Grade 10 Chemistry Notes: Metals, Nonmetals, and Their Production

5.2 General Properties of Metals and Production of Some Metals

• Students should be able to:
  • Mention general properties of metals.
  • Describe the uses of some common metals.
• This section covers general characteristics and extraction of metals, uses of some common metals, alloys, and production of Aluminum, Copper, and Iron.

5.2.1 Properties and Extraction of Metals

• Activity 5.1: Discuss metal properties in groups.
  • A metal which is liquid at room temperature.
  • A metal that exists in the gaseous state at room temperature.
  • A metal which is the best conductor of heat.
  • A metal which is the poorest conductor of heat.
  • A metal which can be cut with a knife.
• About 80% of known elements are metals.

A. Some Physical Properties of Metals

• Lustrous Appearance:
  • Metals have a shining appearance and can be polished.
  • Gold is shining yellow, copper is brown.
  • Iron, aluminum, zinc, and silver are lustrous grey or silvery.
• Malleability:
  • Metals can be beaten into thin sheets.
  • Examples: Aluminum foils, gold and silver ornaments.
• Ductility:
  • Metals can be drawn into wires.
  • Examples: Copper, gold, iron, and silver.
• Hardness and Tensile Strength:
  • Metals are generally hard and have tensile strength, except lithium, potassium, sodium.
• Density:
  • Metals generally have a high density except lithium, potassium, sodium.
• Sonorous:
  • Metals produce a metallic sound when struck (e.g., school bell).

B. Chemical Properties of Metals

• Positive Valency:
  • Metals possess positive valency and have a tendency to lose electrons.
  • $M(g) \rightarrow M^{n+}(g) + ne^-$
• Valence Electrons:
  • Metals have 1, 2, or 3 valence electrons.
• Reducing Agents:
  • Metals are oxidized by losing electrons and act as reducing agents.
• Oxides Formation:
  • They mostly form basic oxides and some amphoteric oxides.
• Chlorides Formation:
  • They form chlorides that are true salts and electrovalent.
• Hydrides Formation:
  • They form hydrides which are ionic, unstable, and reactive.
• Reaction with Acids:
  • They usually replace hydrogen from dilute non-oxidizing acids like $HCl$ and $H<em>2SO</em>4$. Exceptions are copper, silver, and gold.

C. Reactivity Series of Metals

• Activity Series:
  • Arrangement of metals in decreasing order of activity.
  • The most active metal is at the top, and the least active metal is at the bottom.
  • Hydrogen is included for comparison.
• Reaction with Dilute Acids:
  • Metals above hydrogen (potassium, sodium, calcium, magnesium) liberate hydrogen gas when treated with dilute acids.
  • Metals below hydrogen (copper, silver, gold) do not liberate hydrogen.
• Displacement Reactions:
  • A more reactive metal can displace a less reactive metal from its compound or salt solution.
  • Potassium can displace all metals from their salt solutions.
• Reducing Agent Strength:
  • Metals at the top of the reactivity series are strong reducing agents.
  • Metals at the bottom are weak reducing agents.
  • Potassium is the strongest, and gold is the weakest reducing agent.

D. Natural Occurrence and Extraction of Metals

• Uncombined State:
  • Noble metals (Ag, Au, Bi, Cu, Pd, Pt) exist in nature as uncombined or free state.
• Combined State:
  • Active metals (alkali and alkaline earth metals) never exist in uncombined state.
  • They exist in the form of carbonates, halides, oxides, phosphates, silicates, sulphides, and sulphates.
• Minerals:
  • Constituents of the earth’s crust containing metals or their compounds.
  • Examples: Sodium as halite ($NaCl$), potassium as sylvite ($KCl$), magnesium as magnesite ($MgCO<em>3$), calcium as limestone ($CaCO</em>3$).
• Ores:
  • Minerals with a high percentage of a particular metal from which the metal can be profitably extracted.
  • Ores contain impurities (sand and other undesirable materials) called gangue.
• Metallurgy:
  • The science and technology of extracting metals from their ores and compounding alloys.
• Principal Steps in Extraction:
  1. Preparation (concentration) of the ore (e.g., oil floatation, magnetic separation).
  2. Production of the metal (e.g., roasting, calcination).
  3. Purification of the metal (e.g., chemical reduction, electrolytic reduction).
• Extraction by Electrolysis:
  • Most active metals (K, Na, Ca, Mg) are extracted only by electrolysis.

5.2.2 Alloys

• Definition:
  • Mixtures of two or more metals or metals and nonmetals when molten and do not separate when solidified.
• Formation:
  • The constituent elements are melted together and then allowed to cool to form a solid material called alloy.
• Effect of Alloying:
  • Increases hardness and strength.
  • Modifies color and melting point.
  • Decreases electrical conductivity.
  • Increases resistance to corrosion of metals.
• Examples of Alloys:
  • Amalgam: Alloy of mercury and another metal.
  • Gun Metal: Alloy of copper (87%), tin (10%), and zinc (3%).
  • Solder (Fuse Metal): 67% tin and 33% lead, melts at 183°C (lower than tin's melting point of 232°C). Used to join wires and electrical resistances.
• Gold Alloys:
  • Gold is hardened by alloying it with copper and silver.
  • Gold content is expressed in carats or mass percent.
  • Carat: Mass unit of gold in 24 mass units of the alloy.
  • 24 carat: Pure gold.
  • 22 carat: 22 parts of pure gold in 24 parts of the alloy.
• Mass Percent Calculation Examples:
  • Mass % of gold in 24 carat gold = $\frac{24 \times 100}{24} = 100.00$
  • Mass % of gold in 22 carat gold = $\frac{22 \times 100}{24} = 91.67$

5.2.3 Production of Aluminum, Iron, and Copper

A. Aluminum

• Students should be able to:
  • Explain properties, occurrence, and extraction of aluminum.
  • Describe the applications of aluminum.
• Occurrence:
  • Most abundant metal in the earth's crust (about 8%).
  • Third most abundant element after oxygen and silicon.
  • Does not occur in uncombined or free metal state.
  • Main mineral is bauxite (Al2O3
    2H_2O).
  • Other minerals: orthoclase ($KAlSi<em>3O</em>8$), cryolite ($Na<em>3AlF</em>6$), corundum ($Al<em>2O</em>3$), beryl ($Be<em>3Al</em>2Si<em>6O</em>8$), and china clay (Al2Si2O7 2H2O).
• Extraction (Hall–Héroult Process):
  1. Purification of Bauxite:
     • Bauxite is contaminated by silicon dioxide ($SiO_2$), iron oxide, and titanium (IV) oxide.
     • Powdered ore is heated with sodium hydroxide solution to convert silica to soluble silicate.
     • $SiO<em>2(s) + 2NaOH(aq) \rightarrow Na</em>2SiO<em>3(aq) + H</em>2O(l)$
     • Aluminum oxide is converted to soluble sodium aluminate.
     • $Al<em>2O</em>3(s) + 2NaOH(aq) \rightarrow 2NaAlO<em>2(aq) + H</em>2O(l)$
     • Impurities like iron oxides and titanium (IV) oxide remain unaffected and are filtered off.
     • The solution is treated with acid to precipitate aluminum hydroxide.
     • $AlO<em>2^-(aq) + H</em>3O^+(aq) \rightarrow Al(OH)_3(s)$
     • Aluminum hydroxide is collected, washed, dried, and heated to get $Al<em>2O</em>3$.
     • $2Al(OH)<em>3(s) \xrightarrow{Heat} Al</em>2O<em>3(s) + 3H</em>2O(g)$
  2. Electrolysis:
     • Pure aluminum oxide is mixed with cryolite ($Na<em>3AlF</em>6$) to reduce its melting point from 2045°C to 1000°C.
     • The molten cryolite provides a good conducting medium for electrolysis.
     • Graphite electrodes are used as both anode and cathode.
     • Electrode Reactions:
       • Anode: $3C(s) + 6O^{2-} \rightarrow 3CO_2(g) + 12e^-$
       • Cathode: $4Al^{3+}(l) + 12e^- \rightarrow 4Al(l)$
     • Overall Reaction:
       • $4Al^{3+}(l) + 6O^{2-}(l) + 3C(s) \rightarrow 4Al(l) + 3CO_2(g)$
       • $2Al<em>2O</em>3(l) + 3C(s) \rightarrow 4Al(l) + 3CO_2(g)$
     • The anode (graphite electrode) oxidizes to carbon dioxide and must be replaced regularly.
     • Molten aluminum is siphoned from the bottom of the electrolytic cell.
• Physical Properties:
  • Soft, silvery-white, and light metal with a density of 2.7 g/cm3.
  • Melts at 660°C.
  • Can be shaped into wires, rolled, pressed, or cast into different shapes.
  • Good conductor of heat and electricity.
• Chemical Properties:
  • Reactive metal.
  • Reaction with Oxygen:
    • Reacts with atmospheric oxygen to form a thin film of aluminum oxide on its surface ($4Al(s) + 3O<em>2(g) \rightarrow 2Al</em>2O_3(s)$).
    • This film inhibits further reaction with oxygen.
  • Reaction with Nitrogen:
    • Burns in nitrogen gas to form aluminum nitride ($2Al(s) + N_2(g) \rightarrow 2AlN(s)$).
  • Reaction with Dilute Acids:
    • Reacts with dilute acids like $HCl$ and $H<em>2SO</em>4$, forming salts and liberating hydrogen gas.
    • $2Al(s) + 3H<em>2SO</em>4(aq) \rightarrow Al<em>2(SO</em>4)<em>3(aq) + 3H</em>2(g)$
    • $2Al(s) + 6HCl(aq) \rightarrow 2AlCl<em>3(s) + 3H</em>2(g)$
    • Does not react with dilute or concentrated $HNO_3$ due to the formation of a protective oxide layer.
  • Reaction with Chlorine:
    • Burns in chlorine gas to form aluminum chloride ($2Al(s) + 3Cl<em>2(g) \rightarrow 2AlCl</em>3(s)$).
  • Reaction with Sodium Hydroxide Solution:
    • $2Al(s) + 2NaOH(aq) + 6H<em>2O(l) \rightarrow 2NaAl(OH)</em>4(aq) + 3H_2(g)$
• Uses of Aluminum:
  • Light alloys like duralumin (Al, Cu, Mg) are used in the transportation industry (airplanes, ships, automobiles).
  • High thermal conductivity and corrosion resistance make it suitable for household cookware.
  • Manufacture of door and window frames and building roofs.
  • Packaging material in the food industry.
  • Power transmission lines.
  • Thermite Welding:
    • Aluminum displaces iron from iron oxide. $2Al(s) + Fe<em>2O</em>3(s) \rightarrow 2Fe(l) + Al<em>2O</em>3(s)$
    • Mixture of powdered aluminum and iron oxide is called thermite and produces a temperature of about 3000°C.
    • Used in welding rails, propeller shafts, and other steel parts.

B. Iron

• Students should be able to:
  • Explain properties, occurrence, and extraction of iron.
  • Describe the applications of iron.
• Occurrence:
  • Second most abundant metal in the earth's crust after aluminum (about 4.7% of the earth's crust).
  • Never found as a free metal in nature.
  • Exists in the form of compounds (oxides, carbonates, and sulphides).
  • Chief ores are hematite ($Fe<em>2O</em>3$), limonite (Fe2O3
    H2O), magnetite ($Fe</em>3O<em>4$), and siderite ($FeCO</em>3$).
  • Also found as iron pyrites ($FeS_2$), commonly called fool’s gold.
• Extraction (Blast Furnace):
  • Raw materials: iron ore, coke, limestone, and hot air.
  • The furnace is charged with a mixture of iron ore, limestone, and coke at the top, and hot air is blown at the bottom.
  • Reactions in the Blast Furnace:
    1. Oxidation of coke to carbon dioxide: $C(s) + O<em>2(g) \rightarrow CO</em>2(g) + Heat$
    2. Reduction of carbon dioxide to carbon monoxide: $CO_2(g) + C(s) \rightarrow 2CO(g)$
    3. Reduction of iron oxides to metallic iron by carbon monoxide:
       • $3Fe<em>2O</em>3(s) + CO(g) \rightarrow 2Fe<em>3O</em>4(s) + CO_2(g)$
       • $Fe<em>3O</em>4(s) + CO(g) \rightarrow 3FeO(s) + CO_2(g)$
       • $FeO(s) + CO(g) \rightarrow Fe(l) + CO_2(g)$
    4. Decomposition of limestone:
       • $CaCO<em>3(s) \rightarrow CaO(s) + CO</em>2(g)$
    5. Formation of slag:
       • $CaO + SiO<em>2 \rightarrow CaSiO</em>3$
       • Lime + sand calcium silicate
       • (Flux) (Impurity) (Slag)
  • Slag is used for the manufacture of cement.
  • The iron obtained directly from the blast furnace is called pig iron.
  • Pig iron contains about 2% silicon, up to 1% phosphorus and manganese, and traces of sulfur.
  • Pig iron contains high carbon content, typically 3.5 - 4.5%, making it brittle.
• Wrought Iron:
  • The purest form of commercial iron is called wrought iron, obtained by removing most impurities from pig iron.
  • Impure iron is heated with hematite and limestone in a furnace.
  • Increases the purity of the iron to 99.5%.
  • Tough, malleable, and ductile.
• Steel Making from Pig Iron (Purification of Pig Iron):
  • Conversion of pig iron to steel is a purification process where impurities are eliminated through oxidation at high temperatures.
  • Three techniques used: Bessemer converter, Open-hearth Furnace, and Basic Oxygen Process.
    1. Bessemer Converter:
       • Molten pig iron is transferred to a cylindrical vessel with a refractory lining of $MgCO<em>3$ and $CaCO</em>3$.
       • A blast of hot air is blown through the molten metal.
       • Oxygen converts silicon, phosphorus, and sulfur to their oxides, which react with the lining to form a slag.
       • Carbon is oxidized to carbon monoxide.
    2. Open-hearth Furnace:
       • A large, shallow hearth lined with a basic oxide refractory (MgO and CaO).
       • Charged with a mixture of pig iron, $Fe<em>2O</em>3$, scrap iron, and limestone.
       • A blast of hot air and burning fuel is directed over the surface.
       • Impurities are oxidized by the $Fe<em>2O</em>3$ and air, forming carbon dioxide and sulfur dioxide.
       • $C + O<em>2 \rightarrow CO</em>2$
       • $S + O<em>2 \rightarrow SO</em>2$
       • $12P + 10Fe<em>2O</em>3 \rightarrow 3P<em>4O</em>{10} + 20Fe$
       • $3Si + 2Fe<em>2O</em>3 \rightarrow 3SiO_2 + 4Fe$
       • Calcium oxide (from limestone) reacts with oxides of silicon and phosphorus to form slag.
       • $P<em>4O</em>{10} + 6CaO \rightarrow 2Ca<em>3(PO</em>4)_2$
       • $SiO<em>2 + CaO \rightarrow CaSiO</em>3$
       • Impurities manganese and silicon form respective oxides and removed as slag. $(a) Si + O<em>2 \rightarrow SiO</em>2$
    3. Basic Oxygen Process:
       • A mixture of powdered calcium oxide (CaO) and oxygen gas is forced directly into the surface of the molten pig iron.
       • The oxygen reacts exothermically with carbon, sulfur, silicon, phosphorus, and impurity metals.
       • Carbon and sulfur are oxidized to $CO<em>2$ and $SO</em>2$.
       • Oxides of silicon ($SiO<em>2$), phosphorus ($P</em>4O_{10}$), and impurity metals combine with lime (CaO), forming slag.
• Tempering of Steel:
  • A process by which steel is conditioned to a desired hardness by heating and controlled rate of cooling.
  • Some carbon in steel is present as cementite ($FeC_3$).
• Physical Properties of Iron:
  • Gray lustrous, malleable, and ductile metal.
  • Good conductor of heat and electricity.
  • High melting point (1580°C) and high density (7.87 g/cm3).
  • Ferromagnetic metal.
• Chemical Properties of Iron:
  • Reactive metal, but less reactive than group IA and IIA metals.
  • Rusts in the presence of air and moisture to form hydrated iron (III) oxide: 4Fe(s) + 3O2(g) + Moisture \rightarrow 2Fe2O3 xH2O(s)(rust)
  • Reacts with dilute acids to form iron (II) salts and liberating hydrogen gas
    • $Fe (s) + 2HCl (aq) \rightarrow FeCl<em>2 (aq) + H</em>2 (g)$
    • $Fe (s) + H<em>2SO</em>4 (aq) \rightarrow FeSO<em>4 (aq) + H</em>2 (g)$
  • Exhibits different oxidation states (Fe2+ and Fe+3).
  • Solutions of iron (II) compounds are pale-green and solutions of iron (III) compounds are yellowish brown.
  • Heated iron reacts with hydrogen chloride gas to form iron (II) chloride and hydrogen gas.
  • $Fe (s) + 2HCl (g) \rightarrow FeCl<em>2 (g) + H</em>2 (g)$
  • Heated iron reacts with chlorine and sulfur.
    • $2Fe (s) + 3Cl<em>2 (g) \rightarrow 2FeCl</em>3 (s)$
    • $Fe (s) + S (s) \rightarrow FeS (s)$
  • Displaces less active metals from solutions of their salts.
    • $Fe (s) + Cu^{2+} (aq) \rightarrow Fe^{2+} (aq) + Cu (s)$
• Uses of Iron:
  • Widely used metal in construction of buildings and bridges
  • Pig iron is used to make domestic boilers, hot-water radiators, railings, water pipes, castings, and moldings.
  • Wrought iron is used in making nails, sheets, horseshoes, ornamental gates, door knockers, farm machinery etc.
  • Manufacture of alloys such as carbon steels and alloy steels.

C. Copper

• Students should be able to:
  • Explain properties, occurrence and extraction of Copper.
  • Describe the the applications of Copper.
• Occurrence:
  • Occasionally found as native copper, but mainly in compounds (sulphides, oxides and carbonates).
  • Most important sulphide ores are chalcopyrite ($CuFeS<em>2$), chalcocite ($Cu</em>2S$), covellite ($CuS$) and bornite ($Cu<em>5FeS</em>4$).
  • Principal oxide ores are cuprite ($Cu_2O$) and tenorite ($CuO$).
  • Carbonate form:
    • malachite (CuCO3 Cu(OH)2).
• Extraction from Chalcopyrite:
  1. Concentration by Froth Flotation:
     • Crushed and ground sulphide ore is concentrated, changing the concentration from 2 % copper to as high as 30% copper.
  2. Roasting:
     • Concentrated ore is roasted with a limited supply of air (oxygen).
     • $2CuFeS<em>2 (s) + 4O</em>2 (g) \rightarrow Cu<em>2S (s) + 2FeO(s) + 3SO</em>2 (g)$
  3. Smelting:
     • Roasted mixture is smelted by adding limestone and sand to form a molten slag that removes impurities.
     • $CaCO<em>3 (s) + SiO</em>2 (s) \rightarrow CaSiO<em>3 (l) + CO</em>2 (g)$
     • $FeO (s) + SiO<em>2 (s) \rightarrow FeSiO</em>3 (l)$
  4. Reduction:
     • The $Cu_2S$ obtained by roasting chalcopyrite is then reduced by heating it in a limited supply of oxygen.
     • $Cu<em>2S (s) + O</em>2 (g) \rightarrow 2Cu (l) + SO_2 (g)$
     • The copper produced is called blister copper and it has 98.5 – 99.5 % purity.
  5. Electrolytic Refining:
     • Blister copper contains iron, silver, gold and sometimes zinc as impurities.
       • Anode: A thick block of impure copper.
       • Cathode: A thin strip of pure copper.
       • Electrolyte: An aqueous solution of copper sulphate, Small quantity of dilute sulphuric acid is also added to the salt solution to prevent hydrolysis.
         • Oxidation at anode:
           • $Cu (impure metal) \rightarrow Cu^{2+}(aq) + 2e^−$
         • Reduction at cathode:
           • $Cu^{2+}(aq) + 2e \rightarrow Cu ( metal)$
• Physical Properties of Copper:
  • Soft, ductile, malleable, reddish-brown metal with a density of 8.96 g/ cm3.
  • Second to silver in electrical conductivity.
  • Melts at 1086°C and boils at 2310°C.
• Chemical Properties of Copper:
  • Less reactive metal that is why it is found in the native state.
    • Powdered copper, when heated in air forms a black powder of copper (II) oxide, $CuO$.
    • $2Cu (s) + O_2 (g) \rightarrow 2CuO (s)$
  • Does not react with dilute acids like $HCl$ and $H<em>2SO</em>4$.
  • Oxidized by oxidizing acids such as dilute and concentrated nitric acid and hot concentrated sulphuric acid, $H<em>2SO</em>4$.
    • $dilute 3Cu(s) + 8HNO<em>3 (aq) 3Cu(NO)</em>3)<em>2 (aq) + 2NO (g) + 4H</em>2O (l)$
    • $concentrated Cu(s) + 8HNO<em>3 (aq) Cu(NO)</em>3)<em>2 (aq) + 2NO</em>2 (g) + 2H_2O (l)$
    • $Hot and concentrated Cu(s) + 2H<em>2SO</em>4 (aq) CuSO<em>4(aq) + SO</em>2 (g) + 2H_2O (l)$
  • Corrodes in moist air over a long period of time as a result of oxidation caused by a mixture of water, oxygen and carbon dioxide. It turns green, due to the formation of verdigris: a basic copper carbonate (CuCO3 Cu(OH)2) or $Cu<em>2(OH)</em>2CO_3$.
    • 2Cu (s) + H2O (l) + O2 (g) + CO2 (g) \rightarrow CuCO3
      Cu(OH)_2
  • Exhibits different oxidation states. It exists as cuprous ($Cu^+$) and cupric ($Cu^{2+}$) ions.
    • $2Cu^+ (aq) \rightarrow Cu^{2+} (aq) + Cu (s)$
• Uses of Copper:
  • Alloys:
    • bronze (copper and tin).
    • brass (copper and zinc).
  • Electrical industry: electric wires, cables.
  • Copper compounds as pesticides.

5.3 Production of Some Important Nonmetals

5.3.1 General Properties of Nonmetals and Common Uses of Some Nonmetallic Compounds

• Students should be able to:
  • Mention the general properties of non-metals and their uses.
  • Describe some common uses of compounds of nonmetals such as $CO<em>2$, $Na</em>2CO<em>3$, $NH</em>3$, $HNO<em>3$, $H</em>3PO<em>4$, $Ca</em>3(PO<em>4)</em>2$, $SO<em>2$ & $H</em>2SO_4$
  • Describe the occurrence, extraction and uses of nitrogen, phosphorous, oxygen, sulphur and chlorine.
• Nonmetals have opposite characteristics to that of metals.

A. Physical Properties

• State:
  • Solids, liquids, gases.
• Luster:
  • Non-lustrous.
• Malleability and Ductility:
  • Nonmalleable and non-ductile.
• Hardness and Density:
  • Varying hardness and have low density.
• Melting and Boiling Points:
  • Low melting and boiling points.
• Sonorousity:
  • Non-sonorous.
• Conductivity:
  • Poor conductors of heat and electricity.

B. Chemical Properties of Non-Metals

• Reaction with Oxygen:
  • Nonmetals react with oxygen on heating or burning to form their oxides.
• Reaction with Acids:
  • Do not displace hydrogen on reaction with dilute acids.
• Oxide Nature:
  • React with oxygen to form acidic or neutral oxides.
• Hydride Formation:
  • Combine with hydrogen to form stable hydrides.
• Reaction with Water:
  • Do not react with water.
• Electronegativity:
  • Electronegative i.e for negative ions by gaining electrons.
• Oxidizing Agents:
  • Oxidizing agents.

5.3.2 Production of Nitrogen, Phosphorous, Oxygen, Sulphur and Chlorine

A. Nitrogen

• Students should be able to:
  • Explain properties, Occurrence and extraction of Nitrogen.
  • Describe the the applications of Nitrogen.
• Occurrence and Production:
  • Occurs in nature in the elemental form as a diatomic molecule, $N_2$, in atmospheric air (about 80% by volume).
  • In the form of compounds, it exists as sodium nitrate (Chile salt peter, $NaNO<em>3$) and potassium nitrate ($KNO</em>3$) also called saltpetre.
  • Also found in DNA molecules and proteins of all living things.
  • Industrial Production:
    • Impurities (dust, other particles, $CO_2$, water vapor) are removed from air.
    • Air is compressed under high pressure and low temperature.
    • Fractional distillation of liquid air separates nitrogen.
    • Argon distills off the mixture at –186°C, leaving behind a blue liquid of oxygen that boils at –183°C.
  • Laboratory Preparation:
    • Warming an aqueous solution containing ammonium chloride and sodium nitrite.
    • $NH<em>4Cl(aq) + NaNO</em>2(aq) \rightarrow NaCl(aq) + N<em>2(g) + 2H</em>2O(l)$
• Physical Properties of Nitrogen:
  • Colorless, odorless, and tasteless gas.
  • Inert under ordinary conditions due to the strength of the triple bond.
• Chemical Properties of Nitrogen:
  • Reacts with metals of group IA and IIA as well as oxygen at higher temperatures.
  • Reaction with Reactive Metals:
    • Lithium: $6Li (s) + N<em>2 (g) \rightarrow 2Li</em>3N (s)$
    • Calcium: $3Ca (s) + N<em>2 (g) \rightarrow Ca</em>3N_2 (s)$
    • Magnesium: $3Mg (s) + N<em>2 (g) \rightarrow Mg</em>3N_2 (s)$
  • Reaction with Oxygen:
    • $N<em>2 (g) + O</em>2 (g) \rightarrow 2NO (g)$
    • $N<em>2 (g) + 2O</em>2 (g) \rightarrow 2NO_2 (g)$
    • Nitric oxide (NO) forms nitrogen dioxide ($NO<em>2$), a reddish brown gas and dimerizes at low temperatures to give a colorless gas of dinitrogen tetraoxide, $N</em>2O_4$.
    • $2NO<em>2 (g) \rightarrow N</em>2O_4 (g)$
  • Forms oxides, like dinitrogen monoxide, $N<em>2O$, dinitrogen trioxide ($N</em>2O<em>3$) and dinitrogen pentoxide ($N</em>2O_5$).
  • Haber Process:
    • Reacts directly with hydrogen in the Haber process to form ammonia.
    • $N<em>2(g) + 3H</em>2(g) \rightleftharpoons 2NH_3(g) \text{ } Fe/200 - 300atm, 400-600^\circ C$
• Uses of Nitrogen:
  • Used in food packaging to prevent oxidation.
  • To create an inert atmosphere in the production of semiconductors.
  • Liquid nitrogen is used as a refrigerant to preserve bulls’ semen and blood.
  • Major use is in the production of ammonia.

B. Phosphorus

• Students should be able to:
  • Explain properties ,Occurrence and extraction of Phosphorous.
  • Describe the the applications of Phosphorous.
• Occurrence and Extraction:
  • A relatively abundant element, ranking 12th in the earth’s crust.
  • Exists naturally only in the combined state, such as in rock phosphate, $Ca<em>3(PO</em>4)<em>2$, fluoroapatite, $Ca</em>{10}(PO<em>4)</em>6F<em>2$ or 3Ca3(PO4)2
    CaF_2.
  • Allotropic forms: white phosphorus and red phosphorus.
• Physical Properties of Phosphorus:
  • White Phosphorus:
    • Very poisonous, white waxy-looking substance that melts at 44.1°C and boils at 287°C.
    • Density is 1.8 g/cm3.
    • Consists of individual tetra- atomic ($P_4$) molecules and is an unstable form of phosphorus
  • Red Phosphorus:
    • Denser (2.16 g/cm3) and is much less reactive than white phosphorus at normal temperatures.
    • Consists of $P_4$ molecules linked together to form a polymer
• Industrial Manufacture of White Phosphorus:
  • Heating a mixture of crushed rock phosphate, $Ca<em>3(PO</em>4)_2$, silica, SiO2, and coke in an electric furnace.
  • $2Ca<em>3(PO</em>4)<em>2 (s) + 6SiO</em>2 (s) + 10C (s) \rightarrow 6CaSiO<em>3(l) + P</em>4(g) + 10CO (g)$
  • The vaporized phosphorus ($$