant 0.5125 nm.[20]

Liquid ammonia

possesses strong ionising powers reflecting its highof 22. Liquid ammonia has a very high standard

enthalpy change of vapourization(23.35 kJ/mol, cf. water 40.65 kJ/mol, methane

8.19 kJ/mol, phosphine 14.6 kJ/mol) and can therefore be used in laboratories in non-insulated vessels

without additional refrigeration.

It is miscible with water. Ammonia in an aqueous solution can be expelled by boiling.

The aqueous solution of ammonia is basic. The maximum concentration of ammonia in water

(asaturated solution) has a density of 0.880 g/cm3

and is often known as '.880 Ammonia'. Ammonia does

not burn readily or sustain combustion, except under narrow fuel-to-air mixtures of 1525% air. When

mixed with oxygen, it burns with a pale yellowish-green flame. At high temperature and in the presence of

a suitable catalyst, ammonia is decomposed into its constituent elements. Ignition occurs when chlorine is

passed into ammonia, forming nitrogen and hydrogen chloride; if chlorine is present in excess, then the

highly explosive nitrogen trichloride (NCl3) is also formed.

The ammonia molecule readily undergoes nitrogen inversion at room temperature; a useful analogy is

an umbrella turning itself inside out in a strong wind. The energy barrier to this inversion is 24.7 kJ/mol,

and the resonance frequency is 23.79 GHz, corresponding to microwave radiation of a wavelength of

1.260 cm. The absorption at this frequency was the first microwave spectrum to be observed.[21]

Ammonia may be conveniently deodorized by reacting it with either sodium bicarbonate or acetic acid.

Both of these reactions form an odourless ammonium salt.
http://en.wikipedia.org/wiki/Gashttp://en.wikipedia.org/wiki/Lighter_than_airhttp://en.wikipedia.org/wiki/Earth%27s_atmospherehttp://en.wikipedia.org/wiki/Liquidhttp://en.wikipedia.org/wiki/Ammonia#cite_note-FOOTNOTEChisholm1911-11http://en.wikipedia.org/wiki/Ammonia#cite_note-FOOTNOTEChisholm1911-11http://en.wikipedia.org/wiki/Ammonia#cite_note-FOOTNOTEChisholm1911-11http://en.wikipedia.org/wiki/Pearson_symbolhttp://en.wikipedia.org/wiki/Space_grouphttp://en.wikipedia.org/wiki/Ammonia#cite_note-19http://en.wikipedia.org/wiki/Ammonia#cite_note-19http://en.wikipedia.org/wiki/Liquidhttp://en.wikipedia.org/wiki/Ionhttp://en.wikipedia.org/wiki/Dielectric_constanthttp://en.wikipedia.org/wiki/Dielectric_constanthttp://en.wikipedia.org/wiki/Dielectric_constanthttp://en.wikipedia.org/wiki/Enthalpy_of_vaporizationhttp://en.wikipedia.org/wiki/Enthalpy_of_vaporizationhttp://en.wikipedia.org/wiki/Joulehttp://en.wikipedia.org/wiki/Water_(molecule)http://en.wikipedia.org/wiki/Phosphinehttp://en.wikipedia.org/wiki/Misciblehttp://en.wikipedia.org/wiki/Waterhttp://en.wikipedia.org/wiki/Base_(chemistry)http://en.wikipedia.org/wiki/Saturation_(chemistry)http://en.wikipedia.org/wiki/Densityhttp://en.wikipedia.org/wiki/Combustionhttp://en.wikipedia.org/wiki/Oxygenhttp://en.wikipedia.org/wiki/Chlorinehttp://en.wikipedia.org/wiki/Hydrogen_chloridehttp://en.wikipedia.org/wiki/Nitrogen_trichloridehttp://en.wikipedia.org/wiki/Nitrogen_inversionhttp://en.wikipedia.org/wiki/Umbrellahttp://en.wikipedia.org/wiki/Resonance_frequencyhttp://en.wikipedia.org/wiki/Hertzhttp://en.wikipedia.org/wiki/Microwavehttp://en.wikipedia.org/wiki/Wavelengthhttp://en.wikipedia.org/wiki/Microwave_spectroscopyhttp://en.wikipedia.org/wiki/Ammonia#cite_note-Cleeton-20http://en.wikipedia.org/wiki/Ammonia#cite_note-Cleeton-20http://en.wikipedia.org/wiki/Ammonia#cite_note-Cleeton-20http://en.wikipedia.org/wiki/Ammonia#cite_note-Cleeton-20http://en.wikipedia.org/wiki/Microwave_spectroscopyhttp://en.wikipedia.org/wiki/Wavelengthhttp://en.wikipedia.org/wiki/Microwavehttp://en.wikipedia.org/wiki/Hertzhttp://en.wikipedia.org/wiki/Resonance_frequencyhttp://en.wikipedia.org/wiki/Umbrellahttp://en.wikipedia.org/wiki/Nitrogen_inversionhttp://en.wikipedia.org/wiki/Nitrogen_trichloridehttp://en.wikipedia.org/wiki/Hydrogen_chloridehttp://en.wikipedia.org/wiki/Chlorinehttp://en.wikipedia.org/wiki/Oxygenhttp://en.wikipedia.org/wiki/Combustionhttp://en.wikipedia.org/wiki/Densityhttp://en.wikipedia.org/wiki/Saturation_(chemistry)http://en.wikipedia.org/wiki/Base_(chemistry)http://en.wikipedia.org/wiki/Waterhttp://en.wikipedia.org/wiki/Misciblehttp://en.wikipedia.org/wiki/Phosphinehttp://en.wikipedia.org/wiki/Water_(molecule)http://en.wikipedia.org/wiki/Joulehttp://en.wikipedia.org/wiki/Enthalpy_of_vaporizationhttp://en.wikipedia.org/wiki/Enthalpy_of_vaporizationhttp://en.wikipedia.org/wiki/Dielectric_constanthttp://en.wikipedia.org/wiki/Ionhttp://en.wikipedia.org/wiki/Liquidhttp://en.wikipedia.org/wiki/Ammonia#cite_note-19http://en.wikipedia.org/wiki/Space_grouphttp://en.wikipedia.org/wiki/Pearson_symbolhttp://en.wikipedia.org/wiki/Ammonia#cite_note-FOOTNOTEChisholm1911-11http://en.wikipedia.org/wiki/Liquidhttp://en.wikipedia.org/wiki/Earth%27s_atmospherehttp://en.wikipedia.org/wiki/Lighter_than_airhttp://en.wikipedia.org/wiki/Gas


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17 NAZURAH ROSLAN

Properties

of

Ammonia

Colourless gas.

Have a pungent

smell.

Lighter than the air.

Boils at 33.3C

When mixed

with oxygen, it burnswith a pale

yellowish-green

flame.

Ammonia in an

aqueous solution canbe expelled by

boiling.

undergoes nitrogen

inversion at room

temperature

freezes at 77.7C to

white crystals

used in laboratori

in non-insulatedvessels without

additionalrefrigeration.


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18 NAZURAH ROSLAN

Manufacture Of Ammonia

The Haber Process is a method of producing ammonia developed in WWI. The

Germans needed nitrogen to for making their explosives. When the Allies blocked

off all trade routes going to and from Germany, they lost all source of sodium

nitrate and potassium nitrate, their source of nitrogen. They found their source

of nitrogen in the air, which was 80% nitrogen. The chemist Fritz Haber developed

the Haber Process in WWI, which takes molecular nitrogen from the air and

combines it with molecular hydrogen to form ammonia gas, which the chemical

formula is NH3.The equation for the reversible reaction is:

N2(g) + 3H2(g) 2NH3(g) + 92 kJ.

A flow scheme for the Haber Process looks like this:


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19 NAZURAH ROSLAN

AMMONIUM FERTILISER

The catalyst

The catalyst is actually slightly more complicated than pure iron. It has potassium

hydroxide added to it as a promoter - a substance that increases its efficiency.

The pressure

The pressure varies from one manufacturing plant to another, but is always high. Youcan't go far wrong in an exam quoting 200 atmospheres.

Recycling

At each pass of the gases through the reactor, only about 15% of the nitrogen andhydrogen converts to ammonia. (This figure also varies from plant to plant.) By continual

recycling of the non reacted nitrogen and hydrogen, the overall conversion is about98%.

Documents

Chemical For Consumer Form 4