OXYGEN FAMILY

Group 16 of periodic table includes oxygen (O), Sulphur (S), Selenium (Se), Tellurium (Te) and Polonium (Po). These are collectively known as Chalcogens or ore forming elements. The first four members are non – metals. Polonium is having metallic character and is radioactive.

General introduction to Oxygen family

Oxygen

Its name came from oxys + genes, "acid former" was discovered in 1772 by Schelle and independently by Pristeley in 1774.

Oxygen is the most common element in the earth's crust and makes up about 20% of the air we breathe. Oxygen is gas at room temperature and is colourless, odourless and testless. Liquid oxygen has a slight blue colour. Oxygen exists in two allotropes, oxygen in air (O₂) and ozone in air (O₃) ozone is very reactive and used bleaching and disinfectant reagent.

Sulfur

Sulfur was classified as an element in 1777 by Lavoisier. Pure sulfur is testless and odourless with white yellow colour. Sulfur is the tenth most abundant element in the known universe. Sulfur exists in three allotropic forms. Two are crystalline and one is amorphous. The most significant compound of sulfur used in modern industrialized societies is sulphuric acid (H₂SO₄). Sulphur dioxide (SO₂) is used in bleaching and refrigeration. SO₂ gas reacts with water vapour in air to produce a weak acid, sulfurous acid (H₂SO₃) which contributes to the acid rain problem.

Selenium

It was discovered by Berzelius in 1818. Its name came from the Greek work Selene which is used for moon. It is an important semiconductor which is particularly sensitive to light. There are at lest two other allotropes of the element, including a semi – metallic state.

Tellurium (Te)

Discovered by Von Reicherstein in 1782. It is brittle metalloid that is relatively rare. It is named from the Latin tellus for 'earth'. Tellurium can be alloyed with same metals to increase their machinability and is a basic ingredient in the manufacture of blasting caps.

Polonium (Po)

It was discovered in 1898 by Marie Curie and named for her native country Poland. Although radioactive, polonium has a few commercial uses. It is synthesized in the lab rather its recovery from minerals.

GENERAL TRENDS IN PHYSICAL PROPERTIES

Illustration 1. Name two compounds in which oxygen has oxidation state different from – 2. Give oxidation states also.

Solution: (i) OF₂ (Oxidation state of O = +2)

(ii) H₂O₂ (Oxidation state of O = -1)

CHEMICAL PROPERTIES OF GROUP 16

The group 16 elements show a lot of variation in their chemical behaviours. Both O and S are reactive but reactivity decreases down the group and oxygen reacts with almost all the elements.

1. Hydrides

The elements of group 16 form volatile hydrides such as H₂O, H₂S, H₂Se, H₂Te an H₂Po.

H₂O can be prepared by heating H₂ in oxygen while other hydrides may be obtained by the action of acids on metal sulphides, selenoids and fluorides, e.g.

$N{a}_{2}S+H_{2}S{O}_4\to N{a}_{2}S{O}_4+H_{2}S$

$N{a}_{2}Se+2HCl\to 2NaCl+H_{2}Se$

Some Important features of hydrides

(i) Physical nature

H₂O is colourless, odourless liquid which is a very essential part of living organism as well as non – living.

Other hydrides are colourless poisonous gases with very unpleasant smell. The volatile nature of hydrides first increases from H₂O to H₂S and then decreases. The low volatile character of water is due to intermolecular hydrogen bonding.

(ii) Acidity

These hydrides are weakly acidic in nature. The acidic nature increases from H₂O to H₂Te. H₂S is a weak diprotic acid.

$H_{2}S\rightleftharpoons H{S}^{-}+H^{+}$

$HS\rightleftharpoons -S^{2-}+H^{+}$

The distance between central atom and hydrogen increases down the group which favours in easy release of hydrogen as proton.

H₂O < H₂S < H₂Se < H₂Te

Acidity order

(iii) Reducing character

Except water all other hydrides are reducing agents. This character increases from H₂S to H₂Te. This is also due to decrease in M – H bond strength.

Illustration 2. Which hydride has greater bond angle:

$H_{2}O,H_{2}S,H_{2}Seor{H}_{2}Te?$

Solution: As the electro negativity of the central atom decreases, the repulsions between element, hydrogen bond pairs decreases and hence the angle decreases accordingly. Thus, H₂O has the greatest bond angle (104.5°).

$H_{2}S$ (Hydrogen sulphide or sulphurated hydrogen)

It mainly occurs in volcanic eruptions, sewage gases, coal gas and in several spring waters.

Methods of preparation

(a) Laboratory preparation

$FeS+H_{2}S{O}_4\to FeS{O}_4+H_{2}S$

$S{b}_2{S}_3+6HCl\to 2SbC{l}_3+3H_{2}S$

Physical properties

It is colourless gas with unpleasant smell of rotten eggs, slightly heavier than air. It is poisonous in nature.

Chemical properties

(a) Combustibility

It burns with blue flame in $O_2.$

$2H_{2}S+3O_2\to 2H_{2}O+2S{O}_2$

In restricted supply of oxygen, sulphur is formed due to incomplete combustion.

$2H_{2}S+O_2\to 2H_{2}O+2S$

(b) Thermal decomposition

$H_{2}S\underset{{310}^{o}C-{1700}^{o}C}{\to }{H}_2+S$

(c) Reducing nature

(i) $H_{2}S+X_2\to 2HX+S\left(X_2=F_2,C{l}_2,B{r}_{2}or{I}_2\right)$

(ii) $2FeC{l}_3+H_{2}S\to 2FeC{l}_2+2HCl+S$

(iii) $S{O}_2+2H_{2}S\to 3S+2H_{2}O$

(iv) $H_2{O}_2+H_{2}S\to 2H_{2}O+S$

(v) $O_3+H_{2}S\to H_{2}O+S+O_2$

(vi) $H_{2}S{O}_4+H_{2}S\to 2H_{2}O+S{O}_2+S$

(vii) $H_{2}S+HN{O}_3\to 2N{O}_2+S+2H_{2}O$

(d) Acidic nature

$H_{2}S\rightleftharpoons H{S}^{-}+H^{+}$

$H{S}^{-}\rightleftharpoons S^{2-}+H^{+}$

$NaOH+H_{2}S\to NaSH+H_{2}O$

Sodium hydro

Sulphide

$2NaOH+H_{2}S\to N{a}_{2}S+2H_{2}O$

Sodium sulphide

(e) Reactions with metals and metal oxides

$2Na+H_{2}S\to N{a}_{2}S+H_2$

$Cu+H_{2}S\to CuS+H_2$

$CaO+H_{2}S\stackrel{\Delta }{\to }CaS+H_{2}O$

$ZnO+H_{2}S\stackrel{\Delta }{\to }ZnS+H_{2}O$

(f) Action on salts

(i) Sulphides precipitated in acidic medium

$CuS{O}_4+H_{2}S\to CuS+H_{2}S{O}_4$

$Pb{(C{H}_{3}COO)}_2+H_{2}S\to PbS+2C{H}_{3}COOH$

$CdS{O}_4+H_{2}S\to CdS+H_{2}S{O}_4$

(ii) Sulphides precipitated in alkaline medium

$MnC{l}_2+H_{2}S\to MnS+2HCl$

$Ni{(N{O}_3)}_2+H_{2}S\to NiS+2HN{O}_3$

$ZnS{O}_4+H_{2}S\to ZnS+H_{2}S{O}_4$

CoS, NiS; Black, ZnS; white, MnS; flesh coloured

(iii) Sulphides which are precipitated neither in acidic nor in alkaline medium

There are the sulphides of Cr, Al, Mg, Ba, Sr, Ca, K and Na.

(g) Formation of polysulphides

$Ca{(OH)}_2+H_{2}S\to CaS+2H_{2}O$

$CaS+4H_{2}S\to Ca{S}_5+4H_2$

$H_{2}S$is an important laboratory reagent for the detection of basic radicals in qualitative analysis.

Illustration 3. Why hydride of oxygen is a liquid whereas hydride of sulphur is a gas?

Solution: Because in hydride of oxygen, i.e., H₂O there is a tendency towards hydrogen bonding.

2. Halides

The elements of group 16 forms many kind of halides from which di, tetra and hexa halides are common. These halides can be obtained by direct interaction of chalcogens to halogen atoms.

$\frac{1}{8}{S}_8\left(s\right)+3F_2\left(g\right)\to S{F}_6\left(g\right)$

$\frac{1}{4}{S}_8\left(\ell \right)+C{l}_2\left(g\right)\to S_{2}C{l}_2\left(\ell \right)$

$Te\left(s\right)+2C{l}_2\left(g\right)\to TeC{l}_4\left(s\right)$

$Te\left(s\right)+2I_2\left(g\right)\to Te{I}_4\left(s\right)$

Oxygen and sulphur forms a number of halides. eg. ClO₂, OF₂, I₂O₅, Cl₂O₇, S₂Cl₂, SF₄, SF₆ etc.

3. Oxides

This group elements for in a number of oxides.

(i) All elements (except Se) forms monoxide.

(ii) All elements form dioxide with formula MO₂, SO₂ is a gas, SeO₂ is volatile solid.

While TeO₂ and PoO₂ are non – volatile crystalline solids.

(iii) All the elements of this group form MO₃ type trioxides. The best known trioxides is SO₃.

In solid state it can exist either as an cyclic trimer, (SO₃)₃ or as a linear chain cross – linked in to sheets. Selenium trioxide is a cyclic tetramer whereas TeO₂ is a solid with network structure.

OZONE $(O_3)$

Ozone was first observed through a rotten smell by Van Marum in 1758 which was confirmed by Schonbiln in 1840. Sorret established its formula,$O_3$ , and pointed out that the ozone is an allotrope of oxygen.

Ozone prevents the living world from the harmful radiations (UV) coming from sun. Its layer lies 12 to 15 miles above earth's surface. Near earth's surface it is decomposed by dust particles.

Preparation of ozone

Ozone is prepared in laboratory by passing silent electric discharge through dry oxygen. By passing the electric current some of the oxygen molecules dissociate and then atomic oxygen combines with oxygen molecules to form ozone.

$O_2\stackrel{energy}{\to }O+O$

$O_2+O\to O_3$

$3O_2\rightleftharpoons 2O_3$-energy

The mixture thus obtained contains 5 – 10% ozone by volume and the mixture is called as ozonised oxygen.

The apparatus used for this purpose is known as ozoniser.

Recovery of pure ozone from ozonised oxygen

The ozonised oxygen is passed through a spiral cooled by liquid air. Ozone condenses as its condensation temperature is $-{112.4}^{o}C.$ The liquid ozone thus obtained may contain some dissolved oxygen. This liquid is fractionally distilled to get pure ozone.

Some other ozone forming reactions

(i) Fluorine reacts with water at a very low temperature when mixture of ozonised oxygen is formed.

$2F_2+2H_{2}O\to 4HF+O_2$

$2F_2+3H_{2}O\to 6HF+O_3$

(ii) By action of ultraviolet rays, oxygen can be converted into ozone.

$O_2\stackrel{uv}{\to }O+O$

$O_2+O\to O_3$

(iii) The electrolysis of acidified water with high current density and platinum anode yields ozone on the anode. Nascent oxygen discharged at anode combines with molecular oxygen to form ozone.

Physical properties

(i) It is having characteristic rotten smell with pale blue colour.

(ii) It can be liquefied into pale blue liquid at $-{112.4}^{o}C.$ If we reach at$-{249.7}^{o}C$, it converts into violet black crystals.

(iii) It is heavier than air.

(iv) Solubility

It is slightly soluble in water but more soluble in turpentine oil, glacial acetic acid or carbon tetrachloride.

(v) It causes headache and nausea when inhaled in small amount.

Chemical properties

(a) Decomposition

It is unstable and easily decomposes into oxygen at ${300}^{o}C.$

$2O_3\stackrel{\Delta }{\to }3O_2+68Kcal$

$Mn{O}_2,$platinum black, silver, lead dioxide, etc., decompose ozone at ordinary temperature, i.e. they catalyse its decomposition.

(b) Oxidising nature

It acts as a strong oxidising agent due to the case with which it can liberate nascent oxygen.

The potential equation is

$O_3\to O_2+O$

The oxidation potential in acidic medium is +2.07 V.

$O_3+2H^{+}+2e^{-}\to O_2+H_{2}O;E^o=2.07V$

Examples

(i) It oxidises HCl into $C{l}_2,$ HBr into $B{r}_2$ and HI into iodine.

$O_3\to O_2+O$

$\frac{2HCl+O\to H_{2}O+C{l}_2}{2HCl+O_3\to H_{2}O+C{l}_2+O_2}$

(ii) Acidified ferrous sulphate into ferric sulphate

$O_3\to O_2+O$

$\frac{2FeS{O}_4+H_{2}S{O}_4+O\to F{e}_2{(S{O}_4)}_3+H_{2}O}{2FeS{O}_4+H_{2}S{O}_4+O_3\to F{e}_2{(S{O}_4)}_3+H_{2}O+O_2}$

(iii) It liberates iodine from neutral KI solution.

$O_3\to O_2+O$

$\frac{2KI+H_{2}O+O\to 2KOH+I_2}{2KI+H_{2}O+O_3\to 2KOH+I_2+O_2}$

(iv) Lead sulphide (black) is oxidised to lead sulphate (colourless).

$\left[O_3\to O_2+O\right]\times 4$

$\frac{PbS+4O\to PbS{O}_4}{PbS+4O_3\to PbS{O}_4+4O_2}$

Similarly, CuS, ZnS and CdS are oxidised to corresponding sulphates.

(v) It oxidise nitrites into nitrates, sulphites into sulphates, arsenites to arsenates, manganate to permanganate and ferrocyanide to ferricyanide.

$NaN{O}_2+O_3\to NaN{O}_3+O_2$

$N{a}_{2}S{O}_3+O_3\to N{a}_{2}S{O}_4+O_2$

$N{a}_{3}As{O}_3+O_3\to N{a}_{3}As{O}_4+O_2$

sodium sodium

arsenite arsenate

$2K_{2}Mn{O}_4+H_{2}O+O_3\to 2KMn{O}_4+2KOH+O_2$

Potassium Potassium

manganate permanganate

$2K_{4}Fe{(CN)}_6+H_{2}O+O_3\to 2K_{3}Fe{(CN)}_6+2KOH+O_2$

Potassium Potassium

Ferrocyanide ferricyanide

(vi) Moist iodine is oxidised to iodic acid

$\left[O_3\to O_2+O\right]\times 5$

$\frac{I_2+H_{2}O+5O\to 2HI{O}_3}{I_2+H_{2}O+5O_3\to 2HI{O}_3+5O_2}$

Iodic acid

Similarly moist sulphur, phosphorus and arsenic are oxidised to their corresponding oxyacids.

$S+H_{2}O+3O_3\to H_{2}S{O}_4+3O_2$

$2P+3H_{2}O+5O_3\to 2H_{3}P{O}_4+5O_2$

$2As+3H_{2}O+5O_3\to 2H_{3}As{O}_4+5O_2$

(vii) $H_{2}S$ is oxidised to sulphur

$H_{2}S+O_3\to H_{2}O+S+O_2$

(viii) Alkaline KI is oxidised to potassium iodate and periodate.

$KI+3O_3\to \underset{\begin{array}{c}Potassium\\ Iodate\end{array}}{KI{O}_3}+3O_2$

$KI+4O_3\to \underset{\begin{array}{c}Potassium\\ periodate\end{array}}{KI{O}_4}+4O_2$

(ix) Dry iodine is oxidised to yellow powder, $I_4{O}_9.$

$\left[O_3\to O_2+O\right]\times 9$

$\frac{2I_2+9O\to I_4{O}_9}{2I_2+9O_3\to I_4{O}_9+9O_2}$

(x) Silver metal is blackened due to alternate oxidation of the metal and reduction of oxide.

$2Ag+O_3\to A{g}_{2}O+O_2$

$A{g}_{2}O+O_3\to 2Ag+O_2$

(xi) Mercury in the presence of ozone is oxidised to suboxide which dissolves in mercury. It starts sticking to glass and loses mobility. Hence, mercury loses its meniscus in contact with ozone. This is known as tailing of mercury.

$2Hg+O_3\to H{g}_{2}O+O_2$

(xii) Ozone reacts with KOH and forms potassium ozonide, $K{O}_3,$which is an orange coloured solid and contains paramagnetic $O_3^{-}$ ion.

$2KOH+5O_3\to K{O}_3+5O_2+H_{2}O$

In all above reactions, oxygen is evolved. There are few reactions in which whole of the oxygen are used up in the process of oxidation.

1. Oxidises $S{O}_{2}toS{O}_3$

$3S{O}_2+O_3\to 3S{O}_3$

2. Acidified stannous chloride is oxidised to stannic chloride.

$3SnC{l}_2+6HCl+O_3\to 3SnC{l}_4+3H_{2}O$

(c) Bleaching property

It acts as a good bleaching agent due to release of nascent oxygen.

Coloured substance + O $\to colourless$

It bleaches oil, ivory, flour, starch, waxes wood pulp etc.

(d) Formation of ozonides

Ozone forms ozonides when reacts with unsaturated organic compounds containing double bond. The ozonides break to form carbonyl compounds when heated with water. This process is called ozonolysis. $H_2{O}_2$is evolved in most of the cases.

(e) Reaction with peroxide

Reaction of ozone with peroxide results in their mutual reduction with the liberation of oxygen.

$Ba{O}_2+O_3\to BaO+2O_2$

$H_2{O}_2+O_3\to H_{2}O+2O_2$

Structure of ozone

The bond length is intermediate between that for a single bond $(1.48\stackrel{o}{A}asin{H}_2{O}_2)$ and for a double bond$(1.21\stackrel{o}{A}asin{O}_2)$. Ozone is, therefore, considered to be a resonance hybrid of the following two forms.

Tests of ozone

(i) It has a strong rotten smell.

(ii) Metallic mercury loses its fluidity in contact with $O_3.$

(iii) It turns an alcoholic solution of benzidine brown.

(iv) It turns an alcoholic solution of tetramethyl base violet.

(v) It turns starch – iodide paper blue.

4. Oxyacids

The important oxyacids of S, Se and Te are given in the table.

But the oxyacids of sulphur are more important.

Extraction of sulphur

(a) Frasch Process

In this method of extraction, sulphur is extracted from underground deposits. The underground deposits are melted through passing superheated ($\approx$ 440 K) steam down the beds of ore. The molten sulphur is then blown out with the help of compressed air.

(b) From natural gas

Sulphur can be recovered from the natural gas which includes separating out of H_­2S by absorption into monoethanolamine and then converting H₂S in to sulphur by the following sequence of reactions.

$H_{2}S+\frac{1}{2}{O}_2\underset{combusition}{\overset{lowtemperature}{\to }}\frac{1}{8}{S}_8+H_{2}O$

Some amount of SO₂ is also forming during the above reaction by oxidation of H₂S which can be converted in to sulphur by passing more of H₂S at about 303 K in the presence of a catalyst Fe₂O₃.

$2H_{2}S+3O_2\to 2S{O}_2+2H_{2}O$

$2H_{2}S+S{O}_2\underset{303K}{\overset{F{e}_2{O}_{3}catalyst}{\to }}\frac{3}{8}{S}_8+2H_{2}O$

Allotropes of sulphur

Main three types of allotropes of sulphur are discussed here:

(a) Rhombic sulphur

It is most stable and common form of sulphur also called as $\alpha$ - sulphur. It has bright yellow colour. It is insoluble in water and carbon disulphide. Its density is 2.07 gm cm^-3 and exists as S₈ molecules. The 8 sulphur atoms in S₈­ molecule forms a puckered ring as shown below. It is solid in nature.

(b) Monoclinic sulphur

It is another form of sulphur which is stable only above 369 K. It is dull yellow coloured solid, also called $\beta$ - sulphur. It is soluble in CS₂ but insoluble in H₂O.

It slowly changes into rhombic sulphur. It also exist as S₈ molecules which have puckered ring structure. It however, differs from the rhombic sulphur in the symmetry of the crystals.

(c) Plastic sulphur

It is obtained by pouring molten sulphur to cold water. It is amorphous form of sulphur also called as $\gamma$ - sulphur. It is insoluble in water as well as CS₂.

Several other modifications of sulphur containing 6- 20 sulphur atoms per ring have been artificially synthesized. In cyclo S₆, the ring adopts the chain form dimensions.

Sulphuric acid (H₂SO₄)

It is the most important oxyacid of sulphur. It is also called as 'oil of vitriol' because of its large application in industries it is also known as 'king of chemicals'.

Manufacture of sulpuric acid

Sulphuric acid (H₂SO₄) is manufactured by contact process which involves following three steps.

(a) Production of sulphur dioxide

It is carried out by burning sulphur powder or roasting of sulphur rich ores.

$S_8+8O_2\to 8S{O}_2$

$4Fe{S}_2+11O_2\to 2Fe{O}_3+8O_2$

(b) Oxidation of SO₂ to SO₃

$2S{O}_2\left(g\right)+O_2\left(g\right)2S{O}_3;\Delta H=-106.6KJ$

This step is the key step of the whole process so with applying Le – Chatelier's principle we can enhance the forward reaction rate to achieve better yield of SO₃.

(i) Temperature

As the reaction is already exothermic so temperature must be kept low. An optimum temperature is maintained between 673 – 723K.

(ii) Pressure

The catalyst is required due to low temperature, so platinised asbestos or divanadium pentoxide (V₂O₅) is used.

(iv) Purity of gases

To prevent the poisoning of catalyst, the gases must be free from the impurities of As₂O₃, dust particles and moisture.

(v) Excess of oxygen

To have a better yield of SO₃, O₂ is used in excess.

(c) Conversion of SO₃ in H₂SO₄

$S{O}_3+H_{2}S{O}_4\to \underset{Oleum}{H_2{S}_2{O}_7}$

Oleum is then diluted with calculated amount of water

$H_2{S}_2{O}_7+H_{2}O\to 2H_{2}S{O}_4$

The sulphuric acid obtained through contact process is generally of 96 – 98% purity.

Physical Properties of H₂SO₄

(i) Pure H₂SO₄ is colourless but commercial acid is yellow in colour due to presence of some impurities.

(ii) Highly concentrated H₂SO₄ (98%) has a specific gravity of 1.84 and high b.p indicates about the hydrogen bonding as shown below:

(iii) Conc. H₂SO₄ has a great affinity for water. Its dissolution is highly exothermic in water.

(iv) It can burn the skin while comes in contact with it.

Chemical properties of H₂SO₄

(i) Dissociation

On heating strongly it dissociates into SO₃ and H₂O.

$H_{2}S{O}_4\rightleftharpoons H_{2}O+S{O}_3$

(ii) Acidic nature

$H_{2}S{O}_4\left(aq\right)\to H^{+}\left(aq\right)+HS{O}_4^{-}\left(aq\right);K_1=1\times {10}^3$

$HS{O}_{{}_4}^{-}\left(aq\right)\to H^{+}\left(aq\right)+S{O}_4^{-2}\left(aq\right);K_2=1.2\times {10}^{-2}$

Thus it forms two series of salts, sulphates and bisulphates. eg

$NaOH+H_{2}S{O}_4\to \underset{sod.bisulphate}{NaHS{O}_4}+H_{2}O$

$2NaOH+H_{2}S{O}_4\to \underset{sod.sulphate}{N{a}_{2}S{O}_4}+H_{2}O$

(iii) As a dehydrating agent

Due to strong affinity for water, H₂SO₄ acts as a powerful dehydrating agent.

(a) Drying of gases

Certain gases like CO₂, SO₂, Cl₂, HCl etc, can be dried by passing through conc. H₂SO₄.

(b) Charring

Concentrated H₂SO₄ reacts with sugar, wood, paper etc to form black mass of carbon. This phenomenon is called charring.

$C_{12}{H}_{22}{O}_{11}\stackrel{conc.H_{2}S{O}_4}{\to }11H_{2}O+12C$

(c) Dehydration of acids

Acids like formic and oxalic acid get dehydrated according to the equation,

$HCOOH\stackrel{conc.H_{2}S{O}_4}{\to }CO+H_{2}O$

$COOH-COOH\stackrel{conc.H_{2}S{O}_4}{\to }CO+C{O}_2+H_{2}O$

4. Oxidizing action

It is moderately strong oxidizing agent. It is intermediate between phosphoric and nitric acids. Both metals and non – metals are oxidised by conc. H₂SO₄, which is reduced to SO₂.

Some of oxidizing reaction are

(i) $C+2H_{2}S{O}_4\to 2H_{2}O+2S{O}_2+C{O}_2$

(ii) $\frac{1}{8}{S}_8+2H_{2}S{O}_4\to 2H_{2}O+3S{O}_2$

(iii) $P_4+10H_{2}S{O}_4\to 4H_{3}P{O}_4+10S{O}_2+4H_{2}O$

(iv) $2HBr+H_{2}S{O}_4\to 2H_{2}O+S{O}_2+B{r}_2$

(v) $H_{2}S+H_{2}S{O}_4\to 2H_{2}O+S+S{O}_2$

(vi) $Cu+2H_{2}S{O}_4\to CuS{O}_4+S{O}_2+2H_{2}O$

$Pb+2H_{2}S{O}_4\to PbS{O}_4+S{O}_2+2H_{2}O$

5. Reaction with salts

(a) Dilute H₂SO₄ decomposes carbonates, bicarbonates, sulphides, sulphites, thiosulphates and nitrites at room temperatures.

$H_{2}S{O}_4+N{a}_{2}C{O}_3\to N{a}_{2}S{O}_4+H_{2}O+C{O}_2\uparrow$

$H_{2}S{O}_4+2NaHC{O}_3\to N{a}_{2}S{O}_4+2H_{2}O+2C{O}_2\uparrow$

$H_{2}S{O}_4+N{a}_{2}S\to N{a}_{2}S{O}_4+H_{2}S\uparrow$

(b) Salts like chlorides, fluorides, nitrates, acetates, oxalates are decomposed by hot conc. H₂SO₄ liberating their corresponding acids.

$H_{2}S{O}_4+CaC{l}_2\to CaS{O}_4+2HCl\uparrow$

$H_{2}S{O}_4+Ca{F}_2\to CaS{O}_4+2HF$

$H_{2}S{O}_4+2KN{O}_3\to K_{2}S{O}_4+2HN{O}_3$

Unlike fluorides and chlorides, the bromides and iodides are decomposed by conc. H₂SO₄ to give Br₂ and I₂ respectively.

$3H_{2}S{O}_4+2NaBr\to NaHS{O}_4+S{O}_2+B{r}_2+2H_{2}O$

$3H_{2}S{O}_4+2NaI\to NaHS{O}_4+S{O}_2+I_2+2H_{2}O$

Uses of sulphuric acid

(i) As an acid in laboratory and industries.

(ii) In storage batteries and lead accumulators.

(iii) As a dehydrating agent.

(iv) In fertilizers.

(v) In drying of gases.

(vi) In preparing other acids like HCl, HF etc.

(vii) Detergent industries.

(viii) Manufacture of pigments paints and dyestuff.

(ix) Metallurgical applications.