NITROGEN AND ITS COMPOUNDS

Nitrogen forms about 78% by volume of the atmosphere. It can be found in combined form in trioxonitrate (V) of sodium and calcium. It can also be found in organic matters like proteins, urea and the vitamin B compounds.

 

Laboratory preparation of nitrogen

– From air: nitrogen can be obtained from air by removing the other constituents. The CO2 and O2 can be removed by passing the air through caustic soda and heated copper turning, respectively.

The nitrogen obtained by this method contains about 1% by volume of rare gases as impurities and is denser than pure nitrogen.

 

– From ammonium dioxonitrate (III) – NH4NO2: thermal decomposition of NH4NO2 yields pure nitrogen. The NH4NO2 should not be heated directly because it is unstable and decomposes exothermically. Hence, a mixture of NaNO2 and NH4Cl in a ratio of 7:5 is heated to yield NH4NO2 which in turn decomposes to form nitrogen and steam.

NaNO2(aq) + NH4Cl(aq)        →         NH4NO2(aq) + NaCl(aq)

NH4NO2(aq)     →           N2(g) + 2H2O(l)

 

– From ammonium heptaoxodichromate (VI): when (NH4)2Cr2O7 is heated, it decomposes to yield nitrogen. (NH4)2Cr2O7(s) N2(g) + Cr2O3(s) + 4H2O(l)

– From ammonia: When ammonia is oxidized by hot copper (II) oxide, nitrogen is liberated.

2NH3(g) + 3CuO(s) 3Cu(s) + 3H2O(g) + N2(g)

– From dinitrogen (I) oxide: When N2O is passed over red-hot copper, the gas is reduced to nitrogen. N2O(g) + Cu(s) CuO(s) + N2(g)

 

Industrial preparation

In the industries, nitrogen is prepared by the fractional distillation of liquid air. Carbon (IV) oxide is first removed from air, by passing the air through caustic soda before liquefaction. The liquid air (without CO2) is subjected to successive compression and cooling processes. During distillation, nitrogen gas first evolves at -1960C (at s.t.p) and is separated from oxygen which distils at -1830C (at s.t.p). The nitrogen is stored in steel cylinders and sold as liquid nitrogen or as compressed gas.

 

Physical properties of nitrogen

1. It is a colourless, odourless and tasteless gas
2. Pure nitrogen is slightly lighter than air
3. It is only slightly soluble in water
4. It has a boiling point of -1960C and melting point of -2100C.

 

Chemical properties of nitrogen

Nitrogen is a Group V element which forms covalent bonds in it compounds. It can also form ionic bonds with Group I and II elements by accepting 3 electrons to form nitride ion; N3-. The oxidation states vary from -3 to +5 in its various compounds. Nitrogen gas is unreactive under ordinary conditions because of the high bond energy of the triple bond (N≡N) in its atoms. However, at high temperature and pressure, nitrogen combines directly with hydrogen, oxygen, and certain metals.

1. Reaction with metals: Nitrogen combines directly with very electropositive metals to form nitrides, e.g

3Mg(s) + N2(g)   →        Mg3N2(s)

Mg3N2(s) + 6H2O(l)  warm →    3Mg(OH)2(aq) + 2NH3(g)

 

2. Reaction with non-metals: nitrogen combines reversibly with hydrogen to form ammonia. With oxygen, it combines directly at a very high temperature of about 2000°C or in the presence of a high-voltage electric spark to form a small amount of nitrogen (II) oxide. This reaction occurs in the atmosphere during lightning flashes. 

N2(g) + O2(g)    →        2NO(g)

 

Uses of nitrogen

– It is used in the manufacturing of ammonia, cyanide, cyanomide, and carbamide (an important fertilizer)

– Liquid nitrogen is used as cooling agent

– Because of its inert nature, it is used as a carrier gas in gas chromatography, provision of inert condition for certain industrial processes that involve easily oxidizable chemicals e.g. the manufacturing of transistors and annealing of metals.

– It is used as preservative to prevent rancidity in packaged foods

 

COMPOUNDS OF NITROGEN

AMMONIA

Ammonia is an essential chemical. It is a hydride of nitrogen.

 

Laboratory preparation of ammonia

Ammonia is prepared in the laboratory by heating any ammonium salt with a non-volatile base. Ammonium chloride and calcium hydroxide are usually employed because calcium hydroxide is cheap and not deliquescent like sodium hydroxide and potassium hydroxide. The two reactants must be properly grounded to provide maximum surface area for the reaction.

Ca(OH)2(s) + 2NH4Cl(s)          →          CaCl2(s) + 2H2O(l) + 2NH3(g)

The suitable drying agent for ammonia is quicklime (CaO). Silica gel may also be used.

Laboratory preparation of ammonia

 

Industrial preparation of ammonia

The industrial process for the production of ammonia is called Haber process. It involves mixing nitrogen

and hydrogen in the volume ratio 1:3, and passing the mixture:

– Over finely divided iron as a catalyst

– At a temperature of about 4500C and

– A pressure of about 200 atm.

The yield of ammonia is about 15%. The Ammonia is liquefied by cooling, while the unused gases are recirculated for further production.

N2(g) + 3H2(g)     →        2NH3(g)

An ammonia molecule has a pyramidal shape. The presence of a lone pair of electrons and the different electronegativity values of nitrogen and hydrogen make the ammonia molecule polar. Because of this, hydrogen bonds occur readily between ammonia molecules and between ammonia and water molecules.

 

Physical properties of ammonia

1. It is a colourless gas with choking smell
2. It is poisonous when in large quantity because of its effects on respiratory muscles.
3. It is an alkaline gas
4. It is about 1.7 times less dense than air
5. It is a very soluble gas. About 1200 volumes of ammonia dissolves in 1 volume of water at s.t.p.

This is due to the readiness with which it forms hydrogen bonds with water to give aqueous ammonia – NH3.H2O. The aqueous ammonia ionizes slightly to give ammonium ion and hydroxyl ion. The hydroxyl ion gives aqueous ammonia its weak alkaline property.

Chemical properties of ammonia

– Ammonia burns readily in oxygen, but not in air, with a greenish yellow flame to form water vapour and nitrogen;

4NH3(g) + 3O2(g)         →          6H2O(g) + 2N2(g)

In the presence of heated platinum catalyst however, ammonia reacts with air to produce

nitrogen (II) oxide;

4NH3(g) + 5O2(g) platinum →  4NO(g) + 6H2O(l)

 

– As a reducing agent

(i) It reduces heated copper (II) oxdie to metallic copper

2CuO(s) + 2NH3(g)               3Cu(s) + 3H2O(l) + N2(g)

(ii) It reduces chlorine to hydrogen chloride and nitrogen. The hydrogen chloride then reacts with excess ammonia to produce dense white fume of ammonium chloride.

2NH3(g) + 3Cl2(g)     →           6HCl(g) + N2(g)

6NH3(g) + 6HCl(g)     →           6NH4Cl(s)

Overall reaction ; 8NH3(g) + 3Cl2(g)  →         6NH4Cl(s) + N2(g)

The ammonia must be supplied in excess. If however, chlorine is in excess, nitrogen (II) chloride – NCl3 , an explosive and oily liquid will be formed.

NH3(g) + 3Cl2(g)              NCl3(l) + 3HCl(g)

– Reaction with carbon(IV) oxide: at about 1500C and 150 atm, ammonia reacts with carbon (IV)

oxide to produce urea, an important organic compound.

2NH3(g) + CO2(g)                (NH2)2CO(s) + H2O(l)

– Thermal decomposition; at temperature above 5000C or during prolong sparkling, ammonia decomposes to form nitrogen and hydrogen

– As a base: ammonia is a weak base because it can accept proton to form ammonium ion. It

reacts with acids to form ammonium salt.

NH3(g) + H+ (aq)     →         NH4+(aq)

2NH3(g) + H2SO4(aq)     →         (NH4)2SO4(s)

 

– As a precipitating agent: aqueous ammonia precipitates the insoluble hydroxides of metals from the solution of their salts. The following are the common insoluble hydroxides.

Pb2+(aq) + 2OH–(aq)   →        Pb(OH)2(s) – white ppt

Zn2+(aq) + 2OH–(aq)      →     Zn(OH)2(s) – white gelatinous ppt

Fe3+(aq) + 3OH–(aq)      →      Fe(OH)3(s) – reddish brown ppt

Cu2+(aq) + 2OH–(aq)    →        Cu(OH)2(s) – blue ppt

 

Some of these hydroxides, such as Cu(OH)2 and Zn(OH)2, dissolve in excess ammonia to form complex ions.

The complex ion formation is used as a confirmatory test for copper and zinc in qualitative analysis.

Cu(OH)2(s) + 4NH3(aq)        →       Cu(NH3)42+(aq) + 2OH–(aq)

                                                        Deep blue solution

Zn(OH)2(s) + 4NH3(aq)      →         Zn(NH3)2+ + 2OH–(aq)

                                                      Colourless solution

Test for ammonia

1. Action of litmus paper – hold a damp red litmus paper in the jar containing the unknown gas with a choking smell. If the litmus paper turns blue, then the gas is ammonia.
2. Action with hydrochloric acid – dip a glass rod in con. HCl and then insert it in the gas jar containing the unknown gas, white fumes are formed if the gas is ammonia.

 

Uses of ammonia

– It is used for softening temporary hard water

– Aqueous ammonia is used in laundry as solvent for removing grease and oil stains

– Aqueous ammonia is used as a refrigerant, although it is being replaced by less toxic compounds

– It is used in the manufacture of HNO3 and Na2CO3.

– It is used in the manufacturing of nitrogenous fertilizers like (NH4)2SO4, NH4NO3, (NH4)3PO4, and carbamide.

 

Tutorial questions

1. State the test for ammonia
2. Describe the preparation of nitrogen from air
3. State three uses of ammonia
4. Describe any three chemical properties of ammonia
5. Explain what happens when aqueous ammonia is added in excess to a solution of copper(II) salt.

 

 

NITROGEN AND ITS COMPOUNDS

 

Oxides of Nitrogen

Oxides are known for every oxidation state of nitrogen from +1 to +5. These are Dinitrogen(i)oxide, N2O, nitrogen(ii)oxide NO, nitrogen(iii)oxide, N2O3, nitrogen(iv)oxide, NO2 and nitrogen(v)oxide, N2O5.

Assignment

Give the preparation (Laboratory and industry), properties (physical and chemical), uses and test of all the five oxides of nitrogen.

 

Trioxonitrate (V) acid – HNO3

It is an important acid used in the laboratory and in the industries. It was previously known as aqua fortis; meaning strong water.

 

Laboratory preparation of HNO3

HNO3 can be prepared from any trioxonitrate(V) salt. KNO3 and NaNO3 are usually used because they are cheap. When these trioxonitarte (V) salts are reacted with conc. H2SO4, HNO3( a more volatile acid) is displaced.

KNO3(aq) + H2SO4(aq)      →          KHSO4(aq) + HNO3(g)

NaNO3(aq) + H2SO4(aq)    →         NaHSO4(aq) + HNO3(g)

 

Laboratory preparation of HNO3

 

Physical properties of HNO3

– It is a fuming liquid with a sharp, choking smell. The pure HNO3 is colourless but turns yellowish after some time due to the decomposition of some of the acid to yield NO2 which then dissolves in it. 

– Pure HNO3 is miscible with water in all proportions and forms a constant boiling mixture with it at 1210C. This mixture, consisting of about 68% HNO3, is the ordinary conc. HNO3 is found in the laboratory. Pure HNO3 containing 98% acid is known as fuming HNO3.

– Pure and ordinary conc. HNO3 are corrosive and must be handled with care.

 

Chemical properties of HNO3.

1. As an acid

(i) It neutralizes bases and alkalis to form metallic trioxonitrate (V) and water

NaOH(aq) + HNO3(aq)     →         NaNO3(aq) + H2O(l)

CaO(s) + 2HNO3(aq)     →         Ca(NO3)2(aq) + H2O(l)

(ii) it reacts with metallic trioxocarboantes(IV) to liberate carbon(IV) oxide.

CaCO3(s) + 2HNO3(aq)     →      Ca(NO3)2(aq) + H2O(l) + CO2(g)

(iii) Very dilute HNO3 (about 1%) reacts with calcium, magnesium or manganese to liberate hydrogen. Conc. HNO3 does not liberate hydrogen gas when reacted with metals, unlike other acids, because the hydrogen gas formed is immediately oxidized to water by the HNO3.

Mg(s) + 2HNO3(aq)     →        Mg(NO3)2 + H2(g)

 

2. As an oxidizing agent:

HNO3 is a strong oxidizing agent. it can undergo reduction in many ways to form various products like NO2, HNO2, NO, N2O, N2, N2OH, N2H4 and NH4+, depending on the concentration of the acid, the strength of the reducing agent and the temperature.

When dilute HNO3 acts as the oxidizing agent, the trioxonitrate (V) ion, NO3–

– is reduced to NO

e.g. NO3–(aq) + 4H+(aq) + 3e-     →          NO(g) + 2H2O(l)

When conc. HNO3 is the oxidizing agent, the NO3

– is reduced to NO2

NO3–(aq) + 2H+(aq) + e-   →         NO2(g) + H2O(l)

With non-metals: Hot conc. HNO3 oxidizes non-metals to their highest oxides which may then react with water to form the corresponding acids. At the same time, the acid itself is reduce to nitrogen (IV) oxide; 

e.g   C(s) + 4HNO3(aq)  heat  →    CO2(g) + H2O(l) + 4NO2(g)

S(s) + 6HNO3(aq)  heat  →    H2SO4(aq) + 2H2O(l) + 6NO2(g)

With metals:

(i) metals which are mild reducing agents like Cu, Pb, Hg and Ag when reacted with pure HNO3 are oxidized ton their trioxonitrates (V) while the HNO3 itself is reduced to NO2. However, with moderately conc. HNO3 (about 50%) these metals still react the same way but the HNO3 is reduced to NO e.g.

Cu(s) + 4HNO3(aq)   →           Cu(NO3)2(aq) + 2H2O(l) + 2NO2(g)

Pure

3Cu(s) + 8HNO3(aq)    →          3Cu(NO3)2(aq) + 4H2O(l) + 2NO(g)

(ii) Aluminium and iron do not react with the conc. HNO3. This may be due to the formation of a surface coating of oxide which is passive and stops any further reaction by the acid. Conc. HNO3 is thus usually transported in containers lined with aluminium or iron.

(iii) magnesium, zinc and iron reacts with dilute HNO3 to form the corresponding metallic trioxonitrates(V) and ammonium trioxonitrate (V) or nitrogen (I) oxide.

4Zn(s) + 10HNO(aq)     →          4Zn(NO3)2(aq) + 3H2O(l) + NH4NO3(aq)

(iv) Tin reacts with conc. HNO3 to form a hydrated oxide instead of a trioxonitrate (V)

Sn(s) + 4HNO3(aq)    →           SnO2.H2O(s) + H2O(l) + 4NO3(g)

(v) Gold and platinum are not attacked by HNO3 at all.

With reducing agents

(i) HNO3 oxidizes H2S to sulphur, some of which may be further oxidized to H2SO4. The HNO3 is itself reduced to NO2.

H2S(g) + 2HNO3(aq)     →         S(s) + 2H2O(l) + 2NO2(g)

(ii) iron (II) salts are oxidized to iron (III) salts by HNO3. The acid is then reduced to NO, which on

exposure to air is converted to NO2.

6Fe2+(aq) + 8H+(aq) + 2NO3–(aq)   →          6Fe3+(aq) + 4H2O(l) + 2NO(g)

3. Decomposition

HNO3 decomposes slowly at room temperature (especially in the presence of sunlight) and rapidly when heated to yield NO2 and O2.

2. Nitration reactions

HNO3 dissociates in the presence of conc. H2SO4 to form the nitryl cation (nitronium ion), NO2+. The nitronium ion can replace the hydrogen ion from many compounds like benzene, methyl benzene and phenol. The process is called nitration and is very important in the industries.

C6H6(l) + HNO3(aq)                C6H5NO2(l) + H2O(l)

                                                 Nitrobenzene

Uses of HNO3.

– It is used as an acid, oxidizing agent and nitrating agent in the laboratory.

– It is used as a rocket fuel

– It is used as an oxidizing agent in the production of important polymers like nylon and terylene.

– A mixture of hydrochloric acid and HNO3 in a ratio of 3:1 (known as aqua regia) is used as a solvent for gold and platinum.

– It is used as an important starting material in the production of many trioxonitrates (V) and organic nitro-compounds, which are used for the production of fertilizers, dyes, drugs and explosives e.g. methyl-2,4,6-trinitrobenzene – TNT)

 

TRIOXONITRATES (V)

Trioxonitrates (V) can be prepared by any of the general methods used for preparation of salts e.g. – Neutralization of HNO3 with appropriate alkali

– Action of HNO3 on metals, a metal oxide or a trioxocarbonate (IV) Precipitation method cannot be used for the preparation of trioxonitrates (V) because they are all soluble. Crystallization is the only suitable method for recovering the trioxonitrate (V) salt from the aqueous solution because they decompose easily on heating.

Action of heat on trioxonitrates (V)

All trioxonitrate (V) salts decompose on heating but the product depend on the position of the metal in

the electrochemical series. Below are the products of the decomposition of trioxonitrates (V).

Metal Decomposition of trioxonitrate (V)

K

Na

The Trioxonitrates (V) decompose to dioxonitrate (III) and oxygen

2KNO3(s) 2KNO2(s) + O2(g)

Ca

Mg

Zn

Pb

H Cu

The trioxonitrates (V) decompose to give the metallic oxide, oxygen and brown fumes of

nitrogen (IV) oxide, e.g.

2Zn(NO3)2(s)         →        2ZnO(s) + O2(g) + NO2(g)

Hg

Ag

Au

The trioxonitrate(V) decompose to give the metal, oxygen and brown fumes of nitrogen (IV) oxide e.g.

2AgNO3(s)      →           2Ag(s) + O2(g) + 2NO2(g)

Unlike the metallic trioxonitrate (V), ammonium trioxonitrate (V) decomposes on heating to produce nitrogen (I) oxide and water.

NH4NO3(s)        →          N2O(g) + 2H2O(l)

 

Action of H2SO4

All trioxonitrate (V) salts liberate HNO3 when heated with conc. H2SO4.

NO3–(aq) + H2SO4(aq)       →         HSO4–(aq) + HNO3(aq)

 

Formation of a brown ring

When conc. H2SO4 is added slowly down the side of a test tube containing an aqueous solution of a trioxonitrate(V) and FeSO4, the acid sinks to the bottom, and two layers are formed. A brown ring is formed at the junction of the two layers. This formation of brown ring can be used as a test for trioxonitrates (V) but it is unreliable because iodides and bromides also give coloured rings. Metallic radicals, which form insoluble tetraoxosulphate (VI) also interfere with the test.

 

Test for trioxonitrate (V)

– With H2SO4: warm the unknown solid with a little conc. H2SO4 in a test tube. If trioxonitrate (V) is present, some HNO3 will be condensed as oily drops on the upper part of the tube. On further heating, the acid decomposes to yield brown fumes of NO2.

H2SO4(aq) + NO3–(aq)        →          HNO3(l) + HSO4–(aq)

4HNO3(l)      →            4NO2(g) + O2(g) + 2H2O(l)

– With copper turnings: warm the mixture of the unknown solid and copper turnings with some

conc. H2SO4. If a trioxonitrate (V) is present, reddish brown fumes of NO2 will be formed.

H2SO4(aq) + NO3–(aq)   →             HNO3(aq) + HSO4–(aq)

Cu(s) + 4HNO3(aq)      →           Cu(NO3)2(aq) + 2NO2(g) + 2H2O(l)

 

Tutorial questions

1. How do we test for trioxonitrates (V) in the laboratory?
2. With the aid of balanced chemical reactions, show the thermal decomposition of;

(i) sodium trioxonitrate (V)

(ii) copper (II) trioxonitrate (V) and

(iii) Silver trioxonitrate (V)

3. State three uses of trioxonitrate (V) acid

4. Briefly describe the acidic properties of trioxonitrate (V) acid

5. How is HNO3 produced in the laboratory?