2nd term SS 2.pptx

HYDROGEN AND ITS COMPOUND
Hydrogen is the lightest among the
elements making up about 1% of the
earth’s crust. It is mainly found in the
combined states with other elements in
forms such as water, acids, organic
compounds and petroleum products

ELECTRONIC CONFIGURATION AND
POSSIBLE OXIDATION OF HYDROGEN
The atomic number of hydrogen is one.
The electronic configuration of hydrogen
is 1s1. It has oxidation number of +1 and
-1

UNIQUE POSITION OF HYDROGEN IN THE
PERIODIC TABLE
Hydrogen has a single electron in its
outermost shell hence it is conveniently
placed in group 1. Hydrogen can also be
placed in group 7 because it gain one electron
to form hydride ion (H-) similar to halogens
(group 7 elements) forming halides ion.

ISOTOPES OF HYDROGEN
There are three isotopes of Hydrogen namely
1.Hydrogen or protium [1
1H]
2.Deuterium or heavy hydrogen [2
1H or D]
3.Tritium [3
1H or T]
These Isotopes have relative atomic masses of 1,
2 and 3 respectively.
Deuterium forms an oxide D2O similar to water
and it is called heavy water. Tritium is radioactive.

LABORATORY PREPARATION OF
HYDROGEN
Hydrogen can be prepared in the
laboratory in the following ways
1.Action of dilute acid on Zinc
2.Action of cold water on active metals
3.Action of steam on Iron

Action of dilute acid on Zinc
Dilute tetraoxosulphate (VI) acid or hydrochloric
acid reacts with zinc granules to liberate hydrogen
gas.
Zn(s) + H2SO4(aq) ZnSO4 (aq) +H2 (g)
Zn(s) + 2HCl (aq) ZnCl2 (aq) + H2 (g)
The hydrogen gas is dried by passing it through
Calcium chloride or conc. H2SO4 and collected by
downward displacement of air.

Laboratory Preparation of hydrogen

NOTE:
Dilute trioxonitrates (v) acid is not used
in the preparation of hydrogen because
of its strong oxidizing property and it
produces water instead of hydrogen on
reaction with metals.

Action of cold water on active metals
Sodium reacts vigorously with cold water
liberating hydrogen gas while Potassium
reacts explosively with cold water to
liberate hydrogen gas.
Na(s) + H2O(l) → NaOH(aq) + H2(g)
K(s) + H2O(l) → KOH(aq) + H2(g)

Preparation of hydrogen by action of sodium on cold water
l

Action of steam on Iron
Iron reacts with steam at red heat
liberating hydrogen. Triiron tetraoxide is
also formed.
3Fe(s) + 4H2O(g) ⇌ Fe3O4(s) + 4H2(g)

INDUSRIAL PREPARATION OF HYDROGEN
From water gas: This method is called the Bosch
process. In this method, steam is passed over red
hot coke (carbon) in a furnace at about 1100℃ to
produce a mixture of carbon (II) oxide and
hydrogen which is known as water gas
C(s) + H2O(g) CO + H2(g)
water gas

water gas and excess stream are passed
over a catalyst such as Iron (III) oxide
Fe2O3 or Chromium (III) oxide at a
temperature of 450℃. The products are
hydrogen and carbon (iv) oxide as shown
in the equation below.
CO + H2 + H2O (g) ⇌ CO2 (g) +2H2 (g)

From Hydrocarbon e.g. methane
Methane reacts with steam in the
presence of nickel catalyst at 800℃ and
30atm to produce synthesis gas (mixture
of carbon (II) oxide and hydrogen).
CH4(g) + H2O(g) 𝑁𝑖 CO +3H2(g)
synthesis gas

synthesis gas is dissolved in excess steam
in the presence of Iron (III) oxide Fe2O3
at a temperature of 450℃. The products
are hydrogen and carbon (iv) oxide and
the latter is removed using caustic alkalis
as shown in the equation below.
CO + 3H2 + H2O(g) 𝐹𝑒𝟐𝑶𝟑 CO2(g) +4H2(g)

By Electrolytic methods
Hydrogen is obtained as a by-product in
the electrolysis of brine. This method is
very expensive because it involves
electricity.

PHYSICAL PROPERTIES OF HYDROGEN
1. It is a colourless, odourless and tasteless gas
which burns in air with a high pitch sound.
2. It is combustible and when it burns in the
absence of gas, it burns quietly with pale blue
flame.
3. It is the lighest known gas.
4. It is 14.4 times less than air.

5. It exists as a diatomic molecule [H2 ]
6. At high pressure, hydrogen can be
liquefy at a critical temperature of -235℃
7. It is neutral to litmus and is insoluble in
water
8. It does not support combustion.
9. It has a very low boiling point.

CHEMICAL PROPERTIES
The chemical reactions of hydrogen arise from
the fact that it can donate its single electron to
form a positive ion [H+]; it can accept electron
[to give a dublet structure] to form a negative
hydride ion [H-] and it can share electron with
another atom to form a covalent molecule.

1. Ability to accept electrons to form the
negative hydride ion H- .
2Na + H2 → 2NaH
Ca + H2 CaH2
2. Reaction with halogens: hydrogen combines
directly with halogens to form halides.
F2(g) + H2(g) 2HF
Cl2(g) + H2(g) 2HCl (g)

3. Reaction with oxygen: Hydrogen burns
with a pale blue flame as it combines
with oxygen to form steam.
2H2(g) + O2(g) → 2H2O(g)
4. Combination with metals: hydrogen
combines with metals to form ionic
hydrides.
Ca + H2 CaH2

5. With Nitrogen: hydrogen combines directly
with Nitrogen in a reversible reaction to form
Ammonia. This reaction is usually catalyzed.
N2(g) +3H2(g) ⇌ 2NH3(g)
6.As a reducing agent: hydrogen is a strong
reducing agent that reduces oxides of metals
to the metal.
Fe2O3(s) + 3H2(g) → 2Fe(s) +3H2O(g)
ZnO(S) + H2(g) → Zn(s) + H2O(g)

Hydrides
When hydrogen combines with other
elements it forms hydrides
1. The hydrides of alkali and alkaline earth
metals are crystalline solids with high melting
points that conduct electricity when molten.
They react with water to liberate hydrogen
gas.
CaH2 + 2H2O Ca(OH)2 + 2H2

2. Boron and aluminium form complex
covalent hydrides which are important
reducing agents especially in organic
chemistry. Some of them are lithium
tetrahydridoaluminate (iii), LiAlH4 and
sodium tetrahydridoborate (iii), NaBH4

USES OF HYDROGEN
1. It is used in the synthesis of ammonia as shown in
the equation below
N2(g) + 3H2(g) ⇌ 2NH3(g)
The ammonia formed is used in the manufacture of
fertilizers, drugs, plastics, wares, dyes among others.
2. It is used in the hardening of vegetable and animal
oils for the manufacture of Margarine, candles, soap
among others. The hydrogenation reaction occurs at
high pressure and in presence of catalyst.

3. It is used to inflate airships and balloons.
The use of hydrogen in balloons is due to its
low density.
4. In oxy-hydrogen flames, small quantities of
hydrogen are required to produce high
temperature that can melt metals.
5.It is used in the synthesis of methanol and
HCl(aq)

TEST FOR HYDROGEN
If a lighted splint is plunged into a
gas jar of hydrogen, it gives a pop
sound.

OXYGEN AND ITS COMPOUNDS
Oxygen is the most abundant element on
earth. It occurs in nature both in the free
and combined states. Free oxygen
constitutes about 21% by volume of
atmospheric air and about 33% by volume
of dissolved air. In the combined state,
oxygen accounts for nearly 50% by mass of
the earth’s crust, the oceans and the air.

In the combined state, oxygen accounts for
nearly 50% by mass of the earth’s crust, the
oceans and the air. It is present in the
trioxosilicates (IV), trioxocrbonates (IV) and
oxides of both metals and non-metals which
make up rocks and clays. Oxygen makes up
88.9% by mass of water. The human body
contains about two-third by mass of oxygen
in the combined state.

Electronic Structure
Oxygen is an element in group (VIA) of
the periodic table. Its atomic number is
8, and has an electronic configuration of
1S
2 2S
2 2P
4

Bonding Capacity Of Oxygen
Oxygen forms a complete octet configuration by
either accepting two electrons from a donor e.g.
metals to form O2-in an ionic substance e.g.
metallic oxide, or it shares two electrons with
other atoms to form a single covalent substance
e.g. water H-O-H or as a discrete gaseous diatomic
molecule, O2 where the two atoms are joined
together by double covalent bond, O=O.

LABORATORY PREPARATION
There are two general and common methods
for the laboratory preparation of oxygen.
These include
1.Thermal decomposition of potassium
trioxochlorate (v) and hydrogen peroxide
2.Oxidation of hydrogen peroxide

1. Thermal decomposition of potassium trioxochlorate
(v) and hydrogen peroxide: oxygen is prepared by
the thermal decomposition of potassium potassium
trioxochlorate (v) in the presence of Manganese (IV)
oxide as catalyst. The gas is collected over water
and dried by passing it throughanhydrous calcium
Chloride or conc. H2SO4
2KClO3(s) 𝑀𝑛𝑂2
(
𝑠
)
KCl(s) + 3O2(g)

2. Using Hydrogen peroxide: manganese
(IV) oxide is added to Hydrogen peroxide
in the absence of heat. The hydrogen
peroxide decomposes to produce
hydrogen gas and water.
2H2O(l) 𝑀𝑛𝑂2
(
𝑠
)
2H2O(l) + O2(g)

NOTE: hydrogen peroxide can also
decompose in the presence of acidified
potassium tetraoxomanganate (VII). This
reaction is a redox reaction.
Oxygen can also be prepared by the thermal
decomposition of Lead (IV) oxide and Silver
oxide.

INDUSTRIAL PREPARATION
Oxygen is prepared industrially by the
fractional distillation of liquefied air. This
process requires two stages:
1.Liquefaction of air
2.Fractional distillation of liquid air.

Atmospheric air is purified by passing it
through caustic soda to remove Carbon(IV)
oxide, water vapour and dust. The purified air
is compressed at a pressure of about 200 atm,
heated and then cooled. It is then allowed to
expand suddenly, for further cooling. By
successive heating and cooling, the pure air
becomes liquefied at about -200℃ [73k]. This
process is called LIQUEFACTION

FRACTIONAL DISTILLATION OF LIQUID AIR
The liquefied air is then passed into a
fractionating column. On distillation, Nitrogen,
being more volatile, boils out first at - 196℃
[77k], leaving oxygen which boils at -
183℃. The oxygen produced is about 99.5%
pure. The Liquefied oxygen is compressed and
stored in steel cylinders at 100atm for medical
and industrial use.

NOTE: Oxygen, like hydrogen is obtained during the
electrolysis of dilute H2SO4.
PHYSICAL PROPERTIES
1. Pure oxygen is colourless, odourless and tasteless.
2. It is slightly denser than air.
3. It is slightly soluble in water.
4. It is neutral to litmus paper.
5. It boils (turn to gas) at -1830C and solidifies at -2250C

CHEMICAL PROPERTIES
(a) Reaction with metals: Metals such as Na, K,
Ca, Mg, Al, Zn burn brightly in oxygen to produce
basic oxides with the exception of metals such as
silver, Gold and platinum. The oxides dissolve in
water to form alkalis e.g.
4Na(s)+ O2(g) 2Na2O(s)
Na2O(s) + H2O(l) 2NaOH(aq)

(b) Reaction with non-metals: Some burning
non-metals such as carbon, sulphur, and
Phosphorus burns in oxygen to form acidic
oxides or acid anhydrides which when
dissolved in water form acidic solutions. E.g
C(s) + O2(g) CO2(g
CO2(g) + H2O(l) H2CO3(aq)

(c ) combustion of hydrocarbons: oxygen
supports the combustion of hydrocarbons to
produce carbon (IV) oxide and water.
CH4(g) + 2O2(g) CO2(g) +2H2O(l)
(d) Oxidation: oxygen oxidizes carbohydrates
as we breathe it in to release energy and
carbon (IV) oxide.

Test for oxygen
Oxygen is identified by its ability to rekindle a
glowing splint. Nitrogen (I) oxide does this as
well but is distinguished from oxygen by its
pleasant sickly smell while oxygen is
odourless. Oxygen also reacts with
nitrogen(II)0xide to give brown fumes of
nitrogen(IV)oxide which nitrogen(I)oxide does
not do.

ALLOTROPE OF OXYGEN
Ozone is the only allotropic form of oxygen. It exists as
a triatomic molecule O3(g) . It can be prepared from
oxygen by the silent electrical discharge. It exhibits the
following chemical properties:
(a)As an oxidizing agent: it oxidizes lead (II) sulphide
and hydrogen sulphide to tetraoxosulphate (VI)
(b)Ozone is more reactive. It decomposes into oxygen
on heating
(c)It liberates iodine from potassium iodide in acidic
solution.

OXIDES
An oxide is formed when an element combine with
oxygen. They are binary compounds containing two
elements only. They are classified into
1. Acidic oxide
2. Basic oxide
3. Amphoteric oxide
4. Neutral oxide
5. Higher oxide

Acidic oxides: These are formed by non-
metals and they dissolve in water to form
acidic solution. They also react with base to
form salt and water. Examples of acid oxides
are P5O10, NO2, SO2, SiO2 etc. They are called
acid anhydride. E.g.
CO2 +NaOH Na2CO3 + H2O
SO3 + 2KOH K2SO4 + H2O
NOTE: SiO2 is insoluble in water

Basic oxides: They are metallic oxides that are
basic in nature because they react with acids
to form salt and water only. Examples are
K2O, MgO, CaO, Li2O etc. Soluble basic oxides
are called alkali. E.g.
Na2O(g) + H2O(l) 2NaOH(aq)
Na2O(s) + 2HCl(aq) 2NaCl(aq) + H2O(l)

Amphoteric oxides: These are oxides of
metals that behave both like acidic and basic
oxides. They react with both acid and base to
produce salt and water.
E.g. ZnO, Al2O3, PbO etc
ZnO(s) + H2SO4(aq) ZnSO4 + H2O(l)
ZnO + 2NaOH + H2O Na2Zn(OH)4

Neutral oxides: They are neither acidic nor basic in
character. They are neutral to litmus paper. E.g. water,
CO, N2O
Higher oxides: they contain a higher proportion of
oxygen than ordinary oxides. They are classified into
peroxides, dioxides and mixed oxides.
Peroxide oxides: These are higher oxides where O-O
bond is present. They give hydrogen peroxide when
reacted with a dilute acid. E.g. Barium peroxide BaO2,
Calcium peroxide CaO2, Sodium peroxide Na2O2,
Hydrogen peroxide H2O2 (most common).

USES OF OXYGEN
1. It is used in oxy-hydrogen for welding and
cutting of metals. A mixture of hydrogen and
oxygen can burn to produce a temperature of
2,500℃
2. In oxy- ethyne [oxy-acetylene] flame for welding
and cutting of steel. The reaction is highly
exothermic.
3. In respiration by plants and animals.

4. As breathing aids in hospitals, high altitude
flying and sea- dividing.
5. In steel production by Linz- Donawitz [L-D]
process.
6. Liquefied oxygen is used in explosives, and as a
rocket fuel
Hydrogen peroxide-preparation,properties,and
uses
Ozone and ozone layer

HALOGENS
Halogens are found in group (VII) of the
periodic table and consist of fluorine (F),
Chlorine (Cl), Bromine (Br), Iodine (I) and
Astatine (At).
The name halogen is of Greek origin meaning
salt formers because they readily form salts
from metals.

Elements in this group are the most reactive
non-metals. Due to their reactivity, halogens
are not found freely in nature but in combined
state with metals forming salts. Flourine is the
most reactive non-metal in the group.
Chlorine is the most important member and is
similar to Iodine and Bromine. Astatine is
radioactive.

ELECTRONIC CONFIGURATION AND SOME PROPERTIES OF THE HALOGENS
AND GRADATION OF THE PROPERTIES
Element Atomic
Number
Electronic
configuration
Atomic
radius
Ionic
radius
Electro-
negativity
Electro-
affinity
Atomic
Mass
Fluorine [F] 9 [2,7] IS
22S
22P
5 0.072 0.136 4.0 -3.35 19.0
Chlorine
[Cl]
17 [2,8,7] IS
22S
2
2p
63S
23P
5
-0.099 0.181 3.0 -3.61 35.5
Bromine
[Br]
35 (2,8,18,7)
..3S
2
3P
63d
104S
2
-0.114 0.195 2.8 -3.36 80.0
Iodine
[I]
53 (2,8,18,18,7)
....4S
2
4P
64d
10
0.133 0.216 2.5 -3.06 126.9
Astatine
[At]
85 (2,8,18,32,18,7)
..5S
25P
65d
106S
26P
5
---- ---- 2.2 ---- 210

PHYSICAL PROPERTIES OF THE HALOGENS
Element Fluorine[F] Chlorine[Cl] Bromine [Br] Iodine[I]
Atomic number 9 17 35 53
Relative Atomic
Mass
19 25.5 79.9 126.9
Physical state at
20℃
Gas Gas Liquid Solid
Colour Pale yellow Greenish
yellow
Dark-red Black
Density(g/cm3) ------ 1.9 3.2 4.9
Melting point
(℃). Boiling point
(℃)
-220 -101 -7 113
Solubility in
water(g per 100g
of water)
Reacts
readily with
water
0.59 3.6 0.018

Summarily, the halogens have the following
characteristics
1. They are powerful oxidizing agents
2. They are diatomic molecules covalently bonded
together
3. They are reactive non metals
4. They are usually coloured and the colouration
deepens down the group
5. At ordinary temperature, chlorine and fluorine are
gases, Bromine is a liquid while Iodine is a solid.
Astatine is radioactive

CHEMICAL PROPERTIES AND GRADATION DOWN
THE GROUP
The chemical reactivity of the halogens decreases
down the group from fluorine to iodine. The
halogens are good oxidising agents and their
oxidising power decreases from fluorine to iodine.
The halogens try to attain the stable noble gas
configuration either by sharing electrons or by
accepting electrons. The following reactions
illustrate the reactivity trend among the halogens.

1. Halogens can react with metals as shown in the equations below
2Na(s) + F2(g) 2NaF(s)
2. Halogens react with non-metals as shown below
C(s) + 2F2(g) CF4(s)
Xe+ 2F
400.8𝑎𝑡𝑚
XeF4(s)
P4(s) + 6I2(s) 4PI3(s)
P4(s) + 6Cl2(g) 4PCl2(s)
P4(s) + 10Cl2(g) 4PCl5(s)
H2(g) + F2(g) 2HF(g)
H2(g) + Cl2(g) 2HCl(g)
Br2(g) + H2(g) 2HBr(g)
I2(g) + H2(g) 2HI(g)
The order of reacting is F2> Cl2>Br2>I2

The acid strength of the hydrogen halides follow the
order:
HI>HBr>HCl>HF
The stability of Hydrogen halides decreases from
chlorine to iodine: HCl>HBr>HI.
Chlorine is now an electron acceptor and so are other
halogens and in so doing acts as oxidising agents.
2Na(s) + Cl2(g) 2Na+Cl-
(s)
The order of decreasing power as oxidising agents is F2>
Cl2>Br2>I2.

3. More reactive halogen displaces the less reactive
from aqueous solution
Cl2 + 2Br- 2Cl- + Br2
Cl2 + 2I- 2Cl- + I2
Br2 + 2I- 2Br- + I2
Iodine is a weak oxidizing agent.

4. Reaction with water.
2F2(g) + 2H2O(l) 4HF(g) + O2
Cl2(g) + H2O(l) HCl(aq) + HClO(aq)
5. Reaction with Alkalis
3Cl2(aq) + 6NaOH(aq) NaClO3(aq) + 5NaCl +
3H2O(l)
Bromine and iodine react in a similar manner.

USES OF HALOGENS AND THEIR COMPOUNDS
1. Fluorine is used in rocket propulsion and in making
Uranium(iv) fluoride. It is also used in making
fluorocarbon compounds which are used as refrigerant,
aerosol propellant, anaesthetics and fire extinguisher,
fluids and polytetrafluoroethene[PTFE] with the trade
name Teflon in making valves, seals, gaskets, electrical
insulators among others.

2. CHLORINE: Is used as oxidant in the
manufacture of bromine and as bleach and
germicide in treatment of water. Chlorine is
used in production of HCl (hydrogen chloride)
in commercial quantity, in the production of
fabrics and papers and also insecticides. (e.g
DDT)

3. Iodine dissolved in alcohol or potassium
iodide [KI] is used as antiseptic for cuts and
scratches.
4. Bromine finds application in the
manufacture of dyestuffs and in making silver
bromide used in photographic materials.

CHLORINE
Chlorine is the most important element in the Halogen
family.it was discovered in 1774 by Scheele and named
Chlorine(greenish yellow) by Davy in 1810.
Chlorine does not occur free in nature due to its
reactivity. It is found in the combined state as Chlorides.
The most abundant of the chlorides is sodium chloride
and are found both in the sea and as salt deposits.

LABORATORY PREPARATION
Chlorine is usually prepared by the oxidation
of concentrated hydrochloric acid with a
strong oxidizing agent such as manganese (IV)
oxide, potassium tetraoxomanganate (VII) or
Lead (IV) oxide.
MnO2(s) + 4HCl (aq) MnCl2(aq) + 2H2O (l) +
Cl2 (g)

Chlorine is prepared industrially by the
electrolysis of brine and molten metallic
chlorides of sodium, Magnesium or Calcium.
The Kellener solvay cell is specifically used for
this purpose.
The chlorine is then liquefied and stored
under pressure in steel cylinders.

PHYSICAL PROPERTIES OF CHLORINE
1. It is a greenish-yellow gas with an unpleasant
choking smell.
2. It is denser than air.
3. It is slightly soluble in water.
4. It is poisonous.
5. It can be easily liquefied under a pressure of
6atm

CHEMICAL PROPERTIES
1. Displacement reaction
Chlorine displaces other halogens from solutions of
their acids and salts with exception of fluorine.
Cl2(g) + 2Nal(g) 2NaCl(aq) + I2(g)
2. As an oxidizing agent
Chlorine is a strong and powerful oxidizing agent due to
its ability to accept electrons from reducing agents to
form chloride.

i. With trioxosulphate (IV): if chlorine is
bubbled through freshly prepared
trioxosulphate (IV) acid solution the
trioxosulphate (IV) acid is oxidized to
tetraoxosulphate (VI) acid.
H2SO3(aq) +H2O(l) + Cl2(g) H2SO4(aq) + 2HCl(aq)

ii. With Iron (II) salts: chlorine oxidizes green
solution of Iron (II) chloride to yellow Iron (III)
chloride.
2FeCl2(aq) + Cl2(g ) 2FeCl3(aq)
iii. With hydrogen sulphide : Chlorine oxidizes
hydrogen sulphide to yellow sulphur.
Cl2 + H2S 2HCl + S

3. As a Bleaching agent
Chlorine bleaches by oxidation. Chlorine does not
bleach the printer’s ink because it contains carbon.
When chlorine dissolves in water, it forms hydrochloric
acid and oxochlorate (i) acid (chlorine water).
Cl2 + H2O HCl + HOCl
In bleaching oxochlorate (i) acid gives out oxygen atom.
The oxygen atom then bleach the material by oxidation.
This bleaching action is permanent because the material
cannot be re-oxidize by atmospheric oxygen

4. Reaction with slaked lime
When chlorine reacts with slaked lime
(calcium hydroxide), bleaching powder is
formed
Cl2(g) + Ca(OH)2(aq ) CaOCl2.H2O

Test for chlorine
1.With starch iodide paper: chlorine changes
the colour of starch iodide paper to blue due
to displacement of iodine.
Cl2(g) + 2KCI(aq ) 2KCl(aq) + I2(s)
2. With moist litmus paper: chlorine bleaches
moist litmus paper first by turning blue litmus
red because it is an acidic gas. But with red
litmus it bleaches it directly.

Hydrogen chloride
Hydrogen chloride is prepared in the laboratory by the
action of concentrated tetraoxosuphate (vi) acid on
sodium chloride. The gas is dried by passing it through
concentrated tetraoxosulphate (vi) acid and collected by
downward displacement of air.
NaCl(s) + H2SO4(aq) Na2SO4(aq) + HCl(g).
Dissolution of hydrogen chloride in water gives
hydrochloric acid.

PHYSICAL PROPERTIES OF HYDROGEN CHLORIDE
1. It is a colourless gas with sharp irritating smell.
2. It is denser than air.
3. It is very soluble in water.
4. It turns moist blue litmus paper red.
5. It is soluble in non-polar solvent like
methylbenzene.

CHEMICAL PROPERTIES OF HYDROGEN
CHLORIDE
1. It reacts with active metals to liberate
hydrogen gas e.g
Zn(s) + 2HCl(g) ZnCl2(s) + H2(g)
2. It reacts with ammonia to form a dense
white fume of ammonium chloride.
HCl(g) + NH3(g) NH4Cl(s)

3. Hydrochloric acid (obtained when hydrogen chloride
dissolved in water) reacts with metallic trioxocarbonate
(iv) or metallic hydrogentrioxocarbonate (iv) to liberate
carbon (iv) oxide.
2HCl(aq) + Na2CO3(s) 2NaCl(aq) + H2O(l) + CO2(g)
4. Hydrochloric acid (obtained when hydrogen chloride
dissolved in water) precipitates silver chloride from
solution.
AgNO3(aq) + HCl(aq) AgCl(s) + HNO3(aq)

Test for solubility of HCl,Preparation, Properties of HCl

USES OF HYDROGEN CHLORIDE GAS AND
HYDROCHLORIC ACID
1. It is used for the synthesis of vinyl chloride
which is used for the manufacture of plastics.
2. Hydrochloric acid is used for pickling of metals
prior electroplating.
3. Hydrochloric acid is for the synthesis and
analysis of many compounds.

NITROGEN AND ITS COMPOUNDS
Nitrogen and other group VA elements are non-
metals and show two common valences of 3 and 5.
They are electron acceptors and both form several
acidic oxides. They also form similar hydrides and
chloride e.g. Nitrogen and phosphorus form N2O3
and P4O6, N2O5 and P4O10, NH3, PH3, NCl3 and PCl3

Property Nitrogen
N
Phosphorous
P
Arsenic
As
Antimony
Sb
Bismuth
Bi
Atomic number 7 15 33 51 83
Outer electron
configuration
2s22p6 3s23p3 3d104s2 4p3 4d104s25p3 4f15d106s26p3
Atomic mass 14.0067 30.9738 74.9216 121.75 208.980
Boiling point
Melting point
-196
-210
280w
44w
2.34r
610sr
817pgr
5.73gr
1380
630gr
6.7gr
1560
271
9.8
Density(gcm-3) 0.81 1.82w
W = white, gr = grey, p = 28 atmospheres, r = red, s = sublimes

Nitrogen
Nitrogen occurs chiefly as a free element in air, making
up about 78% by volume of the atmosphere. Free
Nitrogen in the air is necessary because it dilutes the
oxygen to the point where combustion, respiration and
oxidation of metals are reasonably slow. Nitrogen
occurs abundantly in the earth’s crust as trioxonitrates
(V) of sodium, calcium and ammonium salts. It can also
be found in organic matter such as proteins, urea and
vitamin B compounds.

LABORATORY PREPARATION OF NITROGEN
Since nitrogen makes up a large percentage of air, it can be
obtained from air by removing the other constituents.
Atmospheric air is passed through aqueous NaOH, in order to
absorb carbon (iv) oxide.
2NaOH(aq) + CO2(g) Na2CO3(aq) + H2O
It is then passed over red-hot copper metal in a combustion
tube in order to remove oxygen.
Cu(s) + O2(g) 2CuO(s)
The residual gas collected over water is Nitrogen
contaminated with about 1% of noble gases and it is denser
than pure nitrogen .

Other methods of preparing nitrogen in the laboratory include:
1. Preparation from Ammonia gas.
CuO(s) + NH3(g) 3Cu(s) + H2O(g) + N2(g)
2. From ammonium dioxonitrate (iii)
NaNO2(aq) + NH4Cl(aq) NH4NO2(aq) + NaCl(aq)
NH4NO2(aq) N2(g) +2H2O(l)
3. From ammonium heptaoxodichromate (V
(NH4)2Cr2O7(s) N2(g) +Cr2O3(s)+ 4H2O(l)
4. From dinitrogen (I) oxide: when dinitrogen (I) oxide is passed over
red-hot copper, the gas is reduced to nitrogen.
N2O(g) + Cu(s) CuO(s) + N2(g)

Industrial preparation of nitrogen
Nitrogen is produced commercially from fractional
distillation of liquefied air.
Carbon (iv) oxide is removed from air. This air is
liquefied by subjecting it to successive compression and
cooling processes.
Nitrogen is collected first after distillation before oxygen
because it has a lower boiling point than oxygen.
Nitrogen is stored in steel in steel cylinders and sold as
liquid nitrogen or as the compressed gas.

PHYSICAL PROPERTIES
1. Pure nitrogen is colourless, odourless and
tasteless
2. It is insoluble in water
3. It is lighter than air.
4. It has no effect on litmus paper

CHEMICAL PROPERTIES
1. Combustion: Nitrogen gas is generally unreactive at ordinary
temperatures and pressures. It does not burn and does not support
burning. It extinguishes a lighted splint.
2. Reaction with metals. Magnesium ribbon burns in air to produce a
mixture of magnesium oxide and magnesium nitride, a white solid.
2Mg(s) + O2(g) MgO(s)
3Mg(s) + N2(g) Mg3N2(s)
Magnesium nitride decomposes on addition of water to liberate
ammonia gas
Mg3N2(s) + 6H2O(l) 3Mg(OH)2(aq) + 2NH3(g)

3. Reaction with non-metals:
(a) under conditions of high temperatures and pressures, and
in the presence of finely divided iron as catalyst, nitrogen
combines with hydrogen, to produce ammonia. The reaction
in reversible.
N2(g) + 3H2(g) ⇋
2NH3(g)
(b) Nitrogen combines with oxygen at high temperature and
pressure to produce Nitrogen (ii) oxide, an unstable colourless
and odourless gas. The reaction is reversible.
N2(g) + O2(g) ⇌ 2NO(g)

USES OF NITROGEN
1. In the Haber process in production of ammonia.
2. Liquid nitrogen is used as a refrigerant.
3. Due to its inertness and because it does not support combustion it
is used for the following purposes:
(a) For prevention of fire
(b) As a diluents: to reduce combustion (nitrogen is responsible for
the low heating value of producer gas, and the low combustion rate of
atmospheric oxygen.
(c) Nitrogen provides inert atmosphere for food processing and
packaging, and during some chemical reactions.

COMPOUNDS OF NITROGEN
OXIDES OF NITROGEN
The common oxides of nitrogen are:-Dinitrogen(i)
oxide, N2O- Nitrogen (ii) oxide, NO. Nitrogen (iv)
oxide, NO2, others are dinitrogen(iii) oxide, N2O3, a
pale blue liquid at room temperature and
dinitrogen(iv) oxide, N2O4, a yellow liquid below
the room temperature and dinitrogen(v) oxide
N2O5 that exists as an unusual white solid at room
temperature.

AMMONIA
Ammonia is a hydride of nitrogen. It is produced in nature
when nitrogenous matter decays in the absence of air. The
decomposition may be due to heat or putrefying bacteria.
Laboratory preparation
By the action of heat on any ammonium salt with a non
volatile base. E.g. ammonium chloride and calcium hydroxide.
NH4Cl(s) + Ca(OH)2(aq) CaCl2(s) + 2H2O(l) + 2NH3(g)
The gas is dried using quicklime.

INDUSTRIAL PREPARATION OF AMMONIA
Ammonia is manufactured from its constituent elements
hydrogen and nitrogen by a process known as Haber process.
A mixture of dried nitrogen and hydrogen in the ratio 1:3 by
volume is subjected to a high pressure of 200 – 250
atmosphere, at about 450℃ and in the presence of finely
divided iron as the catalyst; the reaction is exothermic and
reversible.
N2(g) + 3H2(g)
𝐹𝑒. 450𝑜
𝐶
⇌
250𝑎𝑡𝑚
2NH3(g)

Under these conditions, about 20% ammonia is
produced.
Successive heating and cooling under pressure
liquefy the ammonia gas produced and the
uncombined nitrogen and hydrogen are recycled.

PHYSICAL PROPERTIES
1. It is a colourless gas with choking smell.
2. It is very soluble in water.
3. In large quantities, ammonia is poisonous.
4. It is 1.7 times less dense than air
5. It turns moist red litmus paper red.

CHEMICAL PROPERTIES OF AMMONIA
1. Ammonia burns in oxygen with a greenish-yellow
flame producing steam and nitrogen
4NH3 + 3O2 2N2 + 6H2O
In the presence of platinium- rhodium catalyst,
Ammonia reacts with excess air to form nitrogen (ii)
oxide.
4NH3 + 5O2
𝑃𝑙𝑎𝑡𝑖𝑛𝑖𝑢𝑚−𝑟ℎ𝑜𝑑𝑖𝑢𝑚
4NO + 6H2O
2. Ammonia reduces copper (ii) oxide
2NH3 + 3CuO 3Cu + 3H2O + N2

It also reduces chlorine to produce hydrogen chloride and
nitrogen. Then the hydrogen chloride reacts with the excess
ammonia to produce a dense white fume of ammonium
chloride.
2NH3(g) + 3Cl2(g) 6HCl(g) + 2N2(g)
6NH3(g) + 6HCl(g) 6NH4Cl(s)
Overall equation
8NH3(g) + 3Cl2(g) 6NH4Cl(s) + N2(g)
But if the chlorine is in excess nitrogen (III) chloride an
explosive and oily liquid is produced.
NH3(g) + 3Cl2(g) NCl3(l) + 3HCl(g)

3.Aqueous ammonia precipitates many metallic ions
from solution as insoluble hydroxides. It is useful in the
precipitation of amphoteric hydroxides.
FeSO4(aq)+ 2NH3(aq)+ 2H2O(l) Fe(OH)2(s) + (NH4)2SO4(aq)
green ppt
4. Reaction with carbon (IV) oxide: ammonia reacts with
carbon (IV) oxide at 150℃ and a high pressure of 150
atm to produce urea.
2NH3(g) + CO2(g) (NH2)2CO(s) + H2O(l)
urea

USES OF AMMONIA
1. Ammonia is used in refrigeration, since it can be liquefied easily.
2. In softening water used in laundry, to prevent wastage of soap.
3. As a domestic cleaner- to neutralize as in sweat.
4. In treating insect stings- to neutralize methanoic acid injected by
the insects
5. In the production of plastics by polymerization
6. In the production of ammonium salts. Fertilizer and trioxonitrate(v)
acid.
7. As a precipitating regent for the identification of cations in
solutions.

TEST FOR AMMONIA
Place any ammonia salt in a test tube, add any
alkali and heat the mixture. A colourless gas with
choking smell is given off. The gas turns moist red
litmus paper to blue and produces dense white
fumes with hydrogen chloride gas (from
concentrated HCl reagent bottle. The gas is
ammonia

COMPOUNDS OF NITROGEN (TRIOXONITRATE (V)
ACID )
LABORATORY PREPARATION OF TRIOXONITRATE
(V) ACID
Trioxonitrate (v) acid is prepared by heating solid
sodium trioxonitrate (v) with concentrated
tetraoxosulphate (vi) acid.
NaNO3 + H2SO4 HNO3 + NaHSO4

The apparatus used is completely made of glass
because rubber and cork are attack by trioxonitrate
(v) acid. The trioxonitrate (v) acid produced in this
case is yellow in colour due to the slight
decomposition of the acid by heat to produce
reddish-brown gas (nitrogen (iv)) which then
dissolve in the acid to impart the yellow colour.
4HNO3 H2O + 4NO2 + O2

PHYSICAL PROPERTIES
1. It is corrosive and readily destroys organic
materials.
2. It turns blue litmus paper red.
3. It is a colourless liquid which fumes in air on
exposure.

CHEMICAL PPROPERTIES
1. Reaction as an acid:
(a) It reacts with base to form salt and water
HNO3 + NaOH NaNO3 + H2O
(b) It reacts with trioxocarbonate (iv) to liberate
carbon (iv) oxide
HNO3 + CaCO3 Ca(NO3)2 + H2O + CO2

2. As an oxidizing agent:
(a) with metals: Trioxonitrate (v) acid is a strong oxidizing
agent and the oxidizing power of trioxonitrate (v) acid
depends on its concentration. In concentrated form, it
oxidizes certain metal to trioxonitrate (v) while the acid
is reduced to nitrogen (iv) oxide and in moderately
concentrated form it reduced to nitrogen (ii) oxide.
Cu + 4HNO3 Cu(NO3)2 + 2H2O + 2NO2
3Cu + 8HNO3 3Cu(NO3)2 + 4H2O + 2NO

Lead, mercury and silver react in similar way to copper
with the metal
Magnesium, zinc and iron react with dilute trioxonitrate
(v) acid to give ammonium trioxonitrate (v) e.g
4Zn + 10HNO3 4Zn(NO3)2 + 3H2O + NH4NO3
Aluminium and iron do not react with concentrated
HNO3 due to initial formation of an oxide coating on the
metal, which prevent further reaction.

(b) With non-metals: Hot concentrated
tetraoxosulphate (vi) acid oxidizes non-metals to
their oxides, which may dissolve in water to for the
corresponding acids.
C + 4HNO3 CO2 + 2H2O + 4NO2
S + 6HNO3 H2SO4 + 2H2O + 6NO2
P + 5HNO3 H3PO4 + H2O + 5NO2
I2 + 10HNO3 2HIO3 + 4H2O + 10NO2

USES OF TRIOXONITRATE (V) ACID
1. It is an important raw material for the
manufacture of many useful trioxonitrate (v) salts
and organic nitro-compounds such as those used
for making dyes. Explosives, fertilizers and drugs.
2. It is used for making of nylon and terylene.
3. It is useful oxidizing agent for many purposes in
the laboratory.

TRIOXONITRATE(V) SALTS-
All trioxonitrate (v) salts are decomposed by heat.
(a) Trioxonitrate (v) of sodium and potassium are
decomposed to the dioxonitrate (iii) compound and
oxygen.
NaNO3
ℎ𝑒𝑎𝑡
NaNO2 + O2
(b) Trioxonitrate (v) of Zn, Cu, Fe, Mg, Ca, Pb and Al
are decomposed to the oxide of the metal, nitrogen
(iv) oxide and oxygen.

2Pb(NO3)2
ℎ𝑒𝑎𝑡
2PbO + 4NO2 + O2
(c) Trioxonitrate (v) of Hg and Ag are
decomposed to the metal, oxygen and nitrogen
(iv) oxide because the metal oxides are
unstable to heat
2AgNO3
ℎ𝑒𝑎𝑡
2Ag + 2NO2 + O2

TEST FOR TRIOXONITRATE (V) IONS
BROWN RING TEST
Acidify an unknown solution with dilute
tetraoxosulphate (vi) acid, then add some freshly
prepared iron (ii) tetraoxosulphate (vi) and shake. Then
keep the test tube in a slanting position and carefully
add some concentrated tetraoxosulphate (vi) acid down
the side of the test tube. A brown ring will be formed at
the junction of the two liquid layers. This shows that the
unknown contains trioxonitrate (v) ions.

SULPHUR AND ITS COMPOUNDS
GENERAL PROPERTIES OF GROUP VIA ELEMENT
Group VIA elements usually called oxygen family.
They are made up of oxygen (O), sulphur (S),
Selenium (Se) Tellurium (Te) and Polonium (Po)
Sulphur is the second member of the group and it
is a solid. Members of the group has the following
properties.
1. They have six valence electrons.

2. They gain two electrons usually from group I and
II metals to attain octet structure or form negative
divalent ions e.g. S2, O2
3. They form covalent compounds with non-metals
e.g. H -----O--------H in water and H----S-----H ,
hydrogen sulphide
4. Their oxidation states range from -2 in its
compounds except peroxides like H2O2 and K2O2
where oxygen has -1

ELECTRON STRUCTURE OF SULPHUR
Sulphur is represented with a symbol S, it has
electrons. The electronic configuration is 2, 8,6
or Is2 2s2 2p6 3s23p4. It is a P block element,
thus have its valence electrons in the p orbital

ALLOTROPES AND USES OF SULPHUR
Allotropy is the existence of an element in two or more
different forms in the same physical state. Sulphur
exhibits the phenomenon known as allotropy. The
allotropes of sulphur are Rhombic (or ∝ - sulphur),
Monoclinic or prismatic (𝛽- sulphur), Amorphous
sulphur (S- sulphur), Plastic sulphur
Rhombic and monoclinic sulphur are crystalline in
nature and are actually the important allotropes.

Rhombic sulphur (α- sulphur): free sulphur exists
as allotrope at below 96℃. It has octahedral
structure made up of S8 molecules. The colour is
brightly yellow and has a melting point of 113℃
and density of 2.08gcm-3

Rhombic sulphur is prepared by allowing a saturated
solution of sulphur in carbon (IV) sulphide (carbon
disulphide) in a test tube kept below 95𝑜
C to
evaporate slowly. Octahedral crystals will gradually
deposit. This preparation should be done in a fume
cupboard because of the poisonous and flammability of
carbon disulphide

Monoclinic sulphur (β- sulphur): It is the only stable
allotrope between 96℃ and 119℃ and consists of long,
thin and needle shaped. The colour is amber. At room
temperature it changes to rhombic sulphur crystals.
Monoclinic sulphur has a melting point of 119oC and
density of 1.98gcm-3 .Monoclinic sulphur is obtained by
cooling molten sulphur. Powdered sulphur is heated in a
crucible till it melts into amber -coloured liquid. More
sulphur is added, heated and stirred at the same time.

This process is repeated until the crucible is filled with
molten sulphur. It is then allowed to cool while a hard
crust formed at the top. Piercing one or two holes
through the crust and pouring off the remaining molten
sulphur reveals a needle shaped crystals of monoclinic
sulphur deposited on the sides of the crucible.
Rhombic sulphur
𝑏𝑒𝑡𝑤𝑒𝑒𝑛 96𝑜𝐶 𝑎𝑛𝑑 119𝑜𝐶
⇌
𝑏𝑒𝑙𝑜𝑤 96𝑜 𝐶
monoclinic sulphur

Amorphous sulphur: it is pale – yellow in colour and
has no regular crystalline shape. It is prepared as a
deposit when hydrogen sulphide is bubbled through
water for a long time and the saturation exposed to air.
It can also be prepared by the action of dilute HCl acid
on trioxothiosulphate (v) solution
H2S (g) + O2(g) → 2H2O(g) + S(g)
S2O3
2-
(aq) +2H+
(aq) → 2H2O(l) + SO2(g) + S(s)

Plastic sulphur: this is prepared by heating yellow
sulphur until it boils in a test tube. The boiling
sulphur is poured into cold water is seen to roll up
into yellow ribbons which look like a plastic. It
changes to rhombic sulphur after sometime. It is
therefore said to be unstable.

PHYSICAL PROPERTIES
1. Sulphur is a yellow solid existing in crystalline or
amorphous
2. It is non-metallic and exhibits allotropy.
3. It sublimes to give flowers of sulphur.
4. It is a non-conductor of heat and electricity.

5. When heated in the absence of air, roll sulphur
undergoes the following changes:
(i)at 1150C it forms amber - coloured liquid.
(ii)at about 1150C the liquid becomes dark and
discourse.
(iii)near its boiling point, becomes mobile again and
reddish brown in colour.
(iv)at its boiling point of 440C it gives off a brown
vapour, condensing this vapour on a cold surface gives
flowers sulphur.

CHEMICAL PROPERTIES
(i) Combustion in air: sulphur burns in air with a blue
flame to produce sulphur (iv) oxide.
S(s) + O2(g) → SO2(g)
(ii) Reaction with metals: sulphur combines directly
with metals to give the corresponding anhydrous
sulphides
Fe(s) + S(s ) → FeS(s)
2Cu(s) + S(s) → Cu2S(s)

(iii) Reaction with non- metals: sulphur reacts with
coke in the furnace to produce carbon (IV) sulphide
C(s) +2S(s) → CS2(l)
(iv). As a reducing agent: when powdered sulphur is
warmed with concentrated H2SO4, it is oxidized to SO2,
while its acid is reduced to SO2.
S(s) + H2SO4(aq) → 2H2O(l) + 3SO2(g)

USES OF SULPHUR
1. Used in the production of tetraxosulphate (VI) acid.
2. Used in vulcanizing rubber. It makes raw rubber to become
hard, tough and elastic and hence, suitable for making tyres
3. Used in the production of carbon (iv) sulphide used as a
solvent and insecticides
4. It is used in gun powder and matches, dyes tuffs fungicides,
ointment and germicides
5. Also used in the production of calcium hydrogen
trioxosulphate (iv), Ca(HSO3)2 used in bleaching wood pulp for
making news prints

COMPOUNDS OF SULPHUR
TRIOXOSULPHATE (IV) ACIDS AND ITS SALTS
Trioxosulphate (IV) acid is a dibasic acid which is
obtained by dissolving sulphur (IV) oxide in water.
LABORATRY PREPARATION OF TRIOXOSULPHATE (IV)
ACID
It is prepared by the action of dilute hydrochloric acid
on heated sodium trioxosulphate (iv) to produce
sulphur (iv) oxide which is then dissolved in water.

Equations of the reactions
(a) Na2SO3(s) + 2HCl(aq) → 2NaCl(qg) + SO2 (g) + H2O(l)
(a) SO2 (g) + H2O (l) → H2SO3 (aq)
Sulphur (IV) oxide is the acidic hydride of trioxosulphate
(IV) acid.

PHYSICAL PROPERTIES
1. It is a colourless liquid, which smells strongly of
SO2.
2. It turns blue litmus paper red.
3. It mix readily with water.
4. It has an irritating and choking smell.

CHEMICAL PROPERTIES
1. Trioxosulphate (iv) acid is a weak dibasic acid. In
the presence of a limited amount of sodium
hydroxide an acid salt, sodium hydrogen
trioxosulphate (IV) is formed.
2NaOH (aq) + H2SO3 (aq) → Na2SO3 (aq ) + 2H2O(l)

However, when the alkali, sodium hydroxide, is in
excess, the normal salt sodium trioxosulphate (IV)
is produced. That is, complete neutralisation takes
place.
NaOH(aq )+ H2SO3(aq)→ NaHSO3(aq), + H2O(l)
2. It is oxidized in air to tetraoxosulphate (VI) acid.
H2SO4(aq) + O2(g) →2H2SO4(aq)

3. Reducing properties
It is a strong reducing agent. It decolourises the purple
colour potassium tetraoxomanganate (vii) solution, and
changes the colour of potassium heptaoxodichromate
(VI) from orange to green.
4. Bleaching properties: dyes are bleached by aqueous
solution of trioxosulphate (IV) acid. It bleaches by
reduction. Material bleached by trioxosulphate(iv) acid
and sulphur (iv) oxide temporal, it can be re-oxidized by
atmospheric oxygen.

USES OF TRIOXOSULPHATE (IV) ACID
1. It is used for bleaching.
2. It is also used as a germicide.

TRIOXOSULPHATE (IV) SALTS
Trioxosulphates (IV) are the salts of trioxosulphate (IV)
acid.
LABORATORY PREPARATION OF TRIOXOSULPHATE (IV)
SALTS
1. By the action of trioxosulphate (iv) on excess alkali.
Sodium hydroxide is used.
H2SO3 (aq) + NaOH (aq) → Na2SO3(aq) + 2H2O(l)
This is a neutralization reaction.

2. Precipitation of an insoluble trioxosulphate (IV)
from the solution of its metallic salt by sulphur
(iv) oxide.
SO2 (g) +2H2O (l) +ZNNO3 (ag) → ZnSO3(s) +HNO3 (l)
PHYSICAL PROPERTIES
Most of the trioxosulphates (IV) are insoluble in
water but the trioxosulphate (IV) of calcium;
ammonium, potassium and sodium are soluble in
water.

CHEMICAL PROPERTIES
1. They liberate sulphur (IV) oxide on reaction with
aqueous hydrochloric acid e.g. Na2CO3,
Na2CO3 (aq) + 2HCl (aq) → 2NaCl (aq) + H2O (l) +SO2 (g)
2. When exposed to air trioxosulphate (IV) salts are
slowly oxidised to tetraoxosulphate (VI) salts.

3. On addition of barium chloride solution
into solution of trioxosulphate (IV) white
precipitate of barium trioxosulphate (IV)
which is soluble in dilute hydrogen chloride
(HCl) is formed.
BaCl2 (aq) + Na2SO3 (aq) → BaSO3(s) + 2NaCl (aq)

TEST FOR TRIOXOSULPHATE (IV) SALT
1. On warming with dilute hydrochloric acid sulphur
(IV) oxide is given off.
2. Add solution of barium chloride to the solution of
the substance suspected to be trioxosulphate (IV). Any
white precipitate (barium trioxosulphate (IV)), soluble
in dilute hydrochloric acid confirms the presence of a
trioxosulphate (IV) ion e.g. Na2SO3.
BaCl2 (aq) + Na2SO3 (aq) → 2NaCl (aq) + BaSO3(s)

TETRAOXOSULPHATE (VI) ACID
Tetraoxosulphate (VI) Acid is a chemical compound
with the formula H2SO4.
The industrial or commercial production of H2SO4 is
through a process known as the contact process. The
process involves:
1. The oxidation of sulphur (IV) oxide by air to sulphur
(VI) oxide using a catalyst such as vanadium (V) oxide,
V2O5.

SO2 is produced by burning sulphur or by burning pyrite, FeS2 in air.
S + O2 SO2
2SO2(g) + O2
𝑉2𝑂5
⇌ 2SO3(g)
2. The absorption of sulphur (VI) oxide, SO3 in conc. H2SO4 to form a fuming
liquid called ‘Oleum’- heptaoxosulphate (VI) acid.
H2SO4(aq)+ SO3(g)→ H2S2O7(aq)- oleum

3. The oleum is diluted with correct amount of
water to produce the conc. H2SO4.
H2S2O7(aq)+ H2O(l)→ 2H2SO4(aq)
Note: a direct absorption of SO3 in water is not
done - the reaction is violently exothermic, the
heat evolved will cause the acid to boil producing a
mist of fine drops of the H2SO4 which will fill the
environment.

PROPERTIES OF H2SO4
1. As an acid - H2SO4 is dibasic and ionizes almost
completely in solution, this makes it a strong acid.
Due to it being dibasic, it forms two kinds of salts
with alkalis.
2NaOH(aq)+ H2SO4(aq)→ Na2SO4(aq)+ 2H2O(l) and
NaOH(aq)+ H2SO4(aq)→ NaHSO4(aq)+ H2O(l)

PROPERTIES OF H2SO4
A. As an acid
1. H2SO4 is dibasic and ionizes almost completely in
solution, this makes it a strong acid. Due to it being
dibasic, it forms two kinds of salts with alkalis.
2NaOH(aq)+ H2SO4(aq)→ Na2SO4(aq)+ 2H2O(l) and
NaOH(aq)+ H2SO4(aq)→ NaHSO4(aq)+ H2O(l)

2. Action on trioxocarbonate(IV) : Carbon(IV) oxide is
liberated when H2SO4 is added onto a
trioxocarbonate(IV).
Na2CO3(aq)+ H2SO4(aq)→ Na2SO4(aq)+ H2O(l)+ CO2(g)
On a piece of marble, which has the chemical formula
CaCO3, the reaction obtained is prematurely stopped
due to the formation of the sparingly soluble salt, CaSO4,
which forms a deposit on the surface of the marble.

B. Action of dilute H2SO4 on metals Reactive metals
would displace hydrogen from dilute
tetraoxosulphate(VI) acid.
Zn(s)+ H2SO4(aq)→ ZnSO4(aq)+ H2(g)
Note:* Less reactive metals such as copper will not
displace hydrogen from dilute acid.* Cold concentrated
H2SO4 is not attacked by any metal in the complete
absence of water.

C. Conc. H2SO4 as an oxidizing agent When hot and
concentrated, the acid accepts electrons from reducing
agents such as Cu or Zn. It also oxidizes non-metals, such
as carbon and sulphur, and is reduced to SO2 in the
process.
S(s)+ 2H2SO4(aq)→ 2H2O(l)+ 3SO2(g)
C(s)+ 2H2SO4(aq)→ 2H2O(l)+ 2SO2(g)+ CO2(g)
The SO2 given off is detected using a strip of filter paper
moistened with potassium heptaoxodichromate (VI)
solution which turns orange to green - this is one of the
tests for SO2.

D. As a dehydrating agent: Tetraoxosulphate (VI) acid has
a strong affinity for water. An example is the removal of
rust from iron, and the dehydration of sugar. The sugar
becomes a black mass of carbon. Dehydration is the
removal of elements of water from a substance and
the chemical composition of the substance is changed.
C12H22O11(s) + nH2SO4(l)→ 12C(s)+ 11H2O(l)+ nH2SO4(aq))
H2SO4 is hygroscopic, that is, it absorbs water from the
air and becomes dilute.
NOTE: conc. H2SO4 burns the skin by dehydration.

TEST FOR TETRAOXOSULPHATE (VI) ION
The characteristic test for any soluble
tetraoxosulphate(VI) is the formation of white
precipitate (barium sulphate) when a solution of barium
chloride, acidified with dilute HCl, or when a solution of
barium trioxonitrate (v), acidified with dilute
trioxonitrate (v) acid is reacted with it.
BaCl2(aq)+ Na2SO4(aq)→ 2NaCl(aq)+ BaSO4(s)
Ba(NO3)2(aq)+ Na2SO4(aq)→ 2NaNO3(aq)+ BaSO4(s)- white
precipitate

USES OF TETRAOXOSULPHATE (VI) ACID
1. for the manufacture of fertilizer, such as ammonium
tetraoxosulphate (VI), (NH4)2SO4, and superphosphates.
2. for the manufacture of paints and pigments.
3. for the manufacture of artificial and natural fibre.
4. for the manufacture of metallic tetraoxosulphate (VI),
HCl, HNO3, HF, and plastics.

5. for the manufacture of detergents, dyes, explosives
and drugs.
6. for the extraction of metals, example, in pickling
(cleaning) iron and steel before plating them with tin or
zinc.
7. as electrolyte in lead-acid storage battery.
8. In petroleum refining, where to is used to wash
impurities out of gasoline and other refinery products

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