Group 15 Elements: The Nitrogen Family

NITROGEN FAMILY

The elements of groups 15 are nitrogen (N), phosphorous (P), arsenic (As), antimony (Sb) and bismuth (Bi). The first two elements N and P are non – metal. The next two, arsenic and antimony are metalloids whereas the last element, bismuth is a metal. The general electronic configuration of this group is ns²np³.

General Introduction to the nitrogen family members

Nitrogen (N)

Nitrogen gets its name from a mineral known as "niter" (potassium nitrate), from which it can be prepared. It was discovered by a Scottish physician Daniel Rutherford in 1770.

Nitrogen is a fairly common element in the earth's crust. It is most important constituent of atmosphere, making up to 78.08 %. The greatest use of nitrogen is in ammonia which is used for fertilizer production and to produce nitric acid. Nitrogen gas is used where inert atmosphere is needed, such as in an ordinary light bulb. Nitrogen is also used in protecting historic documents. Nitrogen is also used in storage of food stuff and in explosives.

Nitrogen is absolutely essential to all living organisms. It is important part of all protein molecules which, among other functions, are the building materials in all kinds of cells.

Nitrogen oxides (from combustion) are one of the cause of toxic rain.

Phosphorous (P)

Phosphorous got its name because it glows in the dark. "Phosphorous" is from the greek word phos (light) and phoros (bearer). It was first isolated in 1699 by German physician Henning Brand. He discovered it from urine.

It is a non – metals. As a gas it is colourless, but it occurs in 3 allotropic form in solid state. There are white (or yellow) phosphorous, red phosphorous and black phosphorous. The most common and reactive of which is white (or yellow) phosphorous which looks like a waxy. Transparent solid with a garlic like odour. Red phosphorous is much less reactive. It is red powder and does not dissolve in most liquids.

Black phosphorous is produced by heating white phosphorous in the presence of a mercury catalyst and a seed crystal of black phosphorous. It is least reactive, does not ignite easily.

Commercially, phosphorous compounds are used in the manufacture of phosphorus acid (H₃PO₄) which is found in soft drinks, used in fertilizers. Also phosphorous is used in fireworks, phosphorescent compounds. Phosphorous is used in forming china clay. Also used in glass making.

Arsenic (As)

It was first recognized as an element by alchemists. Although it is not certain where and when arsenic was actually discovered, credit for the discovery usually goes to alchemist Albert the great in 1250 AD.

Arsenic is occasionally found as a free element, but it is usually found in a number of minerals. Elemental arsenic occurs in two solid modifications: Yellow and grey or metallic. It is a steel gray, very brittle, crystalline metalloid. Arsenic does not melt when heated but changes directly in to a vapour (sublimation). Arsenic is stable in dry air, but the surface oxidizes slowly in moist air to give a bronze tarnish and finally a black covering to the element. When heated in air, arsenic combines with oxygen to form arseneous oxide (As₂O₃). A blue flame is produced and the compounds give off distinctive garlic – like odour.

It is necessary ultra trace element for humans, red algae, chickens, rats, goats and pigs. A deficiency of arsenic results in inhibited growth. Most infamous use of arsenic is a poison. Arsenic also used in semi conductor devices along with phosphorus.

Antimony (Sb)

The name antimony is derived from the Latin words "anti + monos" which means "not alone". The symbol for antimony comes from its original Latin name stibium. It was studied by Nicholos Lemery in 17^th century.

The standard state of antimony is solid at 298 K. It is very brittle and flakes easily. Antimony has a bluish – white metallic sheen. It is a poor conductor of heat and electricity. It is known to be severely toxic. It causes damage to the liver and is used in some cases to induce vomiting and sweating.

It is added to iron and iron alloys to increase its hardness, stiffness and reduce corrosion. Antimony sulphides are used in the vulcanization of rubber as well as in pigments.

Bismuth (Bi)

It's name came from the German word Wissmuth which means "white mass". Its symbol is Bi from Bismuth. It was distinguished from tin and lead by Claude Geoffroy in 1753.

Bismuth is the heaviest non – radioactive naturally occurring element. It is a hard, brittle metal with an unusually low melting point (271°C). Bismuth is a white – silvery crystalline, brittle metal with a pinkish tinge. It is rare metal usually arrociated with copper, lead, tin, silver and gold ores.

It unusual property to expand upon solidification makes it useful in type – metal alloys and for coating. Bismuth is used in wide variety of pigments ranging form cosmetics to the coloration of glass.

Illustration 1. Why the first ionization energy of group 15 elements much higher than those of corresponding group 14 elements?

Solution: This is due to increase in nuclear charge and extra stable configuration of the elements of this group. The extra stability of configuration is attributed to the exactly half filled p – orbitals in their valence shells.

Illustration 2. Why the tendency of catenation is less in nitrogen as compared to carbon?

Solution: The lesser tendency of nitrogen to catenate as compared with carbon can be understood in terms of their bond energies, as listed below.

As there is a big drop in the bond energy from carbon to nitrogen, evidently the tendency of catenation decreases enormously.

As we move down the group there is a decrease in covalent character.

N > P > As > Sb > Bi

Nitrogen is chemically less reactive, due to high stability of its molecule, N₂ in which two nitrogen atoms are combined through triple covalent bonds (N $\equiv$ N) from which one is sigma ($\sigma$) and two are pie ($\pi$) bonds, thus posses high bond strength (941.4 KJ mol^-1).

Nitrogen has one special feature that it can form p$\pi$ - p$\pi$ multiple bonds with itself, carbon and oxygen due to its small size. In phosphorus rather p$\pi$ - p$\pi$, d$\pi$ - p$\pi$ is found as in POX₃.

1. Hydrides

All elements of this group form gaseous hydrides of the type MH₃. The stability however decreases down the group for these hydrides.

NH₃ PH₃ AsH₃ SbH₃ BiH₃

Ammonia Phosphine Arsine Stibine Bismuthine

All the hydrides are strong reducing agents and reacts with metal ions to give phosphides, arsenides or antimonides. Phosphine and other hydrides of heavier members of this groups are highly poisonous.

Formation of Hydrides

$C{a}_3{N}_2+6H_{2}O\to Ca{\left(OH\right)}_2+2N{H}_3$

$C{a}_3{P}_2+6H_{2}O\to 2P{H}_3+3Ca{\left(OH\right)}_2$

$Z{n}_3{M}_2\left(s\right)+6HCl\left(aq\right)\to \underset{\left(M=As,SborBi\right)}{2M{H}_3\left(g\right)+3ZnC{l}_2\left(aq\right)}$

Beside these, N₂H₄ another hydride of nitrogen can be prepared by the action of NH₃ and sodium hypochlorite.

$N{H}_3+NaOCl\to N_2{H}_4+NaCl+H_{2}O$

Structure of Hydrides

In these entire hydrides central atom assumes sp³ hybrid state forming four sp³ hybrid orbitals. Bond angle of H – M – H decreases down the group. Down the group thermal stability also decreases.

NH₃ > PH₃ > AsH₃ > SbH₃ > BiH₃

It is because of decrease in M – H bond strength due to increase in the size of central atom. These hydrides behave as reducing agents. And down the group reducing power increases. Boiling point of hydrides increases from PH₃ to BiH₃ but NH₃ has exceptionally high B.P. due to presence of intermolecular hydrogen bonding. The hydrides of group 15, due to availability of lone pair on central atom act as Lewis bases. The basic character decreases down the group. This is due to the decrease in density of electron on the central atom down the group as the volume of central atom increases down the group.

Illustration 3. Which of the following is/are not known?

PCl₃, AsCl₃, SbCl₃, NCl₅, BiCl₅

Solution: (i) NCl₅ (due to absence of d-orbitals)

(ii) BiCl₅ (Bi⁵⁺ being an oxidizing agent oxidizes Cl^- ion to Cl_­2 and Cl^- ion being a reducing agent reduces Bi⁵⁺ to Bi³⁺).

Ammonia (NH₃)

Nitrogen forms three different hydrides with hydrogen

(i) NH₃

(ii) Hydrazine, NH₂NH₂

(iii) Hydrazoic acid, N₃H

Amonia (NH₃) is one of the very important hydride found in traces in atmosphere. The atmospheric ammonia is formed by the bacterial decomposition of nitrogenous matter of plants and animals.

It was isolated in 1774 by priestley by the action of ammonium chloride and lime.

It was named alkaline air. Berthelot, in 1788, pointed out that ammonia is a compound of nitrogen and hydrogen. IN 1800, Davy established its formula NH₃.

Preparations

(i) $N{H}_{4}Cl+NaOHN{H}_3+NaCl+H_{2}O$

$2N{H}_{4}Cl+\underset{\left(slaked\lim e\right)}{Ca{\left(OH\right)}_2}\to 2N{H}_3+CaC{l}_2+2H_{2}O$

(ii) $2N{H}_{4}Cl+\underset{\left(Litharge\right)}{\mathrm{PbO}}\to 2N{H}_3+PbC{l}_2+H_{2}O$

(iii) By reacting nitrides with water, ammonia is obtained

$AlN+3H_{2}O\to Al{\left(OH\right)}_3+N{H}_3$

(iv) By reduction of nitrates and nitrites with Zn and caustic soda

$NaN{O}_3+8H\stackrel{Zn/NaOH}{\to }NaOH+N{H}_3+2H_{2}O$

$NaNO2+6H\stackrel{Zn/NaOH}{\to }NaOH+N{H}_3+H_{2}O$

(v) By hydrolysis of calcium cyanamide

$CaC{N}_2+3H_{2}O\to CaC{O}_3+2N{H}_3$

$N{H}_{4}Cl\to N{H}_3+HCl$

${\left(N{H}_4\right)}_{2}S{O}_4\to 2N{H}_3+H_{2}S{O}_4$

(vi) Urea on treatment with caustic soda

$N{H}_{2}CON{H}_2+2NaOH\to N{a}_{2}C{O}_3+2N{H}_3$

For drying of NH₃ gas, the common dehydrating agents like sulphuric acid or CaCl₂ or P₂O₅ cannot be used as these react with ammonia and hence it is cleared by quick lime (CaO).

$3N{H}_3+H_{2}S{O}_4\to {\left(N{H}_4\right)}_{2}S{O}_4$

$2N{H}_3+CaC{l}_2\to CaC{l}_2.2N{H}_3$

$3N{H}_3+P_4{O}_{10}\to {\left(N{H}_4\right)}_{3}P{O}_4$

Manufacture of NH₃

(i) Haber's process

This is the most important industrial method discovered by Fritz haber. This method involves the direct combination of nitrogen and hydrogen according to the following reaction

$N_2+2\rightleftharpoons H_{2}3N{H}_3+24.0Kcal$

(a) High pressure: Usually a pressure of 200 atmosphere is applied.

(b) Low temperature: The working temperature of 450 – 550°C is maintained.

(c) Catalyst: In order to speed up the reaction, a catalyst is used.

The following catalysts have been proposed for this purpose.

(i) Finely divided iron with some molybdenum as a promoter.

(ii) Finely divided nickel and sodalime deposited over pumic stone.

(iii) Finely divided osmium or uranium.

Process

Nitrogen is obtained through liquification followed by fractional evaporation of liquid air. Hydrogen is obtained by electrolysis of water.

Physical Properties

(i) It is colourless gas with a characteristic pigment odour.

(ii) It is highly soluble in water due to hydrogen bonding. The solubility of ammonia increases with increase of pressure and decrease with increase of temperature.

Chemical Properties

(i) Stability

It is highly stable. It decomposes into nitrogen and hydrogen at red heat or when electric sparks and passed through it.

$2N{H}_3\rightleftharpoons N_2+3H_2$

(ii) Combustion

$4N{H}_3+3O_2\to 2N_2+6H_{2}O$

(iii) Basic nature: It acts as Lewis base due to available lone pair on nitrogen. It accepts proton and form ammonium ion.

[NH₄⁺]

It forms salts with acids

$2N{H}_3+3O_2\to N_2+3H_{2}O$

$2N{H}_3+H_{2}S{O}_4\to {\left(N{H}_4\right)}_{2}S{O}_4\left(ammoniumsulphate\right)$

It's solution behaves as weak base.

$2KMn{O}_4+H_{2}O\to 2KOH+2Mn{O}_2+3\left[O\right]$

(iv) Oxidation

It can be oxidised to nitrogen by passing through over heated CuO or PbO.

$2CuO+2N{H}_3\to 2Cu+N_2+3H_{2}O$

$KCN+AgCN\to K\left[Ag{\left(CN\right)}_2\right]$

Both chlorine and bromine oxidize ammonia

$2N{H}_3+3C{l}_2\to N_2+6HCl$

$Si+2C{H}_{3}Cl\underset{{280}^o-{300}^{0}c}{\overset{Cu}{\to }}{\left(C{H}_3\right)}_{2}SiC{l}_2$

$\underset{\left(excess\right)}{8N{H}_3}+3C{l}_2\to N_2+6N{H}_{4}Cl$

When we take chlorine in excess an explosive substance NCl₃ is formed.

$2N{H}_3+3C{l}_2\to \underset{nitrogentrichloride}{NC{l}_3}+3HCl$

$\bullet$ Hypochlorites and hypobromites oxides ammonia to nitrogen.

$SiC{l}_4+2Mg\to Si+2MgC{l}_2$

$\bullet$ The oxidation of NH­₃ occurs with CaOCl₂ (bleaching powder)

$2N{H}_3+3CaOC{l}_2\to 3CaC{l}_2+N_2+3H_{2}O$

Thus ammonia acts as reducing agent.

The restricted oxidation of NH₃ can be done with air, when the mixture is passed over heated platinum gauge at 700 - 800°C.

$A{l}_4{C}_3+2H_{2}O\to 3C{H}_4+4Al{\left(OH\right)}_3$

This is the Ostwald's process and used for the manufacturing of HNO₃.

(v) Formation of amides

$2Na+\underset{\left(dry\right)}{2N{H}_3}\to 2NaN{H}_2+H_2$

(vi) Reaction of aqueous ammonia

When aqueous ammonia reacts with metals, metal hydroxides are formed which may be precipitated or remain dissolved in the form of the complex compound in excess of NH₄OH.

$AlC{l}_3+3N{H}_{4}OH\to \underset{\left(ppt.\right)}{Al{\left(OH\right)}_3}+3N{H}_{4}Cl$

$Cu{\left(OH\right)}_2+{\left(N{H}_4\right)}_{2}S{O}_4+2N{H}_{4}OH\to \underset{\begin{array}{c}Tetra\min ecoppersulphate\\ \left(Deepbluesolution\right)\end{array}}{\left[Cu{\left(N{H}_3\right)}_4\right]S{O}_4}+4H_{2}O$

$CdS{O}_4+4N{H}_{4}OH\to \underset{\begin{array}{c}cadmiumtetra\min esulphate\\ \left(Colourlesssolution\right)\end{array}}{\left[Cd{\left(N{H}_3\right)}_4\right]S{O}_4}+4H_{2}O$

.

$AgN{O}_3+N{H}_{4}OH\to \underset{\left(Whiteppt.\right)}{\mathrm{AgOH}}+N{H}_{4}N{O}_3$

$AgOH+2N{H}_{4}OH\to \underset{\left(so\mathrm{lub}le.\right)}{\left[Ag{\left(N{H}_3\right)}_2\left(OH\right)\right]}+2H_{2}O$

AgCl also dissolves in NH₄OH solution.

$AgCl+2N{H}_{4}OH\to \underset{\left(dia\min esilverchloride\right)}{\left[Ag{\left(N{H}_3\right)}_2\right]Cl}+2H_{2}O$

$ZnS{O}_4+2N{H}_{4}OH\to Zn{\left(OH\right)}_2+{\left(N{H}_4\right)}_{2}S{O}_4$

$Zn{\left(OH\right)}_2+{\left(N{H}_4\right)}_{2}S{O}_4+2N{H}_{4}OH\to \underset{\left(tetra\min ezincsulpahte\left(so\mathrm{lub}le\right)colourless\right)}{\left[Zn{\left(N{H}_3\right)}_4\right]S{O}_4}+4H_{2}O$

Nickel salt first gives a green ppt when dissolves in excess of NH₄OH.

$NiC{l}_2+2N{H}_{4}OH\to Ni{\left(OH\right)}_2+2N{H}_{4}Cl$

$Ni{\left(OH\right)}_2+2N{H}_{4}Cl+4N{H}_{4}OH\to \left[Ni{\left(N{H}_3\right)}_6\right]C{l}_2+6H_{2}O$

It forms a white ppt with mercuric chloride.

$HgC{l}_2+2N{H}_{4}OH\to \underset{\left(a\mathrm{mi}domercuricchloride\right)}{HgN{H}_{2}Cl}+N{H}_{4}Cl+H_{2}O$

It forms a gray ppt with mercurous chloride.

$HgC{l}_2+2N{H}_{4}OH\to \underset{gray}{\underbrace{Hg+HgN{H}_{2}Cl}}++N{H}_{4}Cl+H_{2}O$

(vii) Reaction with Nessler's reagent

$2KI+HgC{l}_2\to Hg{I}_2+2KCl$

$2KI+Hg{l}_2\to K_{2}Hg{I}_4$

Alkaline solution of K₂HgI₄ is called Nessler's reagent. This gives brown ppt with NH₃ called iodide of Million's base.

$2K_{2}Hg{I}_4+\stackrel{+}{N}{H}_4+3KOH\to \underset{brownppt}{H_{2}NHgOHgI}+7Kl+2H_{2}O$

Uses of Ammonia

1. Liquid ammonia is used in refrigeration.

2. Aqueous NH₃ is used in qualitative as well as qualitative analysis in laboratory.

3. In manufacture of nitric acid (ostwald's process), sodium bicarbonate (solvey process) and ammonium compounds. Ammonium sulphate, ammonium calcium phosphate.

Calcium ammonium nitrate, etc are used as fertilizers. Ammonium nitrate is used in certain explosive.

4. It is used as a clanging agent for removing greese.

5. Ammonia is also used in the manufacture of urea which is an excellent fertilizer of nitrogen.

Structure

NH₃ is a covalent molecule in which nitrogen is in sp³ hybridisation state.

But due to presence of one lone pair it acquires pyramidal structure.

The bond angle is 107.1ºnot 109º28' due to lp – bp repulsion.

Phosphine (PH₃)

It is the hydride of phosphorous and was discovered by Gengembre in 1783.

Laboratory Preparation

It is prepared by boiling yellow phosphorous with a concentrated solution of sodium hydroxide in an inert atmosphere.

$P_4+3NaOH+3H_{2}O\to \underset{sodiumhypophosphite}{3Na{H}_{2}P{O}_2}+P{H}_3$

Phosphorus also forms P₂H₄, beside PH₃.

$P_4+4NaOH+4H_{2}O\to 4Na{H}_{2}P{O}_2+2H_2$

$3P_4+8NaOH+8H_{2}O\to 8Na{H}_{2}P{O}_2+2P_2{H}_4$

A concentrated solution of NaOH is taken in a round bottom flask. Few pieces of yellow phosphorous are dropped in to it. Coal gas, oil gas or carbon dioxide is then bubbled through the flask to displace the air from the apparatus.

The flask is then heated. PH₃ is evolved. It is driven out along with a current of inert gas. As soon as the bubbles of gas come in contact with air, they catch fire spontaneously forming rings of smoke known as vortex rings. This combustion is due to the presence of highly inflammable phosphorous dihydride (P₂H₄)

$2P_2{H}_4+7O_2\to \underset{metaphosphoricacid}{4HP{O}_3}+2H_{2}O$

P₂H₄ can be removed from phosphine by the following methods

(i) By passing the evolved gas through a freezing mixture which condenses P₂H₄.

(ii) By passing the gas through HI. PH₃ is absorbed forming phosphonium iodide. This on treatment with caustic potash gives pure phosphine.

$P{H}_{4}I+KOH\to KI+H_{2}O+P{H}_3$

Physical Properties

(i) It is a colourless gas having unpleasant garlic – like odour or rotten fish odour.

(ii) It is heavier than air and sparingly soluble in water.

(iii) It liquefies at -89°C and solidifies at -134ºC.

(iv) It is poisonous in nature.

Chemical Properties

(a) Decomposition

$4P{H}_3\underset{orelectricspark}{\overset{\Delta }{\to }}{P}_4+6H_2$

(b) Combustibility

$P{H}_3+4O_2\to P_2{O}_5+3H_{2}O$

The spontaneous inflammability of phosphine at the time of preparation is due to the presence of highly inflammable phosphorous dihydride, P₂H₄. This property is used in making Holme's signal.

(c) Action of chlorine

Phosphine burns in the atmosphere of chlorine and forms phosphorous pentachloride.

$P{H}_3+4C{l}_2\to PC{l}_5+3HCl$

(d) Basic nature

Phosphine is neutral to litmus. However, it is a weak base, even weaker than ammonia. It reacts with HCl, HBr or HI to form phosphonium compounds.

$P{H}_3+HCl\to P{H}_{4}Cl$

$P{H}_3+HBr\to P{H}_{4}Br$

$P{H}_3+HI\to P{H}_{4}l$

(e) Action of nitric acid

In contact with nitric acid phosphine begins to burns.

$\left[2HN{O}_3\to H_{2}O+2N{O}_2+O\right]\times 8$

$2P{H}_3+8O\to P_2{O}_5+3H_{2}O$

$2P{H}_3+16HN{O}_3\to P_2{O}_5+16N{O}_2+11H_{2}O$

(f) Addition compounds

If forms addition compounds with anhydrous AlCl₃ and SnCl₄.

$AlC{l}_3+2P{H}_3\to AlC{l}_3.2P{H}_3$

$SnC{l}_4+2P{H}_3\to SnC{l}_4.2P{H}_3$

When PH₃ is passed through cuprous chloride solution in HCl, it forms an addition compound.

$C{u}_{2}C{l}_2+2P{H}_3\to 2CuCl.P{H}_3$

(g) Formation of phosphides

(i) When phosphine is passed through copper sulphate solution, a black ppt of cupric phosphide is formed

$3CuS{O}_4+2P{H}_3\to C{u}_3{P}_2+2H_{2}S{O}_4$

(ii) A black ppt of silver phosphide is formed when phosphine is circulated through silver nitrate solution.

$3AgN{O}_3+P{H}_3\to A{g}_{3}P+3HN{O}_3$

(h) The mixture of PH₃ and N₂O or PH₃ and NO explodes in presence of electric spark.

$P{H}_3+4N_{2}O\to H_{3}P{O}_4+4N_2$

Structure of phosphine

Phosphine is a covalent molecule. It has pyramidal structure like ammonia.

The bond angle H ¾ P ¾ H is 93°.

Uses

(i) For making Holme's signals.

(ii) For making smoke screens.

(iii) For making metallic phosphides.

2. Halides

Two types of halides are available for this group. One is MX₃ and another is MX₅ type.

MX₃: NCl₃, PCl₃, AsCl₃, SbCl₃ and BiCl₃

MX₅: PCl₅, AsCl₅, SbCl₅

(a) Tri halides

(i) All trihalides are covalent except BiF₃ which is ionic.

(ii) Like hydrides these trihalides have pyramidal structure and a central atom is sp³ hybridized.

(iii) These trihalides can be easily hydrolyses by water except NX₃.

$PC{l}_3+3H_{2}O\to H_{3}P{O}_3+3HCl$

$AsC{l}_3+H_{2}O\to A{s}_2{O}_3+6HCl$

$SbC{l}_3+H_{2}O\to SbOCl+2HCl$

$BiC{l}_3+H_{2}O\to BiOCl+2HCl$

(iv) The trihalides of P, As and Sb acts as Lewis acids and combine with Lewis bases.

$P{F}_3+F_2\to P{F}_5$

$Sb{F}_3+2F^{-}\to {\left[Sb{F}_5\right]}^{2-}$

(b) Penta halides

In this central atom attains sp³d hybridization and forms five covalent bonds with five chlorine atoms.

Penta halides have less thermal stability as compared to trihalides. All penta halides act as Lewis acids.

$PC{l}_5+C{l}^{-}\to {\left[PC{l}_6\right]}^{-}$

$SbC{l}_3+C{l}^{-}\to {\left[SbC{l}_4\right]}^{-}$

PCl₅ exists as molecule in gaseous state but in solid it exists as ${\left[PC{l}_4\right]}^{+}{\left[PC{l}_6\right]}^{-}$ and is ionic in nature. PBr₅, PI₅ also exist in the ionic form in solid state.

Illustration 4. Are all the five bonds of PCl₅ equivalent? Justify your answer.

Solution: PCl₅ has trigonal bipyramidal structure. It has three equivalent equatorial and two equivalent axial P – Cl bonds. However, due to greater bond pair – bond pair repulsions, the axial P – Cl bonds are longer and hence different from the three equatorial bonds.

4. Oxyacids

Oxides of N, P and As when dissolves in water to form oxy acids. Oxy acids of Sb and Bi are not stable.

Some important oxy acids

Oxyacids of Nitrogen$HN{O}_3,$Nitric acid

Nitric acid is also known as aqua fortis (meaning strong water) which was given by alchemists.

Methods of preparation

(a) Laboratory preparation

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

(b) Birkeland Eyde process

$N_2+O_2\rightleftharpoons 2NO-43,200calories$

The formation of nitric oxide is favoured by high temperature thus it is maintained at ${3000}^{o}C.$ The nitric oxide is cooled to ${1000}^{o}C$ as to prevent its decomposition. Nitric oxide further combines with$O_2.$

$2NO+O_2\to 2N{O}_2$

The vapours then passed through water when nitric acid is produced.

$2N{O}_2+H_{2}O\to HN{O}_3+HN{O}_2$

$3HN{O}_2\to HN{O}_3+2NO+H_{2}O$

(c) Ostwald's process (Modern process)

In this process the nitric acid is formed first when $N{H}_3$ and air passed over platinum gauze at $750-{900}^{o}C.$

$4N{H}_3+5O_2\underset{750-{900}^{o}C}{\overset{Pt}{\to }}4NO+6H_{2}O+21,600cal$

The nitric oxide is then oxidised to $N{O}_2$by air which is cooled to ${50}^{o}C$ and absorbed in water.

$2NO+O_2\to 2N{O}_2$

$3N{O}_2+H_{2}O\to 2HN{O}_3+NO$

Fuming nitric acid

When $N{O}_2$ is dissolved in conc. nitric acid it forms fuming nitric acid. It is brown in colour. It is obtained by distilling concentrated $HN{O}_3$ with a little starch.

Physical properties

HNO₃ is colourless fuming with pungent smell soluble in water. Nitric acid usually acquires yellow colour due to its decomposition by sunlight into $N{O}_2.$

$4HN{O}_3\stackrel{sunlight}{\to }4N{O}_2+2H_{2}O+O_2$

Chemical properties

(a) It is very strong acid and exhibit usual properties of acids.

$CaO+2HN{O}_3\to Ca{(N{O}_3)}_2+H_{2}O$

$N{a}_{2}C{O}_3+2HN{O}_3\to 2NaN{O}_3+H_{2}O+C{O}_2$

$NaOH+HN{O}_3\to NaN{O}_3+H_{2}O$

(b) Oxidising agent

$2HN{O}_3\to H_{2}O+2N{O}_2+O$

$2HN{O}_3\to H_{2}O+2NO+3O$

(i) Oxidation of non – metals

$S+6HN{O}_3\to H_{2}S{O}_4+6N{O}_2+2H_{2}O$

$C+4HN{O}_3\to H_{2}C{O}_3+4N{O}_2+2H_{2}O$

$2P+10HN{O}_3\to 2H_{3}P{O}_4+10N{O}_2+2H_{2}O$

$I_2+10HN{O}_3\to 2HI{O}_3+10N{O}_2+4H_{2}O$

(ii) Oxidation of metalloids

$2As+10HN{O}_3\to 3H_{3}As{O}_4+10N{O}_2+2H_{2}O$

$Sn+4HN{O}_3\to 2N{O}_2+2H_{2}O+Sn$

(iii) Oxidation of compounds

$S{O}_2+HN{O}_3\to H_{2}S{O}_4+2N{O}_2$

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

$6FeS{O}_4+3H_{2}S{O}_4+2HN{O}_3\to 3F{e}_2{(S{O}_4)}_3+2NO+4H_{2}O$

In this reaction NO is absorbed by ferrous sulphate and a dark brown ring of nitroso ferrous sulphate is formed. This is the brown ring test for nitrates.

(c) Action on metals

$HN{O}_3+metal\to Metalnitrate+H$

$2HN{O}_3+H\to 2N{O}_2+2H_{2}O$

$2HN{O}_3+6H\to 2NO+4H_{2}O$

$2HN{O}_3+10H\to N_2+6H_{2}O$

$2HN{O}_3+16H\to 2N{H}_3+6H_{2}O$

(i) $Mg+2HN{O}_3\to Mg{(N{O}_3)}_2+H_2$

(ii) $Mn+2HN{O}_3\to Mn{(N{O}_3)}_2+H_2$

Iron with dil. $HN{O}_3$

$4Fe+10HN{O}_3\to 4Fe{(N{O}_3)}_2+N{H}_{4}N{O}_3+3H_{2}O$

$4Fe+10HN{O}_3\to 4Fe{(N{O}_3)}_2+N_{2}O+5H_{2}O$

Iron with conc.$HN{O}_3$

$Fe+3HN{O}_3\to Fe{(N{O}_3)}_3+3H$

$\frac{[HN{O}_3+H\to N{O}_2+H_{2}O]\times 3}{Fe+6HN{O}_3\to Fe{(N{O}_3)}_3+3N{O}_2+3H_{2}O}$

conc. Ferric nitrate

(d) Action on organic compounds

Nitration

$HN{O}_3+2H_{2}S{O}_4\rightleftharpoons N{O}_2^{+}+2HS{O}_4^{-}+H_3{O}^{+}$

Illustration 5. In $N{O}_3^{-}$ ion, the number of bond pair and lone pair of electrons on nitrogen atom are

(A) 2, 2 (B) 3, 1

(C) 1, 3 (D) 3, 0

Solution: (D)

Oxyacids of phosphorus

Phosphorus forms two series of oxyacids namely; phosphorus acid and phosphoric acid.

(a) Phosphorus acid series

The series consists of p atom exhibiting oxidation state of +3. They act as reducing agents.

Special Note

(i) Meta is used for acid obtained by loss of H₂O molecule.

(ii) Pyro is used for the acid obtained from two molecules with a loss of H₂O.

(iii) Hypo is used for the acid having lower oxygen content than the parent acid.

Phosphorus acid $(H_{3}P{O}_3)$

Preparation

(i) $P_4{O}_6+6H_{2}O\to 4H_{3}P{O}_3$

(ii) $PC{l}_3+3H_{2}O\to H_{3}P{O}_3+3HCl$

Properties

(i) It is colourless crystalline compound. It melts at ${73}^{o}C.$ It is highly soluble in water.

(ii) $4H_{3}P{O}_3\underset{\Delta }{\overset{{200}^{o}C}{\to }}3H_{3}P{O}_4+P{H}_3$

(iii) $H_{3}P{O}_3\rightleftharpoons H^{+}+H_{2}P{O}_3^{-}\rightleftharpoons 2H^{+}+HP{O}_3^{-}$

$K_1={10}^{-1};K_2=2\times {10}^{-7}$

(iv) It acts as a strong reducing agent.

$H_{3}P{O}_3+H_{2}O\to H_{3}P{O}_4+2H$

It reduces

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

$AgN{O}_3+H\to Ag+HN{O}_3$

$2HgC{l}_2+2H\to H{g}_{2}C{l}_2+2HCl$

$I_2+2H\to 2HI$

$2KMn{O}_4+3H_{2}S{O}_4+5H_{3}P{O}_3\to K_{2}S{O}_4+2MnS{O}_4+5H_{3}P{O}_4+3H_{2}O$

It is dibasic acid. The phosphorus lies in $s{p}^3$ hybrid state.

(b) Phosphoric acid series

This series consist of acids having p in +5 oxidation state.

Meta phosphoric does not exist in monomeric form, but it exists as cyclometa phosphoric acid (HPO₃)₃ or

polymeta phosphoric acid (HPO₃)_n.

Orthophosphoric acid $(H_{3}P{O}_4)$

This acid is commonly called as phosphoric acid.

Preparation

(i) $P_4{O}_{10}+6H_{2}O\to 4H_{3}P{O}_4$

(ii) $PC{l}_5+4H_{2}O\to H_{3}P{O}_4+5HCl$

(iii) Laboratory preparation

$P+5HN{O}_3\to H_{3}P{O}_4+H_{2}O+5N{O}_2$

This reaction is carried out in the presence of a crystal of iodine. Iodine acts as a catalyst.

(iv) Industrial preparation

It is prepared from bone ash.

$\Rightarrow C{a}_3{(P{O}_4)}_2+3H_{2}S{O}_4\to 3CaS{O}_4+2H_{3}P{O}_4$

When bone ash is dissolved in $HN{O}_3$ and lead acetate it forms $P{b}_3{(P{O}_4)}_2.$

$P{b}_3{(P{O}_4)}_2+3H_{2}S\to 3PbS+2H_{3}P{O}_4$

Properties

(a) It is a transparent deliquescent crystalline solid. It melts at ${42.3}^{o}C.$ It absorbs water and forms colourless syrupy mass. It is highly soluble in water.

(b) $2H_{3}P{O}_4\stackrel{{250}^{o}C}{\to }{H}_4{P}_2{O}_7+H_{2}O$

Pyrophosphoric

acid

$H_{3}P{O}_4\underset{{600}^{o}C}{\overset{\Delta }{\to }}HP{O}_3+H_{2}O$

metaphosphoric

acid

$4H_{3}P{O}_4\stackrel{atredhot}{\to }{P}_4{O}_{10}+H_{2}O$

(c) Acidic nature

It is tribasic acid.

$H_{3}P{O}_4\rightleftharpoons H^{+}+H_{2}P{O}_4\rightleftharpoons -2H^{+}+HP{O}_4^{2-}\rightleftharpoons 3H^{+}+P{O}_4^{3-}$

(d) Reaction with $AgN{O}_3$

$3AgN{O}_3+H_{3}P{O}_4\to A{g}_{3}P{O}_4+3HN{O}_3$

yellow ppt.

(e) Reaction with $BaC{l}_2$

$3BaC{l}_2+2H_{3}P{O}_4\to B{a}_3{(P{O}_4)}_2+6HCl$

(f) Reaction with bromides and iodides

$3NaBr+H_{3}P{O}_4\to N{a}_{3}P{O}_4+3HBr$

$3NaI+H_{3}P{O}_4\to N{a}_{3}P{O}_4+3HI$

(g) On heating $H_{3}P{O}_4$ in presence of nitric acid with ammonium molybdate, a canary yellow precipitate of ammonium phosphomolybdate is formed.

$H_{3}P{O}_4+21HN{O}_3+12{(N{H}_4)}_{2}Mo{O}_4\to {(N{H}_4)}_{3}P{O}_{4}12Mo{O}_3+21N{H}_{4}N{O}_3+12H_{2}O$

This reaction is used to test $P{O}_4^{3-}$ ion.

Phosphorus atom lies in $s{p}^3$ hybrid state.

Illustration 6. Name the three series of salts formed when orthophosphoric acid reacts with sodium hydroxide.

Solution: NaH₂PO₄, Na₂HPO₄ and Na₃PO₄.

FERTILIZERS

For the proper development and natural growth of plants they require adequate supply of essential nutrients. Nitrogen, phosphorus and potassium are regarded as essential plant nutrients and are consumed in large quantities. To overcome the deficiency of these nutrients in soil additional materials containing N, P and K are added to it in the form of manures. Manures are of two types

(i) Natural and

(ii) Artificial

Artificial fertilizers are the chemical substances which are added to the soil as to make up the deficiency of essential elements are called fertilizers.

Nitrogenous fertilizers

The important nitrogenous compounds used as fertilizers are :

(i) Ammonium sulphate${(N{H}_4)}_{2}S{O}_4$

Preparation :

${(N{H}_4)}_{2}C{O}_3+Ca{(OH)}_2\to CaC{O}_3+2N{H}_3+2H_{2}O$

$2N{H}_{4}Cl+Ca{(OH)}_2\to CaC{l}_2+2N{H}_3+2H_{2}O$

$N{H}_{4}HS\stackrel{\Delta }{\to }N{H}_3+H_{2}S$

$2N{H}_3+H_{2}S{O}_4\to {(N{H}_4)}_{2}S{O}_4$

(ii) Calcium cyanamide, $CaC{N}_2$+ C (Nitrolim)

Preparation:

$Ca{C}_2+N_2\stackrel{{1100}^{o}C}{\to }CaC{N}_2+C$

Calcium cyanamide

In soil it converts to urea.

$CaC{N}_2+H_{2}O+C{O}_2\to CN.N{H}_2+CaC{O}_3$

cyanamide

$CN.N{H}_2+H_{2}O\to CO{(N{H}_2)}_2$

Urea

(iii) Calcium ammonium nitrate, $Ca{(N{O}_3)}_2.N{H}_{4}N{O}_3(CAN)$

Preparation

$N{H}_3+HN{O}_3\to N{H}_{4}N{O}_3$

$2HN{O}_3+CaC{O}_3\to Ca{(N{O}_3)}_2+H_{2}O+C{O}_2$

$Ca{(N{O}_3)}_2+N{H}_{4}N{O}_3\to Ca{(N{O}_3)}_2.N{H}_{4}N{O}_3$

CAN is hygroscopic so the pellets are, therefore, coated with calcium silicate as to protect from moisture.

(iv) Basic calcium nitrate, $Ca{(N{O}_3)}_2.CaO$

Preparation

$CaC{O}_3+2HN{O}_3\to Ca{(N{O}_3)}_2+H_{2}O+C{O}_2$

It is then mixed with lime.

$Ca{(N{O}_3)}_2+CaO\to Ca{(N{O}_3)}_2.CaO$

Basic calcium nitrate

Being highly deliquescent, it is packed in water proof bags.

(v) Urea, $N{H}_{2}CON{H}_2$

Preparation

$C{O}_2+2N{H}_3\to N{H}_{2}COON{H}_4$

Ammonium carbonate

$N{H}_{2}COON{H}_4\to N{H}_{2}CON{H}_2+H_{2}O$

Out of all, urea is having highest nitrogen content (46.6%). It does not even alter the pH of soil. It can be used for all type of crops.

Phosphatic fertilizers

(i) Calcium superphosphate or superphosphate of lime, $Ca(H_{2}P{O}_4).H_{2}O+2(CaS{O}_4.2H_{2}O)$

It is mainly prepared on a large scale by treating phosphorite or bone ash with calculated quantity of 70% $H_{2}S{O}_4.$

(ii) Nitrophosphate

Nitrophosphate or

Calcium superphosphate nitrate

This fertilizer contains both nitrogen and phosphorus, this makes it better.

(iii) Triple phosphate

This is another form of superphosphate.

$C{a}_3{(P{O}_4)}_2+4H_{3}P{O}_4\to 3Ca{(H_{2}P{O}_4)}_2$

It has the three times the amount of available $P_2{O}_5$ in comparison to superphosphate of lime and hence called triple superphosphate. It has about 42 – 46% of $P_2{O}_5.$

(iv) Phosphatic slag or Thomas slag

It is the by product of steel industry.

$4P+5O_2\to P_4{O}_{10}$

NPK fertilizers

Fertilizers containing N, P and K in suitable adjusted proportions are known as NPK fertilizers. These are obtained by mixing nitrogenous, phosphatic and potash fertilizers in suitable proportion.