17.0     HYDROGEN AND ITS COMPOUNDS

Hydrogen (H), the lightest element with atomic number 1, constitutes 75% of the universe’s mass. It plays key roles in industry, energy (e.g., fuel cells), and forms compounds like water and hydrocarbons. With an electron configuration of 1s¹, hydrogen has oxidation states of +1 and -1. It was named by Lavoisier and first isolated by Cavendish in 1766.

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            17.1     ISOTOPES OF HYDROGEN

Hydrogen has three isotopes: Protium (¹H), the most common (99.98%), with a single proton; Deuterium (²H), with one proton and one neutron (0.015%), used in heavy water and fusion reactors; and Tritium (³H), radioactive with two neutrons, used in hydrogen bombs and fusion, and formed by cosmic rays or in reactors.

            17.2     INDUSTRIAL PREPARATION OF HYDROGEN

• FROM WATER-GAS (Bosch process):
  • Hydrogen can be commercially produced from water gas, a mixture of carbon monoxide and hydrogen. Methane or hot coke reacts with steam to form water gas. The water gas is then treated with more steam and passed over a catalyst (iron(III) oxide) to produce additional hydrogen and oxidize carbon monoxide to carbon dioxide. The carbon dioxide is removed by dissolving it in water, and any trace of carbon monoxide is absorbed by copper(I) methanoate in ammonia solution, yielding pure hydrogen.

CH₄ ​+ H₂​O → CO + 3H₂​​      ;           CO + H₂​O → CO₂​ + H₂​

• FROM THE CRACKING OF PETROLEUM
  • When large hydrocarbon molecules are broken up, hydrogen is formed in large quantity as a by-product e.g. the cracking of butane.

C₄H_{10 (g)} → C₄H_{8 (g)} + H_{2 (g)}

            17.3     LABORATORY PREPARATION OF HYDROGEN

• Action of zinc on dilute acid such as hydrochloric or tetraoxosulphate(VI) acid e.g using H₂SO₄;                                Zn+H₂​SO₄​ → ZnSO₄​+H₂​
• Action of sodium on cold water;        2Na_(s) + 2H₂O_(l) → 2NaOH_(aq) + H_2(g)
• Action of iron on steam;         3Fe_(s) + 4H₂O_(g) ⇌ Fe₃O_4(s) + 4H_2(g)

17.4     PROPERTIES OF HYDROGEN

PHYSICAL PROPERTIES

• It is a colourless, odourless and tasteless gas.
• It is neutral to moist litmus paper.
• It has a very low boiling point of -253^oC.
• It is relatively insoluble in water.
• It is the lightest known substance.

CHEMICAL PROPERTIES

• Reacts with oxygen to form water in a combustion reaction.
• Acts as a reducing agent, donating electrons in various chemical reactions.
• Reacts with many metals to form hydrides, such as sodium hydride (NaH).
• Hydrogen atoms in molecules can exchange with deuterium or tritium in specific reactions.
• Reacts with halogens to form hydrogen halides.

17.5     TEST FOR HYDROGEN

With a lighted wooden splinter, hydrogen gives a pop sound and burns with a pale blue flame. It is highly flammable! The pop is the sound of a small explosion. The test is carried out only with a small quantity of the gas.

            17.7     USES OF HYDROGEN

• It is used in the manufacturing of ammonia, hydrochloric acid and methanol.
• It is used in the hydrogenation of fats and oils to produce margarine and other food products.
• It is used for filling balloons.
• Liquid hydrogen is also used as a rocket fuel.
• It is used in oxy-hydrogen flames to produce high temperature that can melt metals etc.

            17.6     HYDRIDES

Hydrogen hydrides are compounds formed when hydrogen bonds with other elements and are classified into three types: ionic, covalent, and metallic hydrides. Ionic hydrides involve hydrogen bonding with alkali or alkaline earth metals (e.g., NaH), acting as strong reducing agents. Covalent hydrides form with non-metals (e.g., CH₄, NH₃, H₂O) and can exhibit polarity and hydrogen bonding. Metallic hydrides, formed with transition metals (e.g., PdHₓ), store hydrogen in metal lattices, making them useful for energy storage. Hydrides have broad applications, including reducing agents in synthesis, hydrogen generation, and potential fuel cell technologies.

17.8     CHLORINE AND ITS COMPOUNDS; HALOGEN FAMILY

Chlorine is the most important element in the halogen family. It does not occur as a free element due to its reactive nature; rather it is usually found in a combine state as chloride.

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            17.9     LABORATORY PREPARATION OF CHLORINE

• By the oxidation of concentrated HCl with strong oxidizing agents such as MnO₂ or KMnO₄.

MnO_2(s) + 4HCl_(aq) → MnCl_2(g) + H₂O_(l) + Cl_2(g)

• By heating concentrated H₂SO₄ with a mixture of NaCl and MnO₂.

MnO₂ + 2NaCl + 2H₂SO₄ → Na₂SO_{4(aq) ­}+ MnSO_4(aq) + H₂O_(l) + Cl_2(g)

            17.10   INDUSTRIAL PREPARATION OF CHLORINE

Chlorine is industrially manufactured via the electrolysis of brine and the chlorides of molten sodium, magnesium and calcium in the chlor-alkali process. The chlorine is then liquified and stored under high pressure in steel cylinders for industrial use.

            17.11   PROPERTIES OF CHLORINE

PHYSICAL PROPERTIES

• It is a greenish-yellowish gas with an unpleasant chocking smell.
• It is moderately soluble in water.
• It is a poisonous gas.
• It can be liquefied under a pressure of about 6atm.
• It is about 2.5 times denser than air.

CHEMICAL PROPERTIES

• It is very reactive.
• It displaces other halogens from solution of their acids and salts.
• It combines directly with other elements except O₂, N₂, C and the noble gases to form chlorides.
• It is a powerful oxidizing agent.
• It has a bleaching action.

17.12   TEST FOR CHLORINE

• It has an irritating smell and is greenish-yellow in colour.
• It turns blue litmus paper red and bleaches it.
• It turns starch-iodide paper dark-blue.

            17.15   USES OF CHLORINE

• It is a powerful germicide used in water purification, sterilization, disinfection and treatment of sewage.
• It is used as a bleaching agent.
• It is used in the production of plastics, PVC and synthetic rubber.
• It is used to manufacture hydrochloric acid, potassium trioxochlorate(V), sodium trioxochlorate(V) and important organic solvents such as CHCl₃, CCl_4, C₂HCl₃ etc.
• It is used in the manufacture of insecticide D.D.T and domestic antiseptics.

17.13   HYDROGEN CHLORIDE

• It exists as gas at a room temperature and dissolves in water to form hydrochloric acid.
• Hydrogen chloride is usually prepared in the laboratory by action of hot concentrated tetraoxosulphate(VI) acid on sodium chloride. Industrially, it is prepared by the direct combination of hydrogen and chlorine.
• It is a very soluble as gas which dissolves readily in water to form hydrochloric acid; it reacts with many metals to form chlorides.

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17.14   HYDROCHLORIC ACID

• Dilute hydrochloric acid behaves like a typical acid:
  • It reacts with metals to liberate hydrogen.
  • It neutralizes bases.
  • It liberates carbon(IV) oxide from hydrogentrioxocarbonates(IV) and trioxocarbonates(IV).
  • It forms insoluble chlorides with silver or lead salts.
• Concentrated hydrochloric acid is oxidized to chlorine by a strong oxidizing agent.

17.13   CHLORIDES

• Chlorides are prepared by:
  • The action of hydrogen chloride or hydrochloric acid on metals.
  • The action of hydrochloric acid on the oxides, hydroxides and trioxocarbonates(IV).
  • Double decomposition and,
  • Direct combination of metals on chlorine.
• Chlorides react with hot concentrated tetraoxosulphate(VI) acid to liberate hydrogen chloride; if strong oxidizing agent is also present, they liberate chlorine.

TEST FOR SOLUBLE CHLORIDES

• With AgNO₃, a white precipitate of silver chloride is formed which is soluble in excess aqueous ammonia.
  • With Pb(NO₃)₂, a white precipitate of lead(II) chloride is formed which dissolves on heating and reappears on cooling.

17.14   HALOGEN FAMILY

The term “halogen” comes from Greek words meaning “salt formers”. Besides chlorine, the halogen family includes fluorine, bromine, iodine and astatine. Fluorine is the most reactive element among the halogens. The halogens are found in group 7 of the periodic table. They tend to attain a stable chemical structure (octet structure) by receiving an electron. 

Halogens exist in different physical states at room temperature: fluorine and chlorine are gases, bromine is a liquid, and iodine and astatine are solids. Their colors range from pale yellow (fluorine), to greenish yellow (chlorine) to red (bromine) to dark purple/violet (iodine). All halogens are diatomic in their elemental form (e.g., F₂, Cl₂).

Fluorine is the most reactive and electronegative, followed by chlorine, bromine, and iodine, with reactivity decreasing down the group. Astatine is rare and radioactive, with limited applications. Halogens show a progressive gradation in properties down the group because of the increasing complexity of the atoms.

These elements are strong oxidizing agents, with wide-ranging industrial and biological applications. Chlorine is used in water purification and the production of PVC, while fluorine is important for toothpaste (as fluoride) and Teflon. Bromine is used in flame retardants, and iodine plays a critical role in thyroid health and medical imaging.

17.16   OXYGEN AND ITS COMPOUND

Oxygen is a chemical element with the symbol “O” and atomic number 8. It is a diatomic molecule, O₂, and makes up a significant part of Earth’s atmosphere (about 21%). Oxygen is crucial for the respiration of most living organisms, serving as a key component in the production of energy. The electronic configuration of oxygen is 1s² 2s² 2p⁴.

            17.17   LABORATORY PREPARATION OF OXYGEN

• 
  Heating potassium trioxochlorate(V) (KClO₃) with manganese(IV) oxide (MnO₂) as a catalyst. Potassium chlorate decomposes to potassium chloride and oxygen:

• 
  Decomposition of hydrogen peroxide (H₂O₂) using manganese(IV) oxide without heating:

17.18   INDUSTRIAL PREPARATION

In the industry, oxygen is mostly prepared by the liquefaction of air followed by fractional distillation of the liquefied air.  It may also be prepared by the electrolysis of water. The oxygen gas is dried, compressed and stored in a steel cylinder under a pressure of about 100atm.

            17.19   PROPERTIES OF OXYGEN

PHYSICAL PROPERTIES

• It is a diatomic gas, which is colourless, odourless and tasteless.
• It is neutral to moist litmus paper.
• It is slightly soluble in water.

CHEMICAL PROPERTIES

• It is strongly electronegative.
• It has an oxidation state of -2.
• It can form multiple bonds with itself.
• It combines with metals to form basic oxides and with non-metals to form acidic oxides.
• It accepts two electrons from an electropositive element to form a stable octet structure.
• It reacts with most hydrocarbons and compounds of carbon, hydrogen and oxygen to form carbon (IV) oxide and water.

17.20   TEST FOR OXYGEN

The presence of oxygen rekindles a glowing splinter. This is a similar case for N₂O; although there are distinguishable features such as:

• O₂ produces a reddish-brown fume of NO₂ on reaction with NO_.
• O_{2 ­}is an odourless gas while N₂O has a pleasant smell.

            17.24   USES OF OXYGEN

• It is required for respiration.
• It combines with ethyne to produce oxy-ethyne flame.
• Liquid oxygen and fuels are used as propellants for space rockets.
• It is used in the steel industry for the removal of impurities from pig iron.
• It is used in the manufacturing of chemical compounds like H₂SO₄, HNO₃, CH₃COOH etc.

17.21   OXIDES

Oxides are formed when other elements combine with oxygen. They can be classified as:

• ACIDIC OXIDES – Oxides of non-metals; React with water to form acids i.e they have acidic properties. Examples include CO₂. SO₃ etc.
• BASIC OXIDES – Oxides of metals; Some do react with water to form alkali. In general, they possess basic properties. Examples include Na₂O, PbO etc.
• AMPHOTERIC OXIDES – These are metallic oxides that behave both as acidic and basic oxides. Examples include ZnO, Al₂O₃ etc.
• NEUTRAL OXIDES – They are neither acidic nor basic. Examples include H₂O, CO and N₂O.
• PEROXIDES – They are oxides containing a higher proportion of oxygen than the ordinary oxides.

17.22   HYDROGEN PEROXIDE

Hydrogen peroxide is a pale blue liquid that appears colorless in dilute solutions. It is a chemical compound with the formula H₂O₂, consisting of two hydrogen atoms and two oxygen atoms. It is known for its powerful oxidizing properties and is widely used in various industries and households. It is prepared by the action of dilute acid on peroxides of certain metals.

            17.23   OZONE (O₃)

It is an unstable allotrope of oxygen formed from atmospheric oxygen by lighting flashes (also prepared by passing electric discharge through atmospheric oxygen). It has the same chemical properties as oxygen except that it is more reactive. It’s a very powerful oxidizing agent and decomposes to ordinary oxygen on heating. The ozone layer in the atmosphere acts as a protective shield by preventing too much ultraviolet radiation from falling on earth and harming living organisms.

17.25   SULPHUR AND ITS COMPOUND

Sulphur is the second member of group VI in the periodic table. It is a yellow non – metallic solid element which make up 0.1% of earth crust. It occurs in USA, Poland, japan and New Zealand. In other instances, naturally occurring sulphides are iron pyrites (FeS), copper pyrites (CuFeS) Galena (PbS), zinc blende (ZnS), cinnabar (HgS). Sulphur is present in petroleum and some protein building foods.

Sulphur belongs to group 6 of the periodic table and its atomic number is 16. The electronic configuration is 1S² 2S² 2P⁶ 3S² 3P⁴ and has a valency of 2.

            17.26   EXTRACTION OF SULPHUR (Frasch Process)

Sulphur is extracted by Frasch process using Sulphur pump which is made up of three concentric tubes sunk to the bottom of Sulphur deposit. Super-heated water at a temperature of 1600 c and a pressure of 18 atmospheres is forced down the outer tube. The water melts the Sulphur which collects at the bottom. Hot compressed air is forced down the central tube and this forces the molten Sulphur to the surface through the middle tube. The Sulphur at this stage is about 99.5% pure and is allowed to solidify and dry.

            17.27   ALLOTROPES OF SULPHUR

• Rhombic Sulphur (α- Sulphur) – Stable at room temperature.
• Monoclinic/Prismatic Sulphur (β-Sulphur).
• Amorphous Sulphur (δ- Sulphur).
• Plastic Sulphur.

17.28   PROPERTIES OF SULPHUR

PHYSICAL PROPERTIES

• Sulphur is a yellow solid that exists in two forms – amorphous and crystalline.
• It is insoluble in water but soluble in carbon (Iv) sulphide and methylbenzene.
• Its melting point is 119^oC and boiling point is 444^oC.
• The density depends on the allotropic form.
• It is a bad conductor of heat and electricity.

CHEMICAL PROPERTIES

• When sulphur is heated in absence of air, it combines directly with many metals.
• It burns in plentiful supply of oxygen with bright blue flame to form Sulphur (iv) oxide.
• Sulphur reacts with hydrogen at high temperature to form hydrogen sulphide.
• Sulphur is readily oxidized when warmed H₂SO₄ to sulphur (iv) oxide.

17.29   HYDROGEN SULPHIDE

• It is a colourless, acidic gas which smells like rotten eggs.
• It is prepared by the action of a dilute acid on a metallic sulphide. It is a strong reducing agent, which is usually oxidized to sulphur in redox reactions.
• It is used in the analysis of sulphur as it forms insoluble, coloured sulphides with solutions of the salts of many metals. The colour of the sulphides formed, together with their varying solubilities in acids and alkalis, help in the identification of the metal ions.

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TEST FOR HYDROGEN SULPHIDE – Beside its characteristic smell, it turns filter paper moistened with lead(II) ethanoate or lead(II) trioxonitrate(V) black.

            17.30   SULPHUR (IV) OXIDE

• It is prepared by:
  • The burning of sulphur or metallic sulphides.
  • Heating of K₂SO₃ or Na₂SO₃ with dilute H₂SO₄ or HCl acid.
  • Heating of concentrated SO_4­ with copper turnings.
• Properties:
  • An acidic gas with a smell of burning firecrackers.
  • Dissolves in water to form H₂SO₃ acid; neutralizes bases.
  • Combines directly with certain metallic oxides
  • With oxidizing agents, it acts as a strong reducing agent in redox reactions.
  • With reducing agents, it acts as a strong oxidizing agent in redox reactions.
  • It bleaches dyes by its reducing action.

TEST FOR SULPHUR (IV) OXIDE – Besides its characteristic smell, it bleaches straws, flowers, fabrics etc.; turns acidified KMnO₄ from purple to colourless and acidified K₂Cr₂O₇ from orange to green.

            17.31   SULPHUR (VI) OXIDE

It is prepared from sulphur(IV) oxide and oxygen in the presence of a heated catalyst like platinized asbestos or vanadium(V) oxide. It is an acidic gas and dissolves exothermically in water to form tetraoxosulphate(VI) acid.

            17.32   TRIOXOSULPHATE (IV) ACID

It is a weak dibasic acid formed when sulphur(IV) oxide dissolves in water. On exposure to air, it slowly becomes tetraoxosulphate(VI) acid. The trioxosulphate(IV) ion, SO₃^2-, in the acid is responsible for the reducing and bleaching action of sulphur(IV) oxide in aqueous solution.

            17.33   TRIOXOSULPHATES (IV)

• They are prepared by:
  • Neutralization of alkalis with sulphur(IV) oxide.
  • Precipitation of corresponding metallic salt solution by passing in sulphur(IV) oxide.
• Only the trioxosulphates(IV) of NH₃, K, Na and Ca are soluble in water. Trioxosulphates(IV) liberate sulphur(IV) oxide when warmed with acids; and on exposure to air, they are slowly converted to tetraoxosulphate(VI) ions, SO₄^2-.

17.34   TETRAOXOSULPHATE(VI) ACID

• H₂SO₄ is one of the most important chemical compounds. Industrially, it is manufactured by the Contact process. This process involves the catalytic combination of sulphur(IV) oxide and oxygen to form sulphur(VI) oxide which is then dissolved in concentrated tetraoxosulphate(VI) acid to form oleum, H₂S₂O₇. This is then diluted with water to produce 98% tetraoxosulphate(VI) acid.
• Tetraoxosulphate(VI) acid is a strong dibasic acid. It neutralizes alkalis; liberates hydrogen with metals and reacts vigorously with bases and hydroxides. The concentrated H₂SO₄ acid is hygroscopic and behaves as a dehydrating agent. It forms esters when reacted with alcohol. It reacts with reducing agents like metals to liberate hydrogen sulphide, H₂S.
• In the industry, it is used in large quantities for the manufacture of fertilizers, soaps, detergents, and also oxidizes metals releasing hydrogen sulphide. It is also used in the manufacture of pigments, textiles, glass, batteries, and explosives and in petroleum refining and metal manufacturing processes.

17.35   TETRAOXOSULPHATE(VI)

• Tetraoxosulphates(VI) are normal salts prepared by the action of dilute tetraoxosulphate(VI) acid on the metals or the oxides, hydroxides or trioxocarbonates(IV) of metals. Insoluble tetraoxosulphates(VI) are prepared by the double decomposition of the appropriate salts or acids.
• Tetraoxosulphates(VI) form salts or alums. The tetraoxosulphates(VI) of sodium, potassium, and ammonium are stable to heat but the tetraoxosulphates(VI) of metals lower down the electrochemical series tends to decompose to the metallic oxides, sulphur(IV) oxide and sulphur(VI) oxide on heating.

TEST FOR TETRAOXOSULPHATES(VI) – On reacting tetraoxosulphate(VI) with barium chloride, they form a white precipitate of barium tetraoxosulphate(VI), which is insoluble in excess hydrochloric acid.

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17.36   NITROGEN AND ITS COMPOUNDS

Nitrogen which occurs in Group 5 of the periodic table has five valence electrons. It can exist in oxidation state of -3 to +5 in its compounds. It makes up 78% by volume of the atmosphere. Nitrogen occurs in the free form in air; in the combined form as trioxonitrates(V) and ammonium salts; and in organic matter.

            17.37   LABORATORY PREPARATION OF NITROGEN

Nitrogen is prepared in the laboratory by heating ammonium nitrite (NH₄NO₂​) or a mixture of ammonium chloride (NH₄Cl) and sodium dioxonitrite(II) (NaNO₂​).

NH₄NO₂​ ​→ N₂ ​+ 2H₂​O                      ;           NH₄Cl + NaNO₂ → N₂ ​+ NaCl + 2H₂​​O

It can also be prepared by the oxidation of ammonia or reduction of certain oxides of nitrogen.

            17.38   INDUSTRIAL PREPARATION OF NITROGEN

Industrially, nitrogen is prepared by the fractional distillation of liquid air. It involves these processes:

• Liquefaction of air – Air is compressed and cooled to form a liquid.
• Fractional distillation – Liquid air is slowly warmed. Nitrogen (boiling point −195.8^oC) is separated first, followed by oxygen and other gases.

I7.39    PROPERTIES OF NITROGEN

PHYSICAL PROPERTIES

• It is colourless, odourless and tasteless.
• Pure nitrogen is lighter than air.
• Slightly soluble in water.
• Melting point – 210^OC and boiling point is -196^OC.

CHEMICAL PROPERTIES

• It reacts with very electropositive metals to form nitrides.
• It reacts with non-metals like hydrogen and oxygen to form ammonia and oxides respectively.

17.40   USES OF NITROGEN

• It is used industrially to manufacture ammonia, carbamide (an important fertilizer), cyanide, cyanamide etc.
• Liquid nitrogen is used as a cooling agent.
• It is used as preservative in packaged foods to prevent rancidity.
• It is used as a carrier gas in gas chromatography due to its inert nature.

17.41   AMMONIA

• NH₃ is a compound of nitrogen and hydrogen which is gaseous at room temperature. Ammonia is prepared industrially by the Haber process, and in the laboratory by heating of slake lime and ammonium chloride.
• The properties of ammonia are as followed:
  • It is an alkaline gas which forms dense white fumes with hydrogen chloride gas.
  • It burns in oxygen to yield nitrogen, but in the presence of a catalyst, it yields nitrogen(II) oxide.
  • It reduces copper(II) oxide to metallic copper.
  • It reduces chlorine to chlorides.
• Aqueous ammonia is a weak alkali which neutralizes acids; it precipitates hydroxides of metals (except copper and zinc); and it sometimes forms complex salts when in excess.

• TEST FOR AMMONIA
  • It has a characteristic pungent smell.
  • It turns red litmus paper blue (confirmatory test).
  • Forms dense white fumes with hydrogen chloride gas.

17.42   AMMONIUM SALTS

Ammonium salts are prepared by the neutralization of acids with aqueous ammonia. They are white crystal solids which are water soluble. They release ammonia with strong alkalis (test for ammonium salts) and they are also decomposed by heating ammonia, nitrogen or oxides of nitrogen. Ammonium salts are extensively used as fertilizers and to manufacture explosives.

            17.43   OXIDES OF NITROGEN

The important oxides of nitrogen are:

• Nitrogen(I) oxide (N₂O) – Colorless gas, neutral oxide
• Nitrogen(II) oxide (NO) – Colorless gas, turns brown in air due to oxidation to nitrogen(IV) oxide.
• Nitrogen(IV) oxide (NO₂) – Brown gas, acidic oxide, dissolves in water to form nitric acid.

All oxides of nitrogen are all reduced to nitrogen by heated metals.

            17.44   TRIOXONITRATE(V) ACID

• Trioxonitrate(V) acid, also known as nitric acid, is prepared in the laboratory by the action of concentrated tetraoxosulphate(VI) acids on trioxonitrates(V), and industrially by the catalytic oxidation of ammonia with excess air.
• As an acid:
  • It neutralizes bases.
  • It liberates CO₂ when reacted with trioxocarbonates(IV).
  • It liberates hydrogen from magnesium.
• As an oxidizing agent:
  • It oxidizes non-metals to their highest oxides.
  • It oxidizes metals to the oxides or trioxonitrate(V), and is itself reduced to nitrogen(II) oxide.
  • It undergoes redox reaction with reducing agents.
  • It decomposes at room temperature liberating oxygen and nitrogen(IV) oxide.
• HNO₃ is used as a nitrating agent. It is also used in the manufacture of trioxonitrates(V) and oragno-nitrate compounds which are used as fertilizers, dyes and explosives.

17.45   TRIOXONITRATES(V)

• Nitrates are decomposed by heat, giving off oxygen and nitrogen(IV) oxide, except sodium, potassium and ammonium trioxonitrates(V); all are crystalline, water-soluble solids; and they produce trioxonitrate(V) acid when they react with concentrated tetraoxosulphate(VI) acid.

• TESTS FOR TRIOXONITRATE(V)
• It gives a brown ring with iron(II) tetraoxosulphate(VI) and concentrated tetraoxosulphate(VI) acid.
• It gives off nitrogen(IV) oxide with copper turnings in concentrated tetraoxosulphate(VI) acid.
• Produces oily trioxonitrate(V) acid droplets with concentrated tetraoxosulphate(VI) acid.

17.46   NITROGEN CYCLE

The nitrogen cycle is crucial for the cycling and availability of nitrogen in various forms for living organisms. It involves several key processes:

• Nitrogen Fixation – Conversion of atmospheric nitrogen (N₂) into ammonia (NH₃) or ammonium ions (NH₄⁺) by nitrogen-fixing bacteria. This can occur in the soil or in symbiotic relationships with plants (like legumes).
• Nitrification – Ammonia (NH₃) produced by nitrogen-fixing bacteria is converted into nitrites (NO₂⁻) and then nitrates (NO₃⁻) by nitrifying bacteria. This process makes nitrogen available for plant uptake.
• Assimilation – Plants and other organisms take up nitrogen in the form of nitrates (NO₃⁻) or ammonium ions (NH₄⁺) to build proteins and nucleic acids.
• Ammonification – Decomposition of organic nitrogen compounds into ammonia (NH₃) and ammonium ions (NH₄⁺) by decomposer bacteria.
• Denitrification – Conversion of nitrates (NO₃⁻) back into atmospheric nitrogen (N₂) by denitrifying bacteria under anaerobic conditions, completing the cycle.

This cycle is essential for maintaining the balance of nitrogen in ecosystems, ensuring that plants and animals have access to this vital nutrient.

            17.47   CARBON AND ITS COMPOUNDS

Carbon forms the largest number of compounds, next only to hydrogen. Carbon is a non-metallic element and the first member of group 4 of the periodic table. It is greatly abundant in the earth’s crust. Carbon occurs in the free state as well as in the combined state. In its elemental form, carbon occurs in nature as diamond and graphite. Coal, charcoal and coke are impure forms of carbon. In the combined state, carbon is present as carbonate in many minerals, such as hydrocarbons in natural gas, petroleum etc. In air, carbon dioxide is present in small quantities, (0.03%).

            17.48   ALLOTROPES OF CARBON

• The crystalline form – Diamond and Graphite.
• The non-crystalline (amorphous) form – Coal, Coke, Charcoal, Lamp black, Carbon black(Soot) etc.

17.49   COMPARISM BETWEEN DIAMOND AND GRAPHITE

            17.50   PROPERTIES OF CARBON

PHYSICAL PROPERTIES

• They are odourless and tasteless.
• They have high melting point.
• They are insoluble in common solvents like water, alkalis, acids etc.

CHEMICAL PROPERTIES

• Carbon can form long chains and complex structures (e.g., hydrocarbons) due to its ability to bond with itself.
• Carbon burns in oxygen to produce carbon dioxide (CO₂) if there’s sufficient oxygen or carbon monoxide (CO) if oxygen is limited.
• Carbon exhibits oxidation states ranging from -4 (in methane, CH₄) to +4 (in carbon dioxide, CO₂).
• It combines directly with other elements at high temperatures.
• It is a strong reducing agent – reduces metallic oxides to metals, reacts with strong oxidizing agents etc.

17.51   CARBON(IV) OXIDE

• CO₂ is a colourless, odourless gas, which is denser than air and soluble in water to form carbonic acid.
• It is prepared in the laboratory by the action of dilute acids on trioxocarbonates(IV) or hydrogen trioxocarbonates(IV).
• It is used in fire extinguishers, carbonated drinks, photosynthesis in plants etc.
• When trioxocarbonate(IV) salts are heated, they decompose to yield CO₂.

• TEST FOR CARBON(IV) OXIDE – It turns lime water (calcium hydroxide solution) milky and when bubbled in excess, the lime water becomes clear again.

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17.52   CARBON(II) OXIDE

• CO is a colourless, odourless and poisonous gas.
• It is prepared by reducing CO₂ with red-hot carbon or by the dehydration of methanoic or ethanedioic acid by concentrated tetraoxosulphate(VI) acid.
• It can combine with hemoglobin in blood, preventing oxygen transport (leading to carbon monoxide poisoning).
• CO is also produced from incomplete combustion of fuels, exhaust fumes, charcoal fires etc.
• It is a strong reducing agent used in the extraction of metals from their oxides.

17.53   COAL

• Coal is a black or brownish sedimentary rock formed from decomposed plant matter. A wide variety of substance can be obtained from it through a process known as the destructive distillation of coal.
• There are four types of coal – Anthracite coal (hard, highest carbon content), Bituminous coal (medium carbon content), Lignite coal (low carbon content) and Peat coal (soft, least carbon content).
• From the destructive distillation of coal (a chemical process in which decomposition of unprocessed material is achieved by heating it to a high temperature in the absence of air), the following are obtained accordingly:

COAL = coke (solid) → ammoniacal liquor (liquid) → coal tar (liquid) → coal gas (gas)

• Coal is a crucial energy source for electricity generation, industry, and steel production. Its affordability and abundance have driven economic growth, but it contributes significantly to environmental issues like pollution and climate change. This has sparked global efforts to transition to cleaner, more sustainable energy alternatives.

COKE – The solid residue, coke with its 98% carbon by mass burns smoothly without smoke and with high calorific value. It is used in the manufacture of gaseous fuels such as producer gas and water gas. It is also used as a reducing agent in metallurgy where it reduces metallic oxides to their respective metals e.g. in the blast furnace for the extraction of iron from its ore.

GASEFICATION OF COKE – It is a process by which gaseous fuels are produced, when coke is heated and air or water is passed through it. Examples of gaseous fuels are:

• Water gas – Produced when steam is passed over a white-hot coke. It contains 50% of CO and 50% of H₂.

C_(s) + H₂O_(g) → CO_(g) + H_2(g)

• Producer gas – Produced when air is passed over red-hot coke. It contains 67% of N₂ and 33% of CO.

O2 + N_2(g) + 2C_(s) ​→ 2CO_(g) + N₂ + heat

• Synthesis gas – Produced when methane is passed through oxygen.

CH₄ + O₂ → CO + 2H₂

AMMONIACAL LIQUOR – It consists essentially of ammonium compounds and benzene. The ammonium compounds can be used in the manufacture of nitrogenous fertilizers while the benzene is used for the manufacture of pharmaceutical products and as a solvent.

COAL TAR – It is a mixture of more than 200 different substances which can be separated by fractional distillation. Most of these e.g. toluene, phenol and naphthalene are used in the synthesis of important commercial products like dyes, paints, insecticides, drugs, plastics and explosives.

COAL GAS – This is the volatile compound which contains about 50% hydrogen, 30% methane, 10% carbon (II) oxide and small amount of other gases e.g. ethane and hydrogen sulphide.

            17.54   DESTRUCTIVE DISTILLATION OF WOOD

Wood has a higher % of H_2­ and N₂ but lower % of carbon than coal. The four products from the destructive distillation of wood are:

Wood = wood charcoal (solid) → pyroligneous acid (liquid) → wood tar (liquid) → wood gas (gas)

            17.55   CARBON CYCLE

The carbon cycle is the process by which carbon is exchanged between the Earth’s atmosphere, oceans, soil, and living organisms. It is crucial for regulating the Earth’s climate and sustaining life, as carbon is a fundamental element in organic molecules such as carbohydrates, proteins, and fats. The cycle involves several key processes that move carbon through different reservoirs, including the atmosphere, terrestrial ecosystems, oceans, and the Earth’s crust. It shows that carbon is being continuously circulated in nature as carbon(IV) oxide. The quantity of carbon(IV) oxide remains constant due to a balance of processes which use up carbon(IV) oxide and those which liberate it.