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 ns²np^1-6 (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:

1. 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.
2. 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.
3. The single N–N bond is weak. So the catenation tendency is weaker in nitrogen.
4. Due to the absence of d orbitals in its valence shell, the maximum covalency of nitrogen is four
5. 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.

Dinitrogen (N₂)

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

Properties

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).

1. It also used to create an inert atmosphere in metallurgy.
2. Liquid dinitrogen is used as a refrigerant to preserve biological materials, food items and in cryosurgery.

Ammonia

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.

Properties

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 H₂SO_4. 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

                                        (brown ring)

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.

Phosphorus

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

Phosphine (PH₃)

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 Halides

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

Properties

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

Structure:

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

Structure:

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.

Structure:

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.

Structure:

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.

Structure:

6- H₄P₂O₇ [Pyrophosphoric Acid]

It is obtained by heating Phosphoric acid at about 250⁰c.

2 H₃PO₄ → H₄P₂O₇ . It’s a tetra basic acid.

Structure:

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

configurations.

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 16^th 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.

Dioxygen (O₂)

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.

Properties:

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

Uses:

1. oxygen is used in oxyacetylene welding, in the manufacture of many metals, particularly steel.
2. Oxygen cylinders are widely used in hospitals, high altitude flying and in mountaineering.
3. Liquid O₂ is used in rocket fuels.

Oxides

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 (O₃)

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)

Preparation:

1. Sulphur dioxide is formed when sulphur is burnt in air or oxygen: S(s) + O₂(g) → SO₂ (g)
2. In the laboratory it is obtained by treating a sulphite with dilute sulphuric acid. SO₃^2-(aq) + 2H⁺ (aq) → H₂O(l) + SO₂ (g)
3. 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⁺ + SO₄^2- + 4H⁺

 5 SO₂ + 2MnO₄^– + 2H₂O 5 SO₄^2- + 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’.

Manufacture:

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

Properties

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:

1. its low volatility
2. strong acidic character
3. strong affinity for water and
4. 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) + SO₄^2-

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:

1. In the manufacture of fertilizers
2. in petroleum refining
3. in the manufacture of pigments, paints and dyestuff intermediates
4. in detergent industry
5. in metallurgical applications
6. as electrolyte in storage batteries
7. in the manufacture of nitrocellulose products and
8. 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.

Oxidation states

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:

1. Ionisation enthalpy, electronegativity, enthalpy of bond dissociation and electrode potentials are higher for fluorine than expected.
2. Ionic and covalent radii, m.p. and b.p. and electron gain enthalpy are quite lower than expected.
3. Most of the reactions of fluorine are exothermic (due to the small and strong bond formed by it with other elements).
4. F forms only one oxoacid while other halogens form a number of oxoacids.
5. 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:

1. 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

1. 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

  (excess)                                               (excess)

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

Ethene                 1,2-Dichloroethane

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.

1. 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

1. for bleaching wood pulp, bleaching cotton and textiles,
2. in the extraction of gold and platinum
3. in the manufacture of dyes, drugs and organic compounds such as CCl₄, CHCl₃, DDT, refrigerants, etc.
4. in sterilising drinking water and
5. 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

1. in the manufacture of chlorine, NH₄Cl and glucose (from corn starch),
2. for extracting glue from bones and purifying bone black,
3. 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

Interhalogen Compounds

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.

Preparation

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

(equimolar) (excess)

Properties:

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:

Type	Examples	Structure
AX	ClF, BrF, IF, BrCl, BrI	Linear
AX3	ClF3, BrF3, IF3, ICl3, IBr3 etc.	Bent T-shaped
AX5	ClF5, BrF5, IF5	Square pyramidal
AX7	IF7	Pentagonal bipyramidal

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-fluorine compounds

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)

(1:5 ratio)

Xe (g) + 3F₂ (g) 573K, 60-70bar XeF₆(s)

(1:20 ratio)

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)

Structures

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

Xenon-oxygen compounds

1. 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

1. 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.