INORGANIC CHEMISTRY
(COMPOUNDS)
In the Chemical Logic and Arithmetic chapter (section-I), we have handled a lot of inorganic
equations and studied their logic, acquainted with the art of predicting reaction products, knew
several techniques to balance them. In other words, we played with the inorganic equations and
developed friendliness and familiarity with them. In this section, called Inorganic chemistry,
we shall study the chemistry of some common elements (both metals and nonmetals) and some
of their common compounds in an interactive manner. We shall try to know all the important
aspects about them, i.e how they are prepared, what are their properties and uses. Inorganic
chemistry is that branch of chemistry which deals with the study of elements and compounds
which are not derived from living beings i.e plants and animals. Substances derived from
non-living matter like the metals and minerals are called inorganic substances. Most of the
inorganic substances are available in the earth crust and sea in the form of minerals. Some
elements and compounds are found in the atmosphere. Inorganic chemistry deals with hundreds
of chemical equations. While studying an equation, we shall recall our knowledge on chemical
reactions that we have learnt in Chemical Logic and Arithmetic chapter and try to utilise them
in understanding the logic of the reaction. We know that broadly inorganic reactions are of two
types i.e metathesis and redox. In stead of looking to the products of a reaction given in the
text, first you predict the products by applying the logic. Then match with the correct products
given in the text. Some reactions give typical products which evade all logic for prediction that
we have learnt. In such cases, we shall know the products first and then apply the logic as to
know how they are formed and not why they are formed. This will make the study of inorganic
chemistry interesting. Often there is a feeling that inorganic chemistry is boring and monotonous.
Yes, this is true only when we learn it without proper logic and without doing a single experiment
in the laboratory. It is really satisfying and rewarding when we carry out a reaction in a test tube,
and observe the changes. For example, when we heat ammonium dichromate, which is a deep
orange coloured powder, we get green chromic oxide and N2 gas in a manner like a volcano
erupting lava. Hundreds of reactions involve colour changes, evolution of gases, often coloured
gases, sometimes with production of light, heat and sound and sometimes with violent explosion.
It is often spellbinding and amazing to observe many chemical changes, sometimes appearing
like black magic. The whole earth is a playing game with the 91 natural elements that it is having.
We are merely mimicking the nature in a miniature scale in the laboratories and trying to unravel
the mysteries of nature.
In every topic that you shall read in this section, you will find many equations written without the
mention of the nature of reaction, whether metathesis or redox. First try to ascertain this and
if find the reaction to be metathesis then make sure which type of metathesis reaction out of the
ten types discussed in the text to which the reaction belongs. If it is a redox reaction, you shall
have to show the changes in ON. You are advised to do this without fail if you want to develop
some expertise in the subject in the long run. Many of the reactions given in this section, you
already know. This is merely a revision in such cases. Many of the equations given in this
section are not deliberately balanced. You have to balance it by following the methods you have
learnt in the chapter chemical arithmetic before. Occasionally balancing has been done to
remind you about it. Do not memorise any reaction. Write it with proper logic. I am sure, after
reading this section, you will develop further interest to learn more on inorganic chemistry. The
whole purpose of this section, in fact, the whole of the book is to arouse interest in young
beginners like you in the subject. Once you start liking the subject, I will feel that whatever little
effort I have put here to bring this book to light has borne immense fruit. There will be three main
chapters in this section.
(a)
Chemistry of some nonmetallic compounds
(b)
Chemistry of some nonmetals
(c)
Chemistry of some metals and their compounds.
CHEMISTRY OF SOME NONMETALLIC COMPOUNDS
AMMONIA(NH3)
Have you experienced the typical smell of urinals in a toilet? This is the smell of ammonia
produced by the bacterial decomposition of urea present in our urine into ammonia. The smell
of ammonia gas is pungent and irritating.
PREPARATION:
1.
Laboratory Method:
Heating of any ammonium salt with a base(alkali) produces ammonia(Refer Metathesis
reaction of the chapter, Chemical Logic and Arithmetic). Ammonia gas is collected by the
downward displacement of air as it is lighter than air. It is not collected over water as it is
highly soluble in water.
NH4Cl + CaO
NH3 + CaCl2 + H2O
(NH4)2SO4 + NaOH
NH3 + Na2SO4 + H2O
(M)
(M)
M stands for metathesis reactions. We get ammonia gas when any ammonium salt is
heated with any alkali or a base. The equations are not balanced and for a good
practice, the reader is advised to balance the equations themselves. Ammonia gas is dried
by passing over quicklime(CaO) but not over conc. H2SO4 because it reacts with it.
SAQ 1: Why does ammonia not dried by conc. H2SO4 and P2O5.
2
3.
Hydrolysis of Metallic Nitrides
(Refer Metathesis reaction):
Mg3N2 + H 2O
Mg(OH)2 + NH3
AlN + H2O
Al(OH) 3 + NH3
Industrial Method:
(M)
(M)
(i)
Haber's Process:
N 2 + 3 H2
Fe/Mo
200-300atm / 450-550 0C
2 NH3 + 22.4 Kcals (R)
(R) stands for redox reaction. Find out the changes in oxidation numbers (ON) of elements
in this reaction. In this reaction, iron is used as catalyst and Molybdenum as a promoter,
which enhances the activity of the catalyst (Fe). High pressure between 200-300 atm. and
high temperature between 450-5500C are applied. The reaction is exothermic, which means
that heat is liberated during the reaction. The yield of ammonia is good at high pressure and
low temperature according to a principle in physical chemistry called Le Chatelier's principle.
You will study about the principle in higher classes. Although pressure applied in the Haber's
process is high(200-300atm), the temperature applied is not low. The temperature applied(4505500C) is called the optimum temperature. At lower temperature than this, the reaction
becomes very slow and takes a long time. At higher temperature than this, the yield of
ammonia is intolerably poor. So a compromise between time and yield is made by using the
optimum temperature of 450-5500C. Nowadays, a mixture of potassium and aluminium oxide
is used as promotor for Fe catalyst in stead of Mo.
(ii)
Cynamide Process:
20000C
CaO + C
10000C
CaC 2 + N 2
CaNCN + H2O
CaC 2 + CO (calcium carbide)
(calcium cynamide)
CaNCN
CaCO3 + N H3
(R)
(R)
(M)
The first reaction is redox, in which C is reduced from 0(C) to -1 state(CaC2), while the
same carbon is oxidised from 0 to +2(CO) state. In the second reaction, which is also
a redox reaction, the ON of C increases from -1(calcium carbide) to +4 in calcium
cyanamide(CaNCN) and the ON of N decreases from 0(N2) to -3 in CaNCN. The third
reaction is a metathesis reaction in which there is no change in the ON of any element.
PROPERTIES:
Ammonia is pungent smelling colourless gas. It can be cooled to -330C to form a colourless
liquid. At room temperature ammonia can be liquified by the application of high pressure as
it is an easily liquifiable gas(refer chapter gaseous state). On further cooling to -780C, liquid
ammonia changes into solid state(melting point= -780C). It is highly soluble in water. At room
temperature and one atmosphere, one litre of water can dissolve nearly 700 litres of ammonia
gas.
1.
Acidic Properties:
NH3 is a weak base and it dissolves in water to produce NH4OH
NH3 + H2O -------> NH4OH
NH4OH is a weak electrolyte and dissociates to a small extent and so the number
of OH- produced is small. Therefore it is a weak base. Red litmus paper turns blue as it
is basic in nature. Chemistry of NH3 gas is analogous to chemistry of NH4OH in aqueous
solution.
+
NH4 + OH
NH4OH
2.
-
(weak dissociation)
Because of its basic properties, ammonia reacts with any acid to give a salt. Ammonium
hydroxide reacts with acid to give salt and water.
NH3 + HCl -------> NH4Cl;
NH3 + H2SO4 --------> (NH4)2SO4
3NH3 + H3PO4 -----> (NH4)3PO4
NH4OH + HCl -------> NH4Cl + H2O;
NH4OH + H2SO4 -----> (NH4)2SO4+H2O
All these are metathesis reactions.
Reaction with O2(air):
(i)NH3 gas burns in air or oxygen gas to give N2 and H2O
-3
NH 3 +
0
0
-2
N 2 + H2 O
O2
(R)
(ii)NH3 reacts with O2 in presence of Pt catalyst at 8000C to produce nitric oxide(NO).
-3
NH3 +
0
+2
Pt
8000C
O2
-2
NO + H 2O
(R)
Look to the changes in ON of elements in the above reactions. Balance them by ON
method that you have learnt in section-I.
3.
With Chlorine:
(i)
Excess NH 3:
NH3 reacts with Cl2 to produce N2 and HCl. In presence of excess ammonia, HCl
reacts with NH3 to form NH4Cl.
-3
0
NH3 + Cl2
NH3 + HCl
0
-1
N 2 + HCl
NH4Cl
(R)
(M)
Balance each equation and add the two to get the balanced overall equation.
(ii)
Excess Cl 2:
NH3 reacts with excess of Cl2 to give nitrogen trichloride(NCl3)
NH3 + Cl2
--------> NCl3 + HCl
(R)
N is oxidised from -3 to +3 while Cl is reduced from 0 to -1. Although N and Cl have same
electronegativity, due to larger size of Cl, it acquires the -ve ON and N acquires the +ve ON.
4.
As Reducing agent:
Ammonia reduces hot metallic oxides like CuO, PbO to the corresponding elements
and itself gets oxidised to N2.
+2
-3
CuO + NH3
5.
to N2.
0
0
Cu + N2 + H2 O
(M)
With Bleaching Powder:
Ammonia reduces bleaching powder[Ca(OCl)Cl] to produce CaCl2 and itself oxidised
-1
-3
+1
0
(R)
CaCl2 + N 2 + H2O
Ca(OCl)Cl + NH 3
Bleaching powder is calcium hypochlorite chloride(sometimes called calcium oxychloride). It
contains OCl -(hypochlorite) and Cl -(chloride) ions. The ON of Cl in OCl - is +1 and it gets
reduced to -1 in CaCl2 while the ON of N, as usual changes, from -3 to 0.
CHEMISTRY OF NH4OH:
We know that when NH3 dissolves in water produces NH4OH which is a weak base.
(i)
Precipitation reactions:
It brings about precipitation reactions with soluble salts of metals like Al, Cr, Fe, Zn
etc. to produce their insoluble hydroxides. These are all metathesis reactions(double
replacement/precipitation type).
Al2(SO4) 3 + NH4OH
Al(OH)3 + (NH 4)2SO 4
(gelatinous white ppt.)
Fe(OH)3 + NH4Cl
(reddish brown ppt.)
Cr(OH)3 + N H4Cl
FeCl3 + NH4OH
CrCl3 + N H4OH
(green ppt.)
(ii)
Complex Formation:
NH4OH dissolves insoluble substances like AgCl forming complex ions. You shall
learn about complex ions in higher classes. Now be happy with its formula.
AgCl(s) + 2NH4OH(aq) -------> [Ag(NH3)2]Cl(aq) + 2H2O(l)
(Complex compound)
The complex compound is called silver diammine chloride which is soluble in water. Note
that this reaction is a metathesis reaction.
Liquid Ammonia:
Liquid ammonia like water is a good solvent. But it is less polar than water. It dissolves alkali
metals like Li, Na, K etc. and alkaline earth metals like Mg, Ca, Sr etc.to form blue/bronze
coloured solution. This solution is a good conductor of electricity due to the presence of free
electrons. In presence of impurities, the metals react with ammonia to produce H2 gas and
metal amide.
0
+1
Na + NH3
+1
0
NaN H2 + H2
(sodium amide)
(R)
Here Na is oxidised from 0 to +1 and NH3 is reduced to H2(ON of H changes from +1
to 0). In short, sodium amide is called sodamide. Write similar reactions for K and Ca
for practice.
USES:
(i)for the preparation of a large number of fertilizers like urea, calcium ammonium
nitrate and many others.
(ii)for manufacture of nitric acid and many other ammonium compounds.
(iii)for removing grease: the part stained by grease is to be washed with NH4OH
solution.
TESTS:
(i)Forms dense white fumes when a glass rod dipped in conc. HCl is shown to the
gas
due to formation of NH4Cl. NH4Cl is a volatile solid and forms the fumes.
NH3 + HCl -------> NH4Cl(dense white fumes)
(ii)With Nessler's reagent, it gives a reddish brown precipitate. Nessler's reagent is
K2[HgI4]. NH3 forms a precipitate when reacts with Nessler's reagent forming a
compound called iodide of Millon's base.
SAQ 2: What happens when
(a) Ammonium nitrate, ammonium chloride, ammonium sulphate are heated separately
in the solid state to about 3000C
(b)Ammonium nitrite and nitrate are heated separately to a high temperature in liquid
state.
SULPHUR DIOXIDE(SO2)
Sulphur dioxide has a burning sulphur smell which is pungent and suffocating. In fact SO2 is
produced along with a small amount of SO3 when sulphur burns in oxygen or air. It is highly
soluble in water. It is liquified by cooling to -100C(b.p) and solidified by further cooling to
-75.50C(m.p). SO2 gas can also be liquified by the application of high pressure at room
temperature.
PREPARATION:
1.
Burning of Sulphur:
S + O2 ------> SO2
(a small amount of SO3 is also formed) (R)
2.
Reaction of sulphite and bisulphite salts with dilute mineral acids:
When any sulphite or bisulphite salt is treated with any dilute mineral acid at room
temperature, we get bubbles of SO2 gas. No heating is necessary. These are all metathesis
reactions.
Na2SO3 + HCl -------> NaCl + SO2 + H2O
CaSO3 + H2SO4 -------> CaSO4 + SO2 + H2O
KHSO3 + HBr --------> KBr + SO2 +H2O
3.
Laboratory Method: When a mixture of Cu turning and conc. H2SO4 is strongly
heated, we get SO2 gas.
0
+6
Cu + H2SO 4 (conc.)
+2
+4
CuSO 4 + SO2 + H2O
(R)
Note that conc. H2SO4 is always reduced to SO2 i.e from +6(H2SO4) to +4(SO2). Cu is
oxidised from 0 to +2. Balance the above equation by either partial equation method
or ON method.
4.
Roasting of sulphide ores:
When sulphide ores such as iron pyrites(FeS2), zinc
blende(ZnS), galena(PbS) etc. are heated in the presence of air, we get SO2 gas. This is
called roasting of the ore.
FeS2 + O2 -------> Fe2O3 + SO2
(R)
ZnS+ O2 -------> ZnO + SO2
(R)
Show the changes in ON.
SAQ 3: What is the ON of Fe in FeS2. Show the changes in ON occurring in the roasting
of iron pyrite.
PROPERTIES:
Acidic Nature:
SO2 is an acidic oxide and is the anhydride of sulphurous acid(H2SO3) which is a weak acid.
SO2 therefore turns moist blue litmus paper red.
(M)
SO2 + H2O -------> H2SO3
SO2 therefore reacts with alkali to form sulphite and bisulphite(hydrogen sulphite)salts. In
dilute solution, the salt remains mostly as bisulphite. In concentrated solution and in the solid
state, the salt remains as sulphite. Note that bisulphite salts are not commonly isolated in the
solid state as they are unstable.
NaOH + SO2 -------> Na2SO3
(M)
Na2SO3 + SO2 + H2O ------> NaHSO3
(M)
As a Reducing agent:
SO2 is a very strong reducing agent. Oxidising agents like Cl2 water, H2O2, acidified KMnO4,
K2Cr2O7 etc. oxidise it to SO3 which soon reacts with water to produce H2SO4.In all such
reactions S is oxidised from +4(SO2) to +6(H2SO4).
(i) With Cl2 water:
0
+4
Cl2 + H2O + SO 2
-1
+6
HCl + H 2SO4
(R)
We can balance this equation by two ways, either by ON method or by partial equation
method. In the latter method, we can either first break the oxidising agent or the reducing
agent. Let us first take up the partial equation method in which we shall break the oxidising
agent producing nascent oxygen.The nascent oxygen then will react with SO2 to give SO3.
SO3 is unstable and react with water to give its corresponding acid, H2SO4. Then we cancel
nascent oxygen and SO3 from the partial equations to get the balanced equation.
Cl2 + H2O
SO 2 + [O]
SO 3 + H2O
Cl2+ SO2 + 2 H2O
2 HCl + [O]
SO 3
H2SO 4
2HCl + H2SO 4
Alternatively, we can first break SO2 as follow: SO2 + 2H2O ------>H2SO4 + 2[H]
In this method, SO2 reacts with H2O to produce nascent H atom which react with the
oxidising agent to reduce it to its lower oxidation state.
SO2 +2 H 2O
Cl2 + 2 [H]
H2 SO 4 + 2 [H]
2 HCl
SO2 + Cl2 +2 H 2O
2 HCl + H2SO4
You can use any one method to balance redox reactions involving SO2.
(ii)
With H2O2:
H2O2
SO2 + [O]
SO3 + H2O
H2O + [O]
SO3
H2SO4
H2O2
H2SO4
+6 -2
+4
-1
+
SO2
In this reaction, the ON of O changes from -1(H2O2) to -2 in H2SO4 while the ON of S
changes from +4 to +6.
(Reactions given below are not balanced and the reader is advised to balance each
equation by ON or partial equation method for redox reactions and hit and trial method
for metathesis reactions for the sake of practice)
(iii)
With acidified KMnO4:
The pink colour of KMnO4 is discharged when SO2 is bubbled through acidified
KMnO4 solution. Refer section-I if you have forgotten balancing techniques. Here we give
without balancing.
+2
+4
+7
KMnO4 + H2SO 4 + SO 2
+6
(R)
K 2SO4 + MnSO4 + H 2SO 4 + H2O
Look to the changes in ON. KMnO4 always is reduced to +2 state in acidic medium.
SO2 always is oxidised to SO3 finally to H2SO4.
(iv) With acidified K2Cr2O7:
The orange colour of K2Cr2O7 turns green on passing SO2 gas through it.
+6
+3
+4
K2Cr2O 7 + H2SO4 + SO2
(orange)
+6
K 2SO 4 + Cr2 (SO 4)3 + H 2SO4 + H2O
(green)
(R)
K2Cr2O7 is always reduced from +6 to +3 state in acidic medium.
As Oxidising Agent:
In a few reaction, SO2 acts as the oxidising agent.
(i)
With H2S:
-2
+4
H2 S + SO 2
0
S + H2O
(R)
The ON of S changes from -2 to 0 and also from +4 to 0. You already know from
section-I that in a redox reaction, usually the decrease or increase of ON takes place via zero
state. When H2S gas is bubbled through SO2 solution, we get a pale white suspension of
sulphur.
(ii)
With Mg metal:
Mg burns in the atmosphere of SO2, and is oxidised to its oxide.
Mg + SO2 -------> MgO + S
(R)
Mg + SO2 ------> MgO + MgS
(with excess Mg)
(R)
In the first reaction, Mg goes from 0 to +2 and S from +4 to 0. In the second reaction S
goes from +4 to -2, while the Mg from 0 to +2.
(iii) With acidified SnCl2:
+2
SnCl2 + HCl
+
+6
+4
+4
SO2
SnCl4
+
H2SO4
(R)
HCl is always used with SnCl2 for giving more Cl- ions to form SnCl4.
Addition Reactions:
(i)
With O2:
+4
0
SO2 + O2
V2 O 5
+6 -2
SO3
(R)
We get SO3 when SO2 reacts with O2 gas at high pressure and moderately high temperature
in the presence of vanadium pentoxide catalyst.
(ii)
With dry Cl2:
0
+4
SO2
+
Cl2
+6
-1
SO 2 Cl2
(R)
We get sulphuryl chloride(SO2Cl2) when SO2 reacts with dry Cl2 gas.
USES:
(i)
It is used as bleaching agent. It bleaches or removes colour from flower, wool, silk,
straw. wood pulp etc. The bleaching carried out by SO2 is temporary in nature and the
bleached material regains the original colour after being exposed to air(O2) for some time.
The bleaching action of SO2 is due to its reducing property,
SO2 + 2H2O ------> H2SO4 + 2[H].
The nascent H atom is responsible to bleach the colour. The coloured substance loses it colour
due to reduction by SO2. On exposure to air(O2),the reduced material is again oxidised to
the same state thereby getting back the colour.
(ii)
It is used as a disinfectant for its germicidal action. It kills bacteria and moulds. It
is used as a food preservative.
CARBON DIOXIDE(CO2)
Atmosphere contains only 0.03% CO2 by volume. It is produced by the respiration of living
beings and combustion and decomposition of organic matter. It is used up by plants for the
photosynthesis and the net amount of CO2 is supposed to remain constant. But due to large
scale deforestation, the CO2 is used up at a slower rate than being produced. Thus the amount
of CO2 in the earth is slowly increasing. You know that CO2 can retain a large amount of
heat in it from sun rays and thus the temperature of the earth is slowly increasing. This is
called global warming. Measures are being taken throughout the world to fight against this
problem.
PREPARATION:
1.
By treating carbonate or bicarbonate salt with mineral acids:
(Refer metathesis reaction in chapter-I)
CaCO3 + HCl -------> CaCl2 + CO2 + H2O
(M)
K2CO3 + H2SO4 ------> K2SO4 + CO2 + H2O
(M)
Bubbles of CO2 are evolved as soon as a carbonate salt is treated with any mineral acid.
No heating is necessary.
2.
By heating carbonate salt(Other than Na, K, Rb and Cs carbonates):
(Refer Metathesis reaction: Section-I)
CaCO3 -----heat-----> CaO + CO2
Li2CO3 -----heat-----> Li2O + CO2
Na2CO3 ----heat-----> No change
PROPERTIES:
It is a colourless, odourless and tasteless gas. When dissolved in water it gives a very
pleasant taste. You have experienced the unique taste while taking soft drinks i.e beverages
(Pepsi, Limca etc) filled with dissolved CO2 gas. It can liquify easily at room temperature
by applying high pressure (refer gaseous state). It is moderately soluble in water and more
so when dissolved under high pressure.
(i)
Acidic Properties:
It is the anhydride of carbonic acid. Carbonic acid is a dibasic and weak acid.
CO2 + H2O -------> H2CO3
(M)
Carbonic acid is stable in water at room temperature only under pressure. On boiling water
or reducing pressure of CO2, large volume of CO2 escapes out. You must have marked when
you open a soft drink bottle.
(ii)
With alkali:
CO2 being an acidic oxide reacts with alkali(metallic hydroxides) forming their carbonates first
and then with excess CO2 forms its bicarbonates(hydrogen carbonates).
With lime water:
This is the most common test for CO2. When passed through
lime water(calcium hydroxide), it forms milky white precipitate(CaCO3) which subsequently
dissolves forming soluble bicarbonate. All these are metathesis reactions.
CO2 + Ca(OH)2 --------> CaCO3↓(white ppt) + H2O
(M)
CaCO3 + CO2 + H2O -----> Ca(HCO3)2 (soluble)
(M)
With NaOH: With NaOH and KOH, the observation is just the opposite. Carbonate
which is formed first is soluble and the bicarbonate which is formed later is weakly soluble
and gives a white precipitate.
CO2 + NaOH ------> Na2CO3(soluble) + H2O
(M)
Na2CO3 + CO2 + H2O -------> NaHCO3↓(white ppt.)
(Μ)
USES:
(i)
It is used as a refrigerant liquid.
(ii)
It is used for extinguishing fire.
(iii)
It is used in soft drinks as a preservative and to enhance taste.
(iv)
Dry Ice: Solid CO2 is called dry ice. Its temperature is -400C. It is prepared
when liquid CO2 is allowed to expand from high pressure to low pressure through a small
hole. Liquid CO2 is further cooled by this process to produce solid CO2. It is used to test
whether a particular electronic part or aircraft part can tolerate a low temperature of -400C
or not. It is also used as a cooling aid.
CARBON MONOXIDE(CO):
Do you know that air pollution caused by the exhaust gases coming from the automobiles
(cars, trucks, buses etc.) contain some amount of carbon monoxide? Carbon monoxide if
formed by the incomplete combustion of petrol/diesel because of poor quality of petrol/diesel
or the engine.
C8H18(petrol) + O2(insufficient) --------> CO + H2O (Incomplete combustion)
(Complete combustion)
C8H18 + O2(excess) ------------> CO2 + H2O
Carbon monoxide is a poisonous gas unlike carbon dioxide which is nonpoisonous and useful
in many ways. A slightly more quantity of CO inhalation will cause death. This is because
CO forms a stable compound with haemoglobin present in blood which loses its power to take
O2 gas from air. Thus the person dies due to want of O2.
PREPARATION:
(i)
From CO2:
When CO2 is passed over red hot coke(C), we get CO.
CO2 + C --------> CO
(R)
(balance this equation)
What change of ON you notice? C changes from +4 to +2 and C changes from 0 to
+2.
(ii)
Incomplete combustion of organic substances: Any organic substance when
burnt in limited quantity of oxygen, we get CO instead of CO2.
CH4 + O2(limited) -------> CO + H2O
(R)
(Show the changes in ON)
(iii)
Reduction of CO2 by metals: When CO2 is passed over red hot zinc, we
get CO.
+2
+2
+4
0
ZnO + CO
Zn + CO 2
PROPERTIES:
It is a combustible gas like H2 but not a supporter of combustion. It burns with a blue flame
and forms CO2.
CO + O2 -------> CO2 + heat
(R) Find the changes in ON
It is a colourless, odourless and tasteless gas having boiling point -1900C and melting point
-205.10C.
(i)
As a Reducing agent:
It is a powerful reducing agent as it itself easily oxidised to CO2. It reduces metallic oxides
to form the corresponding metals. It also reduces nonmetallic oxide like I2O5 to its element(I2).
+4
0
+2
+3
Fe
Fe 2 O 3 + CO
+
CO 2
(R)
ZnO + CO --------> Zn + CO2
PbO + CO --------> Pb + CO2
+5
0
+2
(ii)
Addition reactions:
methyl alcohol respectively.
0
CO + Cl2
(Find the changes in ON)
(Find the changes in ON)
+4
(R)
I2 + CO 2
I2O 5 + CO
+2
(R)
(R)
CO reacts with Cl2 and H2 to give COCl2(phosgene) and
+4
sunlight
-1
COCl2 (carbonyl chloride or phosgene)
Phosgene was used as a war gas in the first world war.
Here also CO acts as a reducing agent, reducing Cl from 0 to -1.
+2
0
CO + H2
0
100-400 C
ZnO
-2
+1
CH3O H (methyl alcohol)
In this reaction CO acts as an oxidising agent as it oxidises H from 0 to +1 and itself
reduced to -2 state.ZnO is used as the catalyst. Do you know something about methyl
alcohol? The liquor tragedy in which several people died by taking spurious liquor contained
this poisonous methyl alcohol along with ethyl alcohol, which is the original substance of a
liquor. We shall know more about them in our organic chemistry section.
USES:
1.
It is used in the form of water gas[CO+H2] and producer gas[CO+N2] as fuels in
industries.
2.
It is used for extraction of metals as a reducing agent.
MISUSE:
It is the major cause of air pollution.
HYDROGEN SULPHIDE (H2S)
Have you experienced the rotten-egg type of smell produced by hydrogen sulphide? It is an
extremely poisonous gas and one of the pollutants present in the atmosphere which causes
air pollution.
PREPARATION:
1.
Action of metallic sulphides with mineral acids:
when metallic sulphides like FeS, ZnS etc. react with dilute mineral acids e.g HCl, H2SO4
hydrogen sulphide gas is evolved. No heating is necessary. All these are metathesis reactions
(Refer section-I)
ZnS + HCl(dil.) -----------> ZnCl2 + H2S
(M)
FeS + H2SO4(dil)---------> FeSO4 + H2S
(M)
Sb2S3 + HCl -----------> SbCl3 + H2S
(M)
PROPERTIES:
It is a colourless gas and is moderately soluble in water. Its boiling point is -60.40C and melting
point is -85.50C. H2S gas can be liquified even at room temperature by applying high pressure.
(i)
Acidic Properties:
It is a weak dibasic acid and forms two types of
salts with alkali.
NaOH + H2S -------> NaHS(sodium bisulphide) + H2O
(M)
NaSH + NaOH-------> Na2S(sodium sulphide) + H2O
(M)
(2)
As a reducing agent: H2S is a very good reducing agent. With any oxidising
agent (mild or strong), it is oxidised to S. In almost all the redox reactions involving H2S,
it is oxidised to S, the ON changes from -2 to 0. The reader is advised to balance each
equation by partial equation or ON method for practice. For small equations, use ON method
and for difficult ones use the partial equation method. The unbalanced equations are given
below with changes in ON shown in some cases.
(i)With Cl2:
When Cl2 gas is bubbled through saturated H2S solution in water, pale
white suspension of sulphur is formed.
-2
0
H2S + Cl2
-1
0
HCl + S
(R)
(ii)With SO2: When H2S gas is bubbled through aqueous solution of SO2, pale white
suspesnion of sulphur is formed.
+4
-2
0
S + H 2O
H2S + SO 2
(R)
(iii)With FeCl3 solution: When H2S gas is bubbled through FeCl3 solution, the reddish
brown colour of FeCl3 turns green due to formation of FeCl2. Sulphur is also produced.
-2
H2S
+
+3
+2
FeCl3
FeCl2 + S
0
+
HCl
(R)
(iv)With conc. H2SO4: When H2S gas is bubbled through conc. H2SO4, sulphur is
produced and SO2 gas is evolved. You know that conc. H2SO4 is always reduced to SO2 i.e
the ON changes from +6 to +4 and . H2S is always oxidised to S.
+4
-2
+6
0
SO2 + S
H2SO4 + H2S
+
H2O
(R)
(v)With conc. HNO3: When H2S is bubbled through conc. HNO3, sulphur is
produced and the reddish brown NO2 gas is also evolved. You know that conc. HNO3 is
always reduced to NO2, the ON changes from +5 to +4.
+5
+4
-2
0
NO 2 + S + H 2O
HN O3 + H2 S
(R)
(vi)With acidified KMnO4: When H2S is bubbled through acidified KMnO4 solution,
the pink colour of permanganate is discharged and sulphur is produced.
+7
+2
-2
KMnO 4 + H 2SO 4 + H2S
0
K 2 SO 4 + MnSO 4 + S + H2 O
(pink)
(colourless)
(R)
(vii)With acidified K2Cr2O7 : H2S turns the orange colour of dichromate into green.
+6
+3
-2
K 2 Cr 2O7 + H2 SO4 + H 2S
0
K2 SO4 + Cr2(SO 4)3 + S + H2 O
(orange)
(green)
(R)
(viii)With Air: H2S is oxidised by O2 dissolved in water producing pale white suspension
of sulphur. However when H2S is burnt in air or oxygen, it produces SO2. This is the only
reaction where the ON of S is promoted from 0 to +4.
-2
0
H2S + O 2
-2
0
H2 S + O2
(3)
(water)
(burning)
0
-2
(R)
S + H2O
+4
-2
SO 2 + H 2O
(R)
Precipitation(Metathesis) Reactions:
Salt solutions of Cu2+, Pb2+, Hg2+, Bi3+, As3+, Sb3+, Sn2+ , Co2+, Ni2+,Zn2+, Mn2+ give
an insoluble sulphide precipitate when H2S is passed through each of them. The sulphides of
all the metal ions are black, excepting As2S3(yellow), Sb2S3(orange), ZnS(white) and
MnS(buff or flesh colour). All these are double replacement type of metathesis reactions.
Let us write a few equations here and the rest you can practise of your own.
CuSO4 + H2S ------> CuS↓ + H2SO4
(M)
AsCl3 + H2S --------> As2S3↓ + HCl
(M)
Zn(NO3)2 + H2S------> ZnS↓ + HNO3
(M)
USES:
(i) As an important reagent in the laboratory. Have you seen how H2S gas is prepared
and used in the laboratory? It is prepared in a separate apparatus called Kipp's apparatus.
One can get H2S gas continuously from Kipp's apparatus whenever required. Whenever, one
does not need the gas, stopcock can be closed and the reaction stops.
(ii)It is used to prepare many metallic sulphides.
Test: It turns lead acetate paper black.
Pb(CH3COO)2 + H2S -----> PbS(black)+HCH3COO( or CH3COOH)
HYDROGEN CHLORIDE (HCL): HYDROCHLORIC ACID
Hydrogen chloride is a gas at room temperature and when it is dissolved in water produces
hydrochloric acid. You are quite familiar with hydrochloric acid.
PREPARATION:
1.
From NaCl and Conc. H2SO4: (Laboratory Method)
NaCl is available plentily in sea water. When a mixture of NaCl and conc. H2SO4
is heated, HCl gas is produced. HCl gas is collected by the upward displacement of air as
it is heavier than air. It is not collected over water as it is highly soluble in water. It is dried
by passing over conc. H2SO4 or P2O5 but not over CaO as it will react with the latter.
NaCl + H2SO4(conc.) -------> NaHSO4 + HCl (<2000C)
(M)
NaHSO4 + NaCl ------------> Na2SO4 + HCl (>2000C)
(M)
PROPERTIES:
It is a colourless gas with suffocating and pungent odour. Its solution (acid) is sour to taste.
The gas liquifies at -850C(bp) and solidifies at -1110C(mp) at room temperature. It is highly
soluble in water and forms constant boiling mixture(azeotropic mixture) at the composition
of 20.4% of HCl by mass (rest water). This mixture boils at 1100C(Read the SAQ below).
We can get hydrochloric acids having different HCl concentrations by dissolving different
amounts of HCl gas in a fixed quantity of water.
SAQ 4 : What is a constant boiling or azeotropic mixture?
(i)Acid Strength:
Hydrochloric acid is a strong acid. HCl gas is a covalent
compound but undergoes dissociation almost completely to form H+ and Cl- ions. It reacts with
any base or alkali to form salt.
NaOH + HCl -----> NaCl + H2O,
Al(OH)3 + HCl ------> AlCl3 + H2O
NH3 + HCl -----> NH4Cl and so on.
These are all metathesis(neutralisation)reactions.
(ii)Reducing Property: HCl acts as a reducing agent particularly in the presence of
strong oxidising agents like MnO2, PbO2, acidified KMnO4, . It is not as good a reducing agent
as HBr or HI.
+4
-1
MnO 2 + HCl
+2
0
MnCl2 + Cl2 + H2 O
(R)
+7
-1
KMnO 4 + H2SO 4 + HCl
+2
0
K 2SO 4 + MnSO4 + Cl2 + H 2O
(R)
The pink colour of KMnO4 is discharged.
(iii)Precipitation (metathesis) reactions: Metal ions like Hg22+(ous), Pb2+, Ag+ etc.
give insoluble white precipitates of their chlorides when HCl reacts with them.
AgNO3 + HCl -----> AgCl↓ + HNO3,
Hg2(NO3)2 + HCl----> Hg2Cl2↓ + HNO3
Pb(NO3)2 + HCl -----> PbCl2↓ + HNO3
USES:
1.
It forms aqua ragia with conc. HNO3. A mixture of 1 part conc. HNO3 and 3 parts
conc. HCl is called aqua ragia. Aqua ragia can dissolve noble metals like Au, Pt. Aqua ragia
forms nascent Cl atoms which react with the metals forming their soluble chlorides.
HCl + HNO3 --------> NOCl + H2O + [Cl]
Au + 3 [Cl] -------> AuCl3,
Pt + 4 [Cl] -----> PtCl4
TEST: When a glass rod dipped in NH4OH solution shown over the HCl gas, dense white
fumes are produced. (same as NH3 gas, refer the topic ammonia).
SAQ 5 : Predict the following reactions with proper reasoning. Indicate whether the reaction
is metathesis ore redox. Show the changes in ON for redox reactions.
(i)
PbO2 + HBr ------->?
(ii)
H2SO4(conc.) + HI ------>?
(iii)
P4 + Br2 ------->?
(iv)
PBr3 + H2O ------->?
(v)
NaBr + H2SO4(conc.)--->?
(vi)
CuSO4 + HI ------>?
HYDROGEN PEROXIDE (H2O2)
Hydrogen peroxide is a water-like liquid. It looks exactly like water, but stored in dark
coloured bottles as it is decomposed by light to water and oxygen. While water boils at 1000C
and freezes at 00C, hydrogen peroxide boils at 151.40C and freezes at -0.90C. It is heavier
than water(density=1.5gm/cc while that of water=1.0gm/cc). Its viscosity is nearly same as
water which means that it is a thin liquid like water.
PREPARATION:
1
From metallic peroxides: When sodium or barium peroxides react with dil. H2SO4
0
at 0 C temperature, we get hydrogen peroxide.
Na2O2 + H2SO4 -------> Na2SO4 + H2O2
(M)
BaO2 + H2SO4 ----> BaSO4 + H2O2
(M)
PROPERTIES:
1.
Decomposition:
H2O2 decomposes to H2O and O2 by light, or heat or even on standing. The rate of decomposition
is low. In presence of small quantity of catalysts like MnO2, Pt, Au, Ag etc.it decomposes
rapidly. Its decomposition is retarded or inhibited by adding a small quantity of phosphoric
acid or metal stannates.
H2O2 -----------> H2O + O2
(R)
(Show the changes in ON)
2. Acidic Property:
It is a very weak acid. It forms two series of salts with NaOH by the removal of the two
H atoms one by one.
NaOH + H2O2 -----> NaHO2(sodium hydrogen peroxide)+ H2O
(M)
(M)
NaHO2+ NaOH ----> Na2O2(sodium peroxide) + H2O
3. As oxidising agent:
It is a moderately strong oxidising agent. It oxidises the black PbS to white PbSO4, colourless
iodide salt to brown I2 solution, green ferrous salt to reddish brown ferric salt, nitrite salt to
nitrate salt and so on. Study the following redox reactions. Balance each of them by ON or
partial equation method. In partial equation method, break down first H2O2-----> H2O + [O]
and then proceed as usual. In all these reactions, H2O2 is reduced to H2O and the ON of
O changes from -1 in peroxide to -2 in water. We give here the one-line equations.
-2
+6
-1
-2
PbSO 4 + H2 O
PbS + H2O2
(R)
In this reaction S changes from -2 to +6 and this is not usual for sulphur as it changes from
-2 to 0 most commonly.
+2
-1
Fe2(SO 4)3 + H2O
0
-1
-2
+3
-1
FeSO 4 + H2SO 4 + H 2O 2
-2
(R)
I2 + KOH
KI + H 2O2
(R)
H2O2 oxidises Cr from +3 to +6 state in alkaline medium. When green insoluble Cr(OH)3
precipitate is boiled with H2O2 in presence of alkali, yellow solution of chromate salt is
produced.
+3
+6
-1
Cr(OH)3 + H2O 2 + N aOH
-2
Na2 CrO 4 + H 2O
(R)
4. As a reducing agent:
In the presence of strong oxidising agents such as acidified KMnO4, K2Cr2O7 it acts as a
reducing agent and is oxidised to O2 gas. While balancing by partial equation method, you
have to first break H2O2 as H2O2 ------> O2 + 2 [H], and then you will use nascent H to
react with nascent O produced from the oxidising agent. Refer section-I for balancing by
partial equation method involving H2O2 if you have forgotten. Here we give the one-line
equations with logic. The reader is advised to practise the balancing by partial equation or
ON method.
+2
-1
+7
+6
0
K2 SO 4 + MnSO 4 + O2 + H2O
KMnO4 + H2SO 4 + H2O 2
+3
-1
0
K2SO 4 + Cr2(SO 4)3 + O2 + H2O
K 2Cr2O 7 + H2SO 4 + H 2O 2
(R)
(R)
It also reduces Cl2, Br2 to their halides(Cl-, Br-), and NaOCl to NaCl. In all these reactions,
H2O2 is oxidised to O2. The ON of O changes from -1 to 0.
0
-1
Cl2 + H2 O 2
SAQ6 :
(ii)Br2
-1
0
HCl + O 2
(R)
Suggest the products when H2O2 reacts with (i)NaOCl
USES: 1.
2.
3.
It is used as an antiseptic for cleansing wounds,teeth and ears.
As a bleaching agent for cotton and synthetic fibres, wool, hair and other soft
materials.
As a preservative for milk and wine.
NITRIC ACID(HNO3):
Pure nitric acid is a colourless liquid having density 1.52gm/cc at room temperature. But
always it is used as aqueous solution in the form of dilute or concentrated acids.
PREPARATION:
It is prepared in the laboratory by heating any nitrate salt(NaNO3) with conc. H2SO4. Simple
double replacement(metathesis) reaction takes place to form HNO3 which is distilled off from
the mixture to a separate container in which it is cooled to give nitric acid.
NaNO3 + H2SO4(conc.) ------------> NaHSO4 + HNO3
(M)
(M)
NaHSO4 + NaNO3(excess) ---------> Na2SO4 + HNO3
Note that since conc. H2SO4 is less volatile and has a boiling point of 3380C while HNO3
is more volatile and has much less boiling point than conc. H2SO4, it is possible for HNO3
to volatilise out and form the distillate. Na2SO4 or NaHSO4 is left behind as the residue.
Therefore we say that a less volatile acid can displace a more volatile acid from the
salt of the more volatile acid.
MANUFACTURE OF HNO3 BY OSTWALD'S PROCESS:
The process involves three steps.
Step-I: Preparation of NO:
Nitric oxide(NO) is prepared by the catalytic oxidation of NH3.
NH 3 + O 2
Pt
8000C
NO + H 2O
(R) Show the changes in ON
Step-II: Preparation of NO2:
The gas mixture is cooled to 1000C and mixed with O2 to form reddish brown NO2 gas.
+2
0
NO + O2
1000C
+4 -2
NO 2
(R)
Step-III: Preparation of HNO3:
NO2 is mixed with excess air(O2) and the mixture reacts with water to produce HNO3. First
NO2 reacts with O2 to give N2O5 which subsequently reacts with H2O to give HNO3. You
know that the last reaction is a metathesis reaction involving no change in ON.
+4
0
NO2 + O 2
N2 O5 + H2 O
+5 -2
N2O5 (R),
HNO 3 (M)
The acid which is prepared first is dilute and contains more water and less of HNO3. By
distilling this diluted acid for a number of times, the concentration of HNO3 is increased upto
68% by mass. Once this concentration is attained, it is not possible to further concentrate the
HNO3 by distillation process. This is called the constant boiling mixture or azeotropic
mixture(boiling point 120.5C). Further concentration to prepare concentrated HNO3(98% by
mass) is done by other chemical methods of dehydration.
PROPERTIES:
1.
Decomposition:
Usually conc. HNO3 looks yellow because of some dissolved NO2 gas. NO2 gas is produced
by the decomposition of HNO3. This decomposition is slow at room temperature but fast if
temperature is increased
+4
+5 -2
0
(R)
NO 2 + O 2 + H2O
HNO3
2.
Acidic Properties:
It is a monobasic strong acid and dissociates as follows.
HNO3 --------> H+ + NO3It forms nitrate salt when reacts with a base or alkali. Carbonates and bicarbonates,
react with dilute nitric acids to form salt and CO2. These are all metathesis reactions, that
you know from section-I.
CaO + HNO3 --------> Ca(NO3)2 + H2O
(M)
(M)
Na2CO3 + HNO3 ------> NaNO3 + CO2 + H2O
Note that sulphite and sulphide salts react with HNO3 but go for redox reactions, since HNO3
is a powerful oxidising agent. We shall see this later.
3.
As Oxidising Agent:
HNO3 is a versatile oxidising agent. It oxidises many metals, nonmetals and common reducing
agents and itself get reduced to one of the oxides of nitrogen such as NO2, NO, N2O. Which
oxide of nitrogen is produced depends on the concentration of nitric acid used. The following
table gives the information. This table was also given in chemical arithmetic chapter before.
conc. HNO3 ---------> NO2,
dilute HNO3 ---------->N2O
moderately conc. HNO3(50%)-------->NO
very dilute HNO3 ------->NH3----->NH4NO3
A.
WITH NONMETALS:
Conc. HNO3 oxidises nonmetals like carbon, sulphur, phosphorous, I2 etc. as follows and itself
gets reduced to the reddish brown, NO2 gas.
0
1.
With Carbon:
+5
+4
C + HNO 3
+4
CO 2 + NO 2 + H2O
(R)
Carbon is oxidised to its stable oxide(CO2) which does not react with water and escapes out.
Can you balance the equation? You can do so either by ON method or by partial equation
method. Let us balance by ON method. The change in ON in carbon is 4 and we place 4 as the
coefficient of HNO3. The change in ON of N is 1 and we leave Cu as such. Then we go for hit
and trial.
C + 4 HNO 3
CO 2 + 4 NO 2 + 2 H2O
(balanced)
In partial equation method, HNO3 is broken down to oxide of nitrogen, water and nascent oxygen.
Nascent oxygen further reacts with the reducing agent to give the other product.
2 HN O3
2 NO 2 + H2O + [O] X 2
C + 2 [O]
C + 4 HNO3
CO 2
4 N O2 + CO 2 + 2H 2O
So henceforth, the balancing job is to be done by you for better practice.
2.
With Sulphur: Sulphur is first oxidised to SO3 which reacts with H2O to produce H2SO4.
Since SO3 is unstable, reacts with water to form sulphuric acid.
0
+5
+6
+4
(R)
H2 SO 4 + NO 2 + H2O
HN O 3 + S
3.
With Phosphorous:
Phosphorous is first oxidised to P2O5 which is unstable and
react with water to produce the corresponding acid, H3PO4.
0
+5
+5
+4
(R)
H 3PO4 + NO 2 + H2O
HNO3 + P 4
4.
With Iodine:
I2 is first oxidised to I2O5 which reacts with water producing
the corresponding acid, HIO3. I2O5 is unstable like SO3 and P2O5 and gives its acid(iodic acid).
0
+5
+5
+4
(B)With Common Reducing Agents:
follows. It is always reduced to NO2.
1.
With H2S:
Conc. HNO3 react with common reducing as
H2S is oxidised as usual to S.
-2
+5
With SO2:
+4
0
HNO3 + H2 S
2.
(R)
S + N O2 + H2O
SO2 is first oxidised to SO3 which reacts with H2O to form H2SO4.
+5
3.
Acidified FeSO4:
sulphate.
+5
+4
+6
+4
HN O3 + SO 2
H2 SO 4 + NO 2 + H2O
With HBr:
(R)
Green ferrous sulphate is oxidised to reddish brown ferric
+3
+2
+4
Fe2(SO 4)3 + NO 2 + H2O (R)
HNO3 + FeSO 4 + H 2SO 4
4.
(R)
HIO 3 + N O 2 + H2O
HNO3 + I2
HBr is oxidised to Br2.
+5
-1
0
HN O3 + HBr
+4
Br2 + N O2 + H2O
(R)
5.
With KI (or HI):
KI or HI is oxidised to I2. Since the medium is nitric acid, KI
produces KNO3 while HI gives H2O.
+5
-1
HN O3+ KI
0
+4
I2 + KN O3 + NO 2 + H2O
(R)
(C)
WITH METALS:
Metals can react with various forms of nitric acids unlike nonmetals and produce different
oxides of nitrogen depending on the concentration of the acid used as given in the table before.
Conc. HNO3 produces usually NO2(N=+4), moderately conc. HNO3(50%) produces NO(N=+2),
dilute HNO3(20%) produces N2O(N=+1), while very dilute HNO3 produces NH3(N=-3) gas
which being basic in nature reacts with HNO3 to produce ultimately NH4NO3. Sometimes, dilute
HNO3 produces N2(N=0) instead of N2O, if it is little more diluted. We shall take a few metals
and use different forms of HNO3 randomly. You can write equations for other metals in similar
lines. Note that metal remains as its nitrate after oxidation.
1.
With Zn using conc. HNO3:
0
+5
+2
Zn + HN O3 (conc.)
+4
Zn(N O3) 2 + N O2 + H2O
(R)
Can you balance the equation by ON method or by partial equation method? Refer chapter-I for
that if you have forgotten how to balance.
2.
With Cu using mod. conc. HNO3: HNO3 is reduced to colourless gas NO in this case.
0
+2
+5
Cu + HNO3
(mod. conc.)
3.
+2
Cu(NO3 )2 + NO + H2O
(R)
With Fe using dilute HNO3: In this case HNO3 will be reduced to N2O.
0
+2
+5
Fe + HNO 3
+1
Fe(NO 3)2 + N 2O + H2O
(dil.)
(R)
Note that dilute HNO3 would oxidise Fe to +2(not +3) and itself reduced to N2O.
4.
With Sn using very dilute HNO3: Very dilute HNO3 produces NH4NO3 as said before.
0
+5
-3
+2
Sn(NO 3)2 + NH4NO3 + H2O
Sn + HNO3
(very dil.)
(R)
Can you not balance this equation by ON method? Try now. The change in ON of N is 8, so 8 is
to be placed as coefficient of Sn. The change in Sn is 2, so 2 to place before HNO3. Then hit and
trial.
8Sn + 2HNO3 -------> Sn(NO3)2 + NH4NO3 + H2O
After equalising Sn in RHS, we have to modify the coefficient of HNO3 to 18 to see if the
equation could be balanced. We fail to achieve our goal. Then place the coefficient 2 before
NH4NO3 and finally adjust HNO3, then H2O.
8Sn + 18HNO3 -------> 8Sn(NO3)2 + NH4NO3 + 7H2O
( n o t
balanced)
8Sn + 20HNO3 --------> 8Sn(NO3)2 + 2 NH4NO3 + 6H2O
(balanced)
Further simplifying on dividing by factor 2, we get
4Sn + 10HNO3 --------> 4Sn(NO3)2 + NH4NO3 + 3H2O
(balanced)
You also try to balance every equation by partial equation method. Refer section-I for the purpose.
5.
With Sn using conc. HNO3: In this case, Sn is oxidised to +4 state(not to +2 as is the
case with dilute HNO3) and forms metastannic acid(H2SnO3). NO2 gas is produced as usual.
Here you can mark a difference. Stannic nitrate is not formed rather metastannic acid is
formed.
0
(conc.)
6.
+4
+5
Sn + HNO 3
+4
H2 SnO 3 + NO 2 + H 2O
With Ag using conc. HNO3:
0
+5
Ag + HN O3
+1
+4
AgNO 3 + N O2 + H 2O
7.
With Hg using mod. conc. HNO3: Hg will be oxidised to +1(ous) state and HNO3 will
give NO. Conc. HNO3 would oxidise Hg to +2(ic) state.
0
+5
Hg + HNO 3
(mod. conc.)
+1
+2
Hg2(NO 3) 2 + N O + H 2O
SAQ 7 : Predict and balance the following:
(i)
HNO3(conc.) + Cu ------>?
--->?
(iii)
Fe + HNO3(v. dil.)----->?
>?
(ii)
HNO3(mod. conc.)+ Mg ---(iv)
Zn + HNO3(dil)----
(D) Hydrogen Displacement:
We know that metals lying above H in the electrochemical series can displace H2 gas from dilute
acids. However, since HNO3 is a strong oxidising agent which is reduced to NO, N2O or
NH4NO3 depending the concentration of acid, metals lying above H in the series cannot displace
H2 gas even from dilute HNO3. The question of H2 displacement does not arise in case of
metals lying below H in the series like Cu, Ag, Hg etc. However there are two exceptions in the
former category, Mg and Mn, which react with very very dilute(1%) HNO3 to produce H2 gas.
0
+5
Mg + HNO3
+2
0
Mg(NO 3)2 + H2
(v.v dilute)
(R)
same is for Mn
(E) Passivity of Metals:
Metals like Al, Fe, Co, Ni, Cr become passive with conc. HNO3. Passivity means, these metals
will react with conc. HNO3 for sometime liberating NO2 gas and after that the reaction will stop.
The metal is said to be passive. It does not react any more. Some authors believe that this is due
to the formation of a thin layer of metallic oxide which deposits over the surface of the metal in
the initial period of reaction, which become impervious (does not allow) to HNO3 to go inside the
metal. However, to this author, the actual cause is not yet known.
(F) Aqua ragia and solubility of noble metals:
Noble metals like Au and Pt are not soluble even in conc. HNO3. A 3:1 mixture of conc. HCl
and conc. HNO3 called the aqua ragia is able to dissolve these metals. This has been discussed
in the topic hydrochloric acid. But can you say what happens when aqua ragia is heated? The
nascent Cl atoms unite to form Cl2 gas.
HNO3 + HCl ------> NOCl + Cl2 + H2O
(R)
The ON of N changes from +5 to +3 and that of Cl changes from -1 to 0.
USES:
1.
It is used for the manufacture of artificial fertilizers such as ammonium nitrate, potassium
nitrate etc.
2.
Preparation of aqua ragia which dissolves Au and Pt required for jewellery.
Test: On heating it gives reddish brown fumes.
SULPHURIC ACID(H2SO4):
Pure sulphuric acid is a viscous colourless liquid. Commercial conc. H2SO4 that we use in the
laboratory is 98.3% H2SO4(rest 1.7% water). This is a constant boiling mixture(azeotropic mixture)
and boils at 3380C. H2SO4 cannot be prepared in the laboratory like HNO3 and HCl, because
sulphuric acid has low volatility. It cannot be displaced by a more volatile acid from a sulphate
salt.
Manufacture of H2SO4 by Contact Process:
It consists of three steps.
Step-I: Preparation of SO2: Sulphur dioxide is prepared either by burning of sulphur or
roasting of sulphide ore such as iron pyrites(FeS2).
S + O2 ------------> SO2
(R),
(Show the changes
in ON)
+2 -1
+3
0
FeS 2 + O2
-2
+4
Fe2O 3 + SO 2
(R)
Here two oxidations have taken place. Fe is oxidised from +2 to +3 state while S is also oxidised
from -1 to +4 state. O has reduced from 0 to -2 state. Note that often wrongly it is believed that
in FeS2, the ON of Fe is +4. It is not true as the sulphide here is a disulphide having S-S bond and
the two sulphur together carry 2- charge. So the ON per S atom is -1.
Step-II: Preparation of SO3: Sulphur dioxide obtained in step-I is mixed with air(O2)
in presence of V2O5 (Pt was used before) catalyst activated by K2O at a high pressure and at a
temperature of nearly 5000C to produce SO3. The reaction is exothermic.
+4
0
SO 2 + O 2
V 2O 5 (K 2 O)
5000C
+6 -2
SO3 + 98 Kjoules
(R)
Step-III: Preparation of Oleum and H2SO4: SO3 produced in step-III is not directly mixed
with water as it is highly dangerous. The reaction of SO3 with H2O is highly exothermic and the
heat produced suddenly boils off H2SO4 as mist causing catastrophic consequences. Therefore
SO3 is first mixed with conc.H2SO4 to produce H2S2O7(oleum) or pyrosulphuric acid to which
water is then added in required quantities to get sulphuric acid of desired concentration.
SO3 + H2SO4(conc) ------> H2S2O7(oleum)
(M)
(see for yourself)
H2S2O7 + H2O ----------> 2 H2SO4
(M)
PROPERTIES:
1.
Acidic Properties: It is a dibasic strong acid and therefore forms two series salts(acidic
and normal)
NaOH + H2SO4 --------> NaHSO4(acidic salt) + H2O
(M)
NaOH(excess)+ NaHSO4 -------> Na2SO4 + H2O
(M)
Dilute H2SO4 reacts with any alkali or a base to form salt and water. It reacts with a sulphite salt
to produce SO2, a carbonate salt to produce CO2, a sulphide salt to produce H2S. All these are
metathesis reactions involving no change of ON. Try for yourself by formulating one example in
each type.
2
As oxidising agent:
(a)
Dilute H2SO4: Metals lying above H in the electrochemical series displace H2
from dilute H2SO4. Unlike dil. HNO3; dil H2SO4 does not give oxide of sulphur. Metals below H,
do not react with dil. H2SO4.
Zn +H2SO4(dil.)---------> ZnSO4 + H2
(R) Show changes in ON.
Al + H2SO4(dil)---------> Al2(SO4)3 + H2 (R) Show the changes in ON.
Cu + H2SO4(dil) -------> NO REACTION
(since Cu lies below H)
(b)
Conc. H2SO4:
I.
With Metals:
It is a powerful oxidising agents. Metals both below and above H in the electrochemical series
react with it. In this case H2SO4 is reduced to SO2. The ON of S changes from +6 to +4. Look
to these reactions.
ZnSO4 + SO2 + H2 O
+2
+6
0
+4
+2
+6
0
Zn + H2SO4
+4
CuSO 4 + SO 2+ H 2O
Cu + H2SO 4
(R)
(R)
Can you not balance these equations? You are advised to balance these equations either by ON
method or by partial equation method. In partial equation method, H2SO4 has to break down as
a mixture of SO2, H2O and nascent oxygen[O], then to proceed in the usual way(Refer
section-I).
II.
With Nonmetals:
Conc. H2SO4 reacts with nonmetals like C, S, P4 etc. in the same manner as conc. HNO3.
H2SO4 is reduced to SO2. Look to the following reactions after visualising the logic of each
reaction, balance them of your own for practice.
0
+4
+6
0
S
(R)
+4
+6
+
+4
C O 2 + S O 2 + H 2O
C + H2 SO 4
H2SO 4
SO2
H2O
+
Here S is oxidised to SO2, and H2SO4 is reduced to the same SO2. With conc. HNO3, however
S is oxidised to SO3 and then to H2SO4.
P4 will be oxidised to P2O5 which reacts with H2O to produce H3PO4 as usual.
0
+5
+6
+4
H3PO 4 + SO 2 + H2O
P 4 + H2SO4
(R)
III.
With Common Reducing Agents:
Strong reducing agents like HBr(or KBr), HI(or KI) react with conc. H2SO4 in the same manner
as conc. HNO3. Bromide and iodide will produce Br2 and I2 respectively.
+6
+6
H2SO 4
0
-1
H2 SO 4 + HBr
-1
KI
+4
Br2 + SO 2 + H2O
0
(R)
+4
I2 + SO 2 + K2SO4
(R)
3.
Dehydrating Properties:
Conc. H2SO4 is a strong dehydrating agent and removes water from substances like glucose and
other carbohydrates such as cellulose etc. This is called charring reaction, in which carbon and
water are formed and the white substance turns black.
4.
Precipitation reactions:(Metathesis):
Soluble salts of Ba2+ and Pb2+ form their insoluble precipitates when react with dilute H2SO4.
BaCl2 + H2SO4 -------> BaSO4(white ppt) + HCl
(M)
Pb(NO3)2 + H2SO4 ------> PbSO4(white ppt) + HCl
(M)
USES: (i). It is used for the manufacture of HCl and HNO3
(ii)Used in the manufacture of fertilisers such as ammonium sulphate and superphosphate
of lime.