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Introduction In-organic chemistry is they study of the elements of the periodic table and their compounds except for carbon but including its carbonates and oxides. In this report, the uses, importance and properties of various metals and non-metals will be explored in an attempt to better understand them and how they react. 1

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Page 1: Chem Project

Introduction

In-organic chemistry is they study of the elements of the periodic table and their compounds except for carbon but including its carbonates and oxides. In this report, the uses, importance and properties of various metals and non-metals will be explored in an attempt to better understand them and how they react.

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Metals and Non metals

Metals

All solids (except for mercury) Good conductors of electricity and heat Generally shiny when polished Many have a high density Are sonorous and malleable Many are ductile

Reactions of Metals

Many react with Oxygen: They form Basic oxides

Example: 2Cu (s) + O2 (g) –> 2CuO (s)

Metals react with acids: They form a salt and hydrogen

Example: Magnesium + Hydrochloric acid Magnesium Chloride + Hydrogen

Metals are reducing agents: We say that metals are reducing agents because they form positive ions and give out electrons.

Example: Mg Mg2+ + 2e-

Non-Metals

Do not conduct heat or electricity Are not ductile, sonorous or malleable They are mostly weak The do not conduct electricity (exceptions graphite)

Reactions of Non-Metals

Non-metals burn in Oxygen: They form acidic oxides

Example: Sulphur + Oxygen Sulphur dioxide

Non metals are oxidising agents: When they react they form negative ions and take an electron.

Example: S + 2e- S2-

Non-metals don’t react with dilute acids or water.

Non-Metallic structure

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Non metals can be either solid, liquid of gas. They combine with themselves to form molecules of various sizes.

Reactions of Metals with Oxygen

When metals react with oxygen they form basic oxides. The vigour of the reaction between metals and oxygen depends upon their place in the activity series. The higher up in the series, the more vigorous the reaction.

Metals such as potassium and calcium burn very brightly when reacted with oxygen, on the other hand metals such as lead and copper, which are low in the series only get coated with a small layer of metal oxide.

Reactions of metals with water

Like the reaction with oxygen, the vigour depends upon the position of the metal in the series. Metals such as potassium rush around the surface of getting smaller and smaller while a low metal in the series such as iron react with water very slowly.

Reactions of Metals with acids

The vigour of the reactions also depends on the metal’s position in the series, the higher in the series the more vigorous the reaction.

Reduction of Metal Oxides by another metal

Metals will reduce the oxides of any metals lower in the activity series. If the reducing metal is high in the activity series, then the reaction is very exothermic.

Reactions of this sort have been used to extract the metals chromium and manganese, by heating their ores with aluminium powder. This method is called the “Thermit process”.

Displacement reactions

All metals will displace a metal lower in the activity series from solutions.

Example: Magnesium + Copper sulphate Magnesium Sulphate + Copper

Stability of metal salts to heat

The position of metals in the activity series tell how stable some of its salts will be. The ease of decomposition increases down the activity series.

Metals3

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The least reactive elements occur in nature as free elements. The more reactive elements occur as compounds. These compounds, together with impurities, are called ores.

Extraction of Metals from their Ores

Extraction of metals form their ores is a reduction reaction.

Choice of extraction method

The extraction method is chosen based on the reactivity of the metal.

Very reactive metals: They require powerful methods of reduction and are extracted using electrolysis. The cathode is used as the reducing agent.

Less reactive metals: Require less powerful methods and can be extracted from their ores by reduction with carbon and carbon monoxide.

Least reactive metals: Can be extracted by heating ores in air at high temperatures.

Extraction of Copper

The main ore of Copper is CuFeS2.

1. The ore is crushed and the powder is added to a vat of oily, frothy water which is being stirred. This process is froth floatation

2. The purified ore is then dried and roasted in air in a furnace. 3. Lime stone is then added.4. The iron part of the ore reacts with silica.

FeO (s) + SiO2 (s) FeSiO3 (l)

5. The copper sulphate part of the ore is burnt to form sulphur dioxide and copper

CuS (s) + O2 (g) SO2 (g) + Cu (l)

6. The impure molten copper is run off into moulds and is purified by electrolysis.

Electrolysis of Iron and Steel

The main ores are:

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Magnetite Haematite Siderite

The extraction is carried out by the reduction of haematite by carbon monoxide in a blast furnace. (See figure a1)

1. A mixture of haematite, lime stone and coke, called charge, is placed into the top of the furnace.

2. Hot air is blown into the bottom of the furnace3. At stage 1- Coke burns in air forming carbon dioxide

C (s) + O2 (g) CO2 (g) Heat absorbed reduces the temperature to about 1100oC

4. At stage 2- more coke reduces carbon dioxide into Carbon monoxide.

C (s) + CO2 (g) 2CO (g) Heat absorbed reduces the temperature to about 1100oC

5. At stage 3- The carbon monoxide the haematite to iron

Fe2O3 (s) + 3CO (g) 2Fe (s) + 3CO2 (g)

The temperature here is around 700oC

6. The iron moves down the furnace where it melts (MP 1535 oC), the molten iron runs to the bottom off the furnace.

7. Impurities in the ore, mainly silica have to be removed. At temperatures above 850 oC( between stage 2 and 3) the limestone decomposes forming calcium oxide and more Carbon dioxide

8. The silica, being acidic, reacts with the calcium oxide to form molten calcium sulphate or slag. The slag runs to the bottom of the furnace where it floats on the iron.

CaO (s) + SiO2 (s) CaSiO3 (l)

9. The slag and Iron are tapped Off Separately: The iron, called pig iron or cast iron, is impure. It is further purified by blowing

air through it to oxidise the impurities. This produces wrought Iron, most of which is converted to steal by adding calculated quantities of Carbon and other Substances.

The slag is used for road building and fertilisers.

10. Waste gasses are removed from the top of the furnace and burnt to heat air blown at the bottom.

Extraction of Aluminium

Ores:

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Bauxite Cryolite

The extraction of aluminium is carried out by electrolysis in 3 stages:

(See figure a2)

1. The bauxite is Purified2. The purified bauxite is dissolved in molten cryolite at 900° C to separate the

ions. The cryolite is used to lower the Melting, point. 3. The bauxite/cryolite solution is electrolysed using graphite electrodes. At the cathode, Molten Aluminium is collected at the bottom and is tapped

off.

Al3+ + 2e- Al (l)

At the anode, Oxygen gas is evolved.

2O2- O2 (g) + 4e-

Extraction of Sodium

(See figure a3)

Sodium is extracted by the electrolysis of molten sodium chloride

Inside is a graphite anode, which is surrounded by molten sodium chloride.

At the cathode: Na+ + e- Na (l) (reduction)

At the anode: 2 Cl- 2Cl (g) + 2e-(oxidation)

Alloys

An alloy is a mixture of 2 or more metals, a few also contain non-metals. In general, alloys are harder, stronger and more corrosion resistant than the pure metals. Because of this they are usually used in place of pure metals.

Table 1: Some common alloys, their uses and composition

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Alloy Composition Use PropertiesDuralumin 95% Al, 4% Cu,

traces of Mg and Mn

Aircraft construction

Low density- as light as Al but more stronger and corrosion resistant.

Magnalium 70% Al, 30% Mg Same as Duralumin Same as Duralumin Steel Fe and about 1&

carbonGirders, Bridges, etc

Much stronger than iron

Solder 67% Pb, 33% Sn Joining Metals Low melting pointBrass 60% Cun 40% Zn Electrical

connections and machine bearings

Harder and stronger than copper

Bronze 90% Cu, 10% tin Machine parts and electrical Circuits etc.

Harder and stronger then copper

Constantan 75% Cu, 25% Ni For ‘silver’ coins Cheaper and harder than silver.

Corrosion

Corrosion occurs when the surface of a solid, usually a metal, is eaten away by the action of oxygen, moisture and pollutants present in the atmosphere. When a metal corrodes, it is covered with a layer of its oxide. Generally the more reactive a metal, the faster it corrodes.

Corrosion of Iron and steel – rusting

On exposure to the atmosphere, iron and steel corrode to from rust (hydrated Iron (III) oxide) Rust gradually ‘eats’ into the metal.

Methods of preventing rust

1. Coating the object with Paint, grease, plastic or rubber: this prevents contact with air

2. Galvanising: Coating the object with zinc, zinc forms an oxide layer which protects against further corrosion.

3. Tinning: Coating the object with tin4. Sacrificial anode: A more reactive metal is oxidised in preference to iron.

Metals and life

Iron forms a part of the haemoglobin molecule in red blood cells, which is essential to carry oxygen.

Magnesium forms a part of the chlorophyll molecule Calcium compounds are constituents of bones and teeth Sodium and potassium ions are essential for transmission of nerve impulse Many other elements are required in minute quantities by plants and animals.

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Pollution by Metals

Lead in compounds for car exhausts affect enzymes, it causes bran damage and even death.

Mercury in compounds from industries cause nerve disorder and even death Tributyl tin n antifouling paints harm shellfish Many others metals can become harmful in large amounts

Analysis of metallic cations

A number of laboratory tests can be performed to identify a given cations. This is called quantitive analysis and can also be applied to ions. The tests performed can be either indicative tests or confirmatory tests.

Indicative tests: points to a particular conclusion but does not prove the presence of the unknown substance

Confirmatory test: Gives proof of the presence of the unknown substance.

Non- metals

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Water

Water is one of the most important compounds derived from non-metallic elements.

The water Cycle

(See figure a4 for diagram of water cycle)This cycle is made up of a few main parts:

evaporation (and transpiration) condensation precipitation collection

Evaporation: Evaporation is when the sun heats up water in rivers or lakes or the ocean and turns it into vapour or steam. The water vapour or steam leaves the river, lake or ocean and goes into the air.

Condensation: Water vapour in the air gets cold and changes back into liquid, forming clouds. This is called condensation. Precipitation: 

Precipitation: Precipitation occurs when so much water has condensed that the air cannot hold it anymore.  The clouds get heavy and water falls back to the earth in the form of rain, hail, sleet or snow.

Collection:  When water falls back to earth as precipitation, it may fall back in the oceans, lakes or rivers or it may end up on land.  Physical Properties:

tasteless and odourless liquid

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Freezing point: 0oC Boiling point: 100oC When cooled water contracts(but when 4oC is reached it begins to expand

until 0oC)

Chemical Properties:

Water acts as a catalyst: nearly all chemical reactions need water to start. Although water acts as a medium in which chemical reactions take place, it

can also react chemically with another substance. Water has remarkable solvent properties and can dissolve a great number of

substances.

Uses of water

Used for drinking Washing Used in hydro electricity Is essential in industry Used for recreation

See figure a8 for a water works

See figure a9 for a sewage treatments works

How soap helps in washing

Soap is a compound called sodium stearate. At one end is a group of atoms which dissolve in water and on the other end the atoms do not. Soap helps by breaking down the surface tension of pure water. When soap is added, the water soluble end of the molecule dissolves in the water, destroying he surface tension, helping the water to spread.

Hardness of water

When soap dissolves in water it forms lather. If the water is hard then the soap is destroyed. Hardness of water is caused by calcium and magnesium compounds in

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the water. They are: Calcium hydrogen Carbonate, magnesium hydrogen Carbonate, calcium sulphate and magnesium sulphate.

There are two types of hardness, temporary hardness and permanent hardness.

Temporary hardness

This is caused by calcium hydrogen carbonate in the water. When it rains, it reacts with carbon dioxide to form tiny amounts of Carbonic acid.

CO2 + H2O H2CO3

If the acid falls on limestone rock, it slowly dissolves them:

Carbonic acid + Calcium Carbonate Calcium hydrogen carbonate.

This is how the hardness gets into the water. Magnesium Hydrogen Carbonate causes hardness o water in a similar manner. It can be removed by boiling.

Permanent Hardness

This is caused when amounts of calcium and magnesium sulphate dissolve in water. It cannot be removed by boiling but it destroys soap.

Ways of removing hardness of water

Boiling Distillation Addition of washing soda Use of detergents Addition of calcium hydroxide Commercial water softeners

Respiration

Air is a mixture of non metallic compounds. Oxygen, which is a part of air, is vital for breathing. Cellular respiration is the process in which the chemical bonds of energy-rich molecules such as glucose are converted into energy usable for life processes. The overall equation for the oxidation of glucose is: C6H12O6 + 6O2 → 6CO2 + 6H2O + energy. (See figure a5 for diagram of alveoli)

Photosynthesis is the process by which plants use the energy from sunlight to produce sugar. We can write the overall reaction of this process as: 6H2O + 6CO2 ----------> C6H12O6+ 6O2

The composition of air

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The percentage of oxygen in the air

(See figure a6 for apparatus used)

1. A Pyrex tube is packed with copper wire. This copper is special in that it is porous and it has an uneven surface and therefore a large surface area.

2. One syringe is filled with air and the other is empty. The tube is heated at one end and air is pushed to the other syringe. As it passes, the oxygen in the air reacts with the copper forming copper oxide.

3. As the copper goes from pink to black the Bunsen burner is moved to a fresh portion of copper. Eventually all the oxygen will react.

4. The apparatus is left to cool down and the volume of air left is noted.

A typical result would be:

Volume at beginning: 100 cm3

Volume at end: 79 cm3

Volume of air used: 21 cm3

Therefore the percentage of Oxygen present would be 21%Separating Air

Air is separated when it is liquefied by fractional distillation.

1. Carbon dioxide and water vapour are removed first; this is done by cooling the air in a refrigeration plant because they are frozen easily.

2. The remaining gases are compressed to about 150 times atm. As this is done it gets very hot. It is then allowed to cool off. It is then allowed to expand again.

3. This continues until the temperature is as low as 73 K. The liquefied gases can then be removed by frictional distillation.

Uses of separated gasses

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Oxygen: is used in burners, It is used in the removal of impurities in the extraction of some metals. It is also used to support breathing.

Nitrogen: This gas is used in liquid form for freezing food. It is also an important ingredient in the production of ammonia.

Pollution of air

Major primary pollutants produced by human activity include:

Sulphur oxides (SOx) - especially sulfur dioxide, a chemical compound with the formula SO2. SO2 is produced by volcanoes and in various industrial processes. This is one of the causes for concern over the environmental impact of the use of these fuels as power sources.

Nitrogen oxides (NOx) - especially nitrogen dioxide are emitted from high temperature combustion, and are also produced naturally during thunderstorms by electrical discharge.

Carbon monoxide (CO) - is a colourless, odorless, non-irritating but very poisonous gas. It is a product by incomplete combustion of fuel such as natural gas, coal or wood. Vehicular exhaust is a major source of carbon monoxide.

Carbon dioxide (CO2) - a colourless, odorless, non-toxic greenhouse gas also associated with ocean acidification, emitted from sources such as combustion, cement production, and respiration.

Combustion

When things burin in air, they combine with oxygen and give out hot gasses. These are usually seen as flames.

Most substances that burn contains carbon and hydrogen, so steam and water vapour are the two main gasses produced. If air supply is limited, complete combustion may not occur and CO may be formed instead of Carbon dioxide.

For burning, three things are needed: Oxygen, heat and fuel.

Because all three are essential, to stop burning, there are three things to do:

1. Cut off the fuel2. Lower the temperature3. Cut off the air supply.

Hydrogen (H)

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Figure 1: Hydrogen Atom Figure 2: Hydrogen Molecule

General:

First Element in the periodic table and has an Atomic number of 1. One proton, one electron and no neutron. Hydrogen is a diatomic element and forms hydrogen gas (figure 2) H2. Hydrogen has the oxidation states +1 and -1. Atomic Mass: 1.0079 Electronic Configuration is 1 Atomic number 1

Physical Properties of Hydrogen:

1) Hydrogen is odourless, tasteless and colourless. 2) Virtually insoluble in water 3) It is the lightest of all gasses.4) H molecules move more quickly than air molecules.

Chemical properties:

1) Hydrogen burns in air.

The hydrogen burns with a blue flame and produces steam.

2H2 (g) + O2 (g) 2H2O (g)

2) Hydrogen is a good reducing agent

Hydrogen removes Oxygen from metal oxides.

Industrial Sources of Hydrogen:

1) Natural Gas(Steam Reforming)

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In this process Methane (CH4) is mixed with steam and passed over a nickel catalyst and high temperature (as high as 1000oC) and pressure (may be greater than 50 atmospheres). This results in the formation of Hydrogen and carbon monoxide.

CH4 (g) + H2O (g) 3H2 + CO (g)

In the second step, the CO is mixed with more steam under an iron oxide catalyst to form CO2 and more H2.

CO (g) + H2O (g) H2 + CO2 (g)

The Carbon Dioxide is then removed.

2) The Bosch process

This process involves an endothermic reaction between Steam and white hot coke.

C (S) + H2O (g) CO (g) + H2 (g)

3) As a by Product

Uses:

1) Synthesis of Ammonia

N2 (g) + 3H2 (g) 2NH3 (g)

The temperature of this reaction is 4500C and the pressure is 200atm. An iron catalyst is used.

2) Used to manufacture margarine, Methanol etc.

Laboratory Preparations of Hydrogen:

Hydrogen is usually produced when metals react with dilute Hydrochloric acid. (See figures a7 for apparatus used)

Metal + Acid Salt + Hydrogen

Example: HCl (aq) + Zn (s) H2 (g) + ZnCl2 (aq)

The Hydrogen can then be collected over water. To produce dry Hydrogen, The hydrogen is bubbled into the dehydrating agent sulphuric acid and then can be collected upwards into a gas Jar.

Carbon (C)

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Figure 1: Carbon Atom

General:

Carbon is the 6th element in the periodic table Carbon has the oxidation states -4, +2, and +4 Has an atomic mass of 12.011 Electronic configuration is 2-4 Atomic Number 6

Properties:

Insoluble in common solvents Combines with oxygen to form Carbon dioxide Reduces Metal Oxides to the Metal

Example: C (s) + PbO CO (g) + Pb (s)

Reduces Steam Is oxidised by hot sulphuric and Nitric acid

Example: C (s) + H2SO4 (aq) 2H2O + 4NO2 (g) + CO2 (g)

Allotropes

Carbon exists on earth in several forms. It is present in its pure form or in its combined states. The two pure forms of carbon are Diamond and Graphite. These are allotropes of carbon. Even though both graphite and diamond are pure forms of carbon they both have very different properties, this is due to how the atoms arrange themselves.

Diamond Graphite

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Properties of Graphite and Diamond

Graphite DiamondSoft and greasy to touch, The flat sheets of the graphite slide over one another.

Hardest known substance, This is because the tetrahedral structure.

Graphite is black; this is because the layers prevent light from being transmitted.

Colourless, this is because the light that enters diamond is reflected from face to face before going out.

Graphite conducts electricity; this is because it had one free electron which is not used in the bonding process.

Diamond does not conduct electricity because all electrons are used in the bonding process.

Graphite burns slowly in air and forms Coke.

Diamond will burn but because of its strong structure it needs a more energy.

The carbon Cycle

(See figure b1 for diagram of carbon cycle )

1) All green plants take in CO2 in a process called photosynthesis and is used to make carbohydrates.

2) Animals eat plants and take the carbon that the plants have.3) Animals give out CO2 as they breath4) CO2 is also given out when bacteria and fungi respire and also by factories

and houses.

Carbon Dioxide

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Preparation of Carbon Dioxide

Carbon Dioxide can be prepared be either:

1) The action of heat on Carbonates

Carbonate + heat Oxide + CO2

2) The action of heat on Hydrogen Carbonates

Hydrogen Carbonate + heat Carbonate + Water + CO2

3) Action of dilute acid on Carbonates

Acid + Carbonate Salt + Water + CO2

(The apparatus used can be seen in figure b2)

Properties of Carbon Dioxide

Colourless, odourless and tasteless It is heavier than air It is slightly soluble in water It forms dry ice when cooled and pressured through sublimation

Reactions of Carbon Dioxide

Carbon Dioxide does not support combustion Carbon Dioxide is an acidic oxide and forms acids when added to water Carbon Dioxide can be reduced by burning magnesium

Carbon Monoxide

Preparation of Carbon Monoxide

The preparation of carbon monoxide is similar to that of Carbon Dioxide. After the carbon Dioxide is produced, it is then reduced by carbon at 10000C to produce Carbon monoxide (CO2 (g) + C (s) 2CO (g)). The resulting gas is then bubbled through Potassium Hydroxide to remove any unreacted CO2. See figure ()

Properties of Carbon Monoxide

It is Colourless, tasteless and odourless It is insoluble in water It is very Poisonous

Reaction of carbon Monoxide

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Burns in air to form Carbon dioxide Is a reducing agent

Manufacture of Sodium carbonate

(See figure b3 for a flowchart of the process)

1) Limestone (CaCO3) is heated to produce Calcium Oxide and Carbon Dioxide.

CaCO3 + Heat CaO + CO2

2) The CaO is then slaked with water to form Calcium Hydroxide.

CaO + H2O Ca (OH) 2

3) Brine is then mixed with water and ammonia gas to for ammoniated Brine.

NH3 (g) + NaCl (aq) + H2O (l) NH4+

(aq) + Cl- (aq) + Na+

(aq) + OH- (aq)

4) The ammoniated Brine is then mixed with Carbon Dioxide to form NH4Cl and NaHCO3. The NahCO3 is the precipitated out of the solution.

NH4+

(aq) + Cl- (aq) + Na+

(aq) + OH- (aq) + CO2 (g) NH4Cl (aq) + NaHCO3 (s)

5) The Sodium Hydrogen Carbonate is then heated to form sodium carbonate Water and CO2.

2NaHCO3 (s) Na2CO3 (s) + CO2 (g) + H2O (l)

6) The let over ammonium Chloride is mixed with Calcium Hydroxide to fore more ammonia for step 3.

Uses of Carbon Compounds

Used it tools Used in metal extraction as a reducing agent Used in fire extinguishers Used to manufacture glass, water-glass, borax and so on

Silicon (Si)19

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Figure 1: Silicon Atom

Silicon is the second most abundant element on earth.

Silicon is present in the combined state as metallic silicates, silicon dioxide flint and quartz.

Silicon is extracted from Silicon Dioxide using carbon (SiO2 (s) + 2C (s) Si (s) + 2CO (g))

General

Oxidation states -4, +2 and +4 Atomic number 14 Atomic mass 28.086 Electronic configuration is 2-8-4

Physical Properties of Silicon

Is grey-Black Is a brittle solid Has a high Melting point Does not exhibit allotropy

Chemical Properties of Silicon

Fairly Inert at Room Temperature Not attacked by acids except for Hydrofluoric acid Burns in air to form Silicon Dioxide When Powdered combines with chlorine on heating and also decomposes

steam Attacked by concentrated or molten alkalis in the powdered state.

Important Substances

Glass: Made from Silicon Dioxide and Soda-ash

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Silicon Tetrachloride Silicic Acids: Give rise to Silicates Silicates: Salts which are used for many different purposes, example: the

production of silica gel which is used as a drying agent. Silicones: are Polymers containing Silicon-oxide. These have many different

properties. Ceramics: have great durability which is good for making many different

items.

Uses of Silicon

To manufacture glass, concrete and other substances. Used in heat resistant paint and varnishes Used as an abrasive Substances of silicon can be used to control radio frequencies and electronic

Equipment

Nitrogen (N)

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Figure 1: Nitrogen Atom

General

Constitutes for 78.08% of the air in the atmosphere by volume Has a large Varity of oxidation states Atomic Mass is 14.007 Has an electronic configuration of 2-5 Atomic number 7

Physical Properties

Colourless, odourless and tasteless Melting point is -210oC Boling Point is -196oC

Chemical Properties

Chemically Inert Produces Nitrogen (II) oxide when sparked with oxygen Produces ammonia when under pressure with Hydrogen under a iron catalyst Reacts with certain metals at high temperatures

Uses of Nitrogen

Used to provide an inert atmosphere for easily oxidisable substances. Example: in food Packaging, chips and so on.

As a cooling agent when liquid To manufacture ammonia and Nitric acid

The Nitrogen Cycle

(See figure b4 for diagram of the nitrogen cycle)

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Nitrogen fixation: In the soil, as well as in the root nodules of certain plants, nitrogen is "fixed" by bacteria, lightning, and ultraviolet radiation. The "fixing of nitrogen" does not mean nitrogen was broken; a better term might be "fixated," because the bacteria put elemental nitrogen into a form that can be used by living organisms and do not allow it to leave that form and revert to elemental nitrogen.

Nitrification: Certain bacteria take the forms into which nitrogen was fixated and further process it (oxidization). Oxidation provides energy for the nitrogen cycle to take place — the bacteria that live in soil cannot harness energy from the sun. The energy they use during their work in the nitrogen cycle comes from this process.

Denitrification and ammonification. Plants absorb nitrates or ammonium ions from the soil and turn them into organic compounds. Animals obtain nitrogen by consuming plants or other animals. Therefore, the waste products of animals contain nitrogen. Ammonium ions, ammonia, urea, and uric acid all contain nitrogen. So regardless of what form of excretion an animal has, some nitrogen is released back into the ecosystem through excrement. Dead plants and animals are food for decomposing bacteria.

Ammonia

Ammonia is an important compound of Nitrogen and has many different uses.

Manufacture of Ammonia

Ammonia is manufactured by the harbour process. (See figure b5 for a flow chart)

1) Hydrogen is made by steam reforming and then is burned in air to produce Nitrogen.

2) Hydrogen and nitrogen are then passed over an iron catalyst and compressed at about 350 times atm and heated to about 450oC. The reaction is an equilibrium reaction and these conditions help to produce as much product as possible.

3) The ammonia formed is then condensed out and stored as liquid ammonia or as a solution.

Laboratory Preparation of ammonia

Ammonia can be made when ever an ammonium salt reacts with any alkali.

(NH4+ + OH- NH3 (g) + H2O (g)) (See Figure b6 for laboratory preparation of

ammonia)

Physical Properties of ammonia

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Ammonia is colourless and has a choking smell It is lighter than air It is alkaline and turns damp red litmus blue It is extremely soluble in water.

Chemical properties of Ammonia

Ammonia is a reducing agent Ammonia reacts with Hydrogen chloride gas to produce white ammonium

chloride. This is a reversible reaction. Ammonia dissolves in water and forms an alkali called ammonium hydroxide.

Reactions of Ammonium Hydroxide

“Undissolving” ammonium Hydroxide, a concentrated solution of ammonium Hydroxide smells strongly, this is because the ammonia gas is undissolving from the solution.

NH4OH (aq) NH3 (g) + H2O (l)

Ammonium Hydroxide reacts with acids, it forms an ammonium salt and water.

NH4OH (aq) + Acid Ammonium Salt + water

Preparation of insoluble metal hydroxides: Ammonium Hydroxide can be used to make precipitates of insoluble metal hydroxides.

Example: 2NH4OH (aq) + FeCL2 (aq) Fe (OH)2 (S) + 2NH4CL (aq)

These reactions are called precipitation reactions

Nitric Acid

Nitric acid is an extremely important compound and is used for many different reasons such as the production of fertilizers, explosives and so on.

Production of Nitric Acid

See figure b7 for flowchart in the production of nitric acid

1. Ammonia gas and air are first cleaned, mixed and compressed to about 4 atm.

2. They are then passed over a platinum (90%)/ Rhodium (10%) gauze catalyst and heated to about 900oC. (4NH3 (g) + 7O2 (g) 4NO2 (g) + 6H2O (g))

3. The resulting Nitrogen (IV) oxide is cooled and mixed with more air and water to form 65% nitric acid. (4NO2 (g) + O2 (g) + 2H2O (l) 4HNO3 (aq))

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Laboratory Preparation

Nitric acid is produced whenever concentrated sulphuric acid is heated with a nitrate (NO3

- (s) + H2SO2 (l) HSO4

- (aq) + HNO3 (g)) See figure b8 for apparatus used

Example: NaNO3 (s) + H2SO4 (l) NaHSO4 (aq) + HNO3 (g)

Chemical Reactions of Nitric Acid

Concentrated nitric acid is a powerful oxidising agent and will oxidise organic compounds and metals. It reduced metals to salts, water and nitrogen (IV) oxide.

Dilute nitric acid shows typical acidic properties. Is a monobasic acid and forms only normal salts

Nitrates

Reactions of Nitrates

Warming with concentrated sulphuric acid makes nitric acid vapour Nitrates decompose when heated

The extent to which they decompose depends on the reactivity of the nitrate. The more reactive the metal the harder it is to decompose and the less reactive metals decompose easily.

Oxides of Nitrates

Nitrogen (l) oxide (N2O) Nitrous oxide: This has a sweet smell and like oxygen, relights a glowing splint

Nitrogen (ll) oxide (NO) Nitric Oxide: Is colourless, reacts with oxygen to form Nitrogen dioxide. It is formed whenever Nitrogen reacts with oxygen due to a hot spark (example lightening).

Nitrogen (lV) oxide (NO2) nitrogen dioxide: It is a heavy brown gas. It forms when nitrates are heated or when nitric acid reacts with a metal.

Test for nitrates

Nitrates are tested by using the brown ring test. In this test the suspected solution is mixed with Iron (II) sulphate and then concentrated sulphuric acid. A brown ring then forms which proves that the solution is a nitrate.

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Phosphorus (P)

Figure 1: Phosphorus atom

General

Has oxidation states +3, -3 and +5 Has an atomic mass of 30.974 Electronic configuration is 2-8-5 Atomic number 15

Extraction of Phosphorus

See figure b9 for diagram on extraction

An electric furnace is fed with crushed rock phosphate, sand and coke. At the high temperature (1400oC) Produced in the furnace the sand combines with the calcium phosphate displacing the more volatile Phosphorus oxide. The latter is reduced by the carbon to phosphate.

Allotropes of Phosphorus

Phosphorus exists in three allotropes forms, white, Red and black phosphorus.

White phosphorus is a highly reactive and very dangerous. It is simultaneously inflammable in air. It is unstable and slowly changes to Red phosphorus.

Red Phosphorus It has a macromolecular structure which makes it more stable than white phosphorus.

Black Phosphorus is the most stable and most difficult to prepare of the allotropes. It is a flaky macromolecular solid which conducts electricity.

See figure c1 for structure of allotropes

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Table 1: Comparison of the properties of red and white phosphorus

Red Phosphorus White Phosphorus 1 Appearance Opaque red, Brittle

powderColourless transparent waxy solid which turns yellow

2 Toxicity Not poisonous poisonous3 Heat in vacuo Sublimes at 416oC Melts at 44 oC4 Density 2.3 g cm-3 1.8 g cm-3

5 Exposure to air No oxidation Emits a green phosphorescence

6 Solubility in organic solvents

insoluble Soluble

7 Hot Concentrated Sodium Hydroxide

NO reaction Reacts to produce phosphine

8 Heat in air Ignites at 260 oC Ignites at 30 oC if moist

Laboratory Preparation of Phosphorus Chloride

See figure c2 for apparatus used

1. Dry Chlorine is generated and passed through a flask submerged in a freezing mixture.

2. Phosphorus (lll) Chloride is dropped from a funnel into the flask.

This preparation must be done in a fume cupboard because chlorine is poisonous.

Compounds of Phosphorus

Oxides

There are Phosphorus (III) oxide (P4O6) and Phosphorus (V) oxide (P4O10). Both of them are acidic oxides and form acids when they react with water.

Phosphorus (V) Oxide is also a good dehydrating agent.

When Phosphorus (III) oxide is heated in air Phosphorus (V) oxide is formed.

Hydrides

The most important Phosphorus hydride is Phosphorus (III) Hydride. It is colourless, it’s an extremely poisonous gas and its boiling point is -87.7oC. It is only slightly soluble in water and it is a stronger reducing agent that Ammonia.

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Oxyacids

There are six different oxyacids of Phosphorus. Each of these acids can form three series of salts.

Phosphorus acid

Manufacture of Orthophosphorus (V) acid

Orthophosphorus (V) acid is manufactured by the burning o white phosphorus in air and dissolving the product in hot water.

P4 (s) + 5O2 (g) P4O10 (s)

P4O10 (s) + 6H2O (l) 4H3PO4 (aq)

It is prepared in the laboratory by evaporating red phosphorus with concentrated nitric acid, the syrupy liquid formed is evaporated over concentrated sulphuric acid, crystals of Orthophosphorus (V) acid are then formed.

Orthophosphorus (III) acid

It is formed by carefully evaporating and cooling solutions of phosphorus (III) oxide or phosphorus (III) chloride.

It decomposed in heat to form (H3PO3 (aq)) and (PH3 (g)) and is a good reducing agent.

Uses of phosphorus

Used to manufacture Orthophosphorus (V) acid Phosphorus compounds are used in “safety matches” to cause the match to

alight Certain insecticides contain phosphorus compounds

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Oxygen (O)

Figure1: Oxygen Atom Figure 2: Oxygen Molecule

General

Atomic number: 8 Electronic configuration: 2-6 Atomic mass: 15.999 Has many oxidation states

Physical Properties

Colourless, tasteless and odourless Slightly soluble in water Slightly denser than air

Chemical Properties

Oxygen is a neutral gas It is a powerful oxidising agent Oxygen supports combustion Most metals and non metals combine with oxygen to form oxides Oxygen causes rusting Oxygen is essential in aerobic respiration

Laboratory Preparation of Oxygen

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Oxygen can be prepared by:

1. Heating lead (IV) oxide. (2PbO2 (s) 2PbO (s) + O2 (g) )2. Heating dilead (II) lead (IV) oxide (red lead) (2Pb3O4 (s) 6PbO (s) + O2 (g))3. Heating potassium nitrate (2KNO3 (s) 2KNO2 (s) + O2 (g))4. The decomposition of Hydrogen peroxide using manganese (IV) oxide as a

catalyst. (2H2O2 (l) 2H2O (l) + O2 (g))

Oxides

1. Acidic Oxide

These are oxides of non metals which react with water to form acids and react with alkalis forming salt and water. Example sulphur dioxide

SO2 (g) + H2O (l) H2SO3 (aq)

2. Basic oxides

These are oxides of metals which react with acids forming salt and water. Example copper oxide

CuO (s) + H2SO4 (aq) CuSO4 (aq) + H2O (l)

3. Amphoteric Oxide

These are oxides which show properties of both basic and acidic oxides. Example lead oxide

PbO as a basic oxide (PbO (s) + 2HNO3 (aq) + H2O (l))

PbO as an acidic oxide (PbO (s) + 2NaOH (aq) Na2PbO2 + H2O (l))

4. Neutral Oxide

These are oxides of certain non-metals which react with neither an acid nor a base. Example carbon monoxide

5. Peroxides

These are oxides of reactive metals which contain the O2-2 ion. They contain more

oxygen that basic oxides.

Reaction of metals with oxygen

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Metals

potassium when heated, it melts very easily and burns with a lilac flame to leave a white powder: K(s) +O2(g)=KO2(s)

sodium When heated it melts at a slightly higher temperature, and burns vigorously with a yellow flame to leave a pale yellow solid.

2Na(s) +

O2(g)=Na2O2(s)

Calcium When heated it does not melt, but it burns with a brick red flame, to leave a white solid. 2Ca(s) + O2(g)=2CaO(s)

Magnesium When heated, it melts just before it burns. It does so with a blinding white flame, to leave a white ash.

2Mg(s) + O2(g)

=2MgO(s)

Aluminium Silver aluminum powder is heated in oxygen; it glows with a white Flame which is less bright than that of magnesium, to leave a white powder.

4Al(s)+302(g)

=2Al2O3(s)

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Zinc Grey Zinc powder burns slowly with a dull red flame to produce wispy yellow/green flakes which are white when they cool. 2Zn(s)+ O2(g)=2ZnO(s)

Iron Steel wool burn with a bright sparkle leaving solid black lumps. 3Fe(s)

+2O2(g)=Fe3O4(s)

Copper When heated after it corrodes in the air it immediately becomes coated with black powder. 2Cu(s) + O2(g)=2CuO(s)

Mercury Mercury becomes covered with red oxide; this coating disappears if it is heated more strongly. 2 Hg(l) + O2(g)= 2HgO(s).

Reaction of Non-Metals with oxygen

Non-Metals

Description of reaction

Equation

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Carbone Burns slowly with a yellow-white flame to form a colorless gas.

C(s)+O2(g)

=CO2(g)

Phosphorus

It burns very brightly to produce clouds of white smoke which settle out as white solid.

P4(s)+ 5O2(g)=P4O10(s)

Sulphur It melts and catches fire, burning with a blue flame. Produces misty gas with choking smell.

S(s)

+O2(g)=SO2(g)

Sulphur (s)

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Figure 1: Sulphur atom Figure 2: Sulphur molecule

General:

Has oxidation states -2, +4 and +6 Atomic number: 16 Atomic mass: 32.065 Electronic configuration: 2-8-6 Number of Protons/Electrons: 16 Number of Neutrons: 16 Classification: Non-metal

Physical Properties:

Density at 293 K: 2.07 g/cm3 Melting Point: 112.8 °C (385.95 K, 235.04001 °F) Boiling Point: 444.6 °C (717.75 K, 832.28 °F) yellow crystalline solid It is tasteless and odourless is a poor conductor of heat and electricity

Extraction of sulphur  

See figure c3 for the extraction of sulphur

Since sulphur in Free State is found at depths of more than 150 to 300 meters below the earth’s surface, the method of extraction of sulphur differs from other metal or non-metal extractions. Sulphur’s relatively low melting point (115°C) is utilized in this process. This is known as the Frasch process. Here compressed super heated water (at 170°C) is

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pressed into a pipe which reaches up to the sulphur deposits. The sulphur here melts. Introducing hot compressed air through another pipe brings it up. The molten sulphur and water mixture is forced up and is collected in a settling tank. The sulphur is cooled and water is evaporated. The sulphur extracted in this way is more than 99% pure.

Allotropes of Sulphur

Sulphur has 2 allotropes namely rhombic and monoclinic sulphur. 96 oqC is called the transition temperature because below 96 oC only rhombic sulphur is stable and above it only monoclinic sulphur is stable.

Effect of heat on Sulphur

Sulphur melts at 115 oC. As the temperature increases it becomes a runny yellow liquid, this soon becomes like treacle in texture and darker in colour. As the temperature increases it becomes runny again and darker. At 444oC the sulphur boils to give a brown vapour.

Plastic Sulphur

If hot runny liquid sulphur is poured into cold water, it forms a soft pliable substance, rather like chewing gum, this is called plastic sulphur. The sulphur goes hard and brittle and it forms changes to that of rhombic sulphur.

Chemical reactions of sulphur

Sulphur burns in air or oxygen with a blue flame forming sulphur dioxide.( S ( s ) + O2 (g ) SO2 (g ))

Sulphur reacts with metals to form sulphides.(Mg ( s ) + S ( s ) MgS

( s ))

Laboratory Preparation of Sulphur Dioxide

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Sulphur dioxide is prepared by:

Burning Sulphur in air ( S ( s ) + O2 (g ) SO2 (g )) Roasting sulphur containing metal ores in air (4FeS 2 ( s ) + 11O2 (g )

2Fe2O3 ( s ) + 8SO2 (g )) Action of acids on sulphate (IV) (sulphate) (acid + sulphate (IV)

(sulphate) salt + sulphur dioxide + water)

Reactions of Sulphur Dioxide

It is quite soluble in water Reacts with alkalis It is a strong reducing agent

Test for sulphur dioxide

If sulphur dioxide is bubbled through potassium manganate (VII) solution the potassium manganate changer colour from purple to colourless.

If sulphur dioxide is bubbled through potassium dichromate (VI) solution, the solution is reduced from an orange colour to a green colour.

Sulphuric Acid

The manufacture of Sulphuric Acid

See figure c4 for flowchart

Sulphuric acid is manufactured using the contact process.

1) Sulphur dioxide and oxygen are cleaned, dried and heated to 450oC and pressurised to 2-3 atm.

2) They are then passed over a vanadium pentoxide catalyst. (2SO 2 + O2 2SO3)

3) The product is then dissolved in concentrated sulphuric acid forming oleum (H2SO4.SO3 ( l ))

4) The mixture is then carefully diluted to the appropriate concentration of sulphuric acid.

Reactions of sulphuric acid

Concentrated sulphuric acid is not an acid. This is because it structure is covalent and does not contain any free hydrogen ions.

Hot concentrated sulphuric acid is an oxidising agent Cold concentrated sulphuric acid is a dehydrating agent

Sulphates

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Test for sulphates:

1) Some sulphate compounds decompose when heated and give off white fumes of sulphur trioxide. Example copper sulphate

2) Sulphate (VI) compounds can be detected using barium chloride. If a solution of barium chloride is added to a sulphate, an insoluble precipitate of barium sulphate is formed. Example sodium sulphate

Hydrogen Sulphate

Hydrogen sulphide, H2S, is an important impurity in natural gas which must be removed before the gas is used. This is done by an absorption and regeneration process to concentrate the H 2S, followed by a catalytic oxidation (Claus process) using porous catalysts such as Al 2O3 or Fe2O3.

8H2S + 4O2 → S8 + 8H2O

Over the years the Claus process has been improved and a modified process can yield 98% recovery.

Chemical reactions of Hydrogen sulphate

Hydrogen sulphate dissolves slightly in water forming weak acids.

The salts of this acid are called sulphides. There are two ways of making them:

1) Soluble sulphides can be made in hydrogen sulphide is bubbled through a solution of an alkali

2) Insoluble sulphides can be made by precipitation.

Hydrogen sulphide is a reducing agent.

Table 1: Uses of sulphur and some of its compounds

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Chlorine (Cl)

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Figure 1: Chlorine atom Figure 2: Chlorine molecule (Cl2)

General

Has many oxidation states Atomic number: 17 Atomic mass: 35.453 Electronic configuration: 2-8-7

Physical Properties

Colour: green/yellow Has a choking smell It is poisonous Melting point -101 °C Boiling point -34.6 °C Denser than air

Chemical properties

Chlorine is an acidic gas Chlorine reacts with metals forming anhydrous ionic chlorides Chlorine reacts with non-metals forming covalent chlorides It is a powerful oxidising agent In the presence of moisture chlorine acts as a bleaching agent. Chlorine

combines with water forming hydrochloric acid and chloric (I) acid

Chloric (I) acid is the bleaching agent. It is unstable and decomposes in light forming hydrochloric acid and oxygen.

Hydrogen gas will burn in chlorine forming hydrogen chloride fumes.

Industrial Production of Chlorine

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The majority of chlorine is obtained from the electrolysis of brine. It is broken down to give both sodium hydroxide and chlorine gas.

Laboratory Preparation of Chlorine

Chlorine can be made by the oxidation of Hydrochloric acid.

2HCl (aq) + O Cl2 (g) + H20 (l)

1) Using Manganese (IV) oxide as the oxidising agent. When the hydrochloric acid and Manganese (IV) oxide are mixed, Chlorine is evolved. The apparatus used is seen is figure ()

2) Using potassium manganate (VII) as the oxidising agent. The reaction occurs without heating and chlorine is evolved. The apparatus used can be seen in figure ()

Uses of Chlorine

Table 1: Uses of chlorine and some of its compounds

Hydrogen Chloride

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It is a covalent gas mainly used to make Hydrochloric acid.

Physical Properties

It is heavier than air It is colourless when dry It has a choking smell Damp litmus turns red It is very soluble in water

Chemical Reaction

Hydrogen chloride gas reacts with ammonia gas forming white ammonium chloride.

Hydrogen chloride gas dissolves in water forming hydrochloric acid

Manufacture of hydrogen chloride

Chlorine is burned in an atmosphere of hydrogen. The hydrogen gas formed is then immediately dissolved in an absorption tower.

Laboratory Preparation of Hydrogen Chloride

The apparatus can be seen in figure c5

Hydrogen is formed when ever a chloride is reacted with concentrated sulphuric acid. Example: Sodium Chloride + sulphuric acid Hydrogen Chloride + Sodium Hydrogen Sulphate

Tests for chlorides

All chlorides react with concentrated sulphuric acid to form hydrogen chloride gas

E.g. NaCl (s) + H2SO4 (aq) HCl (g) + NaHSO 4 (aq)

Adding silver ions will identify chlorine ions in the solution. This is because the white precipitate silver chloride is formed

Cl- (aq) + Ag+

(aq) AgCl (s)

Diagrams and Illustrations

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Figure a1: diagram of the burst furnace

Figure a2: Electrolysis cell for the extraction of Aluminium

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Figure a3: The cell used in the extraction of sodium

Figure a4: Diagram of the water cycle

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Figure a5: Diagram of the alveoli

Figure a6: Finding the percentage of oxygen in the air

Figure a7: Apparatus used for the preparation of Hydrogen

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Figure a8: Diagram of a water works

Figure a9: Diagram of a sewage treatment plant

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Figure b1: Diagram of the carbon cycle

Figure b2: Preparation and collection of Carbon dioxide

Figure b3: Flowchart showing the production of sodium carbonate

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Figure b4: The nitrogen Cycle

Figure b5: The harbour process

Figure b6: Laboratory Preparation of ammonia

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Figure b7: Flow chart showing the production of nitric acid

Figure b8: Laboratory preparation of Nitric acid

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Figure b9: The extraction of Phosphorus

Figure c1: The structure of phosphorus allotropes

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Figure c2: Laboratory Preparation of Phosphorus (V) chloride

Figure c3: The extraction of sulphur

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Figure c4: Flowchart showing production of sulphuric acid (contact Process)

Figure c5: Manufacture of Hydrochloric acid

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In conclusion, the uses, properties and reactions of various elements in the periodic table were investigated and a better understanding about their reactions, compounds, properties, etc was obtained.

Reference

CXC Chemistry Jacqueline Fergusson Richard Hart Chemistry A concise revision course for CXC Anne

Tindale A-lever Chemistry Forth edition E.N. Ramsden General Chemistry 3rd Edition Umland Bellama CXC Human and Social Biology Phil Gadd Wikipedia The internet

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