Divya

The P Block Elements

The long form of the periodic table is divided into s, p, d and f blocks. Elements are grouped in four blocks depending upon the subshell in which the valence electron of the atom of an element enters, the p block on the periodic table constitutes Groups 13 to 18.

In class XI you have already studied the chemistry of the elements of Groups 13 and 14 along with those of Groups 1 and 2. The p-block elements show a much greater variation In properties than that shown by s- and d-block elements.

The variation in properties of p-block elements is tine in varying atomic sizes, ionisation enthalpies, electronegativities, and electron gain enthalpies. You know that the valence shell electronic configuration of the p-block elements is wsJMp’ h.

The first member of the group differs from the subsequent members due to its small size and nonavailability of d orbitals in its atoms. The p-block elements comprise metals, nonmetals and metalloids, thereby showing diversity In their properties.

Group 15 Elements

This group comprises the elements nitrogen, phosphorus, arsenic, antimony and bismuth. Nonmetnllic character changes to metallic character gradually, on descending the group. Nitrogen and phosphorus are nonmetals, arsenic and antimony are metalloids, while bismuth is a metal. The elements of Group 15 are less metallic than the corresponding elements of Group 14.

Occurrence And Uses

Nitrogen is present to the extent of 78% in the atmosphere. However, it is not very abundant in the earth’s crust as all nitrites and nitrates are water soluble. The major minerals are Indian saltpetre (KNO₃) and Chile saltpetre (NaNO₃). Nitrogen is an essential constituent of proteins and amino acids. Phosphorus is the eleventh element in order of abundance in the earth’s crust.

The common minerals of phosphorus arc fluorapatite [Ca₉(PO₄)₆ , CaF₂] and hydroxyapatite [Ca₉(PO₄)6Ca(OH)₂], which are present in phosphate rocks. Phosphorus is an essential constituent of plants and animals. It is present in bones, teeth and other hard tissues of the animal body. Arsenic, antimony and bismuth are not abundant and occur in trace amounts as sulphides along with other minerals.

Nitrogen is used in iron and steel industries and oil refineries to maintain an inert atmosphere. Liquid nitrogen is used as a refrigerant. Large amounts of nitrogen are used in the manufacture of ammonia and calcium cyanamide.

The major use of phosphorus is in making phosphatic fertilisers. Phosphates are also used in the food industry, in detergents and as pharmaceuticals. Phosphorus is used in safety matches and to make phosphor-bronze, an alloy.

It is also used to make pesticides and organophosphorus compounds. Arsenic is used to dope semiconductors and to alloy with lead to make it harder. Compounds of arsenic are used as pesticides and weedicides. Antimony is used in alloys with tin and lead. It is used to electroplate steel to prevent rusting. Bismuth is used to make low-melting alloys which are used in fuses and fire alarms.

Atomic And Physical Properties

Electronic configuration

The valence shell electronic configuration of these elements is ns²np³. The presence of completely filled s orbitals and half-filled p orbitals confers extra stability to the electronic configuration.

Size

Like in other groups, the atomic and ionic radii of elements increase on moving down the group. There is a substantia] increase in size from nitrogen to phosphorus. However, the increase in size from phosphorus to arsenic and then from antimony to bismuth is less due to the presence of completely filled d or f orbitals.

Ionisation enthalpy

It decreases down the group due to an increase in atomic size. The first ionisation enthalpy is quite high because of the stable ns²  np³ configuration.

Electron gain enthalpy

In Group 15, the electron gain enthalpy of nitrogen is positive whereas that of other elements is negative. The positive electron gain enthalpy of nitrogen is due to its compact atomic size. In the case of other elements, though the electron gain enthalpy is negative, it is not very high as the elements of the group have stable, half-filled p orbitals.

Electronegativity

Electronegativity decreases down the group with an increase in atomic size and metallic behaviour of the elements. However, the difference in electronegativities is less pronounced with increasing atomic weight.

Boiling and melting points

All elements in this group are polyatomic and display allotropy, except nitrogen. There is an increase in boiling point down the group; however, the variation in melting point is not regular.

Chemical Properties

Oxidation states and trends in chemical reactivity

The elements of this group have five electrons in their outermost shell. The elements exhibit a maximum oxidation state of +5. The stability of the +5 oxidation state decreases down the group due to the increasing inert pair effect. The only well-characterised binary compound of bismuth in the +5 oxidation state is BiF₅.

(As you know from your previous class, the inert pair effect is the tendency of s electrons to remain inert, i.e., not to take part in bond formation.) Nitrogen and phosphorus may also form N^-3  and P^-3 species with electropositive metals, by gaining three electrons. For example, magnesium and calcium nitrides (Mg₃N₂ and Ca₃N₂), which constitute N^3- ions.

Tire stability of the -3 oxidation state decreases down the group as the electropositive character of the element increases. Nitrogen displays a wide range of oxidation states, ranging from -3 to +5. The common oxidation states of phosphorus and arsenic are -3, +3 and +5.

In the case of nitrogen, disproportionation reactions are common as exemplified by the case of nitrous acid

3HNO₂ →HNO₃+H₂O+2NO

In the case of phosphorus, such reactions are common

4H₃PO₃→3H₃PO₄+PH₃

You have already studied in your previous class that in a disproportionation reaction, a substance is both oxidised and reduced. In the examples discussed above the reactant contains nitrogen and phosphorous in the +3 oxidation state. This oxidation state is not very stable and a higher stable oxidation state, i.e., +5 state is known for the elements. For the heavier members of the group, the stability of the +5 oxidation state decreases and that of the +3 state increases (inert-pair effect). Hence the tendency for disproportionation decreases.

Bond type

Group 15 elements generally form covalent compounds. However, there is a decrease in the covalent character in the order P > As > Sb > Bi. Antimony and bismuth form tripositive cations due to an increase in metallic character, on moving down the group.

Formation of penta- and hexa-coordinated derivatives

Due to the nonavailability of d orbitals, nitrogen cannot expand its octet. The heavier members of the group use d orbitals to form species-like PCI₅, PF₆^–, and AsF₅.

Tendency to form multiple bonds and catenation

The tendency to form multiple bonds relative to single bonds decreases on descending the group. Nitrogen is diatomic two atoms are bound by a triple bond in a molecule. Nitrogen also forms pπ-pπ bonds with other elements having comparable size and electronegativity (like C and O).

The other members of the group do not form multiple bonds as the orbitals are larger and diffuse and therefore do not allow effective overlap. Phosphorus, arsenic and antimony are tetraatomic, the atoms being linked by single bonds. Recently compounds like R₃P= O and R₃P= CH₂ (R = alkyl group) have been isolated involving dir-pn bonds, i.e., overlap of d orbitals of phosphorus with p orbitals of carbon and oxygen.

The trialkyl and triaryl derivatives of phosphorus and arsenic act as electron-pair donors towards transition metals, where the unshared electron pair on phosphorus and arsenic is donated to the vacant d orbital of the metal. The bond thus formed is strengthened by an overlap of filled d orbitals of the metal with vacant d orbitals of P or As. This is referred to as a dπ-dπ bond.

The single N-N bond is weaker than the single P-Pbond. This is due to the small size of nitrogen and small N-N bond length, which lends to high Intcreledrunic repulsion between the nonbonding electrons. Thus the catenation power of phosphorus Is greater than that of nitrogen. This Is manifested in a large number of allotropes of phosphorus.

Anomalous behaviour of nitrogen Nitrogen, the first member of group 15, differs considerably from the rest of the members. These differences arise due to the small size of nitrogen, ils high electronegativity, its tendency to form stable pn pa bonds and the nonavailability of d orbitals in its valence shell.

Some of the anomalous properties of nitrogen are listed as follows.

• Nitrogen exists as a diatomic gaseous molecule, while the lower members of the group exist as polyatomic solids.
• Nitrogen is inert due to the high strength of the NnN bond. The rest of the members of the group are more reactive.
• Nitrogen forms pπ-pπ bonds with itself and with elements like carbon and oxygen. Consequently, it forms a large number of oxides which are monomeric, unlike those of phosphorus which arc dimeric. Also, it forms species N^–₃ and CN^–.
• The maximum covalency of nitrogen is four. The other elements can expand their octet and form species with coordination numbers five and six.

Reactivity towards hydrogen

The elements of Group 15 form trihydric of the general formula EH3. Their ease of formation and stability decrease down the group. This can be explained in terms of E—H bond enthalpy. On descending the group, the size of E increases so that the orbitals become larger and more diffuse.

Consequently, these orbitals do not undergo effective overlap with the small 1 s orbital of hydrogen which leads to a decrease in bond enthalpy on descending the group. Thus, the reducing power of the hydrides increases down the group.

Apart from ammonia, the other hydrides are toxic gases. The low volatility of ammonia is dm* to intermolecular hydrogen bonding.

The hydrides of Group 15 elements have a pyramidal structure. The central atom is sp3 hybridised having a lone pair of electrons. Due to the repulsion between the lone pair and a bond pair of electrons, the bond angle reduces from 109°27′ (that of a regular tetrahedron) to 107°48’.

Thus the tetrahedral shape is slightly distorted. (The shape of the ammonia molecule is described as pyramidal since one of the tetrahedral positions is occupied by a lone pair.) Distortion due to lone pairs is even more in PH₃, AsH₃ and SbH₃, causing the bond angle to reduce further up to 91.3°.

Hydrides of Group 15 elements donate lone pairs of electrons and thus behave as Lewis bases. The basicity decreases down the group.

Reactivity towards oxygen

Since the Group 15 elements exhibit predominantly two oxidation states +3 and +5, two types of oxides are known— and E205. The +5 oxidation state in bismuth is not stable. Therefore it’s the main oxide! is Bi2Ov Oxides of nitrogen and phosphorus are acidic, those of arsenic and antimony are amphoteric, while that of bismuth is basic.

Thus the basic character of the oxides increases down the group as the metallic character increases. When an element forms two oxides E₂O₃ and E₂O₅, the oxide in the higher oxidation state is more acidic.

Reactivity towards halogens

Trihalides, EX₃, of all elements are known. Of all nitrogen halides, only NF₅ is stable due to the strong N-F bond. The trihalides of other elements are stable and are predominantly covalent, except BiF3. The pentahalides are fewer in number than trihalides.

Nitrogen does not form pentahalides due to the non-availability of d orbitals in its valence shell. The well-characterised pentahalides are PX₅ (X = F, Cl, Br), AsF₅, SbF₅, SbCl₅ and BiF₅. Pentahalides are more covalent than trihalides.

This is because the central atom in the +5 oxidation state has greater polarising power and can polarise the anion considerably. You have studied in class XI that when the degree of polarization is large, the concentration of electrons increases between the two bonded atoms and the covalent character increases.

Reactivity towards metals

Group 15 elements react with metals to form binary compounds in which they exhibit a -3 oxidation state. Nitrogen forms nitrides with lithium (Li3N), alkaline earth metals (Mg₃N₂, Ca₃N₂, etc.) and aluminium (AIN). There are also examples of phosphides (Ca₃P₂), arsenides (Na₃As, Mg₃As₂), antimonides (Zn₃Sb₂) and bismuthides (Mg₃Bi₂).

Nitrogen (N₂)

Nitrogen is the most abundant gas in the atmosphere. It comprises about 78.1% of the atmosphere by volume. It is less abundant in the earth’s crust because of the high solubility of nitrates.

Preparation

Nitrogen is obtained commercially by the fractional distillation of liquefied air. Air is first cooled to remove water vapour and carbon dioxide and then liquefied. On fractional distillation, nitrogen distils out first, leaving behind oxygen.

In the laboratory, nitrogen is obtained by warming an aqueous solution of ammonium chloride and sodium nitrite. This produces the thermally unstable ammonium nitrite, which decomposes to give nitrogen.

⇒ \(\mathrm{NH}_4 \mathrm{Cl}(\mathrm{aq})+\mathrm{NaNO}_2(\mathrm{aq}) \stackrel{\Delta}{\longrightarrow} \mathrm{NaCl}(\mathrm{aq})+\mathrm{NH}_4 \mathrm{NO}_2(\mathrm{aq}) \stackrel{\Delta}{\longrightarrow} \mathrm{N}_2(\mathrm{~g})+2 \mathrm{H}_2 \mathrm{O}\)

Small amounts of nitric acid and nitric oxide are formed as by-products in this reaction and may be removed by passing the gaseous product through a mixture of aqueous sulphuric acid and potassium dichromate.

Other methods of preparation of nitrogen include the oxidation of ammonia by bromine and the thermal decomposition of ammonium dichromate.

⇒ \(8 \mathrm{NH}_3+3 \mathrm{Br}_2 \stackrel{\Delta}{\longrightarrow} \mathrm{N}_2+6 \mathrm{NH}_4 \mathrm{Br}\)

⇒ \(\left(\mathrm{NH}_4\right)_2 \mathrm{Cr}_2 \mathrm{O}_7 \stackrel{\Delta}{\longrightarrow} \mathrm{N}_2+\mathrm{Cr}_2 \mathrm{O}_3+4 \mathrm{H}_2 \mathrm{O}\)

Small quantities of very pure nitrogen can be obtained by the thermal decomposition of sodium azide or barium azide.

⇒ \(2 \mathrm{NaN}_3 \stackrel{\Delta}{\longrightarrow} 3 \mathrm{~N}_2+2 \mathrm{Na}\)

Properties

Nitrogen is a colourless, odourless and nontoxic gas. The gas has low freezing and boiling points and low solubility in water. The two stable isotopes of nitrogen are 14N and 15N.

Nitrogen exists as a diatomic molecule, N =N. It is rather inert at room temperature due to the high bond enthalpy of the N =N bond. However, at elevated temperatures, it becomes increasingly reactive and combines directly with hydrogen oxygen, and some electropositive metals.

⇒ \(\mathrm{N}_2(\mathrm{~g})+3 \mathrm{H}_2(\mathrm{~g}) \underset{\text { catalyst }}{\stackrel{773 \mathrm{~K}}{\rightleftharpoons}} 2 \mathrm{NH}_3(\mathrm{~g})\)

⇒ \(\mathrm{N}_2(\mathrm{~g})+\mathrm{O}_2(\mathrm{~g}) \stackrel{2000 \mathrm{~K}}{\rightleftharpoons} 2 \mathrm{NO}(\mathrm{g})\)

⇒ \(6 \mathrm{Li}(\mathrm{s})+\mathrm{N}_2(\mathrm{~g}) \stackrel{\Delta}{\longrightarrow} 2 \mathrm{Li}_3 \mathrm{~N}(\mathrm{~S})\)

Ammonia

Small amounts of ammonia are present in nature as it is released from the decay of nitrogenous organic matter like urea.

⇒ \(\mathrm{NH}_2 \mathrm{CONH}_2+\mathrm{H}_2 \mathrm{O} \rightarrow\left(\mathrm{NH}_4\right)_2 \mathrm{CO}_3\)

Preparation

Ammonia is obtained by heating ammonium salts with an alkali in the laboratory.

⇒ \(2 \mathrm{NH}_4 \mathrm{Cl}+\mathrm{Ca}(\mathrm{OH})_2 \stackrel{\Delta}{\longrightarrow} \mathrm{CaCl}_2+2 \mathrm{NH}_3+2 \mathrm{H}_2 \mathrm{O}\)

On a large scale, ammonia is prepared by the Haber process.

⇒ \(\underbrace{\mathrm{N}_2(\mathrm{~g})+3 \mathrm{H}_2(\mathrm{~g})}_{4 \text { volumes }} \rightleftharpoons \underbrace{2 \mathrm{NH}_3(\mathrm{~g})}_{2 \text { volumes }} \quad \Delta_{\mathrm{f}} \mathrm{H}^{\ominus}=-46.1 \mathrm{~kJ} \mathrm{~mol}^{-1}\)

One volume of nitrogen combines with three volumes of hydrogen to give two volumes of ammonia. The reaction is reversible, accompanied by a reduction in volume and is exothermic. According to Le Chatelier’s principle, the forward reactions are favoured by high pressure and low temperature.

This reaction is carried out at a pressure of 150-250 Pa and a temperature of 600-700 Kin the presence of a finely divided iron catalyst with small amounts of oxides of potassium, aluminium and molybdenum.

The gases are passed over four beds of catalyst, with cooling between each pass. On each pass about 15% conversion occurs and the unreacted gases are recycled so that eventually an overall conversion of 98% can be achieved.

Properties

Ammonia is a colourless gas at room temperature (freezing point 198.4 K boiling point 239.7 K) and has a strong, characteristic pungent smell. It can be liquefied easily. The molecules of ammonia are extensively associated with hydrogen bonding both in the liquid and solid states.

Thus, it is less volatile than the other Group 15 hydrides. It is highly soluble in water and its aqueous solution is weakly basic owing to the presence of OH“ ions as shown by the following reaction.

⇒ \(\mathrm{NH}_3(\mathrm{~g})+\mathrm{H}_2 \mathrm{O}(\mathrm{l}) \rightleftharpoons \mathrm{NH}_4^{+}(\mathrm{aq})+\mathrm{OH}^{-}(\mathrm{aq})\)

NH₃ and NH₄OH both react with acids forming ammonium salts. These salts are thermally unstable and decompose on heating. If the anion is not oxidising, ammonia is evolved.

⇒ \(\mathrm{NH}_4 \mathrm{Cl} \stackrel{\Delta}{\longrightarrow} \mathrm{NH}_3+\mathrm{HCl}\)

If the anion is oxidising then NH⁺₄ is oxidised to N₂ or N₂O.

⇒ \(\mathrm{NH}_4 \mathrm{NO}_2 \stackrel{\Delta}{\longrightarrow} \mathrm{N}_2+2 \mathrm{H}_2 \mathrm{O}\)

⇒ \(\mathrm{NH}_4 \mathrm{NO}_3 \stackrel{\Delta}{\longrightarrow} \mathrm{N}_2 \mathrm{O}+2 \mathrm{H}_2 \mathrm{O}\)

⇒ \(\left(\mathrm{NH}_4\right)_2 \mathrm{Cr}_2 \mathrm{O}_7 \stackrel{\Delta}{\longrightarrow} \mathrm{N}_2+4 \mathrm{H}_2 \mathrm{O}+\mathrm{Cr}_2 \mathrm{O}_3\)

When aqueous ammonia is added to a metal salt solution, in many cases the metal hydroxide or hydrated oxide is precipitated.

CaCI₂(aq)+2NH₄OH(aq) →Ca(OH)₂(s)+2NH₄CI(aq)
Zn(NO₃)(aq)+2NH₄OH(aq)→Zn(OH)₂(s)+2NH₄NO₃(aq)
2FeCI₃(aq)+3NH₄OH(aq) →FeO₃.xH₂O(s)+3NH₄CI(aq)

In the ammonia molecule, the nitrogen atom is sp3 hybridised. There are three bond pairs and one lone pair. The resultant structure is pyramidal with an unshared electron pair on nitrogen. The bond angle (107.8°) is less than the bond angle associated with sp3 hybridisation (109.5°) due to distortion caused by the lone pair.

The ammonia molecule donates this lone pair to electron-pair acceptors and thus behaves as a Lewis base. It donates these electrons to many transition metal ions, forming complex compounds. When ammonia solution is added to copper sulphate solution, a deep blue colouration is obtained owing to the formation of [Cu(NH3)₄]²⁺ ions

⇒ \(\mathrm{Cu}^{2+}(\mathrm{aq})+4 \mathrm{NH}_3(\mathrm{aq}) \rightleftharpoons\left[\mathrm{Cu}\left(\mathrm{NH}_3\right)_4\right]^{2+}(\mathrm{aq})\)

AgCl(s) dissolves inNH3(aq) to form a complex, [Ag(NH3)2 ]C1. You are familiar with this reaction as it forms the basis of confirming chloride ions in salt analysis.

CI^–(aq)+AgNO₃(aq)→AgCI(s)+NO^–₃(aq)
AgCI(s)+2NH₃(aq)→[Ag(NH₃)₂]⁺(aq)

Uses

Ammonia is widely used in the manufacture of a large number of nitrogenous fertilisers like urea, ammonium nitrate, ammonium sulphate and ammonium phosphate. It is also used to prepare many inorganic nitrogen-containing compounds, the most important being nitric acid. Ammonia and ammonium compounds are used in the preparation of explosives, fibres and plastics.

Ammonia is also used as a refrigerant and in the manufacture of detergents and numerous inorganic and organic chemicals. Synthetic ammonia is the key to the industrial production of most inorganic nitrogen compounds as shown in the given scheme.

Oxides Of Nitrogen

Nitrogen forms a large number of oxides in the oxidation states +1 to +5. The oxides in the lower oxidation states are neutral while those in the higher oxidation states are acidic.

They exhibit pπ- pπ bonding between nitrogen and oxygen and exist as resonance hybrids of different canonical forms. Their formulae, names, methods of preparation and structures are summarized in.

Oxoacids Of Nitrogen

Nitrogen forms several oxoacids, many of which are unstable in the free state and are known only in aqueous solution or as their salts. The principal species are hyponitrous acid (H₂N₂O₂, a weak acid, whose salts are known), hydronitrous acid (H₄N₂O₄, whose sodium salt is known), nitrous acid (HNO₂) unstable and weak though salts are known and nitric acid (HNO₃) which is the most stable and the important one.

Nitric acid, HNO₃

Nitric acid is one of the three most important acids in the modern chemical industry (the others are sulphuric acid and hydrochloric acid).

Preparation In the laboratory nitric add is prepared by heating an alkali metal nitrate with concentrated sulphuric acid.

⇒ \(\mathrm{KNO}_3+\mathrm{H}_2 \mathrm{SO}_4 \stackrel{\Delta}{\longrightarrow} \mathrm{KHSO}_4+\mathrm{HNO}_3\)

It is industrially prepared by the Ostwald process, which involves the catalytic oxidation of ammonia to NO.

The nitric and air cooled and the mixture of gases is absorbed in a counter-current of water. During this process, nitric oxidised to nitrogen dioxide, which dissolved in water to give nitric acid.

⇒ \(2 \mathrm{NO}(\mathrm{g})+\mathrm{O}_2(\mathrm{~g}) \rightleftharpoons 2 \mathrm{NO}_2(\mathrm{~g})\)

⇒ \(3 \mathrm{NO}_2(\mathrm{~g})+\mathrm{H}_2 \mathrm{O}(\mathrm{l}) \rightleftharpoons 2 \mathrm{HNO}_3(\mathrm{aq})+\mathrm{NO}(\mathrm{g})\)

The nitric oxide is recycled and the aqueous nitric acid can be concentrated by distillation to the extent of up to 68% by mass since at this composition a constant-boiling mixture (azeotrope) is formed. Further concentration to 98% may be achieved by dehydrating with concentrated sulphuric acid or phosphorous pentoxide.

Properties Pure nitric acid is a colourless liquid (freezing point 231.4 K, boiling point 355.6 K). It has a specific gravity of 1.504. On exposure to light, it undergoes slight decomposition to N02 and Oz and thus acquires a yellowish-brown colour.

4HNO₃→4NO₂+O₂+2H₂O

It is a strong add and is completely dissociated in aqueous solutions

⇒ \(\mathrm{HNO}_3(\mathrm{aq})+\mathrm{H}_2 \mathrm{O}(\mathrm{l}) \rightleftharpoons \mathrm{H}_3 \mathrm{O}^{+}(\mathrm{aq})+\mathrm{NO}_3^{-}(\mathrm{aq})\)

Nitric add forms a large number of nitrates, which are highly soluble in water. Dilute nitric acid (< 2M) behaves as a typical strong acid. More concentrated aqueous solutions are strongly oxidising and attack most metals except noble metals like gold and platinum.

When nitric acid acts as an oxidising agent, it is reduced and the products of reduction depend upon the concentration of the acid, temperature and the other reactant.

Zinc reacts with dilute nitric acid to give nitrous oxide and with the concentrated acid, it forms nitrogen dioxide. With copper, dilute nitric acid gives nitric oxide and concentrated acid gives nitrogen dioxide.

4Zn+10HNO₃(dil.)→4Zn(NO₃)₂+N₂O+5H₂O
Zn+4HNO₃(conc.)→Zn(NO₃)₂+2NO₂+2H₂O
3Cu+8HNO₃(dil.)→3Cu(NO₃)₂+2NO+4H₂O
Cu+4HNO₃(conc.)→Cu(NO₃)+2NO₂+2H₂O

Metals like aluminium and chromium are rendered passive by concentrated nitric acid due to the formation of a superficial oxide film on the surface of the metal.

Concentrated nitric acid oxidises nonmetals to their corresponding oxides or oxoacids.

C+4HNO₃(conc.)→Cu(NO₃)₂+2H₂O

⇒ \(\mathrm{P}_4+20 \mathrm{HNO}_3 \longrightarrow \underset{\text { Phosphoric acid }}{4 \mathrm{H}_3 \mathrm{PO}_4}+20 \mathrm{NO}_2+4 \mathrm{H}_2 \mathrm{O}\)

S₈+48HNO₃→8H₂SO₄+48NO₂+16H₂O

⇒ \(\mathrm{I}_2+10 \mathrm{HNO}_3 \longrightarrow \underset{\text { logic acid }}{2 \mathrm{HIO}_3}+10 \mathrm{NO}_2+4 \mathrm{H}_2 \mathrm{O}\)

A mixture of concentrated nitric and hydrochloric acids (1 : 3 by volume respectively) is referred to as aqua regia; it is a very powerful oxidising mixture and dissolves metals like gold and platinum. However, it does not dissolve silver; instead, it forms an insoluble chloride with it.

When nitric acid is mixed with concentrated sulphuric acid, the nitronium ion(NO₂) is obtained, which is the active species used in the nitration of organic compounds.

Test for nitrates Nitrates are detected by the brown ring test. A freshly prepared ferrous sulphate solution is added to an aqueous solution of the nitrate followed by the slow addition of concentrated sulphuric acid along the side of the test tube so that it forms the layer at the bottom. A brown ring appears at the interface of the two liquids, which confirms the presence of nitrate ions in the solution. The ferrous ions reduce the nitrate ions to nitrogen monoxide. This reacts with hydrated ferrous ions to form a brown complex.

NO^–+3Fe²⁺+4H⁺→NO+3Fe³⁺+2H₂O

⇒ \(\underset{\text { Hydrated ferrous ion }}{\left[\mathrm{Fe}\left(\mathrm{H}_2 \mathrm{O}\right)_6\right]^{2+}}+\mathrm{NO} \longrightarrow \underset{\text { Brown colour }}{\left[\mathrm{Fe}\left(\mathrm{H}_2 \mathrm{O}\right)_5 \mathrm{NO}^{2+}\right.}+\mathrm{H}_2 \mathrm{O}\)

Structure Nitric acid has a planar structure. The bond parameters of a molecule of nitric acid are as follows.

The nitrate ion has a planar structure with equal bond lengths.

Nitric acid is largely used in the production of ammonium nitrate, which is employed for the production of fertilisers and other nitrates, which are used in the manufacture of explosives. It is used to make cyclohexanone

Which is one of the starting materials to prepare the monomers for the polymers nylon 6, 6 and nylon 6. Another major use is in the preparation of organic nitro compounds like nitroglycerine, nitrocellulose and trinitrotoluene. Minor uses include the pickling of stainless steel and the etching of metals. It is also used as an oxidiser in rocket fuel.

Phosphorus—Allotropic Modification

The main allotropes of phosphorus are white phosphorus, red phosphorus and black phosphorus.

White phosphorus is the most common allotrope obtained by the condensation of gaseous or liquid states. It is a waxy, translucent solid, pale yellow in colour and soluble in carbon disulphide and benzene. It is toxic, spontaneously catches fire and is, therefore, stored underwater. It reacts with moist air and gives out a characteristic faint glow and this is referred to as chemiluminescence.

It consists of discrete tetrahedral P4 units. The P-P bond angle is 60° and there is a considerable angular strain in the molecule. This accounts for the low stability and high reactivity of this allotrope. It readily catches fire in the air, forming the pentoxide.

P₄+5O₂→P₄O₁₀

Red phosphorus is obtained by heating white phosphorus at 573 K in the absence of air for several days. It is less reactive than white phosphorous and does not display chemiluminescence. It has a polymeric structure consisting of tetrahedral P4 units joined together.

Thermodynamically the most stable form is black phosphorus. It exists in two forms—a-black phosphorus and (3-black phosphorus. The a-form is obtained by heating red phosphorus in a sealed tube at 803 K.

It sublimes in the air and consists of opaque, monoclinic crystals. The |3-form is obtained by heating white phosphorus at 473 K under high pressure. It is inert and has a layered structure. Important physical properties of these allotropes are summarized.

Phosphine

Preparation

Phosphine, PH₃, is the hydride of phosphorus. It is obtained by the alkaline hydrolysis of yellow phosphorus or by the reaction of calcium phosphide with water or dilute acid.

⇒ \(\mathrm{P}_4+3 \mathrm{NaOH}+3 \mathrm{H}_2 \mathrm{O} \longrightarrow \mathrm{PH}_3+3 \mathrm{NaH}_2 \mathrm{PO}_2\)
Sodium hypophosphite

Ca₃P₂+6H₂O→2PH₃+3Ca(OH)₂
Ca₃P₂+6H_CI→2PH₃+3CaCI₂

The phosphine obtained is purified by passing it through hydrogen iodide, whereby phosphonium iodide is formed. This treatment with alkali gives pure phosphine.

PH₃+HI→PH₄I
PH₄I+NaOH→PH₃+NaI+H₂O

Phosphine catches fire spontaneously because it contains traces of diphosphine (P₂H₆), which is inflammable.

Properties

Phosphine is a colourless, toxic gas having an unpleasant odour similar to that of rotten fish. It is sparingly soluble in water and the resultant solution undergoes photodecomposition to give red phosphorus and hydrogen.

Pure phosphine is stable in the air but catches fire above 423 K.

⇒ \(\mathrm{PH}_3+2 \mathrm{O}_2 \stackrel{423 \mathrm{~K}}{\longrightarrow} \mathrm{H}_3 \mathrm{PO}_4\)

It is feebly basic and forms phosphonium salts with acids.

PH₃+HX→PH₄X

Itbums in chlorine to give phosphorus trichloride and phosphorus pentachloride

PH₃+3CI₂→PCI₃+3HCI
PH₃+4CI₂→PCI₅+3HCI

When phosphine is bubbled through aqueous solutions of copper and mercury(2) salts, the corresponding phosphides are precipitated.

3CuSO₄+2PH₃(g)→Cu₃P₂(s)+3H₂SO₄(aq)
3HgCI₂(aq)+2PH₃(g)p₂(s)+6HCI(aq)

Uses

The spontaneous combustion of phosphine is used in Holme’s signals in deep seas and oceans for signalling danger to ships. Containers containing calcium phosphide and calcium carbide are pierced and thrown into the sea.

In the presence of water, calcium phosphide hydrolyses to give phosphine which contains traces of inflammable P₂H₄. As stated earlier, diphosphine catches fire spontaneously. This ignites the ethyne produced by the hydrolysis of calcium carbide, and a luminous flame is obtained. Phosphine is also used in smoke screens.

Phosphorus Halides

All the possible phosphorus trihalides, PX₄ (X=F, Cl, Br, I) and phosphorus pentahalides, PX5(X=F, Cl, Br) are known.

Phosphorus trichloride

This is commercially the most important trihalide of phosphorus.

It may be prepared by passing dry chlorine over gently heated white phosphorus.

P₄+6CI₂→4PCI₃

It may also be obtained by treating white phosphorus with thionyl chloride.

P₄+8SOCI₂→4PCI₃+4SO₂+2SO₂CI₂

Phosphorus trichloride is a colourless, pungent-smelling liquid with a boiling point of 349 K. In moist air, i1 undergoes hydrolysis to give fumes of hydrochloric acid.

PCI₃+3H₂O→H₃PO₃+3HCI

Phosphorus trichloride reacts with oxygen to form phosphorus oxy-chloride and with chlorine, it forms phosphorus pentachloride

2PCI₃+O₂→2POCI₃
PCI₃+CI₂→PCI₅

PCI₃ reacts with organic compounds containing —OH group (carboxylic acids, alcohols) as follows.

3RCOOH+PCI₃→3RCOCI+H₃PO₃
3ROH+PCI₃→3RCI+H₃PO₃

The shape of the PCI₃ molecule is pyramidal and the phosphorus atom in the molecule is a sp3 hybridised

Phosphorus pentachloride

This is the most well-characterised pentahalide of phosphorus.

It is prepared by the reaction of chlorine with phosphorus trichloride.

\(\mathrm{PCl}_3+\mathrm{Cl}_2 \stackrel{\mathrm{CCl}_4}{\longrightarrow} \mathrm{PCl}_5\)

It can also be obtained by treating white phosphorus with an excess of dry chlorine or thionyl chloride.

P₄+10CI₂→4PCI₅
P₄+10SO₂CI₂→4PCI₅+10SO₂

PCI₅ is a yellowish-white powder.

It sublimes on heating and decomposes at a higher temperature to give the trichloride and chlorine.

\(\mathrm{PCl}_5 \stackrel{\Delta}{\longrightarrow} \mathrm{PCl}_3+\mathrm{Cl}_2\)

PCI₅ is susceptible to hydrolysis, which is a two-step reaction.

⇒ \(\mathrm{PCl}_5+\mathrm{H}_2 \mathrm{O} \longrightarrow\underset{\begin{array}{c}\text { Phosphorus } \\\text {oxychloride }\end{array}}{\mathrm{POCl}_3}+2 \mathrm{HCl}\)

PCI₅ is an excellent chlorinating agent and converts alcohols and carboxylic acids to their respective chloi derivatives.

R.OH+PCI₅→RCI+POCI₃+HCI
R.COOH+PCI₅→RCOCI+POCI₃HCI

It converts many metals to the corresponding chlorides.

⇒ \(\mathrm{Sn}+2 \mathrm{PCl}_5 \stackrel{\Delta}{\longrightarrow} \mathrm{SnCl}_4+2 \mathrm{PCl}_3\)

In the gaseous and liquid states, PCI₅ has a trigonal bipyramidal structure. The two axial P-Cl bonds are longer (240 pm) than the three equatorial bonds (202 pm). This is because bond-pair-bond-pair repulsions are stronger in the axial atoms than in the equatorial atom.

PCI₅ dimerises in the solid state and exists as an ionic solid, containing the tetrahedral [PCI₄]+ cation and the octahedral [PCI₆] anion.

Oxides Of Phosphorus

Phosphorus forms two oxides—the trioxide (P₄O₆) and the pentoxide (P₄O₁₀). The former is prepared by burning white phosphorus in a limited supply of air, while for the other an excess of air is needed.

\(\mathrm{P}_4+3 \mathrm{O}_2 \stackrel{\Delta}{\longrightarrow} \mathrm{P}_4 \mathrm{O}_6\)

Both the oxides are soluble in water, forming acidic solutions. This is expected as they are oxides o a nonmetal.

P₄O₆+6H₂O→4H₃PO₃
P₄O₁₀+6H₂O→4H₃PO₄

Due to the inability of phosphorus to form pπ-pπ double bonds with oxygen, its oxides are dimeric (P₄O₆ and P₄O₁₀). This is in sharp contrast to the oxides of nitrogen which are monomeric (N₂O₃ and N₂O₅)

Oxoacids Of Phosphorus

Phosphorus forms a large number of oxoacids. Before we discuss the properties and structure of each oxoaci note the following points in this context.

1. In a molecule of phosphorus acid, the phosphorus atom is sp3 hybridised and tetrahedrally surrounded by other atoms.
2. The ionisable acidic hydrogens in the molecules are attached to oxygen atoms, i.e., every acid contains at least one P-OH bond. These oxygens, which are attached to hydrogen, are called hydroxylic oxygens. All oxoacids also contain nonhydroxylic oxygens where the oxygen atom is linked only to phosphor forming the P = O bond.
3. In addition to P-OH and P = O bonds, the acids may contain either P-H or P-P bonds.
4. The presence of P-H bonds in the oxoacids confers reducing and not acidic properties.

The names, formulae, methods of preparation and some salient features of the oxoacids of phosphorus are shown.

It is seen that oxoacids containing phosphorus in their lower oxidation state are reducing. For example, hypophosphorus acid (H3P02) reduces silver salts to the metal.

⇒ \(4 \mathrm{AgNO}_3+2 \mathrm{H}_2 \mathrm{O}+\mathrm{H}_3 \mathrm{PO}_2 \longrightarrow 4 \mathrm{Ag}+4 \mathrm{HNO}_3+\mathrm{H}_3 \mathrm{PO}_4\)

The oxoacids in the +5 oxidation state do not have reducing or oxidising properties. Disproportionation is common for oxoacids containing phosphorus in the +3 oxidation state.

⇒ \(4 \mathrm{H}_3 \stackrel{+3}{\mathrm{PO}_3} \longrightarrow 3 \mathrm{H}_3 \stackrel{+5}{\mathrm{PO}_4}+\stackrel{-3}{\mathrm{P}} \mathrm{H}_3\)

The molecular structures of the oxoacids of phosphorus along with the types of bonds present in a molecule of the various oxoacids are shown in.

Group 16 Elements

Group 16 of the periodic table comprises oxygen, sulphur, selenium, tellurium and polonium. The elements of the group are called chalcogens or ore-forming elements as most metals occur as their oxides or sulphides. Nonmetallic character is maximum in oxygen and sulphur, weaker in selenium and tellurium whereas the short-lived and radioactive polonium is predominantly metallic.

Occurrence And Uses

The amount of oxygen present in dry air is about 20.946% by volume. Oxygen is the most abundant element in the earth’s crust. Most of the combined oxygen is in the form of silicates, oxides and water.

In contrast, the abundance of sulphur in the earth’s crust is only 0.03-0.1%. It is the 16th most abundant element and occurs mostly in the form of sulphide (zinc blende—ZnS, galena—PbS, cinnabar—HgS, and copper pyrites—CuFeS₂) and sulphate (gypsum—CaSO₄ 2H₂O, Epsom salt—MgSO₄ -7H₂0, baryte—BaSO₄) ores. Elemental sulphur is also present as hydrogen sulphide in natural gas, crude oil and volcanic ash.

Sulphur is a constituent of proteins and enzymes and some amino acids, for example, cysteine. The other elements of Group 16 show comparatively low abundance. Selenium and tellurium occur among sulphide ores. The main source of selenium and tellurium is the anode mud obtained during the electrolytic refining of copper. Thorium and uranium minerals contain polonium as a natural decay product.

Oxygen is essential for life—most life processes are based on oxidative metabolism. It is used in various energy generation processes through the combustion of wood and fossil fuels. Rocket fuels have liquid oxygen as the oxidant.

Oxyacetylene flames have very high temperatures and are used in welding. Many chemical industries use oxygen as an oxidant. Mountaineers use oxygen cylinders at high altitudes. Ozone, an allotrope of oxygen, is used as a disinfectant and for water sterilisation. It is also a bleaching agent.

Sulphur is predominantly used for the manufacture of sulphuric acid, which in turn is used in making fertilisers and other chemicals. Elemental sulphur is used in the vulcanisation of rubber and as a disinfectant.

Selenium is used to decolourise glass and as a photoconductor in photocopying machines. Tellurium is used as an additive to steel to increase its ductility. Tellurium and polonium are toxic.

Atomic And Physical Properties

Electronic configuration

The valence-shell electronic configuration of these elements isns²np⁴. They attain a noble-gas configuration by gaining two electrons, thus forming an E²” anion (E = O, S, Se, Te) or by sharing two electrons and thus forming two covalent bonds.

Size

The atomic size of Group 16 elements increases down the group as extra shells of electrons are added. The small atomic radius of oxygen and the absence of d orbitals in its atom are responsible for the distinctive chemical properties and also the high electronegativity of the element.

Ionisation enthalpy

The ionisation enthalpy of the elements decreases down the group as the size of the atom increases. The ionisation enthalpies of the Group 16 elements are strikingly less than those of the corresponding Group 15 elements. The unexpectedly high first ionisation enthalpies of Group 15 elements is due to the extra stability associated with half-filled p orbitals, which is not there in Group 16 elements.

Electronegativity

Oxygen is the second most electronegative element, fluorine being the most. Within the group, the electronegativity decreases on moving down. This indicates that the metallic character increases down the group.

Electron gain enthalpy

The electron gain enthalpy of oxygen is less negative than that of sulphur due to the small atomic size of oxygen. The electron gain enthalpy becomes less negative from sulphur to polonium.

Boiling and melting points

Due to small size, high electronegativity and absence of d orbitals, in an oxygen molecule, pπ-pπ bonds are formed between two oxygen atoms (0=0). Thus oxygen is stable and exists as a diatomic molecule in the gaseous state.

Oxygen too has a triatomic allotrope—ozone. The other elements of the group do not form multiple bonds and exist as polyatomic solids. In fact, sulphur is octa-atomic (S₈). The large difference between the boiling points of oxygen and sulphur and that between their melting points can be explained on the basis of their atomicity (oxygen exists as O₂ and sulphur as S₈). All elements of the group exhibit allotropy.

Chemical Properties

Oxidation states and trends in chemical reactivity

As you already know, the outermost electronic configuration of the elements of Group 16 is ns²np⁴ short of the nearest noble-gas configuration by two electrons. The elements can achieve the nearest noble-gas configuration by gaining or sharing two electrons. Thus, the E^2- ion of Group 16 elements exists with highly electropositive elements. Group 16 elements exhibit variable oxidation states due to the presence of empty d orbitals. Oxygen (which has no d orbitals) generally exhibits a^-2 oxidation state.

In peroxides, however, it exhibits a -1 oxidation state (O^-2₂)• The stability of the -2 oxidation state decreases down the group. Oxygen usually displays a negative oxidation state, except in some binary compounds with fluorine like OF₂ and O₂F₂ (Fluorine is more electronegative than oxygen.) Sulphur, selenium and tellurium show a tendency for covalency with formal oxidation states of +2, +4 and +6 in compounds where they are combined with more electronegative elements like oxygen or halogens.

The reactivity of elements generally decreases down the group. Sulphur combines with all elements except noble gases, nitrogen, tellurium, iodine, platinum and gold.

Bond type

Oxygen readily forms the divalent anion, O^2-. The tendency for the formation of the divalent anion decreases from sulphur onwards due to an increase in the size of the atom and a decrease in electronegativity. Compounds in a positive oxidation state are generally covalent; the covalent character decreases down the group.

Tendency to form multiple bonds and catenation

The tendency to form multiple bonds decreases down the group. Thus oxygen exists as Oz held by 0=0, while the other elements are polyatomic. The bond energy of the oxygen-oxygen double bond, 0=0, is 498 kJ mol-1. This makes the 0=0 bond more than three times as strong as the 0-0 bond (bond energy for 0-0 is 142 kJ mol^-1). In comparison, the S = S bond is less than twice as strong as the S-S single bond (bond energy for S = S is 434 kJ mol -1, and that for S-S is 264 kJ mol^-1 ). The tendency of catenation in sulphur is much higher; this is evident from the large number of allotropes of sulphur.

Anomalous behaviour of oxygen Oxygen differs considerably from the rest of the members of the group. The factors responsible for this anomalous behaviour of oxygen are small size, high electronegativity, nonavailability of d orbitals and the tendency to form pn-multiple bonds.

Some specific differences between the properties of oxygen and the other members of the group are as follows.

• Oxygen is a gas while the other members are solids.
• Oxygen is diatomic while the other members are polyatomic.
• Being highly electronegative, oxygen shows only negative oxidation states of -2 and -1 (except with fluorine) while other elements of the group show positive oxidation states.
• Oxygen has a tendency to form hydrogen bonds.

Reactivity towards hydrogen

Binary hydrides of the general formula H₂E are known as Group 16 elements. Some of the physical properties of their hydrides are summarised. The thermal stability of the hydrides decreases down the group. As the size of the central atom increases, the strength of the covalent H-E bond decreases. Thus both bond enthalpy and thermal stability decrease.

A consequence of decreasing bond enthalpy is the increase in acidic character. In other words, bond cleavage becomes easier and dissociation of H₂E to H⁺and HE becomes easier. The stability of hydrides decreases on descending the group and H₂Te in fact is thermodynamically unstable.

Apart from water, the hydrides are foul-smelling, toxic gases. The high boiling point of water is due to the presence of intermolecular hydrogen bonding. Hydrogen sulphide and the lower hydrides are reducing agents and the reducing power increases down the group. The hydrides are with two lone pairs. angular in shape. The central atom is sp3 hybridised

Reactivity towards oxygen

Group 16 elements mainly form dioxides (EO₂) and trioxides (EO₃). The dioxides are known for all elements, whereas the trioxides are known for sulphur, selenium and tellurium. The oxides are acidic in nature. The reducing power of the dioxides decreases down the group.

Reactivity towards halogens

The elements of Group 16 form a large number of binary compounds with halogens. The three main types of halides are EX₂, EX₄ and EX₆, but other halides are also known. The halogen compounds of sulphur, selenium, tellurium and polonium are called halides.

Oxygen too combines with halogens but its compound with fluorine only is said to be a halide. The compounds of oxygen with other halogens are oxides and not halides. This is so because of the high electronegativity of oxygen, which is exceeded only by fluorine and not chlorine, bromine and iodine.

The highest oxidation states of the Group 16 elements are realised only in combination with the most electronegative fluorine. Also, for a given oxidation state of Group 16 elements, fluoride is the most stable. Among the hexafluorides, SF6 is extremely stable. Its inertness is due to the sterically hindered sulphur in an octahedral structure.

SF₆ is unaffected by water as the protected sulphur atom does not allow hydrolysis, which is a thermodynamically favoured reaction. On the other hand, SeF₆ and TeF₆ are slightly more reactive and TeF6 is hydrolysed, probably due to the large size of the central atom.

The tetrafluorides of sulphur, selenium and tellurium are stable and exist as a gas (SF₄), liquid (SeF₄) and solid (TeF₄) respectively. The central atom in these compounds has sp3 hybridization and thus the molecular structure is trigonal bipyramidal with one lone pair in the equatorial position. This kind of geometry is called see-saw geometry.

 

Dimeric monohalides of the type S₂X₂ (X = F, Cl, Br) and Se₂X₂(X = Cl, Br) are known. However, they are unstable and tend to be disproportionate as follows.

2SeCI₂→SeCI₄+3Se

The hydrolysis is also accompanied by disproportionation.

2Se₂CI₂+2H₂O→H₂Se+3Se+4HCI

Oxygen

Oxygen occurs as dioxygen (O₂) and ozone (O₃). Dioxygen (O₂) exists as a gas and makes up 21 per cent of air in the atmosphere. It makes up 89 per cent by weight of the water in the oceans. Oxygen also occurs as sulphates, carbonates, nitrates, borates, etc. The three stable isotopes of oxygen are ¹⁶O,¹⁷O and ¹⁸O. We will now discuss the preparation and properties of dioxygen.

Preparation

1. The most convenient method of preparation of oxygen in the laboratory is by the thermal decomposition of potassium chlorate in the presence of manganese dioxide as a catalyst.

⇒ \(2 \mathrm{KClO}_3 \underset{420 \mathrm{~K}}{\stackrel{\mathrm{MnO}_2}{\longrightarrow}} 2 \mathrm{KCl}+3 \mathrm{O}_2\)

In the absence of manganese dioxide, the reaction occurs slowly at 670-720 K. Apart from potassium chlorate, potassium nitrate, potassium permanganate and barium peroxide also give dioxygen on heating.

⇒ \(2 \mathrm{KNO}_3 \stackrel{\text { heat }}{\longrightarrow} 2 \mathrm{KNO}_2+\mathrm{O}_2\)

⇒ \(2 \mathrm{KMnO}_4 \stackrel{\text { heat }}{\longrightarrow} \mathrm{K}_2 \mathrm{MnO}_4+\mathrm{MnO}_2+\mathrm{O}_2\)

⇒ \(2 \mathrm{BaO}_2 \stackrel{\text { heat }}{\longrightarrow} 2 \mathrm{BaO}+\mathrm{O}_2\)

2. Oxygen can also be prepared by the thermal decomposition of oxides of less reactive metals (metals placed low in the electrochemical series) like mercury and silver and also from oxides of some metals in their higher oxidation states.

⇒ \(2 \mathrm{HgO} \stackrel{\text { heat }}{\longrightarrow} 2 \mathrm{Hg}+\mathrm{O}_2\)

⇒ \(2 \mathrm{Ag}_2 \mathrm{O} \stackrel{\text { heat }}{\longrightarrow} 4 \mathrm{Ag}+\mathrm{O}_2\)

⇒ \(2 \mathrm{PbO}_2 \stackrel{\text { heat }}{\longrightarrow} 2 \mathrm{PbO}+\mathrm{O}_2\)

⇒ \(2 \mathrm{~Pb}_3 \mathrm{O}_4 \stackrel{\text { heat }}{\longrightarrow} 6 \mathrm{PbO}+\mathrm{O}_2\)

⇒ \(3 \mathrm{MnO}_2 \stackrel{\text { heat }}{\longrightarrow} \mathrm{Mn}_3 \mathrm{O}_4+\mathrm{O}_2\)

3. Another convenient method of preparation of oxygen is by the action of water on sodium peroxide or by the decomposition of hydrogen peroxide in the presence of manganese dioxide.

2Na₂O₂+2H₂O→4NaOH+O₂
2H₂O₂→2H₂O+O₂

4. Commercially, oxygen is obtained by the liquefaction of air. Initially, carbon dioxide and water vapour are removed from the air, and then the air is liquefied and subjected to fractional distillation. Nitrogen with a lower boiling point (77 K) distils out first leaving oxygen behind.

5. Large amounts of oxygen are obtained by the electrolysis of water.

⇒ \(2 \mathrm{H}_2 \mathrm{O} \stackrel{\text { electrolysis }}{\longrightarrow} 2 \mathrm{H}_2+\mathrm{O}_2\)

Hydrogen is liberated at the cathode and oxygen at the anode.

Properties

Dioxygen is a colourless, odourless and tasteless gas. It has a freezing point of 65 K and a boiling point of 90 K. It is slightly soluble in water (30.8 g per dm³ of water at 298 K and atmospheric pressure). This small amount of dissolved oxygen in water sustains aquatic life.

Oxygen has an even number of electrons yet it is paramagnetic. You have learnt about this in your previous class under molecular orbital theory.

Oxygen is a supporter of combustion. It reacts with most metals (except some less reactive metals like gold platinum, and noble gases) and nonmetals to form the respective oxides.

2Ca+O₂→2CaO
2Mg+O₂→2MgO
4AI+3O₂→2AI₂O₃
4Fe+3O₂→2Fe₂O₃
2H₂+O₂→2H₂O
S+O₂→SO₂
C+O₂→CO₂
P₄+5O₂→P₄O₁₀

The bond dissociation enthalpy of the O = O bond in an oxygen molecule is high and to initiate the reactions, initial heating is needed. However, these reactions are exothermic in nature and the heat liberated can sustain the reactions.

Oxygen reacts with a variety of compounds as shown.

⇒ \(2 \mathrm{SO}_2+\mathrm{O}_2 \underset{723 \mathrm{~K}, 2 \mathrm{~atm}}{\stackrel{\mathrm{V}_2 \mathrm{O}_5}{\longrightarrow}} 2 \mathrm{SO}_3\)

This reaction is the basis of the contact process for the manufacture of sulphuric acid.

⇒ \(4 \mathrm{NH}_3+5 \mathrm{O}_2 \underset{500 \mathrm{~K}}{\stackrel{\mathrm{Pt}}{\longrightarrow}} 4 \mathrm{NO}+6 \mathrm{H}_2 \mathrm{O}\)

This reaction is involved in the Ostwald process for the manufacture of nitric acid.

Metal sulphides react with oxygen to form oxides. These oxides may then be reduced to give the metal.

⇒ \(2 \mathrm{ZnS}+3 \mathrm{O}_2 \longrightarrow 2 \mathrm{ZnO}+2 \mathrm{SO}_2\)

Both saturated and unsaturated hydrocarbons cause an excess of oxygen to give carbon dioxide and water.

CH₄ + 2O₂→CO₂ + 2H₂O
H₂C= CH₂ + 3O₂→2CO₂ + 2H₂O
2HC = CH+ 5O₂→4CO₂ + 2H₂O

As these reactions are highly exothermic, hydrocarbons are used as fuels.

Oxides

Oxygen reacts with almost all the elements to form binary compounds called oxides. An element may form more than one oxide. For example, nitrogen forms six oxides.

Metal oxides can be simple (for example CaO, ZnO, Fe₂O₃) where the metal displays one oxidation state or they can be mixed (for example FeO₄, Pb₃O₄). A mixed oxide is made up of two oxides, in which the metal shows two oxidation states. Thus Pb₃O₄ may be considered to be a mixture of PbO₂ and PbO and may also be formulated as PbO₂ 2PbO. Depending on how an oxide behaves chemically, the oxides can be classified as

1. Acidic,
2. Basic,
3. Amphoteric And
4. Neutral.

The oxide that combines with water to give an acidic solution is referred to as acidic. Generally, oxides of nonmetals are acidic. For example,

CO₂+H₂O→H₂CO₃
P₄O₁₀+6H₂O→4H₃PO₄
SO₂+H₂O→H₂SO₃
CI₂O₇+H₂O→2HCIO₄

Acidic oxides react with bases to form salt and water oxides of metals generally basic as they dissolve in water to give a basic solution. for example,

MgO+2H₂O→Mg(OH)₂
CaO+2H₂O→Ca(OH)₂
Na₂O+H₂O→2NaOH

Basic oxides react with acids to form salt and water. Amphoteric oxides are some metallic oxides which show both acidic and basic properties. They react with acids as well as bases to form salts. For example,

ZnO+2HCI→ZnCI₂+H₂O
ZnO+2NaOH+ H₂O→Na[Zn(OH)₄]

When a metal forms oxides exhibiting different oxidation states, then the oxides which have metals in their higher oxidation states display acidic properties; for example, V₂O₅ and Mn₂O₇ are acidic in nature.

Some oxides like CO, N₂O and NO display neither acidic nor basic properties and are called neutral oxides.

Ozone

As already stated elemental oxygen exists in two allotropic modifications, dioxygen (O₂) and trioxygen or ozone (O₃).

Ozone is present in the upper atmosphere at a height of about 20 km from the earth’s surface. It is formed by the action of ultraviolet radiation on oxygen.

⇒ \(3 \mathrm{O}_2 \stackrel{\text { UV light }}{\longrightarrow} 2 \mathrm{O}_3\)

The ozone layer in the atmosphere protects us from the harmful effects of ultraviolet rays.

It has been shown that oxides of nitrogen, particularly nitric oxide, combine rapidly with ozone.

NO+O₃→NO₂+O₂

Nitric oxide emitted by supersonic aircraft is responsible for the slow depletion of the ozone layer. Another threat to the ozone layer is chlorofluorocarbons (CFCs) or freons which are used as refrigerants and as aerosol propellants. These molecules diffuse into the stratosphere and undergo slow photochemical degradation to produce atomic chlorine which reacts with ozone. Ozone being thermodynamically unstable liberates nascent oxygen, which combines with the CIO* free radical.

⇒ \(\mathrm{CFCl}_3 \stackrel{\text { UV rays }}{\longrightarrow} \dot{\mathrm{C}} \mathrm{FCl}_2+\mathrm{Cl}^{\circ}\)

⇒ \(\mathrm{Cl}^*+\mathrm{O}_3 \longrightarrow \mathrm{ClO}^*+\mathrm{O}_2\)

\(\mathrm{ClO}^*+\mathrm{O} \longrightarrow \mathrm{Cl}^*+\mathrm{O}_2\)

Preparation

Ozone can be obtained by the action of a silent electric discharge through pure and dry oxygen.

⇒ \(3 \mathrm{O}_2 \longrightarrow 2 \mathrm{O}_3 ; \Delta H^{\ominus}(298 \mathrm{~K})=+142.6 \mathrm{~kJ} \mathrm{~mol}^{-1}\)

This is an endothermic reaction. It is essential to use a silent electric discharge as it generates less heat; if the temperature of the reaction is allowed to rise continuously then the ozone formed may decompose back to oxygen. This reaction produces only about 10% ozone. The product is actually? a mixture of dioxygen and is called ozonised oxygen.

Properties

Ozone is a pale blue gas which condenses to give a deep blue liquid and a violet-black solid. It has a strong characteristic smell and is heavier than air. Ozone is slightly soluble in water but readily soluble in organic solvents.

Ozone in low concentrations is not so toxic, but it becomes harmful in concentrations above 100 ppm. Then it may cause respiratory problems, headaches and nausea. Ozone is diamagnetic. It is not very stable and decomposes slowly to give dioxygen.

2O₃→3O₂

This reaction is exothermic and therefore the enthalpy of the reaction is negative. The reaction is also associated with an increase in entropy (positive ΔS). Thus the reaction is thermodynamically favoured and the Gibbs energy change (ΔG) has a large negative value. Thus high concentration of ozone can lead to an explosion.

Ozone is a very powerful oxidising agent next only to fluorine in oxidising power. It releases atomic oxygen in the reaction which brings about oxidation.

O₃→O₂+O
2NO₂+O₃→N₂O₅+O₂
S+HO₂+O₃→ H₂SO₄

Ozone oxidises metal sulphides to their respective sulphates.

⇒ \(\mathrm{PbS}+4 \mathrm{O}_3 \longrightarrow \mathrm{PbSO}_4+4 \mathrm{O}_2\)

Thus when ozone is passed through a suspension of lead sulphide (black), the colour changes from black to white, owing to the formation of the white lead sulphate.

Ozone oxidises halogen acids to halogens, potassium iodide to iodine and acidified ferrous salts to ferric salts.

⇒ \(2 \mathrm{HCl}+\mathrm{O}_3 \longrightarrow \mathrm{Cl}_2+\mathrm{H}_2 \mathrm{O}+\mathrm{O}_2\)

⇒ \(2 \mathrm{KI}+\mathrm{H}_2 \mathrm{O}+\mathrm{O}_3 \longrightarrow \mathrm{I}_2+2 \mathrm{KOH}+\mathrm{O}_2\)

⇒ \(2 \mathrm{Fe}^{2+}+2 \mathrm{H}^{+}+\mathrm{O}_3 \longrightarrow 2 \mathrm{Fe}^{3+}+\mathrm{H}_2 \mathrm{O}+\mathrm{O}_2\)

The reaction of ozone with potassium iodide is used for quantitative estimation of ozone as the iodine liberated can be estimated by titrating with sodium thiosulphate, using starch as the indicator.

Ozone has an angular structure. The bond length in ozone is 128 pm which is intermediate between the 0-0 single bond length of 148 pm and the 0 = 0 double bond length of 110 pm. The structure is a resonance hybrid of the two resonating forms, as shown in.

Uses

Ozone is used as a disinfectant and for sterilising water. It is used for bleaching delicate fabrics, oils, ivory, starch, etc. It is used in the industry for the manufacture of potassium permanganate, artificial silk, etc. It finds use in organic chemistry, as an oxidising agent and for carrying out ozonolysis.

Allotropic Modifications Of Sulphur

Sulphur has a strong tendency towards catenation and this is manifested in a large number of allotropes of sulphur. The main allotropes of sulphur are rhombic (α-sulphur) and monoclinic ((β-sulphur).

The rhombic form is stable at room temperature whereas monoclinic sulphur is stable over 369 K. Rhombic sulphur is a bright yellow solid, readily soluble in carbon disulphide. It is soluble in ether, alcohol and benzene too but to a lesser extent. It gets converted to the monoclinic form on slow heating above 369 K. It has a melting point of 385.8 K and a specific gravity of 2.06. It is obtained when a solution of sulphur in carbon disulphide is crystallised.

Monoclinic sulphur is prepared by melting rhombic sulphur in a dish and letting it cool till a superficial crust is formed. Two holes are pierced into the crust and the liquid sulphur lying below the crust (which has not yet solidified) is poured out through one of the holes.

Small needle-like crystals of monoclinic sulphur become visible. Monoclinic sulphur melts at 393 K, has a specific gravity of 1.98 and is soluble in carbon disulphide. Bel this temperature rhombic sulphur is stable and at this temperature, both forms are in equilibrium. This temperature is called the transition temperature.

Both rhombic and monoclinic sulphur consist of S₈ rings, the sulphur atoms joined together in the shape of a crown. The packing pattern varies leading to different symmetry in crystals. Several other ring sizes from S₆ to S₂₀ have been synthesised in the recent past. A form is also known where the S₆ rings are arranged in the chair conformation.

Plastic sulphur is an amorphous form of sulphur obtained when molten sulphur is poured into cold water. In addition to this several chain polymers also exist in the liquid and vapour states. Above 1000 K the main species is S₂, which like O₂ is paramagnetic.

Sulphur Dioxide

When sulphur is burnt in the air, sulphur dioxide is the main product obtained (along with 6-8% sulphur trioxide)

S+O₂→SO₂

In the laboratory sulphur dioxide is prepared by heating copper turnings with concentrated sulphuric acid.

⇒ \(\mathrm{Cu}+2 \mathrm{H}_2 \mathrm{SO}_4 \longrightarrow \mathrm{CuSO}_4+\mathrm{SO}_2+\mathrm{H}_2 \mathrm{O}\)

It is also obtained when a sulphite is treated with dilute sulphuric acid.

⇒ \(\mathrm{Na}_2 \mathrm{SO}_3+2 \mathrm{H}^{+} \longrightarrow 2 \mathrm{Na}^{+}+\mathrm{H}_2 \mathrm{O}+\mathrm{SO}_2\)

Commercially, sulphur dioxide is obtained as a by-product in the roasting of sulphide ores.

⇒ \(2 \mathrm{PbS}+3 \mathrm{O}_2 \longrightarrow 2 \mathrm{PbO}+2 \mathrm{SO}_2\)

⇒ \(4 \mathrm{FeS}_2+11 \mathrm{O}_2 \longrightarrow 2 \mathrm{Fe}_2 \mathrm{O}_3+8 \mathrm{SO}_2\)

Properties

Sulphur dioxide gas has a sharp, choking odour. It can be readily liquefied at room temperature at a pressure of two atmospheres. It is an acidic oxide and dissolves in water to give sulphurous acid.

⇒ \(\mathrm{SO}_2+\mathrm{H}_2 \mathrm{O} \rightleftharpoons \mathrm{H}_2 \mathrm{SO}_3\)

When sulphur dioxide reacts with sodium hydroxide, it forms two salts—sodium hydrogen sulphite (i\ahlb(J3) and sodium sulphite (Na2S03). With an excess of sodium hydroxide, sodium sulphite is formed; this reacts with more sulphur dioxide to form sodium hydrogen sulphite.

⇒ \(2 \mathrm{NaOH}+\mathrm{SO}_2 \longrightarrow \mathrm{Na}_2 \mathrm{SO}_3+\mathrm{H}_2 \mathrm{O}\)

⇒ \(\mathrm{Na}_2 \mathrm{SO}_3+\mathrm{SO}_2+\mathrm{H}_2 \mathrm{O} \longrightarrow 2 \mathrm{NaHSO}_3\)

Sulphur dioxide acts as an oxidising agent as well as a reducing agent. However, its role as a reducing agent is more pronounced. It reduces Cr₂O^2-₇ to Cr³⁺

⇒ \(\mathrm{Cr}_2 \mathrm{O}_7^{2-}+2 \mathrm{H}^{+}+3 \mathrm{SO}_2 \longrightarrow 2 \mathrm{Cr}^{3+}+3 \mathrm{SO}_4^{2-}+\mathrm{H}_2 \mathrm{O}\)

It decolorises potassium permanganate solution and reduces ferric salts to ferrous salts.

⇒ \(2 \mathrm{MnO}_4^{-}+5 \mathrm{SO}_2+2 \mathrm{H}_2 \mathrm{O} \longrightarrow 2 \mathrm{Mn}^{2+}+5 \mathrm{SO}_4^{2-}+4 \mathrm{H}^{+}\)

⇒ \(2 \mathrm{Fe}^{3+}+\mathrm{SO}_2+2 \mathrm{H}_2 \mathrm{O} \longrightarrow 2 \mathrm{Fe}^{2+}+\mathrm{SO}_4^{2-}+4 \mathrm{H}^{+}\)

Sulphur dioxide oxidises hydrogen sulphide to sulphur.

⇒ \(2 \mathrm{H}_2 \mathrm{~S}+2 \mathrm{SO}_2 \longrightarrow 2 \mathrm{H}_2 \mathrm{O}+3 \mathrm{~S}\)

It combines with chlorine in the presence of charcoal as a catalyst to form sulphuryl chloride.

⇒ \(\mathrm{SO}_2+\mathrm{Cl}_2 \stackrel{\mathrm{C}}{\longrightarrow} \mathrm{SO}_2 \mathrm{Cl}_2\)

Sulphur dioxide reacts with oxygen in the presence of vanadium pentoxide as a catalyst to form sulphur trioxide. The reaction, as you know, forms the basis of the contact process for the manufacture of sulphuric acid.

⇒ \(2 \mathrm{SO}_2+\mathrm{O}_2 \stackrel{\mathrm{v}_2 \mathrm{O}_5}{\longrightarrow} 2 \mathrm{SO}_3\)

Sulphur dioxide has an angular structure.

Uses

The main use of sulphur dioxide is in the manufacture of sulphuric acid. It is also used as a bleaching agent for wool and silk, as a disinfectant and in petroleum and sugar refining. Liquid sulphur dioxide is used as a nonaqueous solvent.

Oxoacids Of Sulphur

Sulphur forms a significant number of oxoacids. All are not known in the free state but exist in solution and in the form of salts. Some of the important oxoacids of sulphur and their structures are listed.

While calculating oxidation states, remember that the oxidation state of oxygen normally is -2; however, it is -1 in peroxides.

Sulphuric Acid

It is the most important industrial chemical produced worldwide. It is manufactured by the contact process. shows a flow diagram for the process.

The process involves the following three steps.

Production of SO₂ It is obtained by burning sulphur or by roasting iron pyrites.

S+O₂→SO₂

⇒ \(4 \mathrm{FeS}_2+11 \mathrm{O}_2 \longrightarrow 2 \mathrm{Fe}_2 \mathrm{O}_3+8 \mathrm{SO}_2\)

Oxidation of SO₂ to SO₃ The SO₂ produced contains oxygen with it. The gaseous mixture is purified by passing it first through a dust precipitator and then a washing and cooling chamber. The mixture is then introduced to a drying chamber where concentrated sulphuric acid is sprayed to remove moisture.

Further, the gases are passed through an arsenic purifier containing gelatinous ferric hydroxide, which absorbs arsenic and its compounds present as impurities in the mixture. Now the gases are passed through a testing chamber where they are exposed to a strong beam of light to check whether any impurities are present. Impurities, if present in the mixture, scatter light.

If the gases are found impure, then the initial process is repeated till all the impurities are removed. Now the gaseous mixture is heated to about 720-820 K in a preheater and passed through a catalytic converter where the following reaction occurs.

⇒ \(2 \mathrm{SO}_2+\mathrm{O}_2 \longrightarrow 2 \mathrm{SO}_3 \quad \Delta_r H^{\ominus}=-196.6 \mathrm{~kJ} \mathrm{~mol}^{-1}\)

The catalyst used is V₂O₅ at the optimum condition of a temperature of 720 K and pressure of 2 bar. The reaction is exothermic and accompanied by a decrease in volume; thus the forward reaction can proceed if the temperature is low and pressure is high. However, if the temperature is very low, then the reaction becomes too slow.

Absorption of S03 into H₂SO₄ to give oleum The sulphur trioxide is absorbed in concentrated sulphuric acid to give oleum (H₂S₂O₇).

⇒ \(\mathrm{SO}_3+\mathrm{H}_2 \mathrm{SO}_4 \longrightarrow \mathrm{H}_2 \mathrm{~S}_2 \mathrm{O}_7\)

The oleum is diluted with water to give sulphuric acid. The sulphuric acid thus obtained is generally 96-98% pure.

⇒ \(\mathrm{H}_2 \mathrm{~S}_2 \mathrm{O}_7+\mathrm{H}_2 \mathrm{O} \longrightarrow 2 \mathrm{H}_2 \mathrm{SO}_4\)

Sulphur trioxide cannot be directly dissolved in water because the reaction is highly exothermic and the water evaporates from the mixture due to the heat produced.

Properties

Sulphuric acid is a colourless, viscous liquid with a specific gravity of 1.84 at 298 K. It freezes at 283 K and boils at 590 K. It has a strong affinity for water and its dissolution in water is highly exothermic. Therefore if water is poured into concentrated acid, the heat evolved leads to instant boiling of the water which splashes violently. To prevent this the acid is diluted by slowly adding it to water with constant stirring and not the other way.

Sulphuric acid is a strong dibasic acid. It ionises in two steps:

⇒ \(\mathrm{H}_2 \mathrm{SO}_4+\mathrm{H}_2 \mathrm{O} \rightleftharpoons \mathrm{H}_3 \mathrm{O}^{+}+\mathrm{HSO}_4^{-} \quad\) K_4>10

⇒ \(\mathrm{HSO}_4^{-}+\mathrm{H}_2 \mathrm{O} \rightleftharpoons \mathrm{H}_3 \mathrm{O}^{+}+\mathrm{SO}_4^{2-} \quad K_{a_2}^{\prime}=1.2 \times 10^{-2}\)

You know that the greater the value of the ionisation constant of an acid, the stronger the acid. The larger value of the acid ionisation constant K_a1(K_a1>10) shows that sulphuric acid is dissociated into H⁺ and HSO₄^– to a large extent.

It forms two series of salts—hydrogen sulphates or acid sulphates, for example, NaHS04, and sulphates or normal sulphates, for example, Na₂SO₄.

Sulphuric acid is a strong dehydrating and oxidising agent. It has a low volatility. Some of its main chemical properties are as follows.

Acidic nature We have already mentioned that sulphuric acid is a strong acid and forms two series of salts. The dilute acid reacts with active metals, liberating hydrogen.

⇒ \(\mathrm{M}+\mathrm{H}_2 \mathrm{SO}_4 \text { (dilute) } \longrightarrow \mathrm{MSO}_4+\mathrm{H}_2\)

M=Active metal, for example, Fe, Zn

Metal oxides and carbonates dissolve in dilute sulphuric acid.

⇒ \(\mathrm{ZnO}+\mathrm{H}_2 \mathrm{SO}_4 \text { (dilute) } \longrightarrow \mathrm{ZnSO}_4+\mathrm{H}_2 \mathrm{O}\)

⇒ \(\mathrm{CaCO}_3+\mathrm{H}_2 \mathrm{SO}_4 \text { (dilute) } \longrightarrow \mathrm{CaSO}_4+\mathrm{H}_2 \mathrm{O}+\mathrm{CO}_2\)

Low volatility Sulphuric acid is a strong acid with low volatility. The concentrated acid is used to prepare a more volatile acid in reaction with the corresponding salt. For example, when sulphuric acid is treated with sodium chloride, HC1 is evolved.

⇒ \(2 \mathrm{NaCl}+\mathrm{H}_2 \mathrm{SO}_4 \longrightarrow \mathrm{Na}_2 \mathrm{SO}_4+2 \mathrm{HCl}\)

Sulphuric acid also liberates hydrogen fluoride from a metal fluoride and nitric acid from a nitrate. Strong dehydrating agent Since concentrated sulphuric acid has a strong affinity for water, it is used to dry gases (provided they do not react with it). It removes water from organic compounds. For example, it chars sugar.

⇒ \(\mathrm{C}_{12} \mathrm{H}_{22} \mathrm{O}_{11} \stackrel{\text { conc. } \mathrm{H}_2 \mathrm{SO}_4}{\longrightarrow} 12 \mathrm{C}+11 \mathrm{H}_2 \mathrm{O}\)

Some other examples of its dehydrating action are shown.

⇒ \(\underset{\text { Formic acid }}{\mathrm{HCOOH}} \stackrel{\text { conc. } \mathrm{H}_2 \mathrm{SO}_4}{\longrightarrow} \mathrm{CO}+\mathrm{H}_2 \mathrm{O}\)

⇒ \(\begin{aligned}& \mathrm{COOH} \stackrel{\text { conc. } \mathrm{H}_2 \mathrm{SO}_4}{\longrightarrow} \mathrm{CO}+\mathrm{CO}_2+\mathrm{H}_2 \mathrm{O} \\& \mathrm{COOH}\end{aligned}\)
Oxalic acid

Oxidising agent Hot concentrated sulphuric acid is a moderately strong oxidising agent. It oxidises metals and nonmetals, itself getting reduced to sulphur dioxide.

⇒ \(\mathrm{Cu}+2 \mathrm{H}_2 \mathrm{SO}_4 \text { (conc.) } \longrightarrow \mathrm{CuSO}_4+\mathrm{SO}_2+2 \mathrm{H}_2 \mathrm{O}\)

⇒ \(\mathrm{C}+2 \mathrm{H}_2 \mathrm{SO}_4 \text { (conc.) } \longrightarrow \mathrm{CO}_2+2 \mathrm{SO}_2+2 \mathrm{H}_2 \mathrm{O}\)

⇒ \(3 \mathrm{~S}+2 \mathrm{H}_2 \mathrm{SO}_4 \text { (conc.) } \longrightarrow 3 \mathrm{SO}_2+2 \mathrm{H}_2 \mathrm{O}\)

Uses

Sulphuric Acid Is An Important Industrial Chemical And Is needed in bulk amounts in various processes. It is mainly employed in

1. The manufacture of fertilisers, for example., ammonium sulphate and superphosphate,
2. The manufacture of dyes, paints, pigments, drugs and detergents,
3. Petroleum refining,
4. Manufacture of explosives, for example, dynamite and TNT,
5. Metallurgical applications (used to clean the surface of metals before electroplating and galvanising), and
6. Storage batteries.

Group 17 Elements

This group of the p block comprises fluorine, chlorine, bromine, iodine and astatine, the last member being radioactive. Group 17 elements are called halogens (in Greek, halogen means ‘salt giver’). They are highly reactive nonmetals. The general characteristics of the elements are similar to those of the elements of Groups 1 and 2 of the periodic table. There is also a regular gradation in physical and chemical properties within the group.

Occurrence And Uses

Due to the high reactivity of the halogens they occur naturally as compounds only. They occur as halides; however, iodine, which is easily oxidised, is also found as iodates. The important ores of fluorine are fluorspar (CaF₂), fluorapatite [3Ca₃(PO₄)₂ -CaF₂] and cryolite (Na₃AlF₆).

Chlorine is widely present as sodium chloride (and to a lesser extent as chlorides of potassium, magnesium and calcium) in seawater. The deposits of dried-up seas contain these as well as camalite (KC1- MgCl₂ • 6H₂O). Bromine and iodine are less abundant.

They occur as bromides and iodides in seawater. Iodine is present in certain forms of marine life like seaweeds. Sodium iodate is present as an impurity in Chile saltpetre (NaNO₃) deposits.

Fluorine is used to make UF₆ and SF₆. The former is used for nuclear power generation and the latter is a dielectric. Important compounds obtained from fluorine are freons and Teflon (polytetrafluoroethylene). Nonstick cookware is coated with Teflon. Sodium fluoride is used for fluoridation of water. A concentration of 1 ppm of fluoride in drinking water reduces the chances of dental caries. Tin fluoride is used in fluoride toothpastes.

Chlorine is used to purify drinking water and for bleaching textiles, wood pulp and paper. It is also used for the preparation of industrially important organic compounds like chlorinated hydrocarbons (CHCI₃, CCI₄, feons, DDT) polyvinyl chloride and inorganic compounds like HCI and bleaching powder.

It is also used in the manufacture of dyes and drugs, in the extraction of gold and platinum and in the preparation of some toxic compounds like phosgene (COCl₂), tear gas (CCI₃NO₂) and mustard gas (CICH₂CH₂SCH₂CH₂CI).

Bromine is used in the manufacture of silver bromide and potassium bromide, which are respectively used in photography and as a sedative. Iodine is used for the manufacture of potassium iodide and iodoform.

A solution of iodine in alcohol is called a tincture of iodine and is used as an antiseptic. Common salt is iodised by adding sodium or potassium iodide/iodate to it. This is essential as a lack of iodine in the body leads to goitre.

Atomic And Physical Properties

Some important physical constants of the halogens are listed in. The halogens exist as diatomic molecules and a regular gradation in physical properties is observed as we descend the group.

Electronic configuration

The valence-shell electronic configuration of these elements is ns²np⁵, i.e., only one electron is needed to complete the octet.

Size

The halogens are the smallest atoms in their respective periods in the modern periodic table. This is due to their highly effective nuclear charge. The atomic size increases on moving down the group.

Ionisation enthalpy

Haloeens exhibit high ionisation enthalpies and therefore have little tendency to lose valence-shell electrons and form positive ions. On moving down the group the ionisation enthalpy decreases due to an increase in the atomic size.

Electronegativity

Halogens have very high electronegativity values which decrease on moving down the group, i.e., with an increase in size. Fluorine is the most electronegative element.

Electron gain enthalpy

The halogens have the most negative electron gain enthalpies in their respective periods. This indicates that a lot of energy is evolved when the reaction X -> X” takes place. Halogens need only one electron to acquire a stable noble-gas configuration.

This tendency decreases with an increase in size and the electron gain enthalpy becomes less negative on descending the group. However, the electron gain enthalpy of fluorine is less negative than that of chlorine. This is because the fluorine atom is very small in size and the incoming electron encounters a lot of electron-electron repulsion in a small 2 p subshell.

Physical properties

There is a steady increase in the melting and boiling points of the elements on moving down the group. At room temperature, fluorine and chlorine are gases, bromine is a liquid and iodine is a solid.

Fluorine is light yellow, chlorine is greenish-yellow, and bromine and iodine are reddish-brown and violet respectively. The absorption of visible radiation by atoms of the elements results in the promotion of electrons to higher energy levels. When these electrons return to the lower energy levels, the energy absorbed is emitted—the frequency of which lies in the visible range. Thus elements exhibit specific colours. All halogens are soluble in water.

However, their reaction to water varies. Fluorine and chlorine both react with water to form the respective halides and oxygen. The reaction is a violent case of fluorine.

Bromine and iodine are sparingly soluble in water and energy has to be supplied to make them oxidise water. Halogens dissolve in organic solvents like carbon disulphide, chloroform and carbon tetrachloride to give coloured solutions.

An anomaly is noted in a variation of bond dissociation enthalpy of the halogen molecules. The X-X bond dissociation enthalpy is expected to decrease down the group with the increase in atomic size resulting in less effective atomic overlap. This trend is well noted from chlorine to iodine in the group.

However, the F-F bond dissociation enthalpy is abnormally low. This is because the fluorine atom is very small. Also, the intemuclear distance (F-F) is small in the F₂ molecule. This leads to strong repulsion between the lone pairs of electrons in the small F₂ molecule. This factor is responsible for the high reactivity of fluorine.

Chemical Properties

Oxidation states and trends in chemical reactivity

The halogens exhibit a -1 oxidation state in most of their compounds. This is the only oxidation state displayed by fluorine.

The other halogens may display oxidation states of +1, +3, +5 and +7in some of the compounds. The positive oxidation states are displayed in halogen oxides, oxoacids and interhalogens. Oxidation states of +3, +5 and +7 are realised by the unpairing of the s electron pair and its promotion to the vacant d orbitals. Chlorine and bromine also display oxidation states of +4 and +6 in their oxides.

The halogens are highly reactive and the reactivity decreases down the group.

The high reactivity of halogens is due to the following factors.

1. Low enthalpy of dissociation The X-X bond is weak. Therefore, the diatomic molecule dissociates readily.
2. High oxidising power The halogens have high negative electron gain enthalpies, i.e., they have a strong tendency to acquire electrons and are good oxidising agents. Fluorine is the strongest oxidising agent. Generally speaking, a higher member can displace a lower member of the group from the halide. The following reactions illustrate this.

F₂+2X^–→2F^–+X₂
CI₂+2X→2CI+X₂Br₂^–+I₂
Br₂+2I^–→2Br^–+I₂

The oxidising power of halogens doorcases down the group. This is reflected in their decreasing reduction fv4oiUi.il. in descending the group.

As you know, the reduction potential of an element depends on various factors and is shown in the form of a Born-flavour cycle.

⇒ \(\begin{aligned}
\frac{1}{2} \mathrm{X}_2(\mathrm{~s}) \longrightarrow \frac{1}{2} \mathrm{X}_2(\mathrm{~g}) \stackrel{\frac{1}{2} \Lambda_{\mathrm{dit} t} f^{\mathrm{O}}}{\longrightarrow} & \mathrm{X}(\mathrm{g}) \\
\downarrow & \downarrow \Delta_{\mathrm{eg}} H^{\Theta}
\end{aligned}\)

⇒ \(\mathrm{X}^{-} \text {(aq) } \quad \stackrel{\Delta_{\text {hyd }} H^{\Theta}}{\longleftarrow} \quad \mathrm{X}^{-}(\mathrm{g})\)

The reactions of halogens with water indicate their relative oxidising powers. Fluorine oxidises water. Bromine and chlorine dissolve in water, forming hydrohalic and hypohalous acids. The reaction of iodine with water has a positive Gibb’s energy change, i.e., it is nonspontaneous. In fact, an iodide is oxidised to iodine by oxygen. This is the reverse of the reaction observed with fluorine.

⇒ \(2 \mathrm{~F}_2+2 \mathrm{H}_2 \mathrm{O} \longrightarrow 4 \mathrm{H}^{+}+4 \mathrm{~F}^{-}+\mathrm{O}_2\)

⇒ \(\mathrm{X}_2+\mathrm{H}_2 \mathrm{O} \longrightarrow \mathrm{HX}+\mathrm{HOX}\)

⇒ \(4 \mathrm{I}^{-}+4 \mathrm{H}^{+}+\mathrm{O}_2 \longrightarrow 2 \mathrm{I}_2+2 \mathrm{H}_2 \mathrm{O}\)

Anomalous Behaviour Of Fluorine

Fluorine differs from the other halogens because of its exceptionally small size, low F-F bond dissociation enthalpy and the absence of d orbitals in its valence shell.

Some specific examples of its anomalous behaviour are as follows.

• Fluorine is more reactive than other halogens due to the low F-F bond dissociation enthalpy. Also, due to its high electronegativity, it forms strong bonds with other elements.
• Most reactions of fluorine are exothermic due to the formation of strong bonds with other elements.
• Fluorine is the strongest oxidising agent and may oxidise elements to the highest oxidation state. For example, in SF6 and IF7, the oxidation states of sulphur and iodine are +6 and +7 respectively.
• In HF, fluorine forms strong hydrogen bonds. Thus, HF is a liquid while the other hydrogen halides are gases.
• Fluorine forms only one oxoacid while the other halogen forms more oxoacids.
• Due to the high electronegativity of fluorine, fluorides have greater ionic character than other halides.

Reactivity towards hydrogen

All halogens combine with hydrogen to form hydrogen halides, HX. The reaction with fluorine is violent while that with iodine is slow, indicating that the reactivity of the halogens decreases on descending the group. Hydrogen halides are covalent in the gaseous state, but in aqueous solutions, they function as strong acids. Some properties of hydrogen halides are given in.

They undergo dissociation in water giving acidic solutions.

⇒ \(\mathrm{HX}+\mathrm{H}_2 \mathrm{O} \rightleftharpoons \mathrm{H}_3 \mathrm{O}^{+}+\mathrm{X}\)

⇒ \(\mathrm{HF}<\mathrm{HCl}<\mathrm{HBr}<\mathrm{HI}\)

Various factors contribute to the acid strength and an important factor is bond enthalpy. The bond enthalpy decreases from HF to HI. Thus HF is the weakest acid, and HI is the strongest acid. Intermolecular hydrogen bonding also contributes to the low acidity of HF.

Reactivity towards oxygen

A large number of binary compounds of halogens and oxygen are known. The important compounds are shown in.

Binary compounds of fluorine with oxygen are called oxygen fluorides due to the higher electronegativity of fluorine, whereas the analogous compounds with other halogens are called halogen oxides. As the electronegativity difference between the halogens and oxygen is small, the bonds are essentially covalent.

Oxygen fluorides are strong fluorinating agents. O₂F₂ is used to remove plutonium from spent nuclear fuel by converting it to PuF₆. Apart from OF₂, all oxides have positive Gibbs energy of formation and are unstable with respect to dissociation to the elements.

The oxides of iodine are the most stable followed by the oxides of chlorine. The oxides of bromine are the least stable. This can be explained by taking into consideration both kinetic and thermodynamic factors of the reactions involved. The higher oxides tend to be more stable than the lower ones.

Except for that of iodine, all oxides tend to be explosive. Because of its oxidising nature, CIO₂ is used as a bleaching agent for paper pulp and textiles and as a germicide in water treatment. I₂O₅ quantitatively oxidises carbon monoxide to carbon dioxide and is therefore used in the estimation of carbon monoxide.

Reactivity towards metals

The halogens combine with most metals to form metal halides.

⇒ \(2 \mathrm{M}+n \mathrm{X}_2 \longrightarrow 2 \mathrm{MX}_n\)

Fluorine and chlorine react with practically all metals whereas bromine and iodine combine with most metals except the noble metals. The fluorides are all ionic. This is quite obvious as fluorine is the most electronegative element in the periodic table and when it combines with an electropositive metal, owing to the large electronegativity difference an ionic compound is formed. The ionic character of a metal halide follows the order:

fluoride > chloride > bromide > iodide.

Among metal halides in which the constituent elements are the same, the halide in which the halogen is in the higher oxidation state is more covalent than the one in which the halogen is in the lower oxidation state. Thus TICl₃, PbCl₄, AsCl₃, UF₄ are more covalent thanTICl, PbCl₂, AsCl₃ and UF₄ respectively.

Reactivity towards other halogens

The halogens combine among themselves to form binary compounds called interhalogen compounds. Interhalogens of the type XX’, XX₃, XX’₅ and XX’₇ are known. (Here X is the higher member and X is the lower member of the group.) We will discuss interhalogens later in this chapter.

Chlorine

Chlorine was first prepared by Scheele (1774) by oxidising HCI with MnO₂. Davy (1810) identified it to be an element and coined the name chlorine on account of its colour (chlorosis in Greek means greenish-yellow).

Preparation

Chlorine can be prepared by oxidising HCI with MnO₂

⇒ \(\mathrm{MnO}_2+4 \mathrm{HCl} \longrightarrow \mathrm{Cl}_2+\mathrm{MnCl}_2+2 \mathrm{H}_2 \mathrm{O}\)

In practice, a mixture of common salt and concentrated H₂SO₄ is taken in place of HCI.

⇒ \(\mathrm{MnO}_2+4 \mathrm{NaCl}+4 \mathrm{H}_2 \mathrm{SO}_4 \longrightarrow \mathrm{MnCl}_2+4 \mathrm{NaHSO}_4+2 \mathrm{H}_2 \mathrm{O}+\mathrm{Cl}_2\)

It is also obtained upon the oxidation of HC1 by potassium permanganate.

⇒ \(2 \mathrm{KMnO}_4+16 \mathrm{HCl} \longrightarrow 2 \mathrm{KCl}+2 \mathrm{MnCl}_2+8 \mathrm{H}_2 \mathrm{O}+5 \mathrm{Cl}_2\)

Chlorine is produced commercially by two main processes as follows.

1. The electrolysis of brine (concentrated sodium chloride solution) in the manufacture of NaOH
   ⇒ \(2 \mathrm{NaCl}+2 \mathrm{H}_2 \mathrm{O} \stackrel{\text { electrolysis }}{\longrightarrow} 2 \mathrm{NaOH}+\mathrm{Cl}_2+2 \mathrm{H}_2\)
2. By the oxidation of hydrogen chloride Gatsby oxygen at 723 Kin the presence of copper chloride as a catalyst
   ⇒ \(4 \mathrm{HCl}+\mathrm{O}_2 \longrightarrow 2 \mathrm{H}_2 \mathrm{O}+2 \mathrm{Cl}_2\)

Properties

Chlorine is a pungent-smelling, yellowish-green gas which can be liquefied to a greenish-yellow liquid. It reacts with metals and nonmetals to form chlorides.

⇒ \(2 \mathrm{Na}+\mathrm{Cl}_2 \longrightarrow 2 \mathrm{NaCl}\)

⇒ \(\mathrm{Mg}+\mathrm{Cl}_2 \longrightarrow \mathrm{MgCl}_2\)

⇒ \(2 \mathrm{Al}+3 \mathrm{Cl}_2 \longrightarrow 2 \mathrm{AlCl}_3\)

⇒ \(\mathrm{P}_4+6 \mathrm{Cl}_2 \longrightarrow 4 \mathrm{PCl}_3\)

⇒ \(\mathrm{S}_8+4 \mathrm{Cl}_2 \longrightarrow 4 \mathrm{~S}_2 \mathrm{Cl}_2\)

⇒ \(\mathrm{H}_2+\mathrm{Cl}_2 \longrightarrow 2 \mathrm{HCl}\)

It reacts with compounds containing hydrogen (i.e., hydrides, hydrocarbons, etc.) to form hydrogen chloride.

⇒ \(\mathrm{H}_2 \mathrm{~S}+\mathrm{Cl}_2 \longrightarrow 2 \mathrm{HCl}+\mathrm{S}\)

⇒ \(\mathrm{C}_{10} \mathrm{H}_{16}+8 \mathrm{Cl}_2 \longrightarrow 16 \mathrm{HCl}+10 \mathrm{C}\)

The reactions of chlorine with saturated hydrocarbons are substitution reactions and those with unsaturated hydrocarbons are addition reactions.

⇒ \(\mathrm{CH}_4+\mathrm{Cl}_2 \stackrel{\mathrm{uv}}{\longrightarrow} \mathrm{CH}_3 \mathrm{Cl}+\mathrm{HCl}\)

⇒ \(\mathrm{CH}_3 \mathrm{Cl}+\mathrm{Cl}_2 \stackrel{\text { uv }}{\longrightarrow} \mathrm{CH}_2 \mathrm{Cl}_2+\mathrm{HCl}\)

⇒ \(\mathrm{CH}_2 \mathrm{Cl}_2+\mathrm{Cl}_2 \stackrel{\mathrm{uv}}{\longrightarrow} \mathrm{CHCl}_3+\mathrm{HCl}\)

⇒ \(\mathrm{CHCl}_3+\mathrm{Cl}_2 \stackrel{\mathrm{uv}}{\longrightarrow} \mathrm{CCl}_4+\mathrm{HCl}\)

⇒ \(\mathrm{H}_2 \mathrm{C}=\mathrm{CH}_2+\mathrm{Cl}_2 \longrightarrow \underset{\text { 1,2-dichloroethane }}{\mathrm{C}_2 \mathrm{H}_4 \mathrm{Cl}_2}\)

Chlorine reacts with ammonia to give different products depending on the relative proportions of the reactants. With excess ammonia, the products are nitrogen and ammonium chloride, while with excess chlorine, nitrogen trichloride is formed.

⇒ \(8 \mathrm{NH}_3+3 \mathrm{Cl}_2 \longrightarrow 6 \mathrm{NH}_4 \mathrm{Cl}+\mathrm{N}_2\)

⇒ \(\mathrm{NH}_3+3 \mathrm{Cl}_2 \longrightarrow \mathrm{NCl}_3+3 \mathrm{HCl}\)

Chlorine dissolves in water to give a yellow solution. This is called chlorine water. Soon after the solution is prepared it turns colourless due to the following reaction.

⇒ \(\mathrm{Cl}_2+\mathrm{H}_2 \mathrm{O} \longrightarrow \mathrm{HCl}+\mathrm{HOCl}\)

Hypochlorous acid (HOC1) releases nascent oxygen, which is responsible for the oxidising and bleaching action of aqueous chlorine.

⇒ \(\mathrm{HOCl} \longrightarrow \mathrm{HCl}+\mathrm{O}\)

The bleaching action is permanent. Freshly prepared chlorine water is a strong oxidising agent and oxidises ferrous to ferric.

⇒ \(2 \mathrm{Fe}^{2+}+\mathrm{Cl}_2 \longrightarrow 2 \mathrm{Fe}^{3+}+2 \mathrm{Cl}^{-}\)

It also oxidises sulphur dioxide to sulphuric acid and iodine to iodic acid.

⇒ \(\mathrm{SO}_2+2 \mathrm{H}_2 \mathrm{O}+\mathrm{Cl}_2 \longrightarrow \mathrm{H}_2 \mathrm{SO}_4+2 \mathrm{H}^{+}+2 \mathrm{Cl}^{-}\)

⇒ \(\mathrm{I}_2+6 \mathrm{H}_2 \mathrm{O}+5 \mathrm{Cl}_2 \longrightarrow 2 \mathrm{HIO}_3+10 \mathrm{H}^{+}+10 \mathrm{Cl}^{-}\)

Chlorine in reaction with cold and dilute alkalis produces a mixture of chloride and hypochlorite. However, with hot and concentrated alkalis, chlorine forms chloride and chlorate.

⇒ \(\underset{\text { (dilute) }}{2 \mathrm{NaOH}}+\mathrm{Cl}_2 \stackrel{\text { cold }}{\longrightarrow} \underset{\begin{array}{c}
\text { Sodium } \\
\text { hypochlorite }
\end{array}}{\mathrm{NaOCl}}+\mathrm{NaCl}+\mathrm{H}_2 \mathrm{O}\)

⇒ \(\underset{\text { (concentrated) }}{6 \mathrm{NaOH}}+\mathrm{Cl}_2 \stackrel{\text { hot }}{\longrightarrow} \underset{\begin{array}{l}
\text { Sodium } \\
\text { chlorate }
\end{array}}{\mathrm{NaClO}_3}+5 \mathrm{NaCl}+5 \mathrm{H}_2 \mathrm{O}\)

Both these reactions are disproportionation reactions as chlorine (oxidation state zero) undergoes simultaneous oxidation to hypochlorite (oxidation state +1) or chlorate (oxidation state +3) and chloride (oxidation state -1).

The reaction of chlorine with slaked lime yields bleaching powder.

⇒ \(2 \mathrm{Ca}(\mathrm{OH})_2+2 \mathrm{Cl}_2 \longrightarrow \mathrm{Ca}(\mathrm{OCl})_2+\mathrm{CaCl}_2+2 \mathrm{H}_2 \mathrm{O}\)

It is a mixture of calcium hypochlorite, calcium chloride and calcium hydroxide and is represented as Ca(OCI)₂.CaCI₂.Ca(OH)₂.2H₂O.

Hydrogen Chloride

It is prepared by heating sodium chloride with concentrated sulphuric acid.

⇒ \(\mathrm{NaCl}+\mathrm{H}_2 \mathrm{SO}_4 \stackrel{\Delta}{\longrightarrow} \mathrm{NaHSO}_4+\mathrm{HCl}\)

⇒ \(\mathrm{NaHSO}_4+\mathrm{NaCl} \stackrel{\Delta}{\longrightarrow} \mathrm{Na}_2 \mathrm{SO}_4+\mathrm{HCl}\)

⇒ \(2 \mathrm{NaCl}+\mathrm{H}_2 \mathrm{SO}_4 \stackrel{\Delta}{\longrightarrow} \mathrm{Na}_2 \mathrm{SO}_4+2 \mathrm{HCl}\)

The HCI produced is dried by passing through concentrated sulphuric acid.

The presence of HCI can be detected by holding a glass rod dipped in ammonia at the mouth of the test tube where the reaction has taken place when white fumes of ammonium chloride have evolved.

⇒ \(\mathrm{NH}_3+\mathrm{HCl} \longrightarrow \mathrm{NH}_4 \mathrm{Cl}\)

Properties

Hydrogen chloride is an extremely pungent-smelling, colourless gas which can be liquefied to give a colourless liquid and solidified to give a white solid.

It is highly soluble in water and yields an acidic solution, which is called hydrochloric acid.

⇒ \(\mathrm{HCl}+\mathrm{H}_2 \mathrm{O} \longrightarrow \mathrm{H}_3 \mathrm{O}^{+}+\mathrm{Cl}^{-}\)

Hydrochloric acid is a strong acid. It liberates hydrogen from reacting with active metals.

⇒ \(\mathrm{Zn}+2 \mathrm{HCl} \longrightarrow \mathrm{ZnCl}_2+\mathrm{H}_2\)

⇒ \(\mathrm{Fe}+2 \mathrm{HCl} \longrightarrow \mathrm{FeCl}_2+\mathrm{H}_2\)

In reaction with iron, the product formed is ferrous chloride and not ferric chloride as the evolved hydrogen prevents further oxidation of iron. When salts of weak acids (carbonates, sulphites, thiosulphates, sulphides) are treated with dilute hydrochloric acid, the anion is displaced.

⇒ \(\mathrm{Na}_2 \mathrm{CO}_3+2 \mathrm{HCl} \longrightarrow 2 \mathrm{NaCl}+\mathrm{H}_2 \mathrm{O}+\mathrm{CO}_2\)

⇒ \(\mathrm{Na}_2 \mathrm{SO}_3+2 \mathrm{HCl} \longrightarrow 2 \mathrm{NaCl}+\mathrm{H}_2 \mathrm{O}+\mathrm{SO}_2\)

⇒ \(\mathrm{Na}_2 \mathrm{~S}_2 \mathrm{O}_3+2 \mathrm{HCl} \longrightarrow 2 \mathrm{NaCl}+\mathrm{H}_2 \mathrm{O}+\mathrm{SO}_2+\mathrm{S}\)

⇒ \(\mathrm{Na}_2 \mathrm{~S}+2 \mathrm{HCl} \longrightarrow 2 \mathrm{NaCl}+\mathrm{H}_2 \mathrm{~S}\)

A mixture of three parts of concentrated hydrochloric acid and one part of concentrated nitric acid is called aqua regia Aqua regia combines with noble metals forming soluble chloro complexes.

⇒ \(\mathrm{Au}+4 \mathrm{H}^{+}+4 \mathrm{Cl}^{-}+\mathrm{NO}_3^{-} \longrightarrow\left[\mathrm{AuCl}_4\right]^{-}+\mathrm{NO}+2 \mathrm{H}_2 \mathrm{O}\)

⇒ \(3 \mathrm{Pt}+16 \mathrm{H}^{+}+18 \mathrm{Cl}^{-}+4 \mathrm{NO}_3^{-} \longrightarrow 3\left[\mathrm{PtCl}_6\right]^{2-}+4 \mathrm{NO}+8 \mathrm{H}_2 \mathrm{O}\)

Uses

1. Hydrochloric acid is used as a laboratory reagent
2. It is used in the pharmaceutical industry
3. It is used in the manufacture of chlorine and other chemicals.

Oxoacids Of Halogens

Oxoacids of halogens have oxygen attached to the halogen. They have the general formula HOX(O)_n where n=0, 1, 2, 3. Due to its small size and high electronegativity, fluorine forms only one oxoacid—the unstable HOF. The other halogens form four acids—hypohalous (HOX), halos (HOXO), halic (HOXO₂) and perhalic (HOXO₃). Most of the oxoacids are known solutions or salts. The oxoacids of halogens and their structures are shown in.

Interhalogen Compounds

These are binary compounds of two different halogen atoms having the general formula AX_n where both A and X are halogens and X has a higher electronegativity. The value of n may be 1, 3, 5 or 7. The halogen with the lower e ec onega city is assigned a positive oxidation state.

The total number of halogen atoms is always even as this gives rise to a magnetic species, which is generally more stable. The interhalogens can be prepared by a direct combination of the halogens or by the action of a halogen on an interhalogen compound.

A few examples are as follows.

CI₂+F₂→2CIF
CI₂+3F₂→2CIF₃
Br₂+5F₂→2BrF₃
CIF₃+F₂→CIF₅
IF₅+F₂→IF₇

Lists some interhalogens of the four types together with their physical state and colour at room temperature. In the interhalogens of the type AX₅ and AX₇, A is a large halogen (Br or I) and X is fluorine. This type exists as it is easier to pack a large number of small atoms around a large atom.

Most of the interhalogens are volatile solids or liquids and their physical properties are intermediate between those of the constituent halogens. The stability of interhalogen compounds increases with an increase in the difference in electronegativity between the two halogen atoms.

Interhalogen compounds are predominantly covalent and are generally more reactive than halogens (except fluorine). This is because the A-X bond in interhalogens is less stable than the X-X bond in halogens, except for the F-F bond.

All interhalogens hydrolyse to give a halide and an oxohalide (hypohalite in case of AX, halitein case of AX₃, halate in case of AX₅ and perhalate in case of AX₇). The smaller halogen is converted to the halide and the larger to the oxohalide.

⇒ \(\mathrm{ClF}+\mathrm{H}_2 \mathrm{O} \longrightarrow \mathrm{F}^{-}+\mathrm{OCl}^{-}+2 \mathrm{H}^{+}\)

⇒ \(\mathrm{ICl}+\mathrm{H}_2 \mathrm{O} \longrightarrow \mathrm{Cl}^{-}+\mathrm{OI}^{-}+2 \mathrm{H}^{+}\)

The VSEPR theory is applied to assign structures to the interhalogen compounds. Let us discuss a few types here.

In CIF₃ (AX₃ type), the chlorine atom has seven electrons in the outermost shell, out of which three will be used in a bond formation with three fluorine atoms, leaving behind four unused electrons. Thus, there are two lone pairs and three bond pairs, giving rise to sp3d hybridisation, i.e., a trigonal bipyramidal structure.

The two lone pairs will occupy the equatorial positions to minimise repulsions. The molecule is T-shaped with the axial fluorines slightly bent away from the lone pairs and the T is slightly bent.

In order to quickly assign the hybridisation of the central atom in the compound to deduce its structure, the following steps will prove to be helpful.

1. Count the valence electrons of all the atoms. Add or subtract electrons equal to the charge for anions and cations respectively.
2. If the total number of valence electrons in the compound is less than 8, divide by 2. The quotient gives the number of electron pairs. For two electron pairs the hybridisation is sp, for three it is sp2 and for four it is sp3.
3. If the total number of valence electrons exceeds 8, divide by 8 to get the quotient Qj and remainder R. If R = 0, then Q: gives a number of electron pairs. If R≠O then further divide R by 2 to get Q₂. Add Q1 and Q₂ to get the number of electron pairs and assign hybridisation accordingly (Ch = number of bond pairs and Q₂ =number of lone pairs).

Let us now deduce the hybridisation of chlorine in CIF₃ by following the abovementioned steps in sequence.

1. Total number of valence electrons = 7 + 3 x 7 = 28
2. ⇒ \(Q_1=\frac{28}{8}=3 ; R_1=4\)
3. ⇒ \(Q_2=\frac{R_1}{2}=2 ; \quad ∴ Q_1+Q_2=5\)

There are five electron pairs around the central atom, and the hybridisation is sp3d

In interhalogen compounds of the type AX₅, for example, BrF₅, the hybridisation is deduced as follows.

1. Total number of valence electrons =7+5×7 = 42
2. \(Q_1=\frac{42}{8}=5 ; R_1=2\)
3. \(Q_2=\frac{R_1}{2}=1 ; \quad ∴ Q_1+Q_2=5+1=6\)

There are six electron pairs around the central atom, giving rise to sp³d² hybridisation.

There are four bond pairs and one lone pair, giving rise to a square pyramidal structure. Similarly, it can be shown that IF7 has sp3d3 hybridization, giving rise to a pentagonal bipyramidal structure.

Uses

Interhaloeens are useful as nonaqueous solvents. Many of these undergo self-ionisation, for example.

⇒ \(2 \mathrm{BrF}_3 \longrightarrow \mathrm{BrF}_2^{+}+\mathrm{BrF}_4^{-}\)

C1F₃ and BrF₃ are good fluorinating agents and are used in the enrichment of converted to hexafluoride.

⇒ \(\mathrm{U}+3 \mathrm{ClF}_3 / 3 \mathrm{BrF}_3 \longrightarrow \mathrm{UF}_6+3 \mathrm{CIF} / 3 \mathrm{BrF}\)

Group 18 Elements

The elements of this group are helium, neon, argon, krypton, xenon and radon. Group 18 elements are called inert gases, rare gases or noble gases. The last member, radon, is radioactive and short-lived. Group 18 elements are chemically inert as their valence-shell orbitals are completely filled. They react with few compounds and are therefore called noble gases.

Occurrence And Uses

Apart from radon, the other noble gases are present in the atmosphere to an extent of about 1% (the major component being argon). Helium is commercially obtained from natural gas, the other sources of helium and neon are radioactive minerals like pitchblende and monazite. Radon is obtained by the radioactive decay of 27h Ra.

⇒ \({ }_{88}^{226} \mathrm{Ra} \longrightarrow{ }_{86}^{222} \mathrm{Rn}+{ }_2^4 \mathrm{He}\)

Helium is noninflammable and light and therefore used to fill balloons employed in meteorological observations. A mixture of helium and oxygen is used by deep-sea divers for respiration and is also used to provide relief to asthmatic patients. Helium has the lowest boiling point among all elements and thus is used to carry out research at very low temperatures. It is nonradioactive has high thermal conductivity, and is therefore used as a heat transfer agent in gas-cooled atomic reactors. Liquid helium is used as a coolant in superconducting coils used to build superconducting magnets which form a part of the NMR spectrometers used in MR! systems for clinical diagnosis.

Neon is used in discharge tubes and fluorescent lamps for advertising purposes. Neon bulbs are used in gardens and greenhouses. Neon can carry high currents at high voltages and is used to protect electrical instruments, e.g., voltmeters and rectifiers.

Argon is used to create an inert atmosphere in metallurgical operations carried out at high temperatures, for example., welding. It is used to fill electric light bulbs. Krypton and xenon are used for filling electric bulbs and phototubes. Radon is used in radioactive research and in the treatment of cancer.

Atomic And Physical Properties

Some physical characteristics of the noble gases are summarised in.

The electronic configuration of helium is Is2 while the other members of the group have a complete octet of electrons in their outermost shell, or the valence-shell configurations in these elements are MS2up6.

All the noble gases are monatomic. They have large atomic radii, which are actually van der Waals radii (van der Waals radii are nonbonded radii, which represent the distance of closest approach) and not ionic or covalent radii. The atomic radii increase down the group. They have high ionisation enthalpy due to a stable electronic configuration. They have no tendency to accept electrons, as is reflected by large, positive electron gain enthalpy values.

Noble gases are colourless, tasteless odourless and sparingly soluble in water. Noble-gas atoms are held together by weak van der Waals forces, which can be readily overcome and hence they have low melting and boiling points. As stated earlier, helium has the lowest boiling point among all elements (4.2 K). It has the unique property of diffusing through substances like rubber, glass and plastic.

Chemical Reactivity And Compounds

The stable electronic configuration of the elements leads to the involvement of high energies in the loss or gain of electrons (revealed by high ionisation enthalpy values and positive electron gain enthalpy values). Owing to this the noble gases show very low chemical reactivity.

The reactivity of noble gases was constantly investigated and the first real compound of a noble gas was made in 1962. Bartlett had observed that PtF6 combined with molecular oxygen (Oz) to form 02[PtF6 Since the first ionisation enthalpies of O₂ (O₂ ->O₂⁺, 1175 kJ mol^-1) and Xe (Xe —» Xe+; 1170 kJ mol^-1) are comparable, it was predicted that a similar compound could be formed by making Xe react with PtF₆. In fact, PtF₆ vapours reacted with xenon at room temperature to form a reddish-yellow solid which was incorrectly thought of as xenon hexa fluoroplatinate.

⇒ \(\mathrm{Xe}(\mathrm{g})+\mathrm{PtF}_6(\mathrm{~g}) \longrightarrow \underset{\text { Xenon hexafluoroplatinate }}{\mathrm{Xe}^{+}\left[\mathrm{PtF}_6\right]^{-}(\mathrm{s})}\)

Actually, the product formed was really a more complicated one [XeF]⁺[Pt₂F₁₁]^–. After this, it was found that xenon combined with electronegative elements like fluorine and oxygen to form various compounds. Xenon oxofluorides can be obtained from fluorides. Fluorine and oxygen are strong oxidising agents and highly electronegative and hence can oxidise xenon to display positive oxidation states in compounds.

The ionisation enthalpies of He, Ne and Ar are too high to allow the formation of compounds. The ionisation enthalpy of Kr is lower, and so KrF₂ is known, Rn is radioactive and its compounds have not been isolated; the < presence of RnF₂ has been confirmed by radiotracer techniques.

Xenon-Fluorine Compounds

Xenon forms three fluorides, XeF₂, XeF₄ and XeF₆ by direct combination with fluorine. The fluoride former depends on the prevailing experimental conditions.

In all the reactions involving the formation of a xenon fluoride, the reactants are taken in a nickel container and subjected to high temperature and pressure. Depending on the proportion of reactants, temperature and pressure that prevails during the reaction, different xenon fluorides are formed, which are shown in the following.

⇒ \(\begin{aligned}
& \mathrm{Xe}(\mathrm{g})+\mathrm{F}_2(\mathrm{~g}) \stackrel{673 \mathrm{~K}, 1 \mathrm{bar}}{\longrightarrow} \mathrm{XeF}_2(\mathrm{~g}) \\
& \quad(2: 1 \text { ratio })
\end{aligned}\)

⇒ \(\begin{aligned}
& \mathrm{Xe}(\mathrm{g})+2 \mathrm{~F}_2(\mathrm{~g}) \stackrel{873 \mathrm{~K}, 6-7 \text { bar }}{\longrightarrow} \mathrm{XeF}_4(\mathrm{~g}) \\
& \quad(1: 5 \text { ratio })
\end{aligned}\)

⇒ \(\begin{aligned}
& \mathrm{Xe}(\mathrm{g})+3 \mathrm{~F}_2(\mathrm{~g}) \stackrel{573 \mathrm{~K}(60-70 \mathrm{bar})}{\longrightarrow} \mathrm{XeF}_6(\mathrm{~g}) \\
& \quad(1: 20 \text { ratio })
\end{aligned}\)

XeF₆ can also be made by the fluorination of XeF₄ with oxygen fluoride at low temperature

⇒ \(\mathrm{XeF}_4+\mathrm{O}_2 \mathrm{~F}_2 \stackrel{143 \mathrm{~K}}{\longrightarrow} \mathrm{XeF}_6+\mathrm{O}_2\)

The xenon fluorides are colourless, crystalline solids which sublime at 298 K. They are powerful fluorinating agents. They are susceptible to hydrolysis. Hydrogen fluoride is formed during the hydrolysis of all xenon fluorides.

⇒ \(2 \mathrm{XeF}_2+2 \mathrm{H}_2 \mathrm{O} \longrightarrow 2 \mathrm{Xe}+4 \mathrm{HF}+\mathrm{O}_2\)

⇒ \(6 \mathrm{XeF}_4+12 \mathrm{H}_2 \mathrm{O} \longrightarrow 4 \mathrm{Xe}+2 \mathrm{XeO}_3+24 \mathrm{HF}+3 \mathrm{O}_2\)

XeF₆ undergoes partial hydrolysis as well as complete hydrolysis.

⇒ \(\mathrm{XeF}_6+\mathrm{H}_2 \mathrm{O} \longrightarrow \underset{\substack{\text { Xenon } \\ \text { oxofluoride }}}{\mathrm{XeOF}_4}+2 \mathrm{HF}\)

⇒ \(\mathrm{XeF}_6+2 \mathrm{H}_2 \mathrm{O} \longrightarrow \mathrm{XeO}_2 \mathrm{~F}_2+4 \mathrm{HF}\)

⇒ \(\mathrm{XeF}_6+3 \mathrm{H}_2 \mathrm{O} \longrightarrow \mathrm{XeO}_3+6 \mathrm{HF}\)

The products of partial hydrolysis are xenon oxyfluorides and HF while that of complete hydrolysis is xenon trioxide.

It may be noted here that the oxidation state of xenon in XeF₆, XeOF₄, XeO₂F₂ and XeO₃ is +6 and thus the hydrolysis of XeF₆ whether partial or complete is not a redox reaction, unlike in the other cases, where there is a change in the oxidation state of Xe. [XeF₂ gives Xe, the change in oxidation state is from +2 to zero; XeF₄ has Xe in the +4 oxidation state, it gives Xe (zero oxidation state).]

Xenon fluoride acts as a fluoride acceptor as well as a fluoride donor. In reactions with fluoride ion acceptors, it forms cationic species whereas in reactions with fluoride ion donors, it forms anionic species.

⇒ \(\mathrm{XeF}_2+\underset{\substack{\text { Fluoride } \\ \text { ion acceptor }}}{\mathrm{PF}_5} \longrightarrow[\mathrm{XeF}]^{+}\left[\mathrm{PF}_6\right]^{-}\)

⇒ \(\mathrm{XeF}_4+\underset{\substack{\text { Fluoride ion } \\ \text { acceptor }}}{\mathrm{SbF}_5} \longrightarrow\left[\mathrm{XeF}_3\right]^{+}\left[\mathrm{SbF}_6\right]^{-}\)

⇒ \(\mathrm{XeF}_6+\underset{\substack{\text { Fluoride } \\ \text { ion donor }}}{\mathrm{MF}} \longrightarrow \mathrm{M}^{+}\left[\mathrm{XeF}_7\right]^{-}\)

Xenon fluorides are strong oxidising and fluorinating agents and combine quantitatively with hydrogen.

⇒ \(\mathrm{XeF}_2+\mathrm{H}_2 \longrightarrow \mathrm{Xe}+2 \mathrm{HF}\)

⇒ \(\mathrm{XeF}_4+2 \mathrm{H}_2 \longrightarrow \mathrm{Xe}+4 \mathrm{HF}\)

⇒ \(\mathrm{XeF}_6+3 \mathrm{H}_2 \longrightarrow \mathrm{Xe}+6 \mathrm{HF}\)

The structures of xenon fluorides can be derived from the VSEPR theory. In XeF₂, Xe has eight valence electrons out of which two are involved in bond formation. Thus there are two bond pairs and three lone pairs giving rise to sp³d hybridisation with the lone pairs in equatorial positions.

The resultant structure is linear. On the other hand, in XeF₄ the central atom undergoes sp³d₂ hybridisation. There are two lone pairs and the resultant structure is square planar. XeF₆ has a distorted octahedral structure involving sp³d₃ hybridisation. There is one lone pair present at the centre of one of the triangular faces.

Xenon-Oxygen Compounds

Oxides

Two oxides of xenon—XeO₃ and XeO₄—are known. XeO₃ is more important and is formed during the hydrolysis of XeF₄ and XeF₆. It is highly explosive and a vigorous oxidising agent. It has a pyramidal structure involving sp3 hybridisation and one lone pair.

Oxyfluorides

Xenon oxyfluoride (XeOF₂) is obtained by the partial hydrolysis of XeF₄

⇒ \(\mathrm{XeF}_4+\mathrm{H}_2 \mathrm{O} \longrightarrow \mathrm{XeOF}_2+2 \mathrm{HF}\)

Xenon oxytetrafluoride (XeOF₄) and xenon dioxydifluoride (XeO₂F₂) are obtained during the partial hydrolysis of XeF₆. The structures of the oxyfluorides are shown in XeOF₂ has a T-shaped structure due to sp³d hybridisation and two lone pairs. XeOF₄ has a square-pyramid structure because xenon is sp³d₂ hybridised with one lone pair. In XeO₂F₂ xenon is sp³d hybridised with one lone pair, giving rise to a see-saw geometry.

The P Block Elements Multiple-Choice Questions

Question 1. Which of the following is not an acidic oxide?

1. N₂O
2. N₂O
3. P₄O₆
4. SO₂

Answer: 1. N₂O

Question 2. In the molecular state, phosphorus exists as

1. P
2. P₂
3. P₄
4. Px

Answer: 3. P₄

Question 3. Which of the following is true for PCI₅?

1. All the P-Cl bonds are equivalent.
2. The axial bonds are longer than the equatorial bonds.
3. The axial bonds are shorter than the equatorial bonds.
4. The five P-Cl bond lengths differ from each other.

Answer: 2. The axial bonds are longer than the equatorial bonds.

Question 4. The strongest base among the following is

1. NF₃
2. NH₃
3. PF₃
4. PH₃

Answer: 2. NH₃

Question 5. The hybridisation of P in PCI₃ is

1. sp²
2. sp²d
3. sp³d²
4. sp³

Answer: 4. sp³

Question 6. Which of these is paramagnetic?

1. N₂O
2. N₂O
3. N₂O₃
4. N₂O₄

Answer: 2. N₂O

Question 7. PH3 is evolved when phosphorus is treated with

1. H₂O
2. HCI
3. NaOH
4. None of these

Answer: 3. NaOH

Question 8. Solid PClg is made up of

1. discrete PCI₅ units
2. [PCI₆]⁺[PCI₄]^–
3. [PCI₄]⁺[PCI₆]^–
4. [P₂CI₈]²⁺[P₂CI₁₂]^2-

Answer: 3. [PCI₄]⁺[PCI₆]^–

Question 9. Which of these contains an O-O linkage?

1. H₂SO₃
2. H₂S₂O₃
3. H₂SO₄
4. H₂S₂O₈

Answer: 4. H₂S₂O₈

Question 10. Which of the following is not a diprotic acid?

1. H₂SO₃
2. H₂SO₄
3. H₃PO₃
4. H₃PO₄

Answer: 3. H₃PO₃

Question 11. Which of the following is a reducing agent?

1. SO₂
2. SO₃
3. NO₂
4. CO₂

Answer: 1. SO₂

Question 12. The correct order of stability of Group 15 hydrides is

1. H₂O>H₂S>H₂Se>H₂Te
2. H₂Te>H₂Se>H₂S>H₂O
3. H₂S>H₂O>H₂Se>H₂Te
4. H₂Te>H₂Te>H₂Se>H₂S

Answer: 1. H₂O>H₂S>H₂Se>H₂Te

Question 13. The correct order of bond enthalpy among the following is

1. CI₂<I₂<Br₂<F₂
2. F₂<I₂<CI₂<Br₂
3. F₂<Br₂<CI₂<I₂
4. I₂<F₂<Br₂<CI₂

Answer: 4. I₂<F₂<Br₂<CI₂

Question 14. In the reaction Cl2+ 2X“ —> X2+ 2C1, X is

1. Br orI
2. I only
3. F, Br or I
4. F only

Answer: 1. Br orI

Question 15. The reaction 3C10 —> C103 + 2C1 is an example of

1. Decomposition
2. Disproportionation
3. Oxidation
4. Reduction

Answer: 2. Disproportionation

Question 16. When concentrated sulphuric acid is added to sodium chloride, the gas evolves is

1. SO₂ 
2. SO₃
3. CI₂
4. HCI

Answer: 4. HCI

Question 17. Hydrogen bonding is not present in

1. PH₃
2. NH₃
3. H₂O
4. HF

Answer: 1. PH₃

Question 18. Among hydrogen halides, the correct order of acidity is

1. HF<HCI<HI<HBr
2. HI<HBr<HCI<HF
3. HF<HCI<HBr<HI
4. HI<HBr<HF<HCI

Answer: 3. HF<HCI<HBr<HI

Question 19. The hybridisation of Cl in CIF₃ is

1. sp²d
2. sp²
3. sp³
4. sp³d

Answer: 4. sp³d

Question 20. Which of these is paramagnetic?

1. CI₂
2. S₂
3. N₂
4. Br₂

Answer: 2. S₂

Question 21. The hydrolysis product of XeF4 is

1. Xe+HF
2. XeO₃+O₂+HF
3. Xe+HF+XeO₃+O₂
4. XeOF₄

Answer: 3. Xe+HF+XeO₃+O₂

Question 22. Which of these is a planar molecule?

1. XeO₃
2. XeOF₄
3. XeO₄
4. XeF₄

Answer: 4. XeF₄

Question 23. The hybridisation of Xe in XeF₄ is

1. sp³d
2. sp³d
3. sp²
4. sp³

Answer: 2. sp³d

Question 24. Which among the following is tetrahedral?

1. XeO₄
2. XeF₄
3. XeOF₄
4. XeOF₂

Answer: 1. XeO₄

Question 25. Which among these does not exist?

1. XeF₂
2. NeF₂
3. XeO₃
4. XeF₄

Answer: 2. NeF₂

Question 26. The N2 molecule is isoelectronic with

1. CO^–,CN⁺ and NO⁺
2. CO, CN^–and NO⁺
3. CO⁺,N₂O and O₂^-2
4. O₂⁺O₂^– and CO⁺

Answer: 2. CO, CN^–and NO⁺

Question 27. Which of these metals is rendered passive on treatment with concentrated HNO₃?

1. Cr
2. Ni
3. Cu
4. Fe

Answer: 1. Cr

Question 28. Which of these is a dibasic acid and reducing agent?

1. H₃PO₄
2. H₃PO₂
3. H₃PO₃
4. HPO₃

Answer: 3. H₃PO₃

Haloalkanes And Haloarenes

An alkyl halide (haloalkane) or aryl halide (haloarene) is formed when one of the hydrogens in an alkane or benzene molecule is replaced by a halogen atom (fluorine, chlorine, bromine or iodine). For Example,

Chemists often use R-X as a general formula for alkyl halides; R stands for any alkyl group and X for any halogen. The aryl group in a halobenzene is denoted by the symbol Ar to distinguish it from the alkyl group R.

These classes of compounds find wide applications. For Example, chlorine-containing antibiotics, Chloramphenicol and Aureomycin are very effective in the treatment of bacterial infections. Chloroquine is used as a drug in the treatment of malaria.

Our body produces the iodine-containing hormone, thyroxine, the deficiency of which causes a disease called goitre. These are all aryl halides. A majority of alkyl halides are employed as insecticides and there exists virtually no soil microorganism which can degrade them into nontoxic metabolic products

Classification On The Basis Of The Number Of Halogen Atoms

Haloalkanes or haloarenes are classified as mono-, di- or tri-halogen compounds depending on whether they contain one, two or three halogen atoms. For Example,

Monohalo compounds are further classified on the basis of the hybridisation of the carbon atom to which the halogen atom is attached.

Compounds contalnting an sp³ C—X bond (X = F, Cl, Br, I)

1. Alkyl halides or haloalkanes (R—X) Alkyl halides other than methyl halides are classified as primary (1°), secondary (2°) or tertiary (3°), depending upon the number of carbons attached to the halogen-bearing carbon.

2. Allyl halides In an allyl halide, the halogen atom is attached to that sp -sp-hybridised carbon which is bonded to the carbon forming a carbon-carbon double bond.

3. Benzylic halide In a benzylic halide, the halogen atom is attached to that sp³ -hybridised carbon which is bonded to a benzene ring.

Compounds containing an sp² C—X bond

1. Vinylic halides In a vinylic halide, the halogen atom is bonded to the sp2-hybridised carbon atom that forms a carbon-carbon double bond.

2. Aryl halides In an aryl halide, the halogen atom is attached to the sp2-hybridised carbon atom of an aromatic ring.

Dihalo compounds (dihalides) are classified as vicinal (or vie) dihalides and geminal (or gem) dihalides. Molecules with two halogen atoms connected to the same carbon are called geminal (or gem) dihalides. Molecules with two halogen atoms connected to neighbouring carbons, as in the case of 1, 2-dibromoethane, are called vicinal (or vie) dihalides.

Nomenclature

The common names of alkyl halides are simply the alkyl derivatives of the corresponding hydrogen halides. he /UPAC names are halo derivatives of the corresponding hydrocarbons. In the common names, the prefixes n-, •c- (s-) and text- (/-) indicate normal, secondary and tertiary respectively.

Example Write the IUPAC names of the following.

2-Bromobutane

1-Iodocyclohexene

3-Bromo-2-chloro-4-iodopentane

1-Bromo-2-fluoro-4-iodo-3-methyl butane

4-Bromo-2, 2, 6, 6-tetramethyl-4-heptene-3-one

3-Iodo-4-nitrobenzaldehyde

Example Write the structures of all the possible isomers of the monochloro derivative of methylcyclohexane (excluding stereo structures).

Solution

Methods Of Preparation From alkanes

Cyclic and acyclic hydrocarbons can be chlorinated or brominated in the presence of visible or ultraviolet light. For Example, methane on chlorination yields a mixture of methyl chloride, methylene chloride, chloroform and carbon tetrachloride.

⇒ \(\mathrm{CH}_4+\mathrm{Cl}_2 \stackrel{\text { hv }}{\longrightarrow} \underset{\text { Methyl chloride }}{\mathrm{CH}_3 \mathrm{Cl}}+\mathrm{HCl}\)

⇒ \(\mathrm{CH}_3 \mathrm{Cl}+\mathrm{Cl}_2 \stackrel{\text { hv }}{\longrightarrow} \underset{\text { Methylene chloride }}{\mathrm{CH}_2 \mathrm{Cl}_2}+\mathrm{HCl}\)

⇒ \(\mathrm{CH}_2 \mathrm{Cl}_2+\mathrm{Cl}_2 \longrightarrow \underset{\begin{array}{c}
\text { Trichloromethane } \\\text { (chloroform) }\end{array}}{\mathrm{CHCl}_3}+\mathrm{HCl}\)

⇒ \(\mathrm{CHCl}_3+\mathrm{Cl}_2 \longrightarrow \underset{\begin{array}{c}\text { Tetrachloromethane } \\\text { (carbon tetrachloride) }\end{array}}{\mathrm{CCl}_4}+\mathrm{HCl}\)

Haloalkanes And Haloarenes Class 12 Notes

Mechanism

⇒ \(\mathrm{Cl}: \mathrm{Cl} \stackrel{\text { heat }}{\longrightarrow} \mathrm{Cl}^{\circ}+\mathrm{Cl}^{\cdot}\)
Chlorine-free radical

⇒ \(\mathrm{Cl}^{\cdot}+\mathrm{CH}_4 \longrightarrow \mathrm{HCl}+\underset{\begin{array}{c} \text { Methyl free } \\ \text { radical } \end{array}}{\mathrm{CH}_3}\)

⇒ \(\cdot \mathrm{CH}_3+\mathrm{Cl}: \mathrm{Cl} \longrightarrow \underset{\text { Methyl chloride }}{\mathrm{CH}_3 \mathrm{Cl}}+{ }^{\circ} \mathrm{Cl}\)

⇒ \(\mathrm{CH}_3 \mathrm{Cl}+\cdot \mathrm{Cl} \longrightarrow \underset{{\text { Chloromethyl } \\ \text { free radical }}}{\cdot \mathrm{CH}_2 \mathrm{Cl}+\mathrm{HCl}}\)

⇒ \(\cdot \mathrm{CH}_2 \mathrm{Cl}+\mathrm{Cl}: \mathrm{Cl} \longrightarrow \underset{\begin{array}{c} \text { Methylene } \\ \text { chloride } \end{array}}{\mathrm{CH}_2 \mathrm{Cl}_2+\mathrm{Cl}}\)

⇒ \(\mathrm{CH}_2 \mathrm{Cl}_2+\cdot \mathrm{Cl} \longrightarrow \underset{{\text { Dichloromethyl } \\ \text { free radical }}}{\cdot \mathrm{CHCl}_2}+\mathrm{HCl}\)

⇒ \(\cdot \mathrm{CHCl}_2+\mathrm{Cl}: \mathrm{Cl} \longrightarrow \underset{\text { Chloroform }}{\mathrm{CHCl}_3+{ }^{\circ} \mathrm{Cl}}\)

⇒ \(\mathrm{CHCl}_3+\mathrm{Cl} \longrightarrow \underset{{\text { Trichloromethyl } \\ \text { free radical }}}{\cdot \mathrm{CCl}_3}+\mathrm{HCl}\)

⇒ \(\cdot \mathrm{CCl}_3+\mathrm{Cl}: \mathrm{Cl} \longrightarrow \underset{{\text { Carbon } \\ \text { tetrachloride }}}{\mathrm{CCl}_4}+\cdot \mathrm{Cl}\)

The ease of substitution at various carbon atoms is tertiary > secondary > primary, which is the same as the stability of the various alkyl radicals. For Example, the chlorination of propane yields a mixture of 1-chloropropane and 2-chloropropane.

The bromination of an alkane is similar to chlorination except that the rate of the reaction is slow. The bromination of isobutane gives tert-butyl bromide as the major product.

Fluorohydrocarbons cannot be prepared by the direct fluorination of alkanes, because the reaction is explosive. Fluorine can, however, be introduced by the replacement of chlorine from an alkyl chloride using mercurous fluoride.

2CH₃CI+ Hg₂F₂ →2CH₃F+ Hg₂Cl₂

The iodination of an alkane is reversible and leads to the regeneration of the alkane.

Haloalkanes And Haloarenes Class 12 Notes

⇒ \(\mathrm{R}-\mathrm{H}+\mathrm{I}_2 \rightleftharpoons \mathrm{R}-\mathrm{I}+\mathrm{HI}\)

The reaction is, therefore, not synthetically useful and direct iodination is usually achieved by using a strong oxidising agent (HIO₃ or HNO₃ ) to destroy the HI formed in the process.

HIO₃ + 5HI→3I₂ + 3H₂O

From Alkenes

1. Addition of HX

The addition of hydrogen halides to alkenes gives a variety of alkyl halides. HF, HBr and HI can be added at room temperature, while HCI is added at a higher temperature.

In the case of unsymmetrical olefins, the reaction proceeds in accordance with the Markovnikov rule.

Mechanism The hydrogen halides add on to the olefins by an ionic mechanism.

In the presence of light or peroxide, the addition of hydrogen bromide follows a free-radical mechanism the product is formed according to the anti-Markovnikov rule.

Mechanism

⇒ \(\mathrm{CH}_3-\mathrm{CH}=\mathrm{CH}_2+\mathrm{HBr} \stackrel{\text { peroxide }}{\longrightarrow} \mathrm{CH}_3-\mathrm{CH}_2-\mathrm{CH}_2-\mathrm{Br}\)

⇒ \(\mathrm{HBr} \stackrel{\text { peroxide }}{\longrightarrow} \dot{\mathrm{H}}+\dot{\mathrm{Br}}\)

⇒ \(\mathrm{CH}_3-\mathrm{CH}=\mathrm{CH}_2+\dot{\mathrm{Br}} \longrightarrow \mathrm{CH}_3-\dot{\mathrm{C}} \mathrm{H}-\mathrm{CH}_2-\mathrm{Br}\)

⇒ \(\mathrm{CH}_3-\dot{\mathrm{C}} \mathrm{H}-\mathrm{CH}_2-\mathrm{Br}+\mathrm{HBr} \longrightarrow \mathrm{CH}_3-\mathrm{CH}_2-\mathrm{CH}_2-\mathrm{Br}+\dot{\mathrm{Br}}\)

In contrast to the addition of HBr, the free radical addition of HI and HC1 to olefins normally does not take place.

⇒ \(\dot{\mathrm{X}}+\mathrm{H}_2 \mathrm{C}=\mathrm{CH}_2 \longrightarrow \mathrm{XCH}_2 \dot{\mathrm{C}} \mathrm{H}_2 \quad\left[\begin{array}{ll}
\mathrm{X}=\mathrm{Cl} ; & \Delta H=-26 \mathrm{kcal} \mathrm{mol}^{-1} \\
\mathrm{X}=\mathrm{Br} ; & \Delta H=-5 \mathrm{kcal} \mathrm{mol}^{-1} \\
\mathrm{X}=\mathrm{I} ; & \Delta H=+7 \mathrm{kcal} \mathrm{mol}^{-1}
\end{array}\right]\)

⇒ \(\mathrm{XCH}_2 \dot{\mathrm{C}} \mathrm{H}_2+\mathrm{HX} \longrightarrow \mathrm{XCH}_2 \mathrm{CH}_3+\dot{\mathrm{X}}\left[\begin{array}{ll}
\mathrm{X}=\mathrm{Cl} ; & \Delta H=+5 \mathrm{kcal} \mathrm{mol}^{-1} \\
\mathrm{X}=\mathrm{Br} ; & \Delta H=-11 \mathrm{kcal} \mathrm{mol}^{-1} \\
\mathrm{X}=\mathrm{I} ; & \Delta H=-27 \mathrm{kcal} \mathrm{mol}^{-1}
\end{array}\right]\)

For the addition of HCl, the first step is exothermic and so the addition of the chlorine atom takes place readily but the second step is endothermic so this step is not possible. The initially formed radical is unable to decompose HCI.

For the addition of HBr, both the steps are exothermic and so the addition takes place readily. For the free radical addition of HI, the first step is endothermic. Therefore, no HI will add on to an olefin.

Haloalkanes And Haloarenes Class 12 Notes

2. Addition of halogens

Dihalides are formed by the addition of halogens to alkenes.

From Alcohols

Alcohols are important starting materials for the preparation of alkyl halides. Various reagents such as PCI₅, PBr3, PI3 and SOCl₂ readily displace the alcoholic group to yield the corresponding halides.

⇒ \(\mathrm{CH}_3-\mathrm{CH}_2-\mathrm{OH}+\mathrm{PCl}_5 \longrightarrow \underset{\text { Ethyl chloride }}{\mathrm{CH}_3-\mathrm{CH}_2-\mathrm{Cl}+\mathrm{POCl}_3+\mathrm{HCl}}\)

⇒ \(3 \mathrm{CH}_3-\mathrm{CH}_2-\mathrm{OH}+\mathrm{PI}_3 \longrightarrow \underset{\text { Ethyl iodide }}{3 \mathrm{CH}_3-\mathrm{CH}_2-\mathrm{I}+\mathrm{H}_3 \mathrm{PO}_3}\)

⇒ \(\mathrm{C}_2 \mathrm{H}_5 \mathrm{OH}+\mathrm{SOCl}_2 \longrightarrow \underset{\text { Ethyl chloride }}{\mathrm{C}_2 \mathrm{H}_5 \mathrm{Cl}}+\mathrm{SO}_2+\mathrm{HCl}\)

The three hydrohalogen acids HI, HBr and HC1 also react with alcohols but at different rates i.e., HI > HBr > HCI

⇒ \(\mathrm{CH}_3-\mathrm{CH}_2-\mathrm{OH}+\mathrm{HI} \stackrel{\Delta}{\longrightarrow} \mathrm{CH}_3 \mathrm{CH}_2-\mathrm{I}+\mathrm{H}_2 \mathrm{O}\)

With primary alcohols, hydroiodic acid reacts most readily while hydrochloric acid requires zinc chloride as a catalyst and hydrobromic acid displays an intermediate reactivity

⇒ \(\underset{\text { Ethyl alcohol }}{\mathrm{C}_2 \mathrm{H}_5 \mathrm{OH}}+\mathrm{HCl} \stackrel{\text { anhydrous } \mathrm{ZnCl}_2}{\longrightarrow} \underset{\text { Ethyl chloride }}{\mathrm{C}_2 \mathrm{H}_5 \mathrm{Cl}}+\mathrm{H}_2 \mathrm{O}\)

Mechanism

Furthermore, a tertiary alcohol reacts more rapidly than a secondary alcohol, which in turn reacts faster than a primary alcohol.

The Halogen-Exchange Method

This method is used to prepare alkyl iodides which are otherwise not easily obtained. It involves treating an alkyl chloride or bromide with a solution of sodium iodide.

⇒ \(\underset{(\mathrm{X}=\mathrm{Cl}, \mathrm{Br})}{\mathrm{R})}+\mathrm{X}+\mathrm{NaI} \stackrel{\text { acetone }}{\longrightarrow} \underset{\text { Alkyl iodide }}{\mathrm{R} \mathrm{I}}+\mathrm{NaX}\)

This reaction is called the Finkelstein reaction.

An alkyl fluoride is best prepared by heating an alkyl chloride or alkyl bromide in the presence of mercurous fluoride.

2CH₃ —C1+ Hg₂F_2→2CH₃—F+Hg₂Cl₂

This reaction is known as Swart’s reaction.

By The Hunsdiecker Reaction

Alkyl halides and aryl halides can be prepared by decomposing the silver salt of carboxylic acid using chlorine or bromine in the presence of CC1₄.

⇒ \(\underset{\text { iilver carboxylate }}{\mathrm{RCOOAg}}+\mathrm{Br}_2 \underset{\Delta}{\stackrel{\mathrm{CCl}_4}{\longrightarrow}} \mathrm{R}-\mathrm{Br}+\mathrm{CO}_2+\mathrm{AgBr}\)

The reaction probably proceeds through a free-radical mechanism as shown below.

Physical Properties

Melting and boiling points

The melting and boiling points of alkyl halides are higher than those of alkanes with the same number of carbon atoms. This is because of their greater dipole-dipole and van der Waals attractions.

In alkyl halides having the same alkyl groups, the boiling and melting points increase with the increase in the atomic weight of the halogen. Thus, the boiling point of methyl fluoride is the lowest while that of methyl iodide is the highest.

In the case of isomeric alkyl halides, the boiling point depends upon surface area. The boiling point of the primary halide is the highest and that of the tertiary halide is the lowest.

Specific gravity

Most halides are liquids. The bromides, iodides and polyhalides, in general, have a specific gravity greater than and are, therefore, heavier than water.

Solubility

Alkyl halides are insoluble in water because of their inability to form hydrogen bonds with water.

Haloalkanes And Haloarenes Class 12 Notes

Chemical Reactions

Alkyl halides mainly undergo the following two types of chemical reactions.

1. Nucleophilic substitution reactions
2. Elimination reactions

In addition, alkyl halides undergo some other characteristic reactions which we will discuss subsequently.

Nucleophilic Substitution Reactions

All the halogens are significantly more electronegative than carbon. As a result, a carbon-halogen bond is polar. The majority of reactions that alkyl halides undergo involve heterolytic cleavage of the carbon-halogen bond, with the halogen departing as X-.

In such a case, the halogen is referred to as a departing nucleophile. The order of reactivity of alkyl halides is

R—I>R—Br>R—Cl>>R—F

The strengths of the C—X bonds and dipole moments have been measured and are listed.

It is clearly the easiest to break a C—I bond and the most difficult to break a C—F bond. Iodine is the leaving group and fluorine is a very bad one with the other halogens in between.

In this chapter, we will confine our attention to reactions in which the halogen is attached to a saturated (i.e., sp³ -hybridised) carbon.

The main nucleophilic substitution reactions that alkyl halides (in particular ethyl bromide) undergo are as follows.

1. Hydrolysis The reaction of ethyl bromide with aqueous NaOH gives ethyl alcohol.

C₂H₅Br + NaOH -> C₂H₅OH + NaBr

Mechanism

Solvolysis If the attacking nucleophile is the solvent, the reaction is called a solvolysis reaction.

2. Reaction with potassium hydrosulphide Ethyl bromide on treatment with aqueous alcoholic KSH yields alcohol.

3. Reaction with sodium alkoxide On treatment with sodium ethoxide, ethyl bromide gives diethyl ether.

⇒ \(\mathrm{C}_2 \mathrm{H}_5 \mathrm{Br}+\mathrm{C}_2 \mathrm{H}_5 \text { ŌNa } \rightarrow \mathrm{C}_2 \mathrm{H}_5-\mathrm{O}-\mathrm{C}_2 \mathrm{H}_5+\mathrm{NaBr}Diethyl ether\)

4. Reaction with ammonia On being heated with concentrated ammonia in a sealed tube, ethyl bromide gives a mixture of amines.

⇒ \(\mathrm{C}_2 \mathrm{H}_5-\mathrm{Br}+\mathrm{NH}_3 \rightarrow \underset{\text { Ethylamine }}{\mathrm{C}_2 \mathrm{H}_5-\mathrm{NH}_2}+\mathrm{HBr}\)

⇒ \(\mathrm{C}_2 \mathrm{H}_5-\mathrm{Br}+\mathrm{C}_2 \mathrm{H}_5-\mathrm{NH}_2 \rightarrow \underset{\text { Diethylamine }}{\left(\mathrm{C}_2 \mathrm{H}_5\right)_2 \mathrm{NH}}+\mathrm{HBr}\)

⇒ \(\mathrm{C}_2 \mathrm{H}_5 \mathrm{Br}+\left(\mathrm{C}_2 \mathrm{H}_5\right)_2 \mathrm{NH} \rightarrow \underset{\text { Triethylamine }}{\left(\mathrm{C}_2 \mathrm{H}_5\right)_3 \mathrm{~N}}+\mathrm{HBr}\)

⇒ \(\mathrm{C}_2 \mathrm{H}_5-\mathrm{Br}+\left(\mathrm{C}_2 \mathrm{H}_5\right)_3 \mathrm{~N} \rightarrow \underset{{\text { Tetraethylammonium } \\ \text { bromide }}}{\left(\mathrm{C}_2 \mathrm{H}_5\right)_4 \mathrm{~N}^{+} \mathrm{Br}^{-}}\)

Ammonia is a nucleophilic reagent.

Haloalkanes And Haloarenes Class 12 Notes

Similarly ethylamine, diethylamine and triethylamine act as nucleophilic reagents.

5. Reaction with silver acetate On being heated with the silver salt of carboxylic acid, ethyl bromide yields an ester.

⇒ \(\mathrm{C}_2 \mathrm{H}_5-\mathrm{Br}+\underset{\text { Silver acetate }}{\mathrm{CH}_3-\mathrm{CO} A \stackrel{+}{\rightarrow}} \rightarrow \underset{\text { Ethyl acetate }}{\mathrm{C}_2 \mathrm{H}_5-\mathrm{O}-\mathrm{CO}}-\mathrm{CH}_3+\mathrm{AgBr}\)

In this reaction, the acetate anion is the nucleophilic reagent.

6. Substitution by the cyanide (nitrile) anion On being heated with an alcoholic solution of KCN, ethyl bromide gives ethyl cyanide (ethyl nitrile).

⇒ \(\mathrm{C}_2 \mathrm{H}_5 \mathrm{Br}+\mathrm{AgCN} \longrightarrow \underset{\text { Ethyl isocyanide }}{\mathrm{C}_2 \mathrm{H}_5-\stackrel{+}{\mathrm{N}} \equiv \mathrm{C}+\mathrm{AgBr}}\)

⇒ \(\mathrm{C}_2 \mathrm{H}_5-\mathrm{Br}+\mathrm{KCN} \longrightarrow \underset{\text { Ethyl cyanide }}{\mathrm{C}_2 \mathrm{H}_5-\mathrm{CN}+\mathrm{KBr}}\)

KCN is ionic and provides predominantly cyanide ions in solution. The carbon and nitrogen atoms of the CN: ion are in a position to donate electron pairs. However, the attack takes place mainly through carbon and not through nitrogen because the resulting C—C bond in alkyl cyanide (C—CN) is stronger than the C N bond.

Upon reacting with silver cyanide (AgCN), ethyl bromide gives ethyl isocyanide.

⇒ \(\mathrm{C}_2 \mathrm{H}_5 \mathrm{Br}+\mathrm{AgCN} \longrightarrow \underset{\text { Ethyl isocyanide }}{\mathrm{C}_2 \mathrm{H}_5-\stackrel{+}{\mathrm{N}} \equiv \mathrm{C}+\mathrm{AgBr}}\)

Silver cyanide, AgCN, is essentially covalent. Therefore, the lone pair on the nitrogen is mainly available for covalent bond formation, resulting predominantly in the formation of isocyanides.

7. Reaction with silver nitrite On being heated with silver nitrite, an alcoholic solution of alkyl halide gives a mixture of nitroalkane and alkyl nitrite. These are separated by fractional distillation.

Mechanism The nitrite anion \((\mathrm{O}=\ddot{\mathrm{N}}-\overline{\mathrm{O}})\) is an ambident nucleophile with two different points of linkage. The linkage through oxygen results in the formation of alkyl nitrites while that through the nitrogen atom leads to the formation of nitroalkanes.

8. Reaction with sodium acetylide On treatment with sodium acetylide, ethyl bromide yields a higher alkyne.

For Example,

⇒ \(\mathrm{C}_2 \mathrm{H}_5 \mathrm{Br}+\mathrm{Na}-\mathrm{C} \equiv \mathrm{CH} \rightarrow \mathrm{C}_2 \mathrm{H}_5-\mathrm{C} \equiv \mathrm{CH}+\mathrm{NaBr}1-Butyne\)

Sodamide (NaNH₂ ) reacts with acetylene to give sodium acetylide.

⇒ \(\mathrm{H}-\mathrm{C} \equiv \mathrm{C}-\mathrm{H}+\stackrel{+}{\mathrm{N}} \mathrm{a}-\mathrm{NH}_2 \rightarrow \mathrm{H}-\mathrm{C} \equiv \overline{\mathrm{C}}-\stackrel{+}{\mathrm{N}} \mathrm{a}+\mathrm{NH}_3\)

Mechanisms of nucleophilic substitution reactions

A nucleophilic substitution reaction may be described as one in which a substituent is replaced by a nucleophile

The following general equation represents nucleophilic substitution

The above type of traction can take place in the following two ways.

Possibility 1 The leaving group is ejected first and then the nucleophile comes in.

Possibility 2 The leaving group is ejected at the same time as the nucleophile is attached.

Let us now examine each of the possibilities separately.

According to possibility 1, the leaving group goes out with the bonded pair of electrons, leaving a carbonium ion in the first step. In the second or final step the nucleophile, with its lone pair of electrons, interacts with the carbonium ion to form the product.

1. First step:

2. Second step:

The rate of the overall reaction in a multistep process is dependent on the slow step. In the example cited above, the slow step (rate-determining step) is the first step

1. Requiring more energy than the second step.
2. Which creates a charged species out of a neutral molecule.

Rate of overall reaction α rate of slow step reaction (1)

α[substrate]

The rate of the reaction is dependent on the concentration of the substrate molecule only and is independent of the concentration of the nucleophile because the nucleophile becomes involved only after the rate-determining step. Thus the reaction is unimolecular and is termed S_N1—S stands for substitution, N for nucleophilic and 1 indicates that the reaction is unimolecular.

Example s of S_N1 reactions Tertiary halides usually undergo S_N1 displacements. For Example, f-butyl bromide is hydrolysed to f-butyl alcohol through an unimolecular process.

⇒ \(\left(\mathrm{CH}_3\right)_3 \mathrm{C}-\mathrm{Br} \stackrel{\mathrm{O}}{\longrightarrow}\left(\mathrm{CH}_3\right)_3 \mathrm{C}-\mathrm{OH}+\stackrel{\ominus}{\mathrm{Br}}\)

The rate of the reaction depends only on the concentration of the alkyl halide and is independent of the concentration of the added nucleophile. The reaction occurs according to the following rate equation:

Rate=K₁[(CH₃)₃C—Br]

The mechanism of the reaction involves two steps. In the first step, the halide ionises to form an intermediate carbonium ion, which then readily combines with the nucleophile to form the product.

It might seem that the addition of a nucleophile, i.e., an \(\mathrm{O} \stackrel{\ominus}{\mathrm{H}}\) ion, would affect the rate of the reaction in Step 2. But this step is so much faster than Step 1 that the effective rate measured is that of Step 1, i.e., the ionisation step. Tertiary alkyl halides undergo S_N1 reactions very fast because of the high stability of tertiary carbonium ions. The order of reactivity of alkyl halides in the case of S_N1 reactions is tertiary halide > secondary halide > primary halide.

Allylic and benzylic halides undergo S_N1 reactions in which the intermediate carbonium ion is stabilised by resonance.

There are some nucleophilic substitution reactions in which the substrate (R L) as well as the nucleophile \(: \stackrel{\ominus}{N u}\) take part in the rate-determining step. In such a case, the nucleophile attacks the carbon atoms from the side opposite to that of the leaving group and forms a pentavalent transition state which finally collapses to give the product (R—Nu) and the leaving group \(\text { (: } \stackrel{\ominus}{\mathrm{L}})\)

⇒ \(: \stackrel{\ominus}{N u}+R-L \frac{\text { slow }}{\text { (a) }}[\mathrm{Nu}–\mathrm{R}–\mathrm{L}] \underset{\text { (b) }}{\stackrel{\text { fast }}{\rightarrow}} \mathrm{Nu}-\mathrm{R}-: \stackrel{\ominus}{\mathrm{L}}\)

The rate of the reaction depends upon the concentration of the substrate (R—L) as well as that of the nucleophile \(\text { (: } \stackrel{\ominus}{\mathrm{Nu}})\) both the reactants being involved in the rate-determining step (a).

Rate of reaction α rate of slow step (a)
⇒ \(\propto[R-L][\stackrel{\ominus}{N u}]\)

Since the rate of the reaction is dependent on the concentration of the substrate and that of the nucleophile, the reaction is bimolecular and is termed an S_N2 reaction, 2 indicating that the reaction is bimolecular.

Example s of S_N2 reactions Primary halides usually undergo S_N2 reactions. For Example, methyl bromide is hydrolysed to methyl alcohol upon reacting with an aqueous solution of sodium hydroxide. The hydroxide ion displaces the bromide ion.

[The solid wedge represents the bond above the plane of the paper, the dashed line below the plane of the paper and a straight line represents a bond on the plane of the paper.]

The hydroxide ion attacks the carbon atom bearing the leaving group from behind and the bromide ion

departs from the opposite side. In the transition state, the \(\mathrm{O} \stackrel{\ominus}{\mathrm{H}}\) ion and the \(\stackrel{\ominus}{\mathrm{Br}}\) ion is only partially bonded to the central carbon atom. Finally, a new covalent bond is formed between the hydroxide ion and the carbon atom with the loss of the bromide ion.

Methyl halides undergo S_N2 reactions most rapidly because there are only three small hydrogen atoms Tertiary halides are the least reactive because bulky alkyl groups slow down the approaching nucleophiles. The order of reactivity of alkyl halides in the case of S_N2 reactions is primary halide > secondary halide > tertiary halide.

Stereochemical aspect of nucleophilic substitution reactions

Before discussing the stereochemical aspect of nucleophilic substitution reactions in detail, it is useful to discuss some terms used in optical isomerism such as optical activity, chirality, retention of configuration, inversion of configuration and racemisation.

Plane polarised light and optical activity An ordinary beam of light vibrates in all directions. A plane of polarised light is one which vibrates entirely in one direction. An ordinary beam of light can be converted to a plane polarised light by passing it through a specially constructed prism, called Nicol prism. Certain compounds may rotate the plane polarised light either to the left (anticlockwise) or to the right (clockwise).

Compounds which rotate plane polarised light to the left are said to be laevorotatory or of the 1-form. The degree of rotation of light in such cases is prefixed by a – (minus) sign. If the light is rotated to the right, the compound is said to be dextrorotatory or of the d-form. The degree of rotation of light in such cases is prefixed by a + (plus) sign.

The el¬ and /- isomers of a compound are known as optical isomers and the phenomenon is termed optical isomerism. The optical activity of a compound is identified and estimated by an instrument known as a polarimeter.

Enantiomers, diastereomers and chirality Enantiomers are structures that are not identical but are mirror images of each other.

Stereoisomers that are not mirror images of one another are called diastereoisomers. The term diastereoisomer is sometimes shortened to diastereomer.

1 and 3 form a pair of enantiomers. I and 2, and II and IE, form a pair of diastereomers. Structures are said to be chiral if they cannot be superimposed on their mirror images. A hand cannot be superimposed on its mirror image.

If you hold your left hand up to a mirror, the image looks like a right hand. When we try to superimpose our hands on each other, we simply cannot do so. A hand and a foot are chiral objects whereas a ball and a glass are achiral objects. Achiral structures are superimposable in their mirror images.

A carbon atom carrying four different groups is a chiral carbon. Such a carbon is sometimes also referred to as an asymmetric carbon atom. For Example,

The asymmetric carbon atoms are indicated by a star (*).

Examples of achiral carbon-containing compounds:

Inversion, retention and racemisation To understand these terms consider the following general reaction.

If (2) is obtained the process is called retention of configuration because (1) and (2) have a similar configuration. If (3) is obtained the process is known as inversion of configuration because the configuration of 3 is inverted with respect to (1).

If a 50: 50 mixture of (2) and (3) is obtained then the process is called racemisation. The resulting racemic product is optically inactive as one isomer will rotate plane polarised light in the direction opposite to that in which the other isomer will.

Now let us consider the stereochemical aspect of S_N1 and S_N2 reactions taking the Example of optically active alkyl halides.

In the case of optically active alkyl halides, an S_N1 reaction involves racemisation, i.e., an equimolecular mixture of the d- and l- isomers is produced. This is possible because the carbonium ion formed as an intermediate in

The rate-determining step has an equal chance of being attacked from both sides, giving one isomer having the same configuration and the other having the opposite configuration. This process may be illustrated by citing the Example of the hydrolysis of I-phenyl ethyl bromide.

However, the S_N2 reaction of an optically active compound proceeds with a complete inversion of configuration. This is because the nucleophile will be linked to the opposite side with respect to the leaving group (halide ion) in an optically active alkyl halide.

This process may be illustrated by citing the Example of the hydrolysis of 2-bromooctane. The hydrolysis of the optically active 2-bromooctane with sodium hydroxide gives the optically active 2-octanol with an inverted configuration.

Such reactions are generally accompanied by a change in optical rotation and its direction as well. It is termed as Walden inversion or flipping. The reaction may be likened to an umbrella flipping in a high wind, the attacking nucleophile being analogous to a high wind.

S_N2 vs S_N1

The two mechanisms termed S_N1 and S_N2 are the extremes.

The solvent, substrate structure, nucleophile and, to a lesser extent, the leaving group and reaction temperature, nil contribute to determining the mechanism. We shall briefly examine the effect of some of these parameters.

Solvent A reaction is favoured by an ionising solvent whose dielectric constant is high. Ionising solvents accelerate the rate of an S_N1 reaction while that of an S_N2 reaction is relatively solvent-independent. Substrate structure In S_N2 reactions the order of reactivity of various alkyl halides follows the sequence

CH3 —X > primary > secondary > tertiary

In contrast to S_N2 reactions, the sequence of reactivity of alkyl halides follows the following order in S_N1 reactions.

Tertiary > secondary > primary > CH₃ —X.

Nucleophile While the nature of the nucleophile does not alter the rate of an S_N1 reaction, that of an S_N2 reaction is altered by the nucleophile. The high concentration of a good nucleophile will favour an S_N2 reaction.

Elimination Reactions

A reaction involving the loss of two atoms or groups from a molecule, without there being any substitution b) other atoms or groups, is known as an elimination reaction. This is the reverse of an addition reaction. Most commonly a proton is lost from one carbon whereas a nucleophile is lost from the adjacent carbon; these tw< carbon atoms are usually referred to as |i- and a-carbons respectively.

The reaction is known as a reaction or 1, 2-elimination reaction.

Orientation in an elimination reaction

Sometimes elimination reactions lead to mixtures of alkene products; usually, one alkene is formed as a major product and the other as a minor product. There are two empirical rules governing orientation in those reactions.

1. Saytzev rule (Zaitsev rule) states that an alkyl halide capable of forming a double bond in either direction of the chain preferably yields that alkene in which there are a greater number of alkyl groups attached to the double bond. That is, the more substituted olefin is formed as the major product. For Example, 2-bromobutane gives 2-butene as the major product.

The stability of alkenes follows the order

R₂C=CR₂>R₂C=CHR>RCH=CHR>RCH=CH₂>CH₂=CH₂

The greater stability of the more-substituted alkene is explained on the basis of hyperconjugation.

The number of resonating structures increases in more-substituted alkenes and consequently increases stability.

2. The Hofmann rule states that quaternary ammonium salts yield predominantly the least-substituted alkene.

In this case, the hydroxide ion abstracts a proton from the β-carbon, resulting in the formation of a double bond and expulsion of a tertiary amine.

If a quaternary ammonium salt has ethyl as well as propyl groups attached to the positively charged nitrogen the propyl group shows relatively little tendency to lose a β-hydrogen.

Let us now see how the nature of the nucleophile determines whether a reaction will be a substitution read or an elimination reaction.

Basicity

The more basic the nucleophile the more likely is that elimination will replace substitution as the main reaction of an alkyl halide.

For Example,

Size

One of the best bulkier bases for promoting elimination and avoiding substitution is potassium f-butoxide. The large alkyl substituent makes it hard for the negatively charged oxygen to attack carbon in a substitution reaction but it has no problem attacking hydrogen.

Small nucleophile—substitution

Large nucleophile—elimination

Example Among the following, which alkyl halide would you expect to react more rapidly by the SN2 mechanist Explain.

Example Which of the following SN1 reactions would you expect to take place more rapidly? Explain.

Solution (1), is because water is more polar than CH₃OH.

Example Suggest a reagent (or reagents), and state the type of reaction and mechanism involved for each of the following transformations.

1. CH₃CH₂— Br -> CH₃CH₂N(CH₃)₃

2. CH₃CH₂Br -> CH₃CH₂CN

Example Write the structure of the expected product. Indicate the type of reaction and the mechanism involved.

Solution

Temperature

Temperature has an important role to play in deciding whether a reaction is an elimination or a substitution reaction. Elimination is favoured at high temperatures.

Reaction With Metals

1. Alkyl halides react with metallic sodium in dry ether to form alkanes containing an even number of carbon atoms.

⇒ \(2 \mathrm{C}_2 \mathrm{H}_5-\mathrm{I}+2 \mathrm{Na} \longrightarrow \underset{\text { Butane }}{\mathrm{CH}_3-\mathrm{CH}_2-\mathrm{CH}_3+2 \mathrm{NaI}}\)

This reaction is known as the Wurtz reaction.

Mechanism The reaction is believed to proceed through a nucleophilic attack by the alkyl ion on the alkyl halide.

2. Alkyl halides react with metallic zinc in dry ether to form an alkane containing an even number of carbon atoms.

⇒ \(2 \mathrm{CH}_3 \mathrm{CH}_2 \mathrm{I}+\mathrm{Zn} \longrightarrow \mathrm{CH}_3 \mathrm{CH}_2 \mathrm{CH}_2 \mathrm{CH}_3+\mathrm{ZnI}_2\)

This is called the Frankland reaction.

Mechanism

3. When alkyl or aryl halides are refluxed with magnesium in dry ether, alkyl or aryl magnesium halides (R—Mg—X) are formed. They are known as Grignard reagents. They are organometallic compounds. (The metal magnesium is attached to the carbon.)

⇒ \(\mathrm{R}-\mathrm{X}+\mathrm{Mg} \stackrel{\text { dry ether }}{\longrightarrow} \underset{(\mathrm{R}=\text { alkyl, aryl) }}{\mathrm{R}-\mathrm{Mg}-\mathrm{X}}\)

⇒ \(\mathrm{C}_2 \mathrm{H}_5 \mathrm{Br}+\mathrm{Mg} \underset{\text { ether }}{\stackrel{\text { dry }}{\longrightarrow}} \mathrm{C}_2 \mathrm{H}_5 \mathrm{MgBr}\)

⇒ \(\mathrm{C}_6 \mathrm{H}_5 \mathrm{Br}+\mathrm{Mg} \underset{\text { ether }}{\stackrel{\text { dry }}{\longrightarrow}} \mathrm{C}_6 \mathrm{H}_5 \mathrm{MgBr}\)

Grignard reagents form loose covalent bonds with the solvent ether molecules and form solutions in ether.

Since the electronegativities of magnesium and carbon are very different, the Grignard reagents may 1 regarded as polar compounds which are sources of nucleophilic carbanions R.

⇒ \(\stackrel{\delta-}{\mathrm{R}}-\stackrel{\delta_{+}^{+}}{\mathrm{Mg}}-\mathrm{X} \longrightarrow \stackrel{\ominus}{\mathrm{R}}+\stackrel{+}{\mathrm{M} g X}\)

Example Is (CH₃)₃ CO K+ an organometallic compound? If not, why?

Solution

No, the metal is bonded to an oxygen, not a carbon. This is an example of an alkoxide salt.

Haloarenes

Aryl halides (haloarenes) are those compounds in which the halogen atom (F, Cl, Br, I) is attached directly benzene ring. Unlike alkyl halides, these compounds are characterised by their relative inertness towards r common nucleophiles.

Nomenclature

Simple aryl halides are named by prefixing the halo derivative to the word benzene, for example, iodobenzene. The three disubstituted isomers are differentiated by the use of ortho-, meta- and para-.

Numbering, however, is required when more than two halo atoms or other substituents are attached to the ring.

Iodobenzene m-Dichlorobenzene 2-Bromol, 4-dimethylbenzene

Methods Of Preparation

Halogenation

The chlorination or bromination of benzene is carried out at an ordinary temperature in the presence of iron or a Lewis acid (AICI₃ or FeCI₃ ).

From diazonium salts

This is the most general method used to prepare all types of aryl halides. An aromatic amine is first diazotised. The diazonium salt can then be reacted with a metal halide to obtain the aryl halide. Aryl chlorides and bromides are conveniently prepared by the treatment of a freshly prepared solution of diazonium salt with cuprous chloride or cuprous bromide. This is known as the Sandmeyer reaction.

Aryl iodides are, however, obtained by treating the diazonium salt with aqueous potassium iodide.

For the preparation of aryl fluorides, the solution of benediazonium chloride is treated with fluoroboric acid, HBF₄. The solid benzene diazonium fluoroborate, ArN_2⁺BF_4^–, which precipitates out, is filtered, washed and dried. When heated, the fluoroborate decomposes to give the aryl fluoride.

Hunsdlockor reaction

Aryl bromides are obtained by heating the silver salts of aromatic acids with bromine.

 

Physical Properties

Aryl halides are polar. They are insoluble in water. For aryl halides having the same aryl group but different halogen atoms, the melting and boiling points increase with the increase in molecular weight.

For Example, the melting point of iodobenzene is the highest and that of fluorobenzene is the lowest, p-isomers, being symmetrical, have high melting points. They are better packed in the crystal lattice.

Chemical Reactions

Nucleophilic substitution

Aryl halides do not generally undergo nucleophilic substitution reactions due to the following reasons.

1. Resonance effect In an aryl halide, the electron pair on the halogen atom is in conjugation with the benzene ring.

This gives the C—Cl bond a double bond character and thus the bond length decreases. The shorter the bond length, the greater the bond strength and the lower the reactivity. Thus, under ordinary conditions, a nucleophilic attack does not take place.

2. In an alkyl halide, the carbon attached to the halogen is sp³ -sp³-hybridised. However, in an aryl halide, the carbon attached to the halogen is sp² -sp² -hybridised. The sp² -sp² -hybridised carbon with greater s-character is electronegative. Thus, the breaking of the C—X bond becomes difficult.

However, nucleophilic substitution reactions of aryl halides are possible in extreme conditions

1. Substitution by hydroxyl group On being heated with aqueous sodium hydroxide at 613 K and 300 atm, chlorobenzene gives phenol.

2. Substitution by the amino group The reaction of chlorobenzene with ammonia at high temperature (573 K) and high pressure yields aniline.

3. Substitution by the —CN group On being heated in the presence of cuprous cyanide at 473 K, chlorobenzene gives benzonitrile (phenyl cyanide).

The above reactions 1-3 take place easily when electron-attracting groups such as —NO₂, — C = N and —SO₃ H are present at the ortho- and para-positions. Note the condition of the reactions.

The presence of an electron-withdrawing group at the 0- and p- positions increases the rate. The presence of an electron-attracting group on any position of the ring does not increase the rate equally. No effect on the reactivity of an aryl halide is observed by the presence of an electron-withdrawing group at the m- position.

The mechanism is as follows.

The presence of a nitro group at the 0- and p- positions withdraws the electron cloud from the benzene ring. As a result, the attack of the nucleophile on a halobenzene becomes easier. The resulting carbanion thus formed gets stabilised through resonance.

The negative charge appearing at the 0- and p- positions (see structure X) with respect to the halogen substituent is also stabilised by the nitro group through resonance. In the case of m-nitrobenzene, none of the resonating structures bears a negative charge on the carbon atom bearing the nitro group.

The presence of the nitro group at the m- position therefore does not stabilise the negative charge and hence they do not undergo nucleophilic substitution easily.

Electrophilic substitution

The benzene ring of an aryl halide undergoes the usual electrophilic substitution reactions such as halogenation, nitration, sulphonation and Friedel-Crafts reactions. In an aryl halide, the halogen atom deactivates the ring due to its more electronegative character but at the same time, it is 0- and p- directing as is evident from the resonating structures of aryl halides.

Due to resonance, the electron density increases more at the 0- and p- positions than at the m- position. Tt electrophilic substitution reactions in aryl halides require slightly more extreme conditions as compared to those in benzene.

1. Halogenation

2. Nitration

3. Sulphonation

4. Friedel-Crafts reaction

Example Chlorine is an electron-withdrawing group. Yet it is 0- and p- directing in electrophilic substitution reactions. Why?

Solution:

Chlorine withdraws electrons through the -I-effect and releases electrons through resonance. In electrophilic reactions, chlorine cannot stabilise the following carbocation through the -I-effect.

Through resonance, chlorine in chlorobenzenes makes the ortho- and para- para–intermediates more stable.

Reactions with metals

1. Wurtz-Fittig reaction On being heated with an alkyl halide and metallic sodium in dry ether, an aryl halide yields alkyl benzene. This reaction is called the Wurtz-Fittig reaction.

2. Fittig reaction Two molecules of an aryl halide, on being heated with sodium in the presence of dry ether, give diphenyl. This reaction is called the Fittig reaction.

Polyhalogen Compounds

Dichloromethane (methylene chloride)

Dichloromethane is widely used as a

1. Solvent.
2. Paint Remover.
3. Propellent In Aerosols.

A high concentration of methylene chloride in the air may cause dizziness, nausea and numbness of fingers and toes.

Methylene chloride has to be handled with care. Direct contact with the eyes can bum the cornea and direct contact with the skin may cause intense burning and mild redness of the skin. Methylene chloride also affects our central nervous system.

Trichloromethane (chloroform)

Chloroform is used as a solvent for fats, waxes, resins, etc. Fortified by alcohol, it finds use as a general anaesthetic, but its use is discouraged because of its harmful side effects such as nausea and liver toxicity.

The major use of chloroform is in the production of the refrigerant freon. Inhaling chloroform causes dizziness and headache. Chloroform is slowly oxidised by air in the presence of light to an extremely poisonous gas, carbonyl chloride, also known as phosgene. It is, therefore, stored in closed, dark bottles.

⇒ \(2 \mathrm{CHCl}_3+\mathrm{O}_2 \stackrel{\text { light }}{\longrightarrow} \underset{\text { Phosgene }}{2 \mathrm{COCl}_2}+2 \mathrm{HCl}\)

Triiodomethane (iodoform)

It is used as an antiseptic. Due to its bad smell, it has been replaced by other formulations containing iodine, such as tincture iodine.

Tetrachloromethane (carbon tetrachloride)

Carbon tetrachloride is used

1. As A Fire Extinguisher Under The Name Pyrene.
2. As A Refrigerant.
3. In The Synthesis Of Chlorofluorocarbons.
4. As A Dry-Cleaning Agent.

Exposure to CCI₄ causes liver cancer. It also causes dizziness, nausea and vomiting. 1’his chemical may irritate the eyes on contact.

Freons

The chlorofluorocarbon compounds of methane and ethane are collectively known as freons. They are used extensively as noncorrosive coolants, lubricants and nonpoisonous fire extinguishing liquids. They are also used as aerosol propellants, and for refrigeration and air-conditioning purposes.’

Freons and ozone depletion The ozone (O₃) layer in the stratosphere shields the earth from ultraviolet (UV) radiation that is harmful to living organisms.

Freons often make their way into the troposphere and diffuse unchanged into the stratosphere. They initiate a free-radical chain reaction that can upset the natural ozone balance. The reactions are as follows:

⇒ \(\mathrm{CFCl}_3+h v \rightarrow \dot{\mathrm{C} F C l}{ }_2+\dot{\mathrm{C}} \mathrm{l}\)

⇒ \(\dot{\mathrm{Cl}}+\mathrm{O}_3 \rightarrow \mathrm{ClO}+\mathrm{O}_2\)

⇒ \(\mathrm{ClO}+\mathrm{O} \rightarrow \mathrm{O}_2+\dot{\mathrm{Cl}}\)

In the chain-initiating step, UV light causes the homolytic cleavage of one C—Cl bond of the freon. The chlorine atom destroys thousands of molecules of O₃ before it diffuses out of the stratosphere or reacts with other substances. This reaction, therefore, is responsible for the depletion of the ozone layer.

DDT (p, p’-dichlorodiphenyltrichloroethane)

DDT is used as an insecticide. It is effective against mosquitoes that spread malaria. The use of DDT is banned in many countries because it is harmful and non-biodegradable.

Haloalkanes And Haloarenes Multiple-Choice Questions

Question 1. Which of the following is an aryl halide?

Answer: 4.

Question 2. Which of the following is an allyl halide?

1. CH₂ = CH—Cl
2. HC = C—CH₂CI
3. CH₂ = CH—CHCICH₃
4. CH₂ = CH—CH₂CH₂CI

Answer: 3. HC = C—CH₂CI

Question 3. Which of the following is a vinyl halide?

1. CH₂ = CH—CHB1CH₃
2. CH₃—C(Br) = CH₂
3. CH₂ = CH—CH₂—CH₂—Cl
4. HC H C—Br

Answer: 2. CH₃—C(Br) = CH₂

Question 4. How many isomers (including stereoisomers) are possible for C₄H₉Br?

1. 3
2. 2
3. 4
4. 5

Answer: 4. 5

Question 5. How many isomers does C₅H₁₁CI have?

1. 3
2. 4
3. 6
4. 8

Answer: 4. 8

Question 6. The reaction of ethyl bromide with metallic sodium in dry ether gives

1. Ethane
2. Butane
3. Propane
4. Ethylene

Answer: 2. Butane

Question 7. Which of the following is used as an anaesthetic agent?

1. Chloroform
2. Iodoform
3. Acetylene
4. Methane

Answer: 1. Chloroform

Question 8. On being boiled with alcoholic KOH, ethyl bromide gives

1. Ethyl Alcohol
2. Ethylene
3. Acetylene
4. Methane

Answer: 2. Ethylene

Question 9. The IUPAC name of secondary butyl chloride is

1. 3-chloroquine
2. 1-chloro-2-methylpropane
3. 2-chloroquine
4. 2-chloro-2-methylpropnne

Answer: 2.1-chloro-2-methylpropane

Question 10. The reaction CH₃CH₂Br + OH^–→CH₃CH₂OH + Br^–   is

1. S_N1
2. S_N2
3. Elimination Reaction
4. None Of These

Answer: 2. S_N2

Question 11. The conversion of ethyl bromide to alcohol using aqueous KOH is an Example of

1. An Addition Reaction
2. A Substitution Reaction
3. Dehydrohalogenation
4. An Elimination Reaction

Answer: 2. A Substitution Reaction

Question 12. The products of the reaction (CH₃)₃C— Br + CH₃ONa→ are

1. (CH₃)₃C—ONa + CH₃Br
2. (CH₃)₂C = CH₂ + CH₃OH + NaBr
3. (CH₃)₃C—OCH₃ + NaBr
4. (CH₃)₃C—CH₂OH + NaBr

Answer: 3. (CH₃)₃C—OCH₃ + NaBr

Question 13. The reaction of (CH₃)₃CMgCl with D₂O gives

1. (CH₃)₃CD
2. (CH₃)₃COD
3. (CD₃)₃CD
4. (CD₃)₃OD

Answer: 1. (CH₃)₃CD

Question 14. Which of the following is a refrigerant?

1. COCl₂
2. CCl₄
3. CF₄
4. CF₂Cl₂

Answer: 4. CF₂Cl₂

Question 15. Which of the following has a high dipole moment?

1. CH₃CI
2. CH₂CI₂
3. CHCI₃
4. CCI₄

Answer: 1. CH₃CI

Question 16. On being heated with aqueous NaOH, chloroform gives

1. CH₃COONa
2. HCOONa
3. CH₃OH
4. HCOOH

Answer: 2. HCOONa

Question 17. The main product of

1. 
2. CH₃CH=CHCH₃
3. CH₂=CH—CH₂—CH₃152.00 ne of these
4. No Golden Passport to Hindi for Class 11

Answer: 2. CH₃CH=CHCH₃

Question 18. DDT is

1. An Insecticide
2. Biodegradable
3. Nonbiodegradable
4. None Of These

Answer: 1. An Insecticide

Coordination Compounds And Organometallics

One of the characteristic features of transition elements is their ability to form coordination compounds or complexes. Haemoglobin in blood, chlorophyll in plants, dyes and vitamin B₁₂ are all examples of coordination compounds.

Coordination compounds find manifold applications in chemical analysis, chemical industry, metallurgical processes and medicinal chemistry. The challenging field of bioinorganic chemistry, i.e., the application of inorganic chemistry to living systems, also deals with coordination compounds.

In this chapter, we shall study some aspects of this important class of compounds as well as a new class of compounds—organometallic compounds, which contain metal-carbon bonds.

Coordination Compounds

Coordination compounds are addition compounds and are formed by the addition of stoichiometric amounts of two or more stable compounds. Double salts are also addition compounds but they exist in the crystalline state lose their identity in solution and dissociate into simple ions. In contrast, coordination compounds or complexes retain their identity in solution.

Mohr’s salt [FeSO₄.(NH₄)₂SO₄ -6H₂O] and potash alum [K₂SO₄.AI₂(SO₄)₃ -24H₂O] are double salts. A solution of Mohr’s salt cannot be distinguished from a solution obtained by mixing ferrous sulphate and ammonium sulphate since both contain Fe²⁺, NH₄ and SO₄ ions.

Unlike double salts, coordination compounds do not dissociate in solution. When aqueous ammonia is added to copper(2) sulphate solution, a dark blue colouration is obtained. The addition of alcohol to the solution results in the precipitation of tetraammine copper(2) sulphate as dark blue needle-shaped crystals.

⇒ \(\mathrm{CuSO}_4+4 \mathrm{NH}_3+\mathrm{H}_2 \mathrm{O} \rightarrow \mathrm{CuSO}_4 \cdot 4 \mathrm{NH}_3 \cdot \mathrm{H}_2 \mathrm{O}\)
Tetraammine copper(2) sulphate monohydrate

This is a coordination compound represented as [Cu(NH₃ )₄]SO₄. On dissociation, it forms [Cu(NH₃)₄]²⁺ and SO₂– ions and thereby retains its identity.

Another example of a coordination compound is potassium hexacyanoferrate(2), also called potassium ferrocyanide, K₄[Fe(CN)₆]. This on dissolution produces K⁺ and [Fe(CN)₆]4^– ions. The complex ion does not dissociate further to give Fe²⁺ and CN^–.

Important terms used In the context of coordination compounds

Coordination entity (complex) A coordination entity or complex comprises a central metal atom or ion attached to and surrounded by a group of anions or neutral molecules. These ions or molecules are called ligands and they can donate an electron pair to the central atom or ion, for example [Cu(NH3)₄]²⁺, [Fe(H₂O)₆]²⁺ and [Fe(CN)₆]^4-.

A coordination entity may be cationic if it has a net positive charge, for example, [CO(NF₃)₄Cl₂]⁺, and anionic if it bears a negative charge, for example [Ag(CN)₂]^– or electrically neutral, example [Pt(NH₃)₂Cl₂]. These are also called molecular complexes.

Central atom/ion The metal atom or cation of a coordination entity to which a fixed number of ligands or electron pair donors are attached is called the central atom/ion. For example, the complexes [Cu(NH₃)]²⁺ and [FeCl₆]^3-, the central ions are Cu²⁺ and Fe³⁺ respectively.

Ligands A ligand is a neutral or anionic species capable of donating a pair of electrons to the central metal/ion thus forming a coordinate bond. Ligands may be anions like Cl^–, CN^–, OH^–‘, C₂O_4^–  or neutral molecules like NH₃, H₂O, NH_2, CH₂ NCH₂ NH₂   (ethane-1, 2-diamine) all containing lone pair of electrons.

In a coordination entity, the ligand acts as a Lewis base (electron-pair donor) and the metal atom/ion as a Lewis acid (electron-pair acceptor).

Coordination number The total number of sigma bonds formed by ligands surrounding the metal ion is called the coordination number of the central atom/ion. For example, the coordination number of platinum in [Pt (: NH₃)₄]²⁺ is 4.

The sigma bonding electrons can be represented as dots preceding the donor atom.

Remember that pi-bonds if any present between the ligand and the central atom/ion are not considered while determining the coordination number.

Coordination polyhedron Ligands in the complex are arranged in a fixed geometric pattern. The spatial arrangement of ligands attached around the central metal atom/ion is called the coordination polyhedron. The common shapes of coordination polyhedra are tetrahedral, square planar and octahedral. Square pyramidal and trigonal bipyramidal entities are also known.

Coordination and ionisation spheres The coordination sphere is made up of the central atom/ion and the ligand. In the formula, it is represented as enclosed in square brackets.

In the complex [Co(NH₃)₆]Cl₃, the coordination sphere is made up of Co³⁺ and six NH₃ molecules (ligand).

A cationic or anionic complex needs anions or cations to preserve its electrical neutrality. Such ions are called counter ions. They are not directly attached to the metal ion and are present in the ionisation sphere, for example in [CO(NH₃ )₆ ]CI₃ the ionisation sphere is made up of three C^– ions.

The oxidation number of a central atom The oxidation number of the central metal atom is the charge it would carry if all ligands were parentheses after the name or symbol of the metal.

Some examples of coordination entities, their central atom and the oxidation number are summarised.

Types of ligands Ligands may be classified in two ways: (1) on the basis of the charge they carry, and (2) on the basis of the number of donor atoms they contain.

A ligand may be negative (Cl^–, F^–, SCN^–, etc.), neutral (NH₃, H₂O) or, very rarely, positive (NH₂NH₃⁺).

When a ligand is bound to the central metal ion through one donor atom, it is called a monodentate or unidentate ligand, for example, H₂O, NH₃, Cl^–, CN^–, etc. Ligands having two donor atoms are called bidentate or bidentate ligands. Some examples of bidentate ligands are as follows.

The IUPAC nomenclature has accepted certain short-form notations for some legends. Some of these are ox (oxalate), gly (glycine), EDTA (ethylene diamine tetraacetic acid), en (ethane-1, 2-diamine), etc.

Ligands containing more than two donor atoms are called polydentate. These may be tridentate or tridentate having three donor atoms, tetradentate, having four donor atoms, pentadentate (five) and hexadentate (six). An important hexadentate legend is ethylenediaminetetraacetate (EDTA).

Some ligands may have two donor atoms but the two sites are not used simultaneously. These are called ambidentate ligands, for example, the thiocyanate ion (SCN^–) can ligate through either the sulphur or the nitrogen atom.

Similarly, the nitrite ion (NO₂^–) can coordinate through the nitrogen or the oxygen atom. Chelating ligands and chelates When a bidentate and a polydentate ligand bind a single metal ion using ds two or more (in the case of polydentate) donor atoms, a complex known as a chelate is formed.

Such ligands are mlJed chelating ligands. An example of a chelate is the complex formed between ethylenediamine and Cu²⁺ it is represented as

Chelate complexes form a ring structure.

Chelating ligands form more stable complexes than those formed by monodentate ligands. This is called the chelate effect. The five-membered or six-membered rings are the most stable.

A molecule of a bidentate ligand blocks two coordinate sites whereas a tridentate ligand blocks three coordinate sites, and so on.

Therefore, while calculating the coordination number of a metal ion, knowledge of the type of the ligand is important. For example, the coordination number of iron in [Fe(C₂O₄ )₃]³⁺ is 6 and not 3 as oxalate is a bidentate ligand.

Homoleptic and heteroleptic complexes A homoleptic complex is one in which the metal ion is bound to only one ligand, for example [Co(H₂O)₆]³⁺. However, if a metal is bound to more than one ligand, for example, [Co(NH₃)₄(H₂O)₆]³⁺, it is referred to as a heteroleptic complex.

Principles And Processes Of Isolation Of Elements

That matter is made up of different elements is something you have known for a long time. Clements are present in the earth’s crust and also in water bodies (seas and oceans). Elements naturally occur in the form of compounds. However, some of the less reactive elements like gold, platinum, carbon, sulphur and noble gases are found in the free state.

The naturally occurring compounds of elements present in the earth’s crust are called minerals. Minerals may be obtained by mining. A metal may be present in many minerals, but it may be profitably extracted from only a few of them. The mineral from which a metal can be extracted on a commercial basis Is called an ore. Ores are generally associated with rocks and sand.

The process of obtaining a pure metal from an ore is referred to as metallurgy.

Occurrence and Abundance of Metals

As already stated, metals are generally present in the earth’s crust in the combined state. Being exposed to air and moisture, many metals occur as oxides. The abundance of oxygen and its high electronegativity also contribute to the formation of natural oxides.

Many metals occur as sulphides and carbonates too. However, halides, sulphates, phosphates and silicates are also not so uncommon. Different metals occur in different chemical combinations because of the differences in their chemical properties. For example, as you can see in Table 6.1, calcium occurs as a carbonate whereas beryllium occurs as a silicate. Despite their chemical similarity (both Be and Ca belong to Group 2), the metals occur in different chemical combinations in nature. The carbonate of beryllium is unstable whereas that of calcium is stable. Differences in the ionic sizes and different enthalpies of the formation of compounds also lead to an enormous variety of minerals in nature.

As you can a metal may occur in more than one form but for the purpose of extracting it commercially, only a suitable ore is chosen. Generally, silicate minerals are not chosen because it is difficult to extract metals from them.

The abundance of elements in the earth’s crust or any other region like the oceans and the atmosphere also varies to a great extent. For example, lithium is the thirty-fifth most abundant element in the earth’s crust. The elements belonging to the same group (Group 1)—sodium and potassium—are the seventh- and eighth most abundant elements respectively.

Aluminium is the most abundant metal in the earth’s crust, being present in igneous rocks, clay, and mica and as a constituent of some gemstones. However, it is extracted from a hydrated oxide called bauxite. In other words, bauxite is an ore of aluminium. Likewise, iron is found as oxides, sulphides and carbonate but the most important ores are the oxides. Whenever it is possible to extract metals from them, oxides are preferred to sulphides because the latter have to be roasted and during that process, sulphur dioxide, a polluting gas, is produced.

Extraction of Metals

Metals are generally extracted from their ores by reduction. However, prior to this, the ore is subjected to various operations. Broadly, the three steps involved in metallurgy are as follows.

1. Concentration of the ore
2. Extraction of the crude metal from the concentrated ore
3. Refining of the metal

Different methods and techniques are employed for each stage. The choice of the technique depends on the nature of the ore, the chemical reactivity of the metal to be extracted, the type of impurity present in the ore and the available commercially viable processes.

The big lumps of the ore are crushed into smaller pieces. The crushed ore is then ground into fine powder and this process is called pulverisation. The pulverised ore is then concentrated.

Concentration Of Ores Notes

The ores obtained from the earth’s crust contain impurities (gangue). The removal of these impurities from the ore is called concentration, ore dressing or beneficiation.

Some of the important methods of concentration of ores are discussed.

Hydraulic washing or gravity separation

Hydraulic Separation Method

The method is based on the difference between the specific gravities of the ore and the gangue particles. The powdered ore is washed with an upward stream of running water. The lighter gangue particles are washed away, leaving the heavier ore particles behind. The oxide ores of iron and tin and native ores of gold and silver are concentrated by this method.

Concentration Of Ores In Chemistry

Hydraulic washing of the ore is also done using a Wilfley table. The table is rectangular, sloping at 25°, with ribs on its top. The powdered ore is introduced on the table surface by vibrating hoppers. The water is made to flow over the table surface for washing the ore; the lighter gangue particles are flushed off the table soon whereas the heavier ore particles are retained by the grooves in the table.

Magnetic separation

This method has limited applicability and is employed when either the ore or the gangue is magnetic in nature. The powdered ore is placed on a conveyor belt which passes over two rollers, one of which has an electromagnet in it. As the ore passes over the belt, the magnetic particles get attracted by the magnetic roller and fall below the roller, while the nonmagnetic particles fall separately, away from the roller.

Froth flotation

This process is widely used for the concentration of sulphide ores, for example, those of zinc, copper and lead. The principle behind this technique is that the ore particles are preferentially wetted by oil and that of the gangue by water. The finely powdered ore is suspended in water and certain chemicals called collectors (for example pine oil, xanthates, fatty acids) and froth stabilisers (for example cresols, and aniline) are added. The collectors prevent wetting of the mineral particles by water and froth stabilisers stabilise the froth.

A vigorous stream of air Is blown through the ore-wafer suspension. The ore particles are rendered hydrophobic by collectors so that the air bubbles adhere only to particles of tire ore and carry them to the surface while gangue particles are left behind. This results In the formation of a froth on the surface carrying ore particles. The froth is skimmed off the surface. It is then dried to recover the ore particles.

Sometimes the ore may contain sulphides of two metals. These may be separated from each other by adjusting the relative proportions of oil and water. In some cases, depressants are used, which selectively prevent one type of sulphide from coming with the froth.

A mixture of ZnS and PbS may be separated by using NaCN as the depressant, NaCN forms a computer. zinc and prevents ZnS from being carried with the froth.

⇒ \(4 \mathrm{NaCN}+\mathrm{ZnS} \rightarrow \mathrm{Na}_2\left[\mathrm{Zn}(\mathrm{CN})_4\right]+\mathrm{Na}_2 \mathrm{~S}\)

Leaching (Hydrometallurgy)

This method is employed when the ore is soluble In a suitable solvent One of the most importantappSesfier of leachings is in Bayer’s process, where pure aluminium oxide (Al203) Is obtained from bauxite—the principal ore of aluminium. The main impurities present in the ore are silica and oxides of iron and titanium. The ore is first treated with concentrated sodium hydroxide at 473-523 K at high pressure. Al203 and St02 dissolve as sodmrn. aluminate and sodium silicate respectively, leaving the impurities behind.

⇒ \(\mathrm{Al}_2 \mathrm{O}_3(\mathrm{~s})+2 \mathrm{NaOH}(\mathrm{aq})+3 \mathrm{H}_2 \mathrm{O}(\mathrm{l}) \rightarrow 2 \mathrm{Na}\left[\mathrm{Al}(\mathrm{OH})_4\right](a q)\)

The solution is then filtered and cooled and its pH is lowered by passing carbon dioxide through it, and hydrated aluminium oxide gets precipitated. This process is hastened by seeding the reaction mixture with fresh precipitated hydrated aluminium oxide.

⇒ \(2 \mathrm{Na}\left[\mathrm{Al}(\mathrm{OH})_4\right (\mathrm{aq})+\mathrm{CO}_2(\mathrm{~g}) \rightarrow \mathrm{Al}_2 \mathrm{O}_3 \cdot x \mathrm{H}_2 \mathrm{O}(\mathrm{s})+2 \mathrm{NaHCO}_3(\mathrm{aq})\)

Sodium silicate remains in solution. The hydrated oxide is filtered, dried and then heated to give pure alumina.

⇒ \(\mathrm{Al}_2 \mathrm{O}_3 \cdot x \mathrm{H}_2 \mathrm{O} \stackrel{1473 \mathrm{~K}}{\longrightarrow} \mathrm{Al}_2 \mathrm{O}_3(\mathrm{~s})+x \mathrm{H}_2 \mathrm{O}(\mathrm{g})\)

Another important application of leaching is in the concentration of silver and gold ores by using sodium or potassium cyanide. The finely powdered ore is treated with a dilute solution of NaCN or KCN in the presence of air. The metal forms a soluble complex leaving the impurities behind.

⇒ \(4 \mathrm{M}(\mathrm{s})+8 \mathrm{CN}^{-}(\mathrm{aq})+2 \mathrm{H}_2 \mathrm{O}(\mathrm{aq})+\mathrm{O}_2(\mathrm{~g}) \rightarrow 4\left[\mathrm{M}(\mathrm{CN})_2 \mathrm{~J}^{-}(\mathrm{aq})+4 \mathrm{OH}^{-}\right.\)

The metal may be recovered by treating the solution with powdered zinc, which replaces the metal in the complex.

⇒ \(2\left[\mathrm{M}(\mathrm{CN})_2\right]^{-}(\mathrm{aq})+\mathrm{Zn}(\mathrm{s}) \rightarrow\left[\mathrm{Zn}(\mathrm{CN})_4\right]^{2-}(\mathrm{aq})+2 \mathrm{M}(\mathrm{s})\)

Extraction Of The Crude Metal From The Concentrated Ore

If the metal in the concentrated ore is present in the oxidised state (i.e., in the form of a compound), the die oxide has to be reduced in order to obtain the metal in the free state. As you know, not all compounds can be reduced with equal ease. Generally, oxides can be most readily reduced. So the extraction of metals is usually done from their oxides. Metals found in water-free and oxygen-poor conditions occur as their sulphides and not oxides. In the extraction of such metals, the process of reduction is preceded by conversion of the ore into respective oxide.

Conversion of ore Into metal oxide

Melnls present in ores as hydrated oxides, carbonates and sulphides are converted into their oxides by either calcination or roasting. The choice of the process employed depends upon the chemical composition of the ore.

Calcination This process is used when the ore is in the form of a carbonate or a hydrated oxide. The concentrated ore is heated in a reverberatory furnace below its melting point, either in the absence of air or in a limited supply of air. During the process, the following changes occur.

1. Moisture and volatile impurities present in the ore are removed.
2. Impurities like sulphur, phosphorus and arsenic are removed as their volatile oxides.
3. Hydrated oxides and hydroxides undergo dehydration to give the oxides.
   \(\mathrm{Fe}_2 \mathrm{O}_3 \cdot x \mathrm{H}_2 \mathrm{O} \stackrel{\Delta}{\longrightarrow} \mathrm{Fe}_2 \mathrm{O}_3+x \mathrm{H}_2 \mathrm{O}\)
   Limonite
4. Carbonate ores are converted to oxides by the loss of carbon dioxide.
   \(\mathrm{CaCO}_3 \cdot \mathrm{MgCO}_3 \stackrel{\Delta}{\longrightarrow} \mathrm{CaO}+\mathrm{MgO}+\mathrm{CO}_2\)
   Dolomite
   \(\mathrm{CaCO}_3 \stackrel{\Delta}{\longrightarrow} \mathrm{CaO}+\mathrm{CO}_2\)
   Limestone
   \(\mathrm{ZnCO}_3 \stackrel{\Delta}{\longrightarrow} \mathrm{ZnO}+\mathrm{CO}_2\)
   Calamine
   \(\mathrm{CuCO}_3 \cdot \mathrm{Cu}(\mathrm{OH})_2 \stackrel{\Delta}{\longrightarrow} 2 \mathrm{CuO}+\mathrm{H}_2 \mathrm{O}+\mathrm{CO}_2\)
   Malachite
5. The ore becomes porous.

Roasting This method is employed mostly for sulphide ores. The ore is heated strongly, in an excess of air, at a temperature insufficient to melt it. This is usually done in a reverberatory furnace. In the process of roasting the sulphide ores are converted to oxides. Some examples are as follows.

\(2 \mathrm{PbS}+3 \mathrm{O}_2 \rightarrow 2 \mathrm{PbO}+2 \mathrm{SO}_2\)
Galena
\(2 \mathrm{ZnS}+3 \mathrm{O}_2 \rightarrow 2 \mathrm{ZnO}+2 \mathrm{SO}_2\)
Zinc blende
\(2 \mathrm{Cu}_2 \mathrm{~S}+3 \mathrm{O}_2 \rightarrow 2 \mathrm{Cu}_2 \mathrm{O}+2 \mathrm{SO}_2\)
Copper glance

The sulphur dioxide liberated in the processes is used to manufacture sulphuric added by the contact process.

Reduction of the oxide to the metal

The roasted or the calcined ore contains the metal in the form of its oxide, i.e., the metal is present in the positive oxidation state. To obtain the metal in its free state, the metal oxide has to be reduced. Generally, carbon is used as the reducing agent due to its availability and low cost. However, the process of reduction by carbon requires high temperatures, which is expensive. Another disadvantage is the formation of carbides with metals.

Carbon combines with the oxygen of the metal oxide to give the free metal and carbon monoxide. For example,

⇒ \(\mathrm{ZnO}+\mathrm{C} \stackrel{\Delta}{\longrightarrow} \mathrm{Zn}+\mathrm{CO}\)

The carbon monoxide produced may also bring about the reduction of any unreacted oxide.

⇒ \(\mathrm{ZnO}+\mathrm{CO} \stackrel{\Delta}{\longrightarrow} \mathrm{Zn}+\mathrm{CO}_2\)

In all cases, heating is required. Sometimes, the temperature needed for reduction by carbon is too high to make the process economically viable. Then the reduction is done by another highly electropositive metal. Aluminium i one such metal. It releases a large amount of energy on oxidation to A1203. In the process, the metal to be extracted gets reduced to its free state. The release of a large amount of energy when aluminium reacts with iron(III) oxide is the basis of the thermit process, which you have studied in your earlier class. The temperature attained due to this highly exothermic reaction is enough to melt iron, and it is therefore used in localised welding.

⇒ \(2 \mathrm{Al}+\mathrm{Fe}_2 \mathrm{O}_3 \rightarrow \mathrm{Al}_2 \mathrm{O}_3+2 \mathrm{Fe}\)

TW thermal tivaHwnl of mineral* and ores to bring about physical and chemical transformations enabling extraction of the metal is studied in pyrometallurgy—a branch of metallurgy.

In Older to make the choice of reducing agent to reduce a metal oxide and to know the appropriate temperature m\k\l for reduction, it is helpful to use some basic concepts of thermodynamics, particularly the variation in Gibbs fits’ energy of the reactions.

Thermodynamics in Metallurgy

In the process of extracting a metal from its ore, the main difficulty is generally faced in the reduction of the metal oxide. The Gibbs free energy changes occurring during these processes are of importance in metallurgy. You have already studied that the change in AG (Gibbs free energy) for any reaction at a specified temperature (in Kelvin) is given by the equation:

⇒ \(\Delta G^{\Theta}=\Delta H^{\Theta}-T \Delta S^{\Theta},\)

where AH and AS refer to the enthalpy and entropy changes respectively for that reaction and T is the specified temperature. You already know that the change in Gibbs free energy for a reaction can also be expressed as

⇒ \(\Delta G^\theta=-R T \ln K\)

where K is the equilibrium constant for that reaction. For a reaction to be spontaneous, i.e., for it to proceed in the desired direction, AGe should be negative. A negative value of AGÿ is possible only if AHe is negative and ASÿ is positive. As the temperature increases, the value of TAS° also increases, so that TASe > A and therefore AGe becomes more negative.

Consider a reaction involving the formation of an oxide from a metal and oxygen.

⇒ \(2 x \mathrm{M}(\mathrm{s} \text { or } \mathrm{l})+\mathrm{O}_2(\mathrm{~g}) \rightarrow 2 \mathrm{M}_x \mathrm{O}(\mathrm{s} \text { or } \mathrm{l})\)

[Both (s) and (1) are used in the reaction as both the metal and the oxide may melt at a very high temperature.]

Oxygen is used up in the reaction in its gaseous form. You know that gases have greater randomness than solids and liquids, and therefore a high entropy. So in this reaction, the entropy decreases (as the products are not gases) and therefore ASe is negative. As the temperature increases TAS’0’ becomes more negative and since TASe is subtracted from AH0, AGe becomes less negative.

Thus the magnitude of AGe decreases with an increase in temperature.

Ellingham diagram

Ellingham diagrams show the values of standard Gibbs free energy changes for chemical reactions in graphical form with AGe as the ordinate and temperature as the abscissa. The most common diagrams are for the metal oxide systems, i.e., showing the reaction of a metal with oxygen. However, diagrams are also known for metal sulphides and halides. These diagrams help us to predict the conditions under which a metal ore will be reduced to the metal.

The free energy changes that occur when 1 g mol of a common reactant (oxygen) is used may be plotted as a function of temperature (with AGe decreasing upwards) for a number of reactions of metals to their oxides as shown in. The graph shown is called an Ellingham diagram for oxides. (Similar Ellingham diagrams can be drawn for sulphides and halides as well.) The Ellingham diagram for oxides has the following important features.

1. The graphs for metal-to-metal oxide formation slope upwards as AG° becomes less negative, i.e., free energy change decreases with an increase in temperature.
2. Each plot is a straight line and at the kinks in the lines, the slope changes, indicating a change in the phase of the metal, i.e., it either melts or vaporises, and the entropy of the reaction changes accordingly.

The following are the chemical reactions involved in the formation of the respective oxides

• 2Ag+O₂→Ag₂O
• 2Hg+O₂→2HgO
• 2Ni+O₂→2NiO
• 2Fe+O₂→2FeO
• C+O₂→CO₂
• 2C+ O₂ -> 2CO
• \(\frac{4}{3} \mathrm{Cr}+\mathrm{O}_2 \rightarrow \frac{2}{3} \mathrm{Cr}_2 \mathrm{O}_3\)
• Ti+O₂→TiO₂
• \(\frac{4}{3} \mathrm{Al}+\mathrm{O}_2 \rightarrow \frac{2}{3} \mathrm{Al}_2 \mathrm{O}_3\)
• 2Mg+O₂→2MgO
• 2Ca+O₂→2CaO

Entropy Increases as solid changes to liquid and liquid changes to gas, As you can see in the Slope Of the I lg-1 JgO line changes at the boiling point of mercury (l,e„ 630 K) and that of the Mg-MgO line changes at 1303 K.

3. As the temperature Is raised, a stage will come when the graph for a metal oxide crosses the \(\Delta G^\theta\) = 0 line. If it’s temperature, the free energy of formation of the oxide is negative, i.e., the oxide is stable. Above this temperature, the free energy of formation of the oxide is positive, i,e., the oxide is unstable and it decomposes, Theoretically, all graphs will cross the ΔG = 0 line at some temperature but generally these temperatures are too high and thus thermal decomposition of oxides is not an economically and technologically feasible process for obtaining the metals, In practice, AgO and HgO decompose at attainable temperatures,

4. In a number of extraction processes, one meal Is used to reduce the oxide of another metal, which is to be extracted, If the Oils process is feasible the metal which is used as a reducing agent should form more stable oxide than Ine oxide which Is being reduced, A vertical line drawn on the Ellingham diagram of metal oxides, al any temperature, gives the order of the stability of the metal oxides, which decreases on moving upwards, Any meal can reduce the oxides of those metals which lie above it in the Ellingham diagram of metal oxides, This Is because \(\Delta G^\theta\) becomes more negative by an amount equal to the difference between the two graphs al a certain temperature.

Thus, Al can reduce oxides of Fe, Cr and Ni but It cannot reduce MgO below 1773 K (point of Intersection of 2Mg+O₂ →2MgO and \( \frac{4}{3} \mathrm{Al}+\mathrm{O}_2 \rightarrow \frac{2}{3} \mathrm{Al}_2 \mathrm{O}_3\) lines), At tills temperature, AG(t becomes zero and equilibrium results, and an increase In temperature will make the reaction proceed provided no kinetic barriers exist

Carbon As A Reducing Agent

Carbon is an Inexpensive and easily available reducing agent if combined with oxygen to give two oxides, CO (on partial oxidation) and C02 (on complete oxidation).

⇒ \(\mathrm{C}(\mathrm{g})+\frac{1}{2} \mathrm{O}_2(\mathrm{~g}) \rightarrow \mathrm{CO}(\mathrm{g})\)

⇒ \(\mathrm{C}+\mathrm{O}_2(\mathrm{~g}) \rightarrow \mathrm{CO}_2(\mathrm{~g})\)

In the case of partial oxidation of carbon, 2 volumes of CO are produced for every volume of oxygen consumed; thus ASe for the reaction is positive. Hence, \(\Delta G^\theta\) becomes increasingly negative as temperature increases. We know that \(\Delta G^{\Theta}=\Delta H^{\Theta}-T \Delta S^{\Theta}\), which is the basis for the appearance of the Ellinghum diagram.

We also know that the enthalpy and entropy of a reaction are, to a reasonable approximation, independent of temperature; therefore the slope of a line in the Ellingham diagram should be equal to the ASe of that reaction.

If a graph is plotted between \(\Delta G^\theta\) and T for the formation of carbon monoxide, we will get a line sloping downwards.

In the case of complete oxidation of carbon, the volumes of oxygen consumed and CO_Z produced are almost the same, so there is hardly any change in \(\Delta G^\theta\). Therefore, the graph between AGe and T0 is almost horizontal. As we can see the two lines (one for the formation of CO₂ and the other for the formation of CO) intersect at 983 K.

Below this temperature, the formation of CO₂ is favoured while above this temperature, the formation of CO is favoured. (The more negative the value of \(\Delta G^\\) theta, the more is the reaction favoured.) When the Ellingham diagram for the formation of oxides of carbon is superimposed on the Ellingham diagram for the formation of metal oxides, the line for the formation of carbon monoxide intersects all the metal-metal oxide lines due to its slope in the opposite direction. Hence, carbon can reduce any metal oxide provided the temperature is sufficiently high.

The use of carbon for reduction purposes is not feasible for metal oxides seen towards the bottom of the Ellingham diagram. The reduction of these oxides requires very high temperatures, making the reduction process uneconomical and technologically difficult. One more limitation encountered in carrying out such reduction processes is that the metals may form carbides.

We know that a reaction between two or more elements can be expressed in the form of a chemical equation. The reduction of a metal oxide by carbon or carbon monoxide can be expressed as a combination of various reactions whose AyG values are available in the literature. These reactions can be coupled to determine ArGe. If ArGe for the reaction is negative then the reaction is feasible.

During reduction, the metal oxide decomposes as

⇒ \(\mathrm{M}_x \mathrm{O}(\mathrm{s}) \rightarrow x \mathrm{M}(\mathrm{s} \text { or } \mathrm{l})+\frac{1}{2} \mathrm{O}_2(\mathrm{~g})\)

The reducing agent combines with the oxygen liberated in (a).

⇒ \(\mathrm{C}(\mathrm{s})+\frac{1}{2} \mathrm{O}_2(\mathrm{~g}) \rightarrow \mathrm{CO}(\mathrm{g})\)

⇒ \(\mathrm{C}(\mathrm{s})+\mathrm{O}_2(\mathrm{~g}) \rightarrow \mathrm{CO}_2(\mathrm{~g})\)

The values of CO and C02 are available in the literature. The decomposition of MxO is the reverse of its formation.

⇒ \(x \mathrm{M}(\mathrm{s} \text { or } \mathrm{l})+\frac{1}{2} \mathrm{O}_2(\mathrm{~g}) \rightarrow \mathrm{M}_x \mathrm{O}(\mathrm{s})\)

The AyGe of (d) is known. Thus, ArGe of (a) is the reverse of AyGe of MxO. If Equation (a) is added to Equations (b) and (c) separately then the composite equations respectively are as follows.

⇒ \(\mathrm{M}_x \mathrm{O}(\mathrm{s})+\mathrm{C}(\mathrm{s}) \rightarrow x \mathrm{M}(\mathrm{s} \text { or } \mathrm{l})+\mathrm{CO}(\mathrm{g})\)

⇒ \(\mathrm{M}_x \mathrm{O}(\mathrm{s})+\frac{1}{2} \mathrm{C}(\mathrm{s}) \rightarrow x \mathrm{M}(\mathrm{s} \text { or } \mathrm{l})+\frac{1}{2} \mathrm{CO}_2(\mathrm{~g})\)

These chemical equations represent the chemical reactions occurring between MxO and carbon and ArG” be determined.

While using carbon as a reducing agent, the CO produced during reduction may get oxidised to C02 according to the following equation.

⇒ \(\mathrm{CO}(\mathrm{g})+\frac{1}{2} \mathrm{O}_2(\mathrm{~g}) \rightarrow \mathrm{CO}_2(\mathrm{~g})\)

Thus, CO(g) may also act as a reducing agent as here.

⇒ \(\mathrm{M}_x \mathrm{O}(\mathrm{s})+\mathrm{CO}(\mathrm{g}) \rightarrow x \mathrm{M}(\mathrm{s} \text { or } \mathrm{l})+\mathrm{CO}_2(\mathrm{~g})\)

It has been shown that below 983 K, CO is more readily oxidised to COz than carbon and is a better-reducing agent. However, above 983 K, carbon is a better-reducing agent than CO as the latter becomes more stable and less susceptible to oxidation to C02.

Example The values ofAfG* ofMgO and CO respectively at 127 K and 2273 K are as follows.

⇒ \(\frac{1}{2} \mathrm{Mg}(\mathrm{s})+\mathrm{O}_2(\mathrm{~g}) \rightarrow \mathrm{MgO}(\mathrm{s}) ; \quad \Delta_f G^{\ominus}=-941 \mathrm{~kJ} \mathrm{~mol}^{-1}(1273 \mathrm{~K})=-314 \mathrm{~kJ} \mathrm{~mol}^{-1}(2273 \mathrm{~K})\)

⇒ \(\mathrm{C}(\mathrm{s})+\frac{1}{2} \mathrm{O}_2(\mathrm{~g}) \rightarrow \mathrm{CO}(\mathrm{g}) ; \quad \Delta_f G^{\ominus}=-439 \mathrm{~kJ} \mathrm{~mol}^{-1}(1273 \mathrm{~K})=-628 \mathrm{~kJ} \mathrm{~mol}^{-1}(2273 \mathrm{~K})\)

Indicate whether carbon can be used as a reducing agent at 1273 K and 2273 K.

Solution

The overall reaction for the reduction of magnesium oxide is

⇒ \(\mathrm{MgO}(\mathrm{s})+\mathrm{C}(\mathrm{s}) \rightarrow \mathrm{Mg}(\mathrm{s})+\mathrm{CO}(\mathrm{g}) \quad \Delta_r G^\theta=\Delta_f G_{(\mathrm{C} \rightarrow \mathrm{Co})}^{\ominus}-\Delta_f G_{\left(\mathrm{M}_{\mathrm{g}} \rightarrow \mathrm{Mg}_g \mathrm{O}\right)}^{\ominus}\)

⇒ \(\text { At } 1273 \mathrm{~K}, \Delta_r G^{\ominus}=-439-(-941)=+502 \mathrm{~kJ} \mathrm{~mol}^{-1} \text {. }\)

Since \(\Delta_r G^{\Theta}\) is positive, the reaction is not possible,

⇒ \(\text { At } 2273 \mathrm{~K}, \Delta, G^{\Theta}=-628-(-314)=-314 \mathrm{~kJ} \mathrm{~mol}^{-1}\)

Since \(\Delta_r G^{\Theta}\) is negative, the reaction can proceed.

Example The reaction is endothermic at low temperatures but becomes exothermic above the melting point of magnesium. Example giving a reason.

⇒ \(\mathrm{MgO}(\mathrm{s})+\mathrm{C}(\mathrm{s}) \rightarrow \mathrm{Mg}(\mathrm{s})+\mathrm{CO}(\mathrm{g})\)

Solution

The entropy of the liquid state is higher than the entropy of the Solid State. Thus above the melting point of magnesium, the entropy of the products is high, i.e., AS is more positive and thus the reaction becomes exothermic, i.e., ΔG becomes negative.

Under what conditions can magnesium be used for the reduction of chromium oxide?

Solution

The equations involved in the process are

⇒ \(\frac{4}{3} \mathrm{Cr}+\mathrm{O}_2 \rightarrow \frac{2}{3} \mathrm{Cr}_2 \mathrm{O}_3\)

⇒ \(2 \mathrm{Mg}+\mathrm{O}_2 \rightarrow 2 \mathrm{MgO}\)

The overall reaction is

⇒ \(\frac{2}{3} \mathrm{Cr}_2 \mathrm{O}_3+2 \mathrm{Mg} \rightarrow 2 \mathrm{MgO}+\frac{4}{3} \mathrm{Cr}\)

According to the Ellingham diagram (Figure 6.5), the reaction will be possible at any temperature above the point of intersection of the Cr203 and MgO curves.

Example The thermite reaction is exothermic, but does not proceed at room temperature and needs initiation by heating. However, once the reaction has started, further heating is not needed. What do you think might be the reason for this?

⇒ \(2 \mathrm{Al}+\mathrm{Cr}_2 \mathrm{O}_3 \rightarrow 2 \mathrm{Cr}+\mathrm{Al}_2 \mathrm{O}_3\)

Solution

The reaction has a very high activation energy as both reactants are in the solid state. Thus initiation is needed. As the reaction is exothermic, the reaction proceeds on its own, once the activation energy is attained.

Reduction Of Sulphides

Although carbon reduces oxides quite effectively, it is not so effective when used directly on sulphides of metals. The reason why carbon reduces oxides at high temperatures is that the AG vs. T plot for CO has a negative slope due to the positive AS value, i.e., the formation of CO becomes more feasible at high temperatures. In the case of sulphides, the reaction should be

⇒ \(\mathrm{MS}(\mathrm{s})+\frac{1}{2} \mathrm{C}(\mathrm{s}) \rightarrow \mathrm{M}(\mathrm{s})+\frac{1}{2} \mathrm{CS}_2(\mathrm{~g})\)

There is no compound formed by the reaction between carbon and sulphur which is analogous to CO (i.e., CS) with a negative slope. The Age vs. T plot for CS2 is almost horizontal.

⇒ \(\frac{1}{2} \mathrm{C}(\mathrm{s})+\mathrm{S}(\mathrm{g}) \rightarrow \frac{1}{2} \mathrm{CS}_2(\mathrm{~g})\)

That is, the plot cannot intersect the plots of metal sulphides. In other words, carbon does not have much affinity for sulphur and thus sulphides are first roasted to oxides and then reduced.

Limitations of the Ellingham diagram The study of the \(\Delta G^{\Theta}\) vs. T plots merely suggests whether or not a reaction is possible. It can only indicate whether or not thermodynamically a reducing agent can be used. It totally ignores the kinetic aspect of a reaction, i.e., the rate of a reaction.

A reaction is thermodynamically feasible if ArG° I; is negative; however, sometimes the rate of the reaction may be so low that despite the reaction being thermodynamically possible, it does not occur. The study of Fllingtwm diagrams does not include the pÿlMtity of any other competing reaction. For example, carbon cannot be used as a reducing agent for some metals which form carbides.

It is also implicitly assumed that the reactants and products are In equilibrium, which Is not necessarily true. Sometimes, if the reactants and products are all solids, the reaction is stow and becomes possible if there is a change in phase, i.e., a compound melts.

This is because the randomness increases in tire liquid state resulting in an increase in ASe. For instance, the reduction of ZnO and MgO by carbon becomes more spontaneous after the melting point of Zn or Mg as this results in greater A$i%.

Some of the common extraction processes are discussed as follows.

Extraction Of Iron From Its Oxides

Iron is extracted from its oxides in a blast furnace. The furnace is a tall refractory-lined cylindrical structure that is charged at the top. The refractory material generally used is MgO. (Refractory materials are those that can withstand high temperatures without melting or softening).

The calcined ore mixed with coke (a reducing agent) and limestone (slag forming) is referred to as charge, which is fed from the top of the blast furnace and a blast of hot air is introduced from the bottom. Coke bums to produce heat and CO.

⇒ \(2 \mathrm{C}(\mathrm{s})+\mathrm{O}_2(\mathrm{~g}) \rightarrow 2 \mathrm{CO}(\mathrm{g})\)

Since the reaction is exothermic, the temperature at the lower part of the furnace is very high. As the gases move up, they meet the descending charge and the temperature drops gradually, on moving towards the top of the furnace, resulting in a temperature gradient in the furnace. The iron oxide is reduced to iron mainly by CO and some reduction takes place by coke.

Knowing the thermodynamics of the process helps us to understand how the reduction occurs and why the blast furnace is chosen for extracting iron from its oxide.

The following is the main reduction step in the process.

⇒ \(\mathrm{FeO}(\mathrm{s})+\mathrm{C}(\mathrm{s}) \rightarrow \mathrm{Fe}(\mathrm{s} / \mathrm{l})+\mathrm{CO}(\mathrm{g})\)

This is the overall reaction obtained by coupling two reactions. One is the reduction of metal oxide to metal and the other is the oxidation of carbon to CO.

⇒ \(\mathrm{FeO}(\mathrm{s}) \rightarrow \mathrm{Fe}(\mathrm{s})+\frac{1}{2} \mathrm{O}_2(\mathrm{~g}) ; \Delta G_1\)

⇒ \(\mathrm{C}(\mathrm{s})+\frac{1}{2} \mathrm{O}_2(\mathrm{~g}) \rightarrow \mathrm{CO}(\mathrm{g}) ; \quad \Delta \mathrm{G}_2\)

Δ, G for the overall reaction is

⇒ \(\Delta_r G=\Delta G_2+\Delta G_1\)

The reaction will proceed when ArG is negative. If we consider the Ellingham diagram showing the change in AG for the formation of oxides then we find that the AG plot for Fe ->• FeO is sloping upwards while that for Q QQ is sloping downwards. The two lines intersect at about 1073 K. Above this temperature the C -> CO line falls below the Fe -» FeO line and a reduction of FeO occurs.

As discussed earlier in this section, there is a temperature gradient in the furnace. We have also studied earlier that below 983 K, carbon monoxide is a better-reducing agent than carbon. The other oxides of iron (Fe203 and Fe30< ) get reduced at relatively lower temperatures by CO, Thus the reactions occurring in the upper parts of the blast furnace/ which are at lower temperatures than that at the bottom, are as follows.

3Fe₂O₃+CO→2Fe₃O₄+CO₂
Fe₃O₄+CO→3FeO+CO₂
Fe₂O₃+CO→2FeO+CO₂

At a higher temperature in the furnace, coke is the reducing agent

CO₂+C →2CO
FeO+CO→Fe+CO₂

Adding these reactions, we get

Limestone is added to the charge to form slag; thus, it acts as a flux It decomposes to form CaO. This combines with infusible impurities like silica present in the concentrated ore to form an easily fusible material called slag.

⇒ \(\mathrm{CaO}+\mathrm{SiO}_2 \rightarrow \mathrm{CaSiO}_3\)

Flux Cangue Slag

Slag melts at the temperature that prevails in the furnace. It is insoluble in the molten metal and being lighter floats on the surface of the metal therefore it can be skimmed off, and molten iron gets collected at the bottom of the furnace. Slag protects the molten iron from being oxidised.

The iron thus obtained is called pig iron. It contains impurities like carbon (4%) and smaller amounts of sulphur, phosphorus, silicon and manganese. Another form of iron is called cast iron. It has a lower carbon content and is very hard and brittle. The purest form of iron is wrought iron.

It is malleable in nature and is prepared by heating cast iron in the presence of air in a reverberatory furnace lined with haematite. Limestone is added as flux. In the presence of air, impurities like sulphur, silicon and phosphorus are oxidised to the respective oxides. These combine with CaO (obtained by the thermal decomposition of limestone) to form slag. Haematite oxidises carbon to carbon monoxide.

Fe₂O₃+3C→2Fe+3CO

Thus all impurities from cast iron are removed

Extraction Of Copper From Copper(I) Oxide

Copper(I) oxide is not very stable and can be readily reduced by carbon. (In the Ellingham diagram the Cu→ Cu₂ Oline lies above the C→C0₂and C -> CO lines.)

Generally copper is present as the sulphide along with iron sulphide in the ore. The sulphide ore is first roasted to give the oxide.

2Cu₂S+3O₂→2Cu₂O+2SO₂

Iron, if present as a sulphide, is also converted to its oxide.

2FeS+3O₂→2FeO+2SO₂

The roasted ore is then heated in a reverberatory furnace, using silica as flux. FeO combines with silica to form iron silicate slag, which floats on the surface and therefore can be easily skimmed off leaving behind copper matte, which consists of Cu2O and unreacted Cu2Sand FeS.

⇒ \(\mathrm{FeO}+\mathrm{SiO}_2 \rightarrow \underset{\text { Slag }}{\mathrm{FeSiO}_3}\)

The molten matte is fed into a silica-lined converter through which hot compressed air is blown, which causes partial oxidation. The remaining FeS is converted to FeO, which forms a slag with silica.

2FeS+3O₂→2FeO+2SO₂

⇒ \(\mathrm{FeO}+\mathrm{SiO}_2 \rightarrow \underset{\text { Slag }}{\mathrm{FeSiO}_3}\)

The supply of hot compressed air is turned off after some time and self-reduction of the oxide and sulphide occurs.

2Cu₂O+Cu₂S→6Cu+SO₂

The molten copper obtained is 99% pure and is called blister copper. The term blister comes from the (fact that as the melt solidifies, the dissolved sulphur dioxide, oxygen and nitrogen escape giving rise to blisters on the surface of the metal.

Extraction Of Zinc From Zinc Oxide

The main ore of zinc is zinc blende (ZnS), which is first concentrated and then roasted to form the oxide. The zinc oxide is made into briquettes with coke and clay and heated In vertical retorts at 1673 K. (The temperature range for the reduction of ZnO is higher than that in the two cases discussed above.) The coke reduces the oxide to zinc.

ZnO+C→Zn+CO

The zinc obtained distils off because its boiling point is lower than the temperature of the retort. The gaseous zinc is prone to reoxidation; therefore it is collected by sudden cooling.

Not all elements can be obtained from their ores by the reduction of their oxides by carbon. Some of them which are active metals are reduced by electrolysis whereas those which are not so active can be reduced by adding some reducing element.

Electrolytic Reduction

Active metals like alkali metals, alkaline earth and aluminium cannot be obtained by the reduction of their oxides by carbon. This is because their oxides are very stable and can be only reduced at extremely high temperatures. However, at such high temperatures, the metals may form carbides.

Example Of Magnetic Separation

Such active metals are obtained by the electrolysis of their fused salts. The relation between Gibbs energy and electrode potential is as follows.

⇒ \(\Delta G^{\Theta}=-n F E^{\Theta}\)

where n is the number of electrons involved in the redox process, Ee is the electrode potential of the redox couple and F is the Faraday constant. Highly reactive metals have large negative Ee values and it is difficult to carry out their reduction chemically.

In electrolysis, the fused salt dissociates to form M”+ ions, which migrate to the cathode and get deposited and discharged (lose their charge by gaining electrons) there. Only such materials are used as electrodes which do not react with the metal.

An important example is the electrolysis of A1203 to give aluminium. Purified A1203 is mixed with cryolite (Na3AlF6) or calcium fluoride, which lowers the melting point of the mixture and makes it more conducting. The fused mixture is electrolysed in a cell containing a graphite-lined steel cathode and graphite anode. Aluminium is liberated at the cathode and oxygen at the anode.

At cathode      AI₂O₃→2AI³⁺+3O^2-

AI³⁺+3e^–→AI (I)

The oxygen combines with the carbon anodes to give CO₂ and CO.

At anode C(s)+O^2-→CO₂(g)+2e^–
C(s)+2O^2-→CO₂(g)+4e^–

Thus, the carbon anodes are progressively burnt away.

Metal Displacement

You have studied in your previous class that when a more reactive metal like iron is brought in contact with the solution of a less reactive metal like copper the former goes into the solution and the latter precipitates out.

Cu²⁺(aq)+Fe(s) →Fe²⁺(aq)+Cu(s)

This principle may be applied in metallurgy. Low-grade copper ores may be leached using or subjected to the action of certain bacteria. To this solution, scrap iron is added, which displaces copper. (Leaching with bacteria is known as bioleaching.)

Extraction Based on Oxidation

So far you have studied extraction based on reduction. Nonmetals are generally extracted by oxidation. A common example is the extraction of chlorine from seawater. The following reactions show the oxidation of chlorine.

2CI^–(aq)+2H₂O(I) →2OH^–(aq)+H₂(g)+CI₂(g)

The \(\Delta G^{\ominus}\) for this reaction is +422\(E^\theta=-2.2 \mathrm{~V}\) and £e‘= -2.2 V, You know that for a reaction to become spontaneous£e (the electrode potential) should be positive. To make the net potential positive (in case of negative potential values) an external emf greater than the standard potential value of the die reaction has to be applied. Thus, in tins case the minimum potential needed is +2.2 V. If molten NaCl is electrolysed then sodium metal is produced at the cathode.

The extraction of gold and silver by leaching the native ores with sodium cyanide also involves oxidation in the first step.

4Au(s)+8CN^–(aq)+2H₂O(aq)+O₂(g)→4[Au(CN)₂]^–(aq)+4OH^–

This is because the oxidation state of gold in the elemental state is zero whereas in the cyanide complex is +1 (i.e., the electron is lost). In the next step, gold is recovered from the cyanide complex by displacement by zinc.

2[Au(CN)₂]^–(aq)+Zn(s)→2Au(s)+[Zn(CN)₄]^2-(aq)

In this step, zinc acts as a reducing agent as it reduces gold from Au (I) to Au.

Refining Techniques

A metal obtained by extraction is not pure. It is contaminated with various impurities, a few examples of which are as follows.

1. Other metals also present in the ore are obtained by simultaneous reduction of their oxides.
2. Metalloids like silicon and nonmetals like phosphorus are obtained by the reduction of their oxides.
3. Unreacted oxides and sulphides of the metals.
4. Substances used in the furnace like flux and slag.

To obtain the pure metal, it is important to subject the crude metal to an appropriate refining technique so that the impurities are removed. The choice of the refining technique depends on the nature of the metal, the nature of impurities present and the purpose for which the pure metal is needed.

Some common methods of refining are as follows.

Distillation

This method is used when the metal being purified has a low boiling point, like mercury, cadmium and zinc. The impure metal is heated in an iron retort and the vapours are chilled to obtain the pure metal. The less volatile impurities are left behind.

Liquation

This technique is used when the melting point of the metal is lower than that of the impurities. Theerode metal is heated on a sloping surface. It melts and flows down the Scope, from where it is collected. Tin and lead are purified by this method.

Electrolytic Refining

This method is used in the refining of metals like copper,anc. aluminium, gold. silver and nickel. A block of the pure metal is made of the anode, while a thin strip of fee pure metal is made of fee cathode. The electrodes are suspended in an electrolyte.

An electrolyte is a Squid that conducts electodtj’ due to the presence of ions. Electrolytes are molten ionic compounds or solutions of ionic salts or compounds that ionise in solution. In electrolytic refining, a fee solution of an ionic salt of fee metal to be extracted is used as a fee electrolyte.

On passing current, fee electrolytes dissociate?, fee metal Ions deposit on the cathode as fee pure metaL while an equivalent amount of fee metal dissolves from fee anode and passes into solution as fee metal ion. The reactions involved in the fee process are as follows.

Anode: M→Mⁿ⁺+ne^–

Cathode: Mⁿ⁺+ne^–→M

The impurities present may either be more electropositive than the metal being refined, or less electropositive. An optimum voltage is applied such feat fee more electropositive metal impurities remain in solution, while fee less electropositive ores remain un-ionised and deposit under fee anode as anode mud.

In the refining of copper, the electrolyte is copper sulphate solution, the fee anode is a block of blister copper, while the fee cathode is a thin strip of pure copper. Impurities like selenium, tellurium, silver, gold, platinum and antimony form fee anode mud.

Zone Refining

This technique is employed to get metals of very high purity. Germanium, silicon, gallium, indium and boron, which are used in semiconductors, are purified by this method. This method is based on the fact that the solubility impurities are different in the solid and liquid states of a material.

Generally, the solubility is more in the liquid state than in the solid state. Thus, when molten metal is allowed to cool, the impurities remain in the melt, while the pure metal crystallises out.

The impure metal is taken in the form of a rod and a mobile circular heater is fitted at one end of the rod.  The portion of the metal rod in contact with the heater melts. The heater is slowly moved forward and the molten zone moves along with the heater.

As the heater moves forward, pure metal crystallises out of the melt and impurities pass on to the adjacent molten zone. The process is repeated several times; the heater is moved in fee same direction. The impurities get concentrated on one end of the rod and this end is cut off to obtain fee pure metal.

Vapour Phase Refining

This method is used to purify certain metals. The basic requirement of the process is that the metal to be purified should form a volatile compound on being heated with a reagent; furthermore, the volatile compound should undergo thermal decomposition at a high temperature to give the pure metal. There are two important processes based on this technique —the Mond process and the Van Arkel method.

Mond process

This method is used to obtain highly pure nickel. Nickel oxide is heated with water gas to 323 K. The hydrogen reduces nickel oxide, leaving behind impure nickel.

NiO(s)+H₂(g) →Ni(s)+H₂O(g)

The impure nickel is heated in a current of carbon monoxide at 330-350 K and forms a volatile carbonyl complex, tetracarbonyl nickel. This complex undergoes thermal decomposition at 450-470 K to give pure nickel.

⇒ \(\underset{\text { Impure }}{\mathrm{Ni}}+4 \mathrm{CO} \rightarrow \mathrm{Ni}(\mathrm{CO})_4\)

Van Arkel method

This method is used to prepare highly pure samples of titanium and zirconium. These metals are difficult to extract because they react readily with air, oxygen, nitrogen and hydrogen at elevated temperatures. The respective oxides cannot be reduced by C or CO as the oxides form carbides. The crude metal is heated in an evacuated vessel in the presence of iodine to form the volatile tetraiodide—Til4 or Zrl4.

The iodide is then brought in contact with an electrically heated tungsten filament when it decomposes to give the pure metal which gets deposited on the filament. The iodine liberated in the process is reused.

⇒ \(\underset{\text { Impure }}{\mathrm{Ti}(\mathrm{s})}+2 \mathrm{I}_2(\mathrm{~g}) \stackrel{523 \mathrm{~K}}{\longrightarrow} \mathrm{TiI}_4(\mathrm{~g}) \stackrel{1800 \mathrm{~K}}{\longrightarrow} \underset{\text { Pure }}{\mathrm{Ti}(\mathrm{s})}+2 \mathrm{I}_2(\mathrm{~g})\)

⇒ \(\mathrm{Zr}(\mathrm{s})+2 \mathrm{I}_2(\mathrm{~g}) \stackrel{870 \mathrm{~K}}{\longrightarrow} \mathrm{ZrI}_4(\mathrm{~g}) \stackrel{2073 \mathrm{~K}}{\longrightarrow} \mathrm{Zr}(\mathrm{s})+2 \mathrm{I}_2(\mathrm{~g})\)

An essential requirement of the process is that the metal should not be volatile at the temperature at which the decomposition of the iodide takes place.

Chromatographic Methods

As you know from your previous class, chromatography is a useful process to separate the constituents of a mixture. It is based on the principle that when a mixture is moved through an adsorbent, different components of the mixture are adsorbed to different extents.

Chromatography involves two phases—a stationary phase and a mobile phase. Many different chromatographic processes are available depending on the choice of the two phases. A few examples are paper chromatography, column chromatography and gas chromatography.

In column chromatography, an adsorbent like silica or A1203 is packed in a glass column similar to a burette. This is the stationary phase. A mixture of compounds (in soluble form) is introduced from the top of the column. As the mixture trickles down the column, the different components are adsorbed in different areas of the column and are seen as coloured bands.

The mixture to be purified is dissolved in a suitable solvent and poured through the top of the column. As the liquid percolates through the adsorbent, the different components are adsorbed at different regions of the column.

The strongly adsorbed component is adsorbed readily and is present towards the top of the column and the weakly adsorbed component towards the bottom. The adsorbed components are then removed by using a suitable reagent.

This process is called elution and the reagent used is called the eluate. The weakly adsorbed component is eluted first followed by the strongly absorbed ones. This method is suitable for elements which are available in very small amounts and whose impurities are chemically similar to the element being purified.

Zirconium and hafnium can be separated from each other by passing a solution of their chlorides (NrCI₄ and HfCI₄) in anhydrous methanol through a silica column followed by elution with 1.9 M HCI whereby NrCI₄ is eluted. HfCI₄ is later eluted by using 3.5 M H₂SO₄.

Uses of Aluminium, Copper, Iron and Zinc

Aluminium is a very useful metal due to its properties like high electrical conductivity, lightness, resistance to corrosion, and the fact that it can be recycled. Aluminium wires are used as electrical conductors. Aluminium powder is used for making paints and lacquer.

Aluminium foil is used for wrapping photographic films, medicines, and food items like sweets and chocolates. Aluminium metal is used to extract manganese and chromium from its oxides.

Being corrosion-resistant, the metal is used to make utensils. Alloys of aluminium are light and strong and find a variety of uses. The greatest use of aluminium alloys is in construction.

Copper is a very good conductor of heat and electricity and is highly malleable. It is used to make electrical wires and cables. Alloys of copper like brass (with zinc) and bronze (with tin) are very useful.

Copper compounds like basic copper hydroxide (also known as Bordeaux mixture) are used to prevent potato blight (a fungal disease).

Iron is a widely used metal and many industries are based on iron. The uses of iron are manifold. Cast iron is an important engineering material with a wide range of applications. For example, it is used in making pipes, tanks, railway coaches, machines and car parts.

It is also used to manufacture wrought iron and steel. Examples of items produced from wrought iron include nuts, bolts, rivets, horseshoes, boiler tubes, agricultural implements and furniture. There are various types of steel like nickel steel, chrome steel and stainless steel.

Steel is used in making a variety of things, from the thinnest surgical needle to immense ships. Modern structures like skyscrapers and bridges are based on a steel framework. Nickel steel is used in ship and aircraft building, for automobiles and as support in reinforced concrete and railroad tracks.

Chrome steel is used in making cutting tools while stainless steel is used in making utensils, surgical equipment, and cycle and automobile parts.

Zinc is an important constituent of alloys like brass, solder, bronze and German silver. A large proportion of zinc is used to galvanize metals like iron to prevent corrosion. Zinc metal is used in dry batteries. Zinc oxide is used as a paint pigment and in pharmaceutical products. Zinc sulphide is used for making luminous dials.

Principles And Processes Of Isolation Of Elements Multiple-Choice Questions

Question 1. The impurities present in an ore are called

1. Slag
2. Flux
3. Mineral
4. Gangue

Answer: 1. Gangue

Question 2. Which of these metals does not occur in the native state?

1. Gold
2. Platinum
3. Lead
4. Silver

Answer: 2. Platinum

Question 3. Which of these is an ore of copper?

1. Malachite
2. Calamine
3. Haematite
4. Bauxite

Answer: 1. Malachite

Question 4. Froth floatation is used to concentrate

1. Native Ores
2. Oxide Ores
3. Sulphide Ores
4. None Of These

Answer: 3. Sulphide Ores

Question 5. Which of these is a depressant?

1. Xanthate
2. Pine Oil
3. Aniline
4. Sodium Cyanide

Answer: 4. Sodium Cyanide

Question 6. Which of these is used as a collector in froth flotation?

1. Aniline
2. Pine Oil
3. Cresol
4. None Of These

Answer: 2. Pine Oil

Question 7. Ti02 is separated from bauxite by leaching with

1. Acid
2. Alkali
3. Sodium Cyanide
4. Water

Answer: 2. Alkali

Question 8. Which of these ores is roasted?

1. ZnCO₃
2. CaCO₃. MgCO₃
3. ZnS
4. AI₂O₃.xH₂O

Answer: 2. CaCO₃. MgCO₃

Question 9. Which of these is used as a flux in the extraction of iron from its oxide?

1. C
2. CO
3. Fe₂O₃
4. CaO₃

Answer: 4. CaO₃

Question 10. The purest form of iron is

1. Wrought Iron
2. Cast Iron
3. Pig Iron
4. Steel

Answer: 1. Wrought Iron

Question 11. Which metal can be purified by distillation?

1. Iron
2. Copper
3. Lead
4. Zinc

Answer: 4. Zinc

Question 12. In the extraction of aluminium, the anode is made of

1. Aluminium
2. Cryolite
3. Carbon
4. Iron

Answer: 3. Carbon

Question 13. Copper is refined by

1. Distillation
2. Liquation
3. Zone Refining
4. Electrolysis

Answer: 4. Electrolysis

Question 14. Which of these is purified by zone refining?

1. Ge
2. Pb
3. A1
4. Fe

Answer: 1. Ge