Non metals and their compounds

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NON METALS AND THEIR COMPOUNDS
Non-metals these are elements which react with the electrons by gaining to form negatively
charged ions
Where e = electron.
Example
There are only 17 non-metals (excluded H) arranged at right in the periodic table. They are
electronegative elements with high electronegativity and ionization energy. They have a
tendency to accept electron and form anions.
Characteristics of non-metals
1. They are oxidizing agents. This is because they are electron acceptor or they gain
electrons. The substance which gains electrons is said to be oxidizing agent.
2. They have low melting point and boiling point.
3. They have low density.

4. They gain electrons during chemical reactions.
5. They have low physical strength.
6. Their oxides are acidic in nature.
Chemical Properties of non-metals
1. Non-metals react with oxygen to form non-metallic oxides, which are acidic in nature.
Hence non-metallic oxides form an acidic solution when dissolved in water and turn
litmus solution red. For example, carbon is oxidized to form carbon dioxide, which is
acidic in nature.
Non metal Its oxide
Chlorine –Cl HCl or HOCl
Carbon- C CO2, CO
Sulphur – S SO2, SO3
Nitrogen – N NO2, N2O4
 Most of non-metal oxides form gas with acidic, the neutral or basic
Common properties of gases
GASES SMELL NATURE SOLUBILITY COLOUR
Ammonia NH3 Urine smell
Basic or
alkaline
Soluble Colorless
Carbon dioxide CO2 Odorless Acidic Soluble Colorless
Carbon monoxide CO Odorless Neutral Insoluble Colorless
Sulphur dioxide SO2
Irritating chocking
smell
Acidic Soluble Colorless
Chlorine Cl2
Irritating chocking
smell
Acidic Soluble Greenish yellow
Hydrogen sulphide H2S
Rotten eggs or
cabbage
Acidic Soluble colorless
Nitrogen dioxide NO2 Irritating chocking Acidic Soluble Reddish brown

smell yellow
Nitrogen monoxide NO Odorless neutral Insoluble Colorless
Hydrogen chloride HCl
Irritating chocking
smell
acidic Soluble Colorless
2. Non-metals are good oxidizing agents and are oxidized in almost all of their reactions. Like
aluminum is oxidized with bromine to form aluminum bromide.
3. Non-metals have a tendency to oxidize metals.
4. They can easily oxidize those compounds with which they react.
5. Less electronegative non-metals like carbon & hydrogen can act as a reducing agent for some
compounds like ferric oxide and copper (II) oxide.
6. Generally, no reaction takes place between non-metals and acids. But non-metals react with
base’s to form salts. For example, chlorine reacts with calcium hydroxide to form bleaching
powder.
OXIDIZING PROPERTY OF NON METAL
An oxidizing agent is a substance which accepts electrons; a reducing agent is a substance which
donates electrons. Oxidizing and reducing agents can be identified in redox reactions. The
elements with low electronegativity (metals) tend to form ions by losing electrons (oxidation)
and so can act as reducing agents; the elements with high electronegativity (non-metals) tend to
form ions by gaining electrons (reduction) and so can act as oxidizing agents. The strongest
reducing agents are found in Group 1 whilst the strongest oxidizing agents come from Group 7.
The electrochemical series indicates the effectiveness of oxidizing and reducing agents.
One of the main properties of non metal is to gain electron. None metal have 4 to 6 electron in
outer most shell.

example
 Nitrogen:2:5
 Oxygen :2:6
 Chroline:2:8:7
To become stable the outer most shell must be filled. Hence it is for non metal to gain an electron
to become stable
example
 Nitrogen gain 3 electron = 2:8
 Oxygen gain 2 electron = 2:8
 Chlorine gain 1 electron = 2:8:8
Any chemical substance gain electron is oxidizing agent.Alsoall non metal is OXIDIZING
AGENT
The number of electron gained per atom signifies a valency. A Cl gain 1 and oxygen gain 2
hence valency of Cl is 1 and for oxygen is 2.
1. OXIDIZING POWER OF HALOGENS
These are elements found in the group VII of the periodic table and are salt producers.
Oxidizing power of halogens decrease as one move down the group, this is because of tendency
of electronegativity decrease.
Examples of halogens are; chlorine, bromine, fluorine and Iodine
All halogens name end with suffix -ine
Properties of halogens
1. Are salt producers.
2. Are most electronegative elements. With fluorine being the most electronegative and
Iodine the least in the group. The electronegativity decreases downwards the group.
3. Some of them are gases, solids and others liquids at room temperature.
Halogen state at room temperature
Fluorine Gas
Chlorine Gas
Bromine Liquid

Iodine Solid
4. They are colored.
Halogen Color
Chlorine Greenish-yellow
Fluorine Purple
Bromine Brown-reddish.
Iodide Brown
5. They are strong oxidizing strong oxidizing agents i.e. they gain or accept electrons.
DISPLACEMENT REACTION INVOLVING HALOGENS
In displacement reaction one element take part of another in a compound. A more reactive non
metal displace a less reactive non metal from a solution of one of its salt.
In group VII are more reactive halogen displaces a less reactive halogen from another halide
salt.
For example;Chlorine displace bromine or iodine
Chlorine
Chlorine is the greenish yellow gas which turns dump blue litmus red and bleaches it. Chlorine is
a chemical element with symbol Cl and atomic number 17. Chlorine is in the halogen group (17)
and is the second lightest halogen following fluorine. The element is a yellow-green gas under
standard conditions, where it forms diatomic molecules. Chlorine has the highest electron
affinity and the fourth highest electronegativity of all the reactive elements. For this reason,
chlorine is a strong oxidizing agent. Free chlorine is rare on Earth, and is usually a result of
direct or indirect oxidation by oxygen.
The most common compound of chlorine, sodium chloride (common salt), has been known since
ancient times. Around 1630 chlorine gas was first synthesized in a chemical reaction, but not
recognized as a fundamentally important substance.
Preparation of chlorine from potassium permanganate KMnO4 and hydrochloric acid
When potassium permanganate reacted with concentrated HCl, chlorine gas is formed, which is
collected by the downward displacement.

Consider the diagram below.
KMnO4(aq) + HCl(aq) → 2KCl + 2MnCl2(aq) + 8H2O(l) +5Cl2(g)
Preparation of the chlorine gas, hydrogen sulphide (H2S) and carbon monoxide in the fume
chamber/ cupboard
Normally these gases are prepared in the fume chamber for the following reasons.
Because some of them are toxic and poisonous this can cause death to occur. So when the gases
are prepared into the laboratory; windows and doors should be left open to allow good
ventilation of the air in the laboratory, when the poisonous/toxic gases are produced, they can
escape freely from the laboratory to the atmosphere to allow the leaked gaseous molecules to
accumulate in the laboratory.
The Test for Chlorine Gas, Cl2 (g)
1.Chlorine gas is green-yellow in color.
2.Chlorine gas has a pungent choking smell.
3.Chlorine gas turns moist litmus paper from blue to red and moist universal indicator paper
to red - it is acidic. After turning red both papers are then bleached white.
4.Chlorine gas will put out a lit splint.
Specific Test for Chlorine Gas
Chlorine is the only gas that has a bleaching affect

Properties of chlorine
1. It is a greenish-yellow gas.
2. It has an irritating shocking smell
3. It is soluble in water.
4. It is slight denser than air.
5. It turns blue litmus red then bleach.
Chlorine gas is soluble in water.
When chlorine gas is dissolved in water, it dissolves to form hydrochloric acid (HCl) and
hypochlorous acid (HOCl)
HOCl – hypochlorous acid
HCl – hydrochloric acid
 Where HCl can kill bacteria and HOCl can bleach or decolorize.
The acid decomposes to form oxygen and hydrochloric acid.
Chlorine gas is the bleaching agent.
Bleaching agent- is the substance which can decolorize or bleach or remove the original color of
another substance, E.g.
 Blue into colorless
 Red into colorless
 Yellow to colorless
How chlorine bleaches colored matter such as water, clothes, litmus.
E.g. when the chlorine gas is dissolved in water it forms two acids Hypochlorous acid and
hydrochloric acid. hypochlorous acid is unstable so it decompose slowly to form hydrochloric
acid and oxygen , the formed oxygen combines with the colored matter to such that they undergo
oxidation then bleaching occur such that the substance becomes colorless.

Chlorine gas or compounds added to the swimming pool and drinking water.
This is because chlorine gas dissolves in the swimming pool or drinking water to form HCl
which kills bacteria and pathogens and at the same time decolorizes water thus makes water to
become safe for using for human beings.
The addition of chlorine in water is called chlorination of water. The aim is to kill bacteria and
decolorize water.
Water can be clean but not safe for drinking.
Because it may contain bacteria and microorganisms which cause diseases and can’t be seen by
our naked eyes and so water should be boiled or chlorinated.
Chemical properties of chlorine gas
Chlorine shows the following chemical properties.
1. Reaction with metal.
2. Reaction with water.
3. Reaction with alkali
E.g. NaOH, KOH etc
(a) Reaction with cold dilute alkali
KOCl is sodium hypochlorous.
(b)Reaction with hot concentrated alkali.

NaClO3 is sodium chlorate.
4. Reaction with ammonia gas
DISPLACEMENT REACTIONS OF CHLORINE
Chlorine gas can show the displacement reaction such that those elements below, it would
displace.
The potassium iodide solution turns to brown due to the formation of iodine gas.
Effects of sunlight into chlorine water
Chlorine water is the unstable compound, when the heat or sunlight become into contact with
chlorine water (HOCl), the chlorine water decomposes then produces oxygen and hydrochloric
acid. Oxygen gas produced escape to atmosphere.
Consider the diagram below,

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Commercial uses of chlorine gas
1. It is used in the production of common salt, e.g. NaCl.
2. Used in the chlorination of water, this is the addition of chlorine in water so as to kill the
micro-organisms in water.
3. Used in the bleaching action. E.g. Bleaching of water, some of the clothes can be
bleached using chemical compounds made up of chlorine. Such as water guard.
4. Used to make compounds used in the preparation of oxygen gas. E.g. potassium
chlorate KCLO3 can be prepared from chlorine gas
5. (PVC)… materials used to make water pipes and plastics.
(PVC – polyvinyl chloride)
6. Used in disinfecting.
7. Used to bleach wood pulp in paper making.
8. Used to make chemical such as carbon tetrachloride used in dry cleaning.
HYDROGEN CHLORIDE GAS
This is the compound hydrogen chloride has the chemical formula HCl. At room temperature, it
is a colorless gas, which forms white fumes of hydrochloric acid upon contact with atmospheric
humidity. Hydrogen chloride gas and hydrochloric acid are important in technology and industry.
Hydrochloric acid, the aqueous solution of hydrogen chloride, is also commonly given the
formula HCl has the chemical formula HCl. At room temperature, it is a colorless gas, which
forms white fumes of hydrochloric acid upon contact with atmospheric humidity. Hydrogen
chloride gas and hydrochloric acid are important in technology and industry. Hydrochloric acid,
the aqueous solution of hydrogen chloride, is also commonly given the formula HCl

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Preparation of hydrogen chloride gas
Where
H2SO4 (aq) is concentrated sulphuric acid.
NaCl is table salt.
NaHSO4 is sodium bisulphate.
HCl is hydrogen chloride.
A balanced chemical equation for the reaction
The gas cannot be collected over water because it is soluble in water. Thus, it is collected by
downward delivery method.
Test of hydrogen chloride gas
The gas forms white denser fumes of ammonium chloride with ammonium gas.
The physical properties of chlorine
1. It is soluble in water.
Where H3O+ Is oxonium ion

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2. It is colorless
3. It has a pungent smell.
4. It is denser than air.
5. It turns blue litmus to red since it is acidic.
Uses of hydrogen chloride gas
These are some of the uses for hydrogen chloride gas:
1. Most hydrogen chloride is used in the production of hydrochloric acid.
2. Hydrochlorination of rubber
3. Production of vinyl and alkyl chlorides
4. Chemical intermediate in other chemical production
5. Used in toilet bowl cleaner (The Works)
6. Use as babbitting flux
7. Treatment of cotton
o Delinting
o Separation from wool
8. Used in semiconductor industry (in pure grade)
o Etching semiconductor crystals
o Converting silicon to SiHCl3 for purification of silicon
Sulphur and sulphur extraction
1. Sulphur ( S= 2:8:6)
Sulphur is a chemical element with the symbol S and atomic number 16. It is an abundant,
multivalentnon-metal. Under normal conditions, sulfur atoms form cyclic octatomic molecules
with chemical formula S8. Sulphur element is a bright yellow crystalline solid when at room
temperature. Chemically, sulphur can react as either an oxidacing or a reducing agent. It oxidizes
most metals and several nonmetals, including carbon, which leads to its negative charge in most
organosulphur compounds, but it reduces several strong oxidants, such as oxygen and fluorine.
Sulphur is the solid yellow powdered non metal, yellowish in color. Sulphur is found within the
soil in the place called Sulphur bed.
The following countries lead in the extraction of sulphur; Mexico, Italy, Japan, France and USA.
Sulphur can be found combined with ores such as;

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1. Galena - Pbs
2. Iron Pyrite -FeS2
3. Zinc blend – ZnS
Some compounds of sulphur have unpleasant smell (odors). Sulphur compounds are responsible
for the smell such as that of onion, cabbage. The substance which causes tears during cutting of
onions, also the smell associated with rotten eggs is caused by the sulphur compounds.
Compounds of sulphur
They include;
1. Sulphur dioxide. (SO2)
2. Sulphuric acid.( H2SO4)
3. Sulphur trioxide. (SO3)
4. Hydrogen Sulphide( H2S)
Hydrogen sulphide gas – H2S have smell of the rotten eggs or cabbages. At the same time the gas
is poisonous and toxic.
Allotropy and allotropes of sulphur
Allotropy is the existence of an element in various different forms with different physical
properties but same chemical properties.
Allotropes are the different forms of the same element with the same chemical properties but
different physical properties e.g. rhombic, monoclinic, plastic, amorphous are all allotropes of
sulphur.

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Rhombic sulphur
1. Exists at 960C
2. Transparent white crystal
3. Stable at 960C
4. Has shape of octahedral
5. Prepared by heating sulphur
With methyl benzene (toluene)
Monoclinic sulphur
1. exists above 960C
2. yellow translucent
3. stable above 960C
4. has octahedral shape
5. prepared by heating sulphur
Relationship between rhombic and monoclinic sulphur (at transition temperature)
Transition temperature is the temperature which one allotrope can change to another allotrope
either by raising or lowering the temperature.
Rhombic and monoclinic sulphur are related to the temperature.
Preparation of plastic sulphur
Plastic sulphur- is the allotrope of sulphur prepared by heating sulphur then cooling it and
mixing it with cold water. Sulphur heated then melts and becomes molten after that mixed with
cold water, sulphur contact becomes very hard than any allotrope.
NB: plastic sulphur is not considered to be actual allotrope of sulphur because it is harder than
others.

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Amorphous sulphur prepared from chemical compound
Amorphous sulphur is the sulphur allotrope of sulphur produced from the chemical reaction of
compounds containing sulphur.
Commercial uses of plastic sulphur
Used in the synthesis of plastic materials such as jars, trough, plastic water pipes, basin, pen-
tubes containing ink.
EXTRACTION OF SULPHUR BY FRASCH PROCESS
The Frasch process is a method to extract sulphur from underground deposits. It is the only
economic method of recovering sulphur from elemental deposits.[1] Most of the world's sulphur
was obtained in this way until the late 20thcentury, when sulphur was recovered from petroleum
and gas sources (recovered sulphur) became more commonplace (see Claus process).
In the Frasch process, superheated water is pumped into the sulphur deposit; the sulphur melts
and is extracted. The Frasch process is able to produce high purity sulphur.
As from 2011, the only operating Frasch mines worldwide are in Poland and since 2010 in
Mexico. The last mine operating in the United States was closed in 2000. A Frasch mine in Iraq
was closed in 2003 due to the U.S. invasion of Iraq.
Process
In the Frasch process, three concentric tubes are introduced into the sulphur deposit. Superheated
water (165 °C, 2.5-3 MPa) is injected into the deposit via the outermost tube. Sulphur (melting
point 115 °C) melts and flows into the middle tube. Water pressure alone is unable to force the
sulphur into the surface due to the molten sulphur's greater density, so hot air is introduced via
the innermost tube to froth the sulphur, making it less dense, and pushing it to the surface.

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The sulfur obtained can be very pure (99.7 - 99.8%). In this form, it is light yellow in color. If
contaminated by organic compounds, it can be dark-colored; further purification is not economic,
and usually unnecessary. Using this method, the United States produced 3.89 million tonnes of
sulfur in 1989, and Mexico produced 1.02 million tonnes of sulfur in 1991.
The Frasch process can be used for deposits 50–800 meters deep. 3-38 cubic meters of
superheated water are required to produce every tonne of sulfur, and the associated energy cost is
significant. A working demonstration model of the Frasch process suitable for the classroom has
been described.
CHANGES WHICH OCCUR WHEN SULPHUR IS HEATED IN THE ABSENCE OF
AIR
When sulphur is heated in the absence of air, it undergoes changes as shown below.
(a) Sulphur at 1600C
Sulphur melts and becomes mobile yellow liquid.
(b) Sulphur at 2000C
Sulphur becomes dark, viscous and the mobility of sulphur increases more.
(c) Sulphur at 4440C
Sulphur boils to become into vapour or gaseous form.
1. Chemical properties of sulphur
Consider the reactions below.

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(Sulphur dioxide gas formed turning yellow potassium dichromate to green)
Commercial uses of sulphur
1. Manufacture of gun powder.
2. Manufacture of sulphuric acid through contact process.
3. Manufacture of matches.
4. Manufacture of plastic flower.
5. Manufacture of detergents e.g. powder soap, puff.
6. Manufacture of rubber through vulcanization.
Vulcanization is the hardening of rubber or is the process of heating sulphur with rubber to make
it hard.
7. Manufacture of drugs and skin ointment.
8. Production of bleaching compound
Sulphuric acid and contact process
The contact process is the process used to manufacture concentrated sulphuric acid (H2SO4)in
large scale in the industries.
E.g. concentrated sulphuric acid – is the acid containing few (small) amounts of water molecules.
Sulphuric acid, its industrial manufacture, uses and reactions

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Sulphuric acid is made in large quantity industrial because it has many important uses. The
production of sulphuric acid takes place in four stages.
Stage 1
Production of sulphur dioxide
Sulphur dioxide gas can be produced by three methods.
1. Burning sulphur in air. The sulphur from Frasch process is burned in air. The equation is
2. Hydrogen sulphide from crude oil is burned in air. The equation is;
3. Sulphide ores are roasted in air to extract metals from the ores. Sulphur dioxide is also
produced, the equation is ;
Stage 2
The contact process
The sulphur dioxide gas produced in stage 1 is mixed with air. The air and sulphur dioxide pass
through a dust precipitate to remove any dust impurities and then the gases are passed through a
drier. It is essential that this purification process takes place because impure gases will poison the
vanadium (V) oxide catalyst.
The sulphur dioxide gas produced in stage one is mixed with air. Sulphur dioxide is converted
into sulphur trioxide in a converter. Like the Haber process, this is an equilibrium reaction and
goes in both directions. A yield of 95% is obtained by using the following conditions.
 Catalyst vanadium (V) oxide(vanadium pentaoxide)
 Temperature – 4500C
 Pressure 1 atmosphere
This reaction is exothermic, which means that as sulphur trioxide is formed energy is released. If
the temperature rises above 4500C, the yield of sulphur trioxide decreases. The heat exchanger
maintains the temperature at 4500C.

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Stage 3
Production of oleum
If sulphur trioxide is dissolved in water, sulphuric acid is formed.
Oleum: H2S2O7 - yellow heavy chemical.
Definition; is the heavy yellow liquid formed when sulphur trioxide gas – SO3 reacts with acid
H2SO4
NOTE: Oleum have strong affinity to water, so absorbs a lot of water than concentrated
sulphuric acid formed.
Diagrammatic representation of the contact process
NB:
Do not add water to an acid but acid to water for safe dilution.
Because heat absorbed from the surrounding environment which cause water spray to boil which
scatter. Then it can burn at the same time with denser fumes that will be formed. Normally acids
have strong affinity to water. Reaction with water is endothermic.
Reactions of dilute sulphuric acid

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Reaction of sulphuric acid with marblechips - CaCO3 (egg shell)
Insoluble calcium sulphate formed, this cover acid to the top hence prevent further (more)
reaction to occur or continue, then the reaction stops. This is why we can’t prepare CO2 by
reacting marble chips and acid.
Commercial uses of sulphuric acid (H2SO4)
1. Used to make detergents for washing.
2. Used to make paints for house painting.
3. Used to make dyes for making clothes
4. Used in the synthesis of solvents. E.g. thinner for diluting paints.
5. Used in pickling – the removal of the oxide layer on the metal surface e.g. before
galvanization.
6. Used in the lead accumulator battery as the electrolyte.
7. Used to manufacture fertilizers e.g. SA ( ammonium sulphate)
8. Used to make plastic materials.
NITROGEN AND ITS COMPOUNDS
Nitrogen, symbol N, is the chemical element of atomic number seven and electronic
configuration (2:5). At room temperature, it is a gas of diatomic molecules and is colorless and
odorless. Nitrogen is a common element in the universe, estimated at about seventh in total
abundance in our galaxy and the Solar System. On Earth, the element is primarily found as the
gas molecule; it forms about 78% of Earth's atmosphere. The element nitrogen was discovered as
a separable component of air, by Scottish physician Daniel Rutherford, in 1772.

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Nitrogen gas as the diluents gas
Nitrogen gas is said to be the diluents gas. This is because it reacts with oxygen gas supporting
combustion, the reaction between oxygen and nitrogen help to reduce the rate of combustion,
rusting. This shows that if nitrogen in the atmosphere was absent, then rusting and combustion
could occur faster on the Earth’s surface.
Inert/ unreactive noble property of nitrogen
Nitrogen gas – N2 is unreactive gas and can resemble or be similar to noble gas due to the truth
that covalent bond in nitrogen is long and strong to break down hence nitrogen becomes
unreactive gas. The bond is triple covalent bond. E.g. N2
LABORATORY PREPARATION OF NITROGEN
Apparatus – delivery tubes, furnace, beaker, trough, Bunsen burner, 2 wash bottles Chemicals –
Air, water, sodium hydroxide Procedure Either: Nitrogen is prepared from the air by removing
oxygen and carbon dioxide. Water is used to push air through sodium hydroxide solution (caustic
soda solution) which removes carbon dioxide.
The remaining gas is passed over heated copper turnings to remove Oxygen.
Nitrogen is collected over water as it is insoluble in water.

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PROPERTIES OF NITROGEN
Physical properties
(a) it is colorless gas without smell
(b) it is a reactive gas
(c) it does not burn / doesn’t support combustion
(d)it is neither acidic nor basic
(e)It is less denser than air
(d)It is insoluble in air
Chemical properties
Nitrogen is inert unlike Oxygen; it reacts under special conditions for example
1. It reacts with some metals at very high temperatures forming nitrides e.g. Calcium and
magnesium.
1. It reacts with hydrogen to form ammonia (Haber’s process)
Commercial uses of nitrogen
1. Used in the storage of petroleum to avoid the exploding of petrol.
2. Used in the preparation of ammonia gas.
Common drying agent in laboratory for gas
 Concentration of sulphuric acid
 Anhydrous calcium chloride.
 Calcium oxide (CaO)
Ammonia gas
Ammonia, or azane, is a compound of nitrogen and hydrogen with the formula NH3. It is a
colorless gas with a characteristic of pungentsmell. Ammonia contributes significantly to the
nutritional needs of terrestrial organisms by serving as a precursor to food and fertilizers.
Ammonia in both directly or indirectly is also a building-block for the synthesis of many
pharmaceuticals and is used in many commercial cleaning products. Although in wide use,

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ammonia is both caustic and hazardous. The gas form white denser fumes with hydrogen
chloride gas.
PREPATION OF AMMONIA
Ammonia is prepared by heating a mixture of calcium hydroxide and ammonium chloride.
Ca(OH)2(s) + 2NH4Cl(s) → CaCl2(s) + 2H2O(L) + 2NH3(g)
The tube in which ammonia is generated is fixed in a slanting position to prevent the water
formed from running back and crack the whole tube. Concentrated sulphuric acid and anhydrous
calcium chloride are not used to dry ammonia because they react with it. The gas is passed
through fresh quicklime (solid calcium oxide lumps) to effectively dry it. Ammonia is collected
by upward delivery as it is lighter than air
Compounds of nitrogen
1. Nitric acid ( HNO3)
2. Nitrogen monoxide (NO)
3. Dinitrogen Monoxide (N2O)
4. Nitrogen dioxide.
Physical properties, chemical properties and uses of ammonia
Physical Properties of Ammonia
 Ammonia is a colorless gas.

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 It has a pungent odor with and an alkaline or soapy taste. When inhaled suddenly, it
brings tears into the eyes.
 It is lighter than air and is therefore collected by the downward displacement of air.
 It is highly soluble in water: One volume of water dissolves about 1300 volumes of
ammonia gas. It is due to its high solubility in water that the gas cannot be collected over
water.
 It can be easily liquefied at room temperature by applying a pressure of about 8-10
atmosphere.
 Liquid ammonia boils at 239.6 K (- 33.5°C) under one atmosphere pressure. It has a high
latent heat of vaporization (1370 J per gram) and is therefore used in refrigeration plants
of ice making machines.
 Liquid ammonia freezes at 195.3 K (-77.8°C) to give a white crystalline solid.
Chemical properties of ammonia
They include;
1.The combustion of ammonia proceeds with difficulty but yields nitrogen gas and water.
4NH3 + 3O2 + heat → 2N2 + 6H2O
2. Ammonia readily dissolves in water with the liberation of heat.
NH3 + H2O ⇔ NH4
+ + OH−
3. Ammonia gets oxidized to nitrogen, when passed over heated metal oxides.
Uses of ammonia
 In the manufacture of rayon and urea
 In the manufacture of fertilizers such as urea diammonium phosphate, ammonium nitrate,
ammonium sulphate etc.
 In ice plants, as a refrigerant
 In furniture industry, as a cleansing agent for furniture and glass surfaces.

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 In the manufacture of nitric acid by Ostwald's process.
 In the manufacture of sodium carbonate by Solvay's process.
EXPERIMENT TO SHOW THE REDUCING PROPERTIES OF AMMONIA TO
COPPER OXIDE- CuO
Chemical equation for the reaction
Fountain experiment to demonstrate solubility of the gases e.g. NH3, HCl
Fountain experiment this is an experiment used to show the solubility of the gases.
The ammonia fountain is a type of chemical demonstration. The experiment consists of
introducing water through an inlet to a container filled with ammonia gas.Ammonia dissolves
into the water and the pressure in the container drops. As a result more water is forced into the

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container from another inlet creating a fountain effect. The demonstration introduces concepts
like solubility and the gas laws at entry level.
A different gas of comparable solubility in water, such as hydrogen chloride, can be used instead
of ammonia.
If the ammonia is replaced by a liquid vapor, such as water vapor, at a pressure higher than its
room-temperature vapor pressure, a similar effect is produced. In this case, the reduction in
pressure in the container is due to condensation of the vapor as the container cools to room
temperature
Nitric acid
Nitric acid(HNO3), also known as aqua Fortis and spirit of niter, is a highly corrosive strong
mineral acid. The pure compound is colorless, but older samples tend to acquire a yellow cast
due to decomposition into oxides of nitrogen and water. Most commercially available nitric acid
has a concentration of 68%. When the solution contains more than 86% HNO3, it is referred to
as fuming nitric acid. Depending on the amount of nitrogen dioxide present, fuming nitric acid is
further characterized as white fuming nitric acid or red fuming nitric acid, at concentrations
above 95%.
Nitric acid is the primary reagent used for nitration - the addition of a nitro group, typically to an
organic molecule. While some resulting nitro compounds are shock- and thermally-sensitive
explosives, a few are stable enough to be used in munitions and demolition, while others are still
more stable and used as pigments in inks and dyes. Nitric acid is also commonly used as a strong
oxidizing agent.
Preparation of Nitric Acid
Laboratory preparation
Nitric acid is usually prepared by heating potassium nitrate with concentrated sulphuric acid.
This heating is done in a glass retort and vapors of nitric acid are condensed in a receiver, which
is cooled by water. The reaction is:
Chemical
1. Potassium nitrate: 30 g
2. Sulfuric acid: 98% 35mL
Procedure

27
1. Place 30 g potassium nitrate into a round-bottomed flask.
2. Pour 35 ml 98% sulfuric acid and place a stir bar into the flask.
3. Set up a simple distillation apparatus and start heating the round-bottomed flask with a hot
plate with stirring. Soak the collecting vessel in a cold-water bath
To ensure our yield of anhydrous nitric acid is higher, adding some excess sulfuric acid is very
helpful. It makes the final product less water and produces more nitric acid. So weadd equation 2
sulfuric acids into equation 1 potassium nitrate to produce anhydrous nitric acid.
Industrial Preparation
On a commercial scale, nitric acid is manufactured through the Ostwald's process - the process of
catalytic oxidation of ammonia.
Ostwald's process
The conversion of ammonia into nitric acid in this process is done through the following steps:
Step1
Oxidation of ammonia to nitric oxide
Ammonia is oxidized by air in the presence of Pt catalyst at 800°C to give nitric oxide.
Step 2

28
Oxidation of NO to NO2
The nitric oxide is oxidized by air at temperature below 100°C, to give nitrogen dioxide (NO2)
Step 3
Formation of nitric acid
Nitrogen dioxide is then converted to nitric acid by absorbing NO2 in water, in the presence of
air.
Properties of nitric acid (HNO3)
1. It can attack the rubber materials.
2. It is unstable so can decompose at room temperature.
3. It is yellowish in color, due to the dissolved nitrogen dioxide, NO2 in the acid.
4. It is the strong oxidizing agent
5. It is a fuming acid.
Chemical properties of nitric acid
1. Reaction with non metals.
CO2 Gas turns lime water to milky
And extinguishes fire and a reddish brown gas is evolved.
2. Reaction with metal.

29
3. Reactions with Basic Oxides
4. Reaction with Bases (Hydroxides)
5. Reaction with Carbonates and Hydrogen Carbonates
6. Reaction with Metals
Nitric acid usually does not behave as an acid, with metals to form the corresponding salt and
liberate hydrogen.
However, magnesium and manganese are the only two metals, which react with cold and very
dilute (1%) nitric acid to evolve hydrogen.
7. Reaction with Metallic Sulphites

30
Uses of nitric acid
The important uses of nitric acid are as follows:
1) Nitric acid plays a significant role in the manufacture of various products such as:
 Explosives like trinitrotoluene (T.N.T.) nitro glycerin, gun cotton, ammonia etc.
 Fertilizers such as calcium nitrate, ammonium nitrate etc.
 Nitrate salts such as calcium nitrate, silver nitrate, ammonium nitrate.
 Dyes, perfumes, drugs etc. from coal tar products.
 Sulphuric acid by Lead Chamber process.
2) It is used in the purification of silver, gold, platinum etc.
3) Nitric acid is used in etching designs on copper, brass, bronze ware etc.
4) It is used to prepare "aqua regia" to dissolve the noble elements.
5) It is used as a laboratory reagent.
NITROGEN MONOXIDE (NO)
Nitric oxide, or nitrogen oxide, also known as nitrogen monoxide, is a molecule with chemical
formula NO. It is a free radical and is an important intermediate in the chemical industry. Nitric
oxide is a by-product of combustion of substances in the air, as in automobile engines, fossil fuel
power plants, and is produced naturally during the electrical discharges of lightning in
thunderstorms. This is the colorless gas which turns reddish-brown on the exposure to air. The
gas has strong affinity to oxygen, so can react with oxygen from the atmosphere then turn to
reddish brown.
Properties of nitrogen monoxide
1. It is colorless
2. It is neutral to litmus.
3. It is odorless.
4. It is insoluble in water.
5. It is denser than air.

31
Preparation of nitrogen monoxide in laboratory
The gas can be prepared when nitric acid react with moderate reactive metal.
The delivery tube is passed over the water to remove dissolved impurities (nitrogen dioxide)
Chemical properties of nitrogen monoxide
They include;
The gas reacts with oxygen gas readily to form nitrogen dioxide gas.
Commercial uses of the nitrogen monoxide
1. Used to prepare nitrogen dioxide gas.
2. Used to prepare nitric acid in Haber process.
Nitrogen dioxide
Nitrogen dioxide is the chemical compound with the formula NO2. It is one of several nitrogen
oxides. NO2 is an intermediate in the industrial synthesis of nitric acid, millions of tons of which
are produced each year. This reddish-brown toxic gas has a characteristic sharp, biting odor and
is a prominent air pollutant. Nitrogen dioxide is a paramagnetic, bent molecule with C2 point
group symmetry.
Preparation of nitrogen dioxide gas

32
The gas can be prepared by heating lead nitrate Pb (NO3)2
Properties of nitrogen dioxide gas
1. Soluble in water and form two acid nitrous acid- HNO2 and nitric acid.
2. Have irritating chocking smell.
3. Its acidic gas turns dump blue litmus red.
4. It is reddish brown yellow gas.
5. It is less dense than air.
Chemical properties of nitrogen dioxide
They include;
1. React with base
2. Action of heat.
On strong heating the gas nitrogen dioxide decompose to produce nitrogen monoxide and
oxygen.
Commercial uses of nitrogen dioxide

33
1. Used in preparation of nitric acid
2. Used in manufacturing of the salt.
Differences between nitrogen monoxide and nitrogen dioxide
Nitrogen dioxide Nitrogen monoxide
It is soluble in water It is insoluble in water
It is reddish brown yellow in color It is colorless
It has irritating chocking smell It has no pungent smell
It is acidic in nature It is neutral in nature
It change dump blue litmus to red It doesn’t change litmus.
Don’t burn air It burn in air

Non metals and their compounds

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