Class 12 Chemistry Chapter 7 THE P-BLOCK ELEMENTS
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NCERT Notes for Class 12 Chemistry Chapter 7 THE P-BLOCK ELEMENTS
Class 12 Chemistry Chapter 7 THE P-BLOCK ELEMENTS
The elements in which the last electron enters in the valence p-sub shell are called the p-block elements. They include elements of the groups 13 to 18. Their general outer electronic configuration is ns2np1-6 (except He which has 1s2 configuration).They includes metals, non-metals and metalloids.
Group 15 Elements
Group 15 includes nitrogen (N), phosphorus (P), arsenic (As), antimony (Sb) and bismuth (Bi). As we go down the group, the metallic character increases. Nitrogen and phosphorus are non-metals, arsenic and antimony metalloids and bismuth is a typical metal. The valence shell electronic configuration of these elements is ns2np3. The s orbital in these elements is completely filled and p orbitals are half-filled, making their electronic configuration extra stable.
Covalent and ionic radii increase down the group. There is a considerable increase in covalent radius from N to P. However, from As to Bi only a small increase in covalent radius is observed. This is due to the presence of completely filled d or f orbitals in heavier members. Ionisation enthalpy decreases down the group due to gradual increase in atomic size. Because of the extra stable half-filled p orbitals and smaller size, the ionisation enthalpy of the group 15 elements is much greater than that of group 14 elements.
Oxidation states and trends in chemical reactivity
The common oxidation states of these elements are –3, +3 and +5. The tendencies to exhibit –3 oxidation state decreases down the group due to increase in size and metallic character. The last member of the group, bismuth does not form any compound in –3 oxidation state. The stability of +5 oxidation state decreases and that of +3 state increases (due to inert pair effect) down the group. Nitrogen exhibits + 1, + 2, and + 4 oxidation states also when it reacts with oxygen. Phosphorus also shows +1 and +4 oxidation states in some oxoacids. Nitrogen is restricted to a maximum covalency of 4 since only four orbitals (one s and three p) are available for bonding.
Anomalous properties of nitrogen
Nitrogen differs from the rest of the members of this group due to its smaller size, high electro negativity, high ionisation enthalpy and non-availability of d orbitals. Some of the anomalous properties shown by nitrogen are:
- Nitrogen has the ability to form pπ–pπ multiple bonds with itself and with other elements like C and O. Other elements of this group do not form pπ–pπ bonds.
- Nitrogen exists as a diatomic molecule with a triple bond (one s and two p) between the two atoms. So its bond enthalpy is very high. While other elements of this group are poly atomic with single bonds.
- The single N–N bond is weak. So the catenation tendency is weaker in nitrogen.
- Due to the absence of d orbitals in its valence shell, the maximum covalency of nitrogen is four
- N cannot form dπ–pπ bond. While Phosphorus and arsenic can form dπ–dπ bond with transition metals and with C and O.
Hydrides of Group 15 Elements
All the elements of Group 15 form hydrides of the type EH3 (where E = N, P, As, Sb or Bi). The hydrides show regular gradation in their properties. The bond dissociation enthalpy of E – H decreases from NH3 to BiH3. So the thermal stability decreases from NH3 to BiH3 and the reducing character increases.
Ammonia is only a mild reducing agent while BiH3 is the strongest reducing agent amongst all the hydrides. Basicity decreases in the order NH3 > PH3 > AsH3 > SbH3 > BiH3. The melting point of these hydrides increases from top to bottom. This is due to increase in the atomic size of the central atom which increases the van der Waal’s force of attraction. NH3 has the highest melting and boiling points due to inter molecular hydrogen bonding. All these hydrides have pyramidal geometry.
Q1. Though nitrogen exhibits +5 oxidation state, it does not form pentahalides. Give reason.
Nitrogen with n = 2, has s and p orbitals only. It does not have d orbitals to expand its covalence beyond four. That is why it does not form pentahalide.
Q2. PH3 has lower boiling point than NH3. Why?
Unlike NH3, PH3 molecules are not associated through inter molecular hydrogen bonding in liquid state. That is why the boiling point of PH3 is lower than NH3.
Preparation: Dinitrogen is produced commercially by the liquefaction and fractional distillation of air.
In the laboratory, dinitrogen is prepared by treating an aqueous solution of ammonium chloride with sodium nitrite.
NH4CI(aq) + NaNO2(aq) → N2(g) + 2H2O(l) + NaCl (aq)
It can also be obtained by the thermal decomposition of ammonium dichromate.
(NH4)2Cr2O7 →Heat N2 + 4H2O + Cr2O3
Very pure nitrogen can be obtained by the thermal decomposition of sodium or barium azide.
Ba(N3)2 → Ba + 3N2
Dinitrogen is inert at room temperature because of the high bond enthalpy of N≡ N bond. At higher temperatures, it directly combines with some metals to form ionic nitrides and with non-metals to form covalent nitrides.
6Li + N2 →Heat 2Li3N
3Mg + N2 →Heat Mg3N2
It combines with hydrogen at about 773 K in the presence of a catalyst (Haber’s Process) to form ammonia:
Dinitrogen combines with dioxygen at very high temperature (at about 2000 K) to form nitric oxide N2 + O2 → 2 NO
Uses: 1. The main use of dinitrogen is in the manufacture of ammonia and other industrial chemicals containing nitrogen (e.g., calcium cyanamide).
- It also used to create an inert atmosphere in metallurgy.
- Liquid dinitrogen is used as a refrigerant to preserve biological materials, food items and in cryosurgery.
Preparation: In laboratory, ammonia is obtained by treating ammonium salts with caustic soda (NaOH) or slaked lime.
(NH4)2SO4 + 2NaOH → 2NH3 + 2H2O + Na2SO4
2NH4Cl + Ca(OH)2 → 2NH3 + 2H2O + CaCl2
On a large scale, ammonia is manufactured by Haber’s process.
N2(g) + 3H2(g) → 2NH3(g)
In accordance with Le Chatelier’s principle, high pressure of about 200 atm, a temperature of about 773 K and the catalyst such as iron oxide with small amounts of K2O and Al2O3 are employed to increase the rate of this reaction.
Ammonia is a colourless gas with pungent smell. It is highly soluble in water because of its ability to form inter molecular hydrogen bond with water. Liquid ammonia has high melting and boiling points because of inter molecular hydrogen bonding.
The ammonia molecule has a trigonal pyramidal geometry. It has three bond pairs and one lone pair of electrons.
Its aqueous solution is weakly basic due to the formation of OH– ions.
NH3(g) + H2O(l) → NH4+ (aq) + OH– (aq)
As a weak base, it precipitates the hydroxides of many metals from their salt solutions. For example, 2FeCl3 (aq )+ 3NH4OH (aq ) → Fe2O3.xH2O (s) + 3NH4Cl (aq)
ZnSO4 (aq)+ 2NH4OH (aq) → Zn(OH)2 (s) + (NH4)2SO4 (aq)
The presence of a lone pair of electrons on the nitrogen atom of the ammonia molecule makes it a Lewis base. It donates the electron pair and forms complex compounds with Cu2+, Ag+ etc. So it is used for the detection of these metal ions.
Cu2+ (aq) + 4 NH3(aq) → [Cu(NH3)4]2+(aq)
(blue) (deep blue)
Ag+ ( aq) + Cl− (aq) → AgCl (s)
(colourless) (white ppt)
AgCl (s) + 2NH3 (aq ) → [Ag (NH3)2]Cl (aq)
(white ppt) (colourless)
Uses: Ammonia is used i) to produce various nitrogenous fertilizers (ammonium nitrate, urea, ammonium phosphate and ammonium sulphate) ii) in the manufacture of nitric acid iii) liquid ammonia is used as a refrigerant.
Oxides of Nitrogen
Nitrogen forms a number of oxides in different oxidation states. They are:
1- Nitrous Oxide [Nitrogen (I) Oxide]: It is prepared by heating ammonium nitrate.
NH4NO3 → N2O + 2 H2O
It is a colourless, neutral gas. Its structure is:
2- Nitric Oxide [Nitrogen (II) Oxide]: It is prepared by treating sodium nitrite with acidified ferrous sulphate.
2 NaNO2 + 2FeSO4 + 3H2SO4 → Fe2(SO4)3 + 2NaHSO4 + 2H2O + 2NO
It is a colorless neutral gas. Its structure is:
3- Dinitrogen trioxide [Nitrogen (III) oxide]: It is prepared by treating nitric oxide with dinitrogen tetroxide
It is a blue solid with acidic nature. Its structure is:
4- Nitrogen dioxide [Nitrogen (IV) oxide]: It is prepared by heating lead nitrate at about 673K.
It is an acidic brown gas. Its structure is:
5- Dinitrogen tetroxide [Nitrogen (IV) oxide]: It is prepared by cooling NO2.
It is a colourless solid or liquid with acidic character. Its structure is:
6- Dinitrogen pentoxide [Nitrogen (V) oxide]: It is prepared by nitric acid with phosphorus pentoxide.
4 HNO3 + P4O10 → 4 HPO3 + 2 N2O5
It is a colourless solid with acidic character. Its structure is:
Nitric Acid (HNO3)
Preparation: In the laboratory, nitric acid is prepared by heating KNO3 or NaNO3 and concentrated H2SO4 in a glass retort.
NaNO3 + H2SO4 → NaHSO4 + HNO3
On a large scale it is prepared by Ostwald’s process. It involves three steps:
1- The catalytic oxidation of NH3 by atmospheric oxygen in presence of platinum/ rhodium gauge (wire) catalyst.
4 NH3(g) + 5 O2(g) Pt/Rh gauge catalyst, 500K & 9 bar 4NO(g) + 6 H2O(g)
2- The nitric oxide is converted to NO2
2NO(g) + O2(g) → 2 NO2 (g)
3- Absorption of nitrogen dioxide in water to get nitric acid.
3 NO2(g) + H2O(l) → 2 HNO3(aq) + NO(g)
The aqueous HNO3 can be concentrated by distillation up to 68% by mass. Further concentration to 98% can be achieved by dehydration with concentrated H2SO4. 98% HNO3 is known as fuming nitric acid.
Properties: It is a colourless liquid. In the gaseous state, HNO3 exists as a planar molecule with the structure as shown below:
In aqueous solution, nitric acid behaves as a strong acid giving hydronium and nitrate ions.
HNO3(aq) + H2O(l) → H3O+(aq) + NO3– (aq)
Concentrated nitric acid is a strong oxidising agent and attacks most metals except noble metals such as gold and platinum. The products of oxidation depend upon the concentration of the acid, temperature and the nature of the material undergoing oxidation.
3Cu + 8 HNO3(dilute) → 3Cu(NO3)2 + 2NO + 4H2O
Cu + 4HNO3(conc.) → Cu(NO3)2 + 2NO2 + 2H2O
Zinc reacts with dilute nitric acid to give N2O and with concentrated acid to give NO2.
4Zn + 10HNO3(dilute) → 4 Zn (NO3)2 + 5H2O + N2O
Zn + 4HNO3(conc.) → Zn (NO3)2 + 2H2O + 2NO2
Some metals (e.g., Cr, Al) do not dissolve in concentrated nitric acid because of the formation of a passive film of oxide on the surface.
Concentrated nitric acid also oxidises non–metals and their compounds. Iodine is oxidised to iodic acid, carbon to carbon dioxide, sulphur to H2SO4, and phosphorus to phosphoric acid.
I2 + 10HNO3 → 2HIO3 + 10 NO2 + 4H2O
C + 4HNO3 → CO2 + 2H2O + 4NO2
S8 + 48HNO3(conc.) → 8H2SO4 + 48NO2 + 16H2O
P4 + 20HNO3(conc.) → 4H3PO4 + 20 NO2 + 4H2O
Brown Ring Test: It is a test used for the detection of nitrates. The test is carried out by adding dilute ferrous sulphate solution to an aqueous solution containing nitrate ion, and then carefully adding concentrated sulphuric acid along the sides of the test tube. A brown ring at the interface between the solution and sulphuric acid layers indicate the presence of nitrate ion in solution.
NO3– + 3Fe2+ + 4H+ → NO + 3Fe3+ + 2H2O
[Fe (H2O)6 ]2 + NO → [Fe (H2O)5 (NO)]2+ + H2O
Uses: It is used i) in the manufacture of ammonium nitrate for fertilizers and other nitrates for use in explosives and pyrotechnics. ii) for the preparation of nitroglycerin, trinitrotoluene and other organic nitro compounds.iii) in the pickling of stainless steel, etching of metals and as an oxidiser in rocket fuels.
The allotropic forms of phosphorus:
Phosphorus exists mainly in three allotropic forms – white (yellow) phosphorus, red phosphorus and black phosphorus
1- White phosphorus: It is a translucent white waxy solid. It is poisonous, insoluble in water but soluble in carbon disulphide and glows in dark (chemiluminescence). It dissolves in boiling NaOH solution in an inert atmosphere giving PH3 (phosphine).
P4 + 3NaOH + 3H 2 O → PH3 + 3NaH 2 PO2 (sodium hypophosphite)
White phosphorus is less stable and therefore, more reactive. This is because in white phosphorus, the P-P-P bond angles are only 60°. So it has greater angular strain and highly unstable.
It readily catches fire in air to give dense white fumes of P4O10.
P4 + 5O2 → P4O10
It consists of discrete tetrahedral P4 molecule
2- Red phosphorus: It is obtained by heating white phosphorus at 573K in an inert atmosphere for several days. Red phosphorus has iron grey lustre. It is odourless, non-poisonous and insoluble in water as well as in carbon disulphide. Chemically, red phosphorus is much less reactive than white phosphorus. It does not glow in the dark. It contains polymeric chains of P4 tetrahedra.
3- Black phosphorus: It has two forms- α-black phosphorus and β-black phosphorus. α-black phosphorus is formed when red phosphorus is heated in a sealed tube at 803K. It does not oxidise in air. β-Black phosphorus is prepared by heating white phosphorus at 473K under high pressure. It does not burn in air up to 673K
Preparation: It is prepared by the reaction of calcium phosphide with water or dilute HCl.
Ca3P2 + 6H2O → 3Ca(OH)2 + 2PH3
Ca3P2 + 6HCl → 3CaCl2 + 2PH3
In the laboratory, it is prepared by heating white phosphorus with concentrated NaOH solution in an inert atmosphere of CO2.
P4 + 3NaOH + 3H2O → PH3 + 3NaH2PO2
( sodium hypophosphite)
Properties: It is a colourless gas with rotten fishy smell and is highly poisonous. It is slightly soluble in water. The solution of PH3 in water decomposes in presence of light giving red phosphorus and H2. When absorbed in copper sulphate or mercuric chloride solution, the corresponding phosphides are obtained.
3CuSO4 + 2PH3 → Cu3P2 + 3H2SO4
3HgCl2 + 2PH3 →→Hg3 P2 + 6HCl
Like NH3,Phosphine is weakly basic and gives phosphonium compounds with acids.
PH3 + HBr → PH4Br
Uses: Phosphine is technically used to produce Holme’s signal. Containers containing calcium carbide and calcium phosphide are pierced and thrown in the sea. The gases evolved burn and serve as a signal. It is also used in smoke screens.
Phosphorus forms two types of halides- PX3 and PX5
Phosphorus trichloride (PCl3)
Preparation: It is obtained by passing dry chlorine over heated white phosphorus.
P4 + 6Cl2 → 4PCl3
It is also obtained by the action of thionyl chloride with white phosphorus.
P4 + 8SOCl2 → 4PCl3 + 4SO2 + 2S2Cl2
It is a colourless oily liquid and hydrolyses in the presence of moisture.
PCl3 + 3H2O → H3PO3 + 3HCl
It reacts with organic compounds containing –OH group such as CH3COOH, C2H5OH.
3CH3COOH + PCl 3 → 3CH3COCl + H3PO3
3C2H5OH + PCl 3 → 3C2H5Cl + H3PO3
Structure: It has a pyramidal shape as shown, in which phosphorus is sp3 hybridized.
Phosphorus Pentachloride (PCl5)
Preparation: Phosphorus pentachloride is prepared by the reaction of white phosphorus with excess of dry chlorine.
P4 + 10Cl2 → 4PCl 5
It can also be prepared by the action of SO2Cl2 on phosphorus.
P4 + 10SO2Cl 2 → 4PCl5 + 10SO2
PCl5 is a yellowish white powder and in moist air, it hydrolyses to POCl3 and finally gets converted to Phosphoric acid.
PCl5 + H2O →POCl3 + 2HCl
POCl3 + 3H2O → H3PO4 + 3HCl
When heated, it sublimes but decomposes on strong heating.
PCl5 → PCl3 + Cl2
It reacts with organic compounds containing –OH group to give chloro derivative.
CH3COOH + PCl5 → CH3COCl + POCl3 +HCl
C2H5OH + PCl5 → C2H5Cl + POCl3 +HCl
In gaseous and liquid phases, it has a trigonal bipyramidal structure. The three equatorial P–Cl bonds are equivalent, while the two axial bonds are longer than equatorial bonds. This is due to the fact that the axial bond pairs suffer more repulsion as compared to equatorial bond pairs.
In the solid state it exists as an ionic solid, [PCl4]+[PCl6]– in which the cation, [PCl4] + is tetrahedral and the anion, [PCl6] – is octahedral.
Oxoacids of Phosphorus: Phosphorus forms a number of oxoacids.
1- H3PO2 [Hypophosphorus Acid (Phosphinic Acid)]
It is prepared by heating white phosphorus with concentrated NaOH solution followed by passing through cation exchange resin.
P4 + 3NaOH + 3H2O → PH3 + 3NaH2 PO2 (sodium hypophosphite)
NaH2PO2 + H+ -Resin → H3PO2 + Na+-Resin
It is a strong reducing agent due to the presence of a P-H bond. It is monobasic even though it contains three hydrogen atoms. This is because the hydrogen atoms directly bonded to the P atom will not dissociate.
2- H3PO3 [Orthophosphorus Acid (Phosphonic Acid)]
It is prepared by the action of water on P2O3
P2O3 + H2O → H3PO3
It is dibasic because of the presence of two –OH groups.
3- H4P2O5 [Pyrophosphorus Acid]
It is prepared by the action of H3PO3 on PCl3
PCl3 + 5H3PO3→ 3 H4P2O5 + 3HCl
It is also dibasic because of the presence of two –OH groups.
4- H4P2O6 [Hypophosphoric Acid]
It is prepared by the action of an alkali on red Phosphorus followed by passing through cation exchange resin.
2P + NaOH + H2O→ Na2H2P2O6
Na2H2P2O6 + 2H+ – Resin → H4P2O6 + 2Na+– Resin
It is a tetra basic acid.
5- H3PO4 [Orthophosphoric Acid]
It is obtained by the action of water on phosphorus pentoxide (P4O10) P4O10 + 6 H2O → 4 H3PO4
It is also called Phosphoric acid. It’s a tribasic acid and has a tetrahedral shape.
6- H4P2O7 [Pyrophosphoric Acid]
It is obtained by heating Phosphoric acid at about 2500c.
2 H3PO4 → H4P2O7 . It’s a tetra basic acid.
7- (HPO3)n [Metaphosphoric acid]
It is obtained by heating phosphorus acid with Br2 vapours in a sealed tube.
H3PO3 + Br2 → HPO3 + 2HBr
Structure: It exists as a trimer or a polymer as follows:
The oxoacids of phosphorus in +3 oxidation state undergo disproportionation (i.e. simultaneously oxidised and reduced). For example, orthophophorous acid (or phosphorous acid) on heating disproportionates to give orthophosphoric acid (phosphoric acid) and phosphine.
4H3PO3 → 3H3PO4 + PH3
Group 16 Elements
The members of this group are oxygen (O), sulphur (S), selenium (Se), tellurium (Te) and polonium (Po). They are also called chalcogens (means ore producing). Oxygen and sulphur are non-metals, selenium and tellurium are metalloids, while polonium is a radioactive metal.
Ionisation enthalpy of these elements decreases down the group. It is due to increase in size.
However, the elements of this group have lower ionisation enthalpy values compared to those of Group15 elements. This is due to the fact that Group 15 elements have extra stable half- filled p orbitals electronic
Oxygen atom has less negative electron gain enthalpy than sulphur because of the compact nature of its shells due to which the electronic repulsion is greater.
.Oxidation states: The elements of Group 16 exhibit a number of oxidation states (-2,+2,+4 & +6). The stability of -2 oxidation state decreases down the group. Since electronegativity of oxygen is very high, it shows only –2 oxidation state (except in the case of OF2 where its oxidation state is + 2). Other elements of the group exhibit + 2, + 4 & + 6 oxidation states. But + 4 and + 6 are more common. Sulphur, selenium and tellurium usually show + 4 oxidation state in their compounds with oxygen and + 6 with fluorine. Down the group, the stability of + 6 oxidation state decreases and that of + 4 oxidation state increases (due to inert pair effect).
Hydrides of 16th group elements
All the elements of Group 16 form hydrides of the type H2E (E = S, Se, Te, Po). Their acidic character increases from H2O to H2Te. This is due to the decrease in bond (H–E) dissociation enthalpy down the group. So the thermal stability also decreases down the group. All the hydrides except water possess reducing property and this character increases from H2S to H2Te.
Preparation: (i) By heating chlorates, nitrates and permanganates.
(ii) By the thermal decomposition of the oxides of metals low in the electrochemical series and higher oxides of some metals.
2Ag2O(s) → 4Ag(s) + O2(g); 2Pb3O4(s) → 6PbO(s) + O2(g)
2HgO(s) → 2Hg(l) + O2(g) ; 2PbO2(s) → 2PbO(s) + O2(g)
(iii) By the decomposition of Hydrogen peroxide (H2O2) in presence of manganese dioxide.
2H2O2(aq) → 2H2O(1) + O2(g)
(iv) On large scale it can be prepared from water or air. Electrolysis of water leads to the release of hydrogen at the cathode and oxygen n at the anode. It is also obtained by the fractional distillation of air.
Dioxygen directly reacts with metals and non-metals (except with some metals like Au, Pt etc and with some noble gases).
e.g. 2Ca + O2 → 2CaO
P4 + O2 → P4O10
4 Al + 3O2 → 2Al2O3
C + O2 → CO2
- oxygen is used in oxyacetylene welding, in the manufacture of many metals, particularly steel.
- Oxygen cylinders are widely used in hospitals, high altitude flying and in mountaineering.
- Liquid O2 is used in rocket fuels.
Oxides are binary compounds of oxygen with other elements. There are two types of oxides – simple oxides (e.g., MgO, Al2O3 ) and mixed oxides (Pb3O4, Fe3O4)
Simple oxides can be further classified on the basis of their acidic, basic or amphoteric character. An oxide that combines with water to give an acid is called acidic oxide (e.g., SO2, Cl2O7, CO2, N2O5).
Generally, non-metal oxides are acidic but oxides of some metals in higher oxidation states also have acidic character (e.g., Mn2O7, CrO3, V2O5 etc.).
The oxide which gives an alkali on dissolved in water is known as basic oxide (e.g., Na2O, CaO, BaO). Generally, metallic oxides are basic in nature.
Some metallic oxides exhibit a dual behaviour. They show the characteristics of both acidic and basic oxides. Such oxides are known as amphoteric oxides. They react with acids as well as alkalies.
E.g.: Al2O3, Ga2O3 etc.
There are some oxides which are neither acidic nor basic. Such oxides are known as neutral oxides.
Examples of neutral oxides are CO, NO and N2O.
Ozone is an allotropic form of oxygen.
Preparation: When a slow dry stream of oxygen is passed through a silent electric discharge, oxygen is converted to ozone. The product is known as ozonised oxygen.
3 O2(g)→ 2 O3(g); ∆H= +142 kJ/mol
Since the formation of ozone from oxygen is an endothermic process, a silent electric discharge should be used, unless the ozone formed undergoes decomposition.
Properties: Pure ozone is a pale blue gas, dark blue liquid and violet-black solid. Ozone has a characteristic smell.
Ozone is thermodynamically unstable with respect to oxygen since its decomposition into oxygen results in the liberation of heat (∆H is negative) and an increase in entropy (∆S is positive). So the Gibbs energy change (∆G) for this process is always negative (∆G = ∆H – T∆S).
Due to the ease with which it liberates nascent oxygen (O3 → O2 + O), it acts as a powerful oxidising agent.
For e.g., it oxidises lead sulphide to lead sulphate
PbS(s) + 4O3(g) → PbSO4(s) + 4O2(g)
Oxides of nitrogen (particularly nitric oxide) combine very rapidly with ozone and deplete it. Thus nitrogen oxides emitted from the exhaust systems of supersonic jet aeroplanes, slowly depleting the concentration of the ozone layer in the upper atmosphere.
NO(g) + O3(g) NO2(g) + O2(g)
Estimation of ozone: When ozone reacts with an excess of potassium iodide solution buffered with a borate buffer, iodine is liberated. The liberated iodine can be titrated against a standard solution of sodium thiosulphate. This is a quantitative method for estimating O3 gas.
Structure: O3 has an angular structure. It is a resonance hybrid of the following two forms:
Uses: It is used as a germicide, disinfectant and for sterilising water. It is also used for bleaching oils, ivory, flour, starch, etc. It acts as an oxidising agent in the manufacture of potassium permanganate.
Allotropes of Sulphur
Sulphur forms a large number of allotropes. Among these yellow rhombic (α-sulphur) and monoclinic (β -sulphur) forms are the most important. The stable form at room temperature is rhombic sulphur, which transforms to monoclinic sulphur when heated above 369 K.
1- Rhombic sulphur (α-sulphur)
It is prepared by evaporating the solution of roll sulphur in CS2. It is insoluble in water but readily soluble in CS2.
2- Monoclinic sulphur (β-sulphur)
It is prepared by melting rhombic sulphur in a dish and cooling, till a crust is formed. Two holes are made in the crust and the remaining liquid is poured out. On removing the crust, colourless needle shaped crystals of β-sulphur are formed. It is stable above 369 K and transforms into α-sulphur below it. At 369 K both the forms are stable. This temperature is called transition temperature.
Both rhombic and monoclinic sulphur have S8 molecules. The S8 ring in both the forms is puckered and has a crown shape.
Sulphur Dioxide (SO2)
- Sulphur dioxide is formed when sulphur is burnt in air or oxygen: S(s) + O2(g) → SO2 (g)
- In the laboratory it is obtained by treating a sulphite with dilute sulphuric acid. SO32-(aq) + 2H+ (aq) → H2O(l) + SO2 (g)
- Industrially, it is produced by roasting of sulphide ores. 4 FeS2(s) + 11 O2(g) 2 Fe2O3(s) + 8 SO2(g)
Properties: Sulphur dioxide is a colourless gas with pungent smell and is highly soluble in water.
With water, it forms a solution of sulphurous acid which is a dibasic acid and form two types of salts with alkalies – normal salt (sulphite) and acid salt (bisulphate or hydrogen sulphite).
SO2(g) + H2O(l) → H2SO3(aq)
With sodium hydroxide solution, it forms sodium sulphite, which then reacts with more sulphur dioxide to form sodium hydrogen sulphite.
2NaOH + SO2 → Na2SO3 + H2O
Na2SO3 + H2O + SO2 → 2NaHSO3
SO2 is oxidised to sulphur trioxide by oxygen in the presence of vanadium pentoxide (V2O5) catalyst.
2SO2 + O2 → 2SO3
Moist sulphur dioxide behaves as a reducing agent. It converts iron(III) ions to iron(II) ions and decolourises acidified potassium permanganate(VII) solution (This used as a test for SO2).
2Fe3+ + SO2 + 2H2O 2Fe2+ + SO42- + 4H+
5 SO2 + 2MnO4– + 2H2O 5 SO42- + 4H+ + 2Mn2+
Structure: SO2 has an angular shape. It is a resonance hybrid of the following two canonical forms:
Uses: Sulphur dioxide is used (i) in refining petroleum and sugar (ii) in bleaching wool and silk and (iii) as an anti-chlor, disinfectant and preservative (iv) for the production of Sulphuric acid, sodium hydrogen sulphite and calcium hydrogen sulphite (v) Liquid SO2 is used as a solvent to dissolve a number of organic and inorganic chemicals.
Oxoacids of sulphur
Sulphur forms a large no. of oxoacids like Sulphurous acid (H2SO3), Dithionous acid (H2S2O4), Sulphuric acid (H2SO4), Pyrosulphuric acid (Oleum, H2S2O7), Peroxomonosulphuric acid (Caro’s acid, H2SO5), Peroxodisulphuric acid (Marshell’s acid, H2S2O8) etc. structure of some oxoacids are:
Sulphuric Acid (H2SO4)
The most important oxoacid of sulphur is sulphuric acid which is also known as the ‘King of Chemicals’.
Sulphuric acid is manufactured by the Contact Process which involves three steps:
( i) burning of sulphur or sulphide ores in air to generate SO2.
S(s) + O2(g) → SO2 (g) Or, 4 FeS2(s) + 11 O2(g) 2 Fe2O3(s) + 8 SO2(g)
(ii) conversion of SO2 to SO3 by the reaction with oxygen in the presence of a catalyst (V2O5) 2SO2 + O2 → 2SO3
(iii) absorption of SO3 in H2SO4 to give Oleum (H2S2O7).
SO3 + H2SO4 → H2S2O7
(iv) Dilution of oleum with water gives H2SO4 of the desired concentration.
H2S2O7 + H2O → 2H2SO4
Sulphuric acid is a colourless, dense, oily liquid. It dissolves in water with the evolution of a large quantity of heat. Hence, for diluting the acid, the concentrated acid must be added slowly into water with constant stirring.
Chemical properties: The chemical reactions of sulphuric acid are due to the following reasons:
- its low volatility
- strong acidic character
- strong affinity for water and
- its ability to act as an oxidising agent.
In aqueous solution, sulphuric acid ionises in two steps.
H2SO4(aq) + H2O(l) → H3O+(aq) + HSO4–
HSO4–(aq) + H2O(l) → H3O+(aq) + SO42-
So it is dibasic and forms two series of salts: normal sulphates and acid sulphates.
Because of its low volatility sulphuric acid can be used for the manufacture of more volatile acids from their corresponding salts.
2 MX + H2SO4 → 2 HX + M2SO4 (where X = F, Cl, NO3 etc. and M is a metal)
Concentrated sulphuric acid is a strong dehydrating agent and drying agent. Many wet gases can be dried by passing them through sulphuric acid. Sulphuric acid removes water from organic compounds
e.g.: C12H22O11 + H2SO4 → 12C + 11H2O
Hot concentrated sulphuric acid is a moderately strong oxidising agent. It oxidises both metals and non- metals and the acid itself reduces to SO2.
Cu + 2 H2SO4(conc.) → CuSO4 + SO 2 + 2H 2O
S + 2H 2SO4(conc.) → 3SO2 + 2H2O
C + 2H2SO4(conc.) → CO2 + 2 SO2 + 2 H2O
Uses: The important uses of Sulphuric acid are:
- In the manufacture of fertilizers
- in petroleum refining
- in the manufacture of pigments, paints and dyestuff intermediates
- in detergent industry
- in metallurgical applications
- as electrolyte in storage batteries
- in the manufacture of nitrocellulose products and
- as a laboratory reagent.
Group 17 Elements
Fluorine (F), chlorine (Cl), bromine (Br), iodine (I) and astatine (At) are the members of Group 17. They are collectively known as the halogens (means salt producers). They are highly reactive non-metallic elements. All these elements have seven electrons in their outermost shell (ns2np5) and so they do not readily lose their electron. So they have very high ionisation enthalpy.
Halogens have maximum negative electron gain enthalpy in the corresponding periods. This is due to the fact that the atoms of these elements have only one electron less than stable noble gas configurations. Electron gain enthalpy of these elements decreases down the group. However, the negative electron gain enthalpy of fluorine is less than that of chlorine. It is because, in fluorine the incoming electron goes to the 2p subshell, but in Cl it enters in to the 3p subshell. Due to the compactness of 2p subshell compared to 3p subshell, the electron – electron repulsion is greater in fluorine than in chlorine. So F does not easily gains electron.
Halogens have very high electronegativity. The electronegativity decreases down the group.
Fluorine is the most electronegative element in the periodic table.
All halogens have characteristic colour. For example, F2 has yellow, Cl2-greenish yellow, Br2-red and I2, violet colour. This is due to absorption of radiations in visible region which results in the excitation of outer electrons to higher energy level.
The bond dissociation enthalpy of F2 is low. This is due to the relatively large electron-electron repulsion among the lone pairs in F2 molecule.
All the halogens are highly reactive. They react with metals and non-metals to form halides. The reactivity of the halogens decreases down the group.
Halogens are strong oxidising agents since they readily accept electron. F2 is the strongest oxidising halogen and it oxidises other halide ions in solution or in the solid phase.
All the halogens exhibit –1 oxidation state. Chlorine, bromine and iodine also show + 1, + 3, + 5 and + 7 oxidation states in their oxides, oxy acids and in inter halogen compounds. Due to the absence of vacant d orbitals and the maximum electronegativity, fluorine exhibits only –1 oxidation state.
Anomalous behavior of fluorine
Due to the small size, highest electronegativity, low F-F bond dissociation enthalpy, and non availability of d orbitals in valence shell, fluorine shows properties different from other halogens.
Some of the anomalous properties of fluorine are:
- Ionisation enthalpy, electronegativity, enthalpy of bond dissociation and electrode potentials are higher for fluorine than expected.
- Ionic and covalent radii, m.p. and b.p. and electron gain enthalpy are quite lower than expected.
- Most of the reactions of fluorine are exothermic (due to the small and strong bond formed by it with other elements).
- F forms only one oxoacid while other halogens form a number of oxoacids.
- Hydrogen fluoride is a liquid due to strong hydrogen bonding. While the hydrogen halides of other elements are gases.
Hydrides of halogens
Halogens react with hydrogen to give hydrogen halides which dissolve in water to form hydrohalic acids. The acidic strength of these acids varies in the order: HF < HCl < HBr < HI. The stability of these halides decreases down the group due to decrease in bond dissociation enthalpy from HF to HI. Chlorine (Cl2)
Preparation: It can be prepared by any one of the following methods:
- By heating manganese dioxide with concentrated hydrochloric acid.
MnO2 + 4HCl → MnCl2 + Cl2 + 2H2O
HCl can be replaced by a mixture of common salt and concentrated Conc.
H2SO4 4NaCl + MnO2 + 4H2SO4 → MnCl2+ 4NaHSO4 + 2H2O + Cl2
(ii) By the action of HCl on potassium permanganate.
2KMnO4 + 16HCl → 2KCl + 2MnCl2 + 8H2O + 5Cl2
Manufacture of chlorine
- Deacon’s process: By oxidation of hydrogen chloride gas by atmospheric oxygen in the presence of CuCl2 (catalyst) at 723 K.
4HCl+O2 ⎯⎯⎯ CuCl2⎯→2Cl2 +2H2O
(ii) Electrolytic process: Chlorine is obtained by the electrolysis of brine solution (concentrated NaCl solution). During electrolysis chlorine is liberated at the anode.
Properties: It is a greenish yellow gas with pungent and suffocating odour. It is soluble in water. It reacts with a number of metals and non-metals to form chlorides.
2Al + 3Cl2 → 2AlCl3; P4 + 6Cl2 → 4PCl3
2Na + Cl2 → 2NaCl; S8 + 4Cl2 → 4S2Cl2
2Fe + 3Cl2 → 2FeCl3 ;
With excess ammonia, chlorine gives nitrogen and ammonium chloride whereas with excess chlorine, nitrogen trichloride (explosive) is formed.
8NH3 + 3Cl2 → 6NH4Cl + N2; NH3 + 3Cl2 → NCl3 + 3HCl
With cold and dilute alkalies chlorine produces a mixture of chloride and hypochlorite but with hot and concentrated alkalies it gives chloride and chlorate.
2NaOH + Cl2 → NaCl + NaOCl + H2O
(cold and dilute)
6 NaOH + 3Cl2 → 5NaCl + NaClO3 + 3H2O
(hot and conc.)
With dry slaked lime it gives bleaching powder.
2Ca(OH)2 + 2Cl2 → Ca(OCl)2 + CaCl2 + 2H2O
Chlorine reacts with hydrocarbons and gives substitution products with saturated hydrocarbons and addition products with unsaturated hydrocarbons.
CH4 + Cl2 ⎯⎯UV⎯→ CH3Cl + HCl
Methane Methyl chloride
C2H4 + Cl2 ⎯⎯⎯→ C2H4Cl2
Chlorine water on standing loses its yellow colour due to the formation of HCl and HOCl. Hypochlorous acid (HOCl) so formed is unstable and dissociates to give nascent oxygen which is responsible for oxidising and bleaching properties of chlorine.
- It oxidises ferrous to ferric, sulphite to sulphate, sulphur dioxide to sulphuric acid and iodine to iodic acid.
2FeSO4 + H2SO4 + Cl2 → Fe2(SO4)3 + 2HCl
Na2SO3 + Cl2 + H2O → Na2SO4 + 2HCl
SO2 + 2H2O + Cl2 → H2SO4 + 2HCl
I2 + 6H2O + 5Cl2 → 2HIO3 + 10HCl
It is a powerful bleaching agent; bleaching action is due to oxidation.
Cl2 + H2O → 2HCl + [O]
Coloured substance + [O] → Colourless substance
It bleaches vegetable or organic matter in the presence of moisture. Its bleaching action is permanent.
Uses: It is used
- for bleaching wood pulp, bleaching cotton and textiles,
- in the extraction of gold and platinum
- in the manufacture of dyes, drugs and organic compounds such as CCl4, CHCl3, DDT, refrigerants, etc.
- in sterilising drinking water and
- preparation of poisonous gases such as phosgene (COCl2), tear gas (CCl3NO2), mustard gas (ClCH2CH2SCH2CH2Cl).
Hydrogen Chloride (HCl)
Preparation: It is prepared in the laboratory, by heating sodium chloride with concentrated sulphuric acid.
NaCl + H2SO4 ⎯⎯420⎯K⎯→ NaHSO4 + HCl
NaHSO4 + NaCl ⎯⎯823⎯K⎯→ Na2SO4 + HCl
Properties: It is a colourless and pungent smelling gas. It is extremely soluble in water and ionises as: HCl + H2O → H3O+ + Cl–
Its aqueous solution is called hydrochloric acid, which is a strong acid in water. It reacts with NH3 and gives white fumes of NH4Cl.
NH3 + HCl → NH4Cl
When three parts of concentrated HCl and one part of concentrated HNO3 are mixed, aqua regia is formed which is used for dissolving noble metals, e.g., gold, platinum.
Au + 4 H+ + NO3– + 4Cl– →AuCl4– + NO + 2H2O
3Pt + 16H++ 4NO3– + 18Cl– → 3PtCl6 2- + 4NO + 8H2O
Hydrochloric acid decomposes salts of weaker acids like carbonates, hydrogen carbonates, sulphites, etc.
Na2CO3 + 2HCl → 2NaCl + H2O + CO2
NaHCO3 + HCl → NaCl + H2O + CO2
Na2SO3 + 2HCl → 2NaCl + H2O + SO2
Uses: It is used
- in the manufacture of chlorine, NH4Cl and glucose (from corn starch),
- for extracting glue from bones and purifying bone black,
- in medicine and as a laboratory reagent.
Oxoacids of Halogens
Due to high electronegativity and small size, fluorine forms only one oxoacid, HOF known as fluoric
(I) acid or hypofluorous acid. The other halogens form several oxoacids like Hypohalous acid (HOX), halous acid (HOXO) , halic acid(HOXO2) and perhalic acid (HOXO3). They are stable only in aqueous solutions or in the form of their salts.
Chlorine forms 4 types of oxoacids – hypochlorous acid (HOCl), Chlorous acid (HOClO or HClO2), Choric acid (HOClO2 or HClO3) and perchloric acid (HOClO3 or HClO4). The structures of them are
When two different halogens react with each other, interhalogen compounds are formed. They can be assigned general compositions as AX, AX3, AX5 and AX7, where both A and X are halogens. A is larger and more electropositive than X. As the size of the central atom (A) increases, the stability of the compound also increases.
The interhalogen compounds can be prepared by the direct combination or by the action of halogen on lower interhalogen compounds.
Cl2 + F2 437K 2ClF I2 + 3Cl2 2ICl3
Cl2 + 3F2 573K 2ClF3
Br2 + 3F2 2BrF3
I2 + Cl2 2ICl
Br2 + 5F2 2BrF5
These are all covalent molecules and are diamagnetic in nature. They are volatile solids or liquids except CIF which is a gas at 298 K. Their physical properties are intermediate between those of constituent halogens. The interhalogen compounds are more reactive than halogens (except fluorine).
This is because A–X bond in interhalogens is weaker than X–X bond in halogens except F–F bond. The types of inter halogen compounds and their structures are as follows:
ClF, BrF, IF, BrCl, BrI
ClF3, BrF3, IF3, ICl3, IBr3 etc.
ClF5, BrF5, IF5
Uses: These compounds can be used as non aqueous solvents. Interhalogen compounds are very useful fluorinating agents.
Group 18 Elements
Group 18 consists of six elements- helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe) and radon (Rn). All these are gases and chemically unreactive. So they are called inert gases or noble gases.
All noble gases have general electronic configuration ns2np6 (except helium which has 1s2). Due to stable electronic configuration these gases have very high ionisation enthalpy and electron gain enthalpy.
Even though these elements are chemically inert, Kr and Xe form some compounds with oxygen and fluorine under special conditions.
In noble gases, there is only weak van der Waals force of attraction. So they have low melting and boiling point.
Xenon forms three binary fluorides, XeF2, XeF4 and XeF6 by the direct reaction of elements under suitable conditions.
Xe (g) + F2 (g) 673K, 1 bar XeF2(s)
(xenon in excess)
Xe (g) + 2F2 (g) 873K, 7 bar XeF4(s)
Xe (g) + 3F2 (g) 573K, 60-70bar XeF6(s)
XeF6 can also be prepared by the interaction of XeF4 and O2F2 at 143K. XeF4 +O2F2 →XeF6 +O2
XeF2, XeF4 and XeF6 are colourless crystalline solids. They are powerful fluorinating agents. They are readily hydrolysed even by traces of water. For example, XeF2 is hydrolysed to give Xe, HF and O2.
2XeF2 (s) + 2H2O(l) → 2Xe (g) + 4 HF(aq) + O2(g)
XeF2 and XeF4 have linear and square planar structures respectively. XeF6 has seven electron pairs (6 bonding pairs and one lone pair) and thus, have a distorted octahedral structure
- XeO3: It is obtained by the hydrolysis of XeF4 and XeF6 with water.
6XeF4 + 12 H2O → 4Xe + 2XeO3 + 24 HF + 3 O2
XeF6 + 3 H2O → XeO3 + 6 HF
- XeOF4 & XeO2F2: Partial hydrolysis of XeF6 gives oxyfluorides, XeOF4 and XeO2F2.
XeF6 + H2O → XeOF4 + 2 HF
XeF6 + 2 H2O → XeO2F2 + 4HF
XeO3 is a colourless explosive solid and has a pyramidal molecular structure. XeOF4 is a colourless volatile liquid and has a square pyramidal molecular structure.
Uses of noble gases
Helium is used in filling balloons for meteorological observations. It is also used in gas-cooled nuclear reactors. Liquid helium is used as cryogenic agent for carrying out various experiments at low temperatures. It is used as a diluent for oxygen in modern diving apparatus because of its very low solubility in blood.
Neon is used in discharge tubes and fluorescent bulbs for advertisement display purposes. Neon bulbs are used in botanical gardens and in green houses.
Argon is used to provide an inert atmosphere in high temperature metallurgical processes and for filling electric bulbs. It is also used in the laboratory for handling substances that are air-sensitive.
Xenon and Krypton are used in light bulbs designed for special purposes.