5.3 Production of Some Important Nonmetals

5.3 Production of Some Important Nonmetals

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

A. Physical Properties of Nonmetals:

• Exist as solids, liquids, or gases.

• Non-lustrous (lack shine).

• Non-malleable and non-ductile (brittle).

• Varying hardness with generally low density.

• Low melting and boiling points.

• Do not exhibit metallic luster.

• Softer than metals.

• Non-sonorous (do not produce a ringing sound).

• Poor conductors of heat and electricity.

B. Chemical Properties of Nonmetals:

• React with oxygen when heated or burned, forming oxides.

• Do not displace hydrogen in reactions with dilute acids.

• Form acidic or neutral oxides upon reacting with oxygen.

• Combine with hydrogen to create stable hydrides.

• Generally do not react with water.

• Electronegative, gaining electrons to form negative ions.

• Serve as oxidizing agents.

5.3.2 Production of Nitrogen, Phosphorus, Oxygen, Sulfur, and Chlorine

A. Nitrogen

Occurrence and Production:

• Predominantly found as a diatomic molecule (N_2) in the atmosphere, making up about 80% by volume.

• Exists in compounds such as sodium nitrate (NaNO3) and potassium nitrate (KNO3), found in DNA and proteins.

• Industrial production involves:

  • Removing impurities from air.

  • Compressing air under high pressure and low temperature to eliminate CO_2 and water vapor.

  • Fractional distillation of liquid air to separate nitrogen from oxygen.

Physical Properties of Nitrogen:

• Colorless, odorless, and tasteless gas.

• Inert under normal conditions due to the strength of the triple bond (N≡N).

Chemical Properties of Nitrogen:

• Reacts with metals from groups IA and IIA at elevated temperatures to form nitrides. Examples of Reactions:

  • With lithium: 6Li (s) + N2 (g) → 2Li3N (s)

  • With calcium: 3Ca (s) + N2 (g) → Ca3N_2 (s)

  • With magnesium: 3Mg (s) + N2 (g) → Mg3N_2 (s)

• Combines with oxygen at high temperatures to form various oxides:

  • N2 (g) + O2 (g) → 2NO (g)

  • N2 (g) + 2O2 (g) → 2NO_2 (g)

Uses of Nitrogen:

• Food packaging to prevent oxidation.

• Creating inert atmospheres in semiconductor manufacturing.

• Liquid nitrogen as a refrigerant.

B. Phosphorus

Occurrence and Extraction:

• Exists naturally only in combined forms (e.g., rock phosphate Ca3(PO4)_2).

• Common allotropes include white phosphorus (P_4) and red phosphorus.

Physical Properties of Phosphorus:

• White Phosphorus: Poisonous, waxy, melts at 44.1°C, boils at 287°C, consists of unstable P_4 molecules.

• Red Phosphorus: Denser (2.16 g/cm^3), less reactive, forms a polymeric structure.

Industrial Production:

• White phosphorus is produced by heating a mixture of rock phosphate, silica, and coke in an electric furnace:

  • 2Ca3(PO4)2 (s) + 6SiO2 (s) + 10C (s) → 6CaSiO3 (l) + P4 (g) + 10CO (g)

• Collected under water due to high reactivity.

Preparation of Red Phosphorus:

• Obtained by heating white phosphorus in sunlight for several days.

Chemical Properties and Uses of Phosphorus

Reactions with Oxygen

• Limited oxygen: Forms tetraphosphorus hexoxide (P4O6)

P4 (s) + 3O2(g) → P4O6(s)

• Excess oxygen: Forms tetraphosphorus decoxide (P4O{10})

P4 (s) + 5O2 (g) → P4O{10}(s)

Reactions with Water

• P4O6 dissolves in water to form phosphorous acid (H3PO3)

P4O6 (s) + 6H2O (l) → 4H3PO_3 (aq)

• P4O{10} dissolves in water to form orthophosphoric acid (H3PO4)

P4O{10} (s) + 6H2O (l) → 4H3PO_4 (aq)

Reactions with Chlorine

• Limited chlorine: Forms phosphorus(III) chloride (PCl_3)

P4 (s) + 6Cl2 (g) → 4PCl_3 (s)

• Excess chlorine: Forms phosphorus(V) chloride (PCl_5)

P4 (s) + 10Cl2(g) → 4PCl_5 (s)

Industrial Production and Storage of Phosphorus

White Phosphorus Production

• Heating a mixture of:

• Crushed rock phosphate (Ca3(PO4)_2)

• Silica (SiO_2)

• Coke (C) in Electric furnace
  • Electric furnace reaction: 2Ca3(PO4)2 (s) + 6SiO2 (s) + 10C (s) → 6CaSiO3 (l) + P4 (g) + 10CO (g)

Condensation and Storage

• Vaporized phosphorus (P_4) is condensed and collected
• White phosphorus is stored under water due to spontaneous ignition in the presence of oxygen

Red Phosphorus Production

• Heating white phosphorus in sunlight for several days

Solubility

• Red phosphorus:

• Insoluble in water

• Soluble in carbon disulfide (CS_2)

Stability of CS_2 Solutions

• Solutions of P4 in CS2 are reasonably stable
• Ignition occurs when CS_2 evaporates

Summary

• White phosphorus is produced industrially through an electric furnace process.
• Red phosphorus is prepared by heating white phosphorus in sunlight.
• White phosphorus is highly reactive and must be stored under water, while red phosphorus is less reactive and can be stored without water.
• Red phosphorus is soluble in carbon disulfide, but these solutions are unstable and ignite upon evaporation of CS_2.

Uses of Phosphorus

• Red phosphorus:

• Making matches
  • White phosphorus:

• Production of phosphoric acid and other phosphorus compounds converted into acids and salts to be used such as

• Fertilizers

• Baking powder

• Chemical industries

• Fireworks, smoke bombs, rat-poisons, tracer bullets

• Essential for plant growth

Summary

• Phosphorus exhibits various chemical reactions, primarily involving oxidation and halogenation.
• It has diverse industrial and practical applications, including fertilizer production and pyrotechnics.
• Phosphorus is crucial for biological processes in plants.

C) Oxygen: Occurrence and Production

Abundance and Occurrence

• Most abundant element on Earth (46.6% by weight of the Earth's crust)
• Found in compounds (oxides, silicates, carbonates, phosphates)
• Elemental form in atmospheric air (20% by volume)

Industrial Production

• Fractional distillation of liquid air

Allotropes

• Diatomic (O2): Common form • Triatomic (O3): Ozone

Ozone Formation

• Natural production in the stratosphere
• Two-step process involving:

• Solar ultraviolet radiation breaking apart O_2 molecules

• Reaction of three O2 molecules to form two O3 molecules

Significance of Ozone

• Removes harmful radiation from sunlight
• Protects life on Earth's surface

Summary

• Oxygen is abundant in the Earth's crust and atmosphere.
• It exists as two allotropes: O2 and O3 (ozone).
• Ozone formation in the stratosphere is crucial for protecting life from harmful radiation.

Physical and Chemical Properties of Oxygen

Physical Properties

• Colorless, odorless, tasteless gas
• Changes from gas to liquid at -182.96°C
• Liquid has a bluish color
• Solidifies at -218.4°C
• Density: 1.429 grams per liter (denser than air)

Chemical Properties

• Reactive element, combines with most elements to form oxides
• Oxidation of Metals: Forms basic oxides (e.g., MgO, CaO)
Example: Oxidation: 2Mg (s) + O2 (g) → MgO (s) • Oxidation of Non-Metals: Forms acidic oxides (e.g., SO2, P4O10) Example: Non-Metal Oxidation: S8 (s) + 8O2 (g) → 8SO2 (g)
• Combustion Support: Necessary for burning substances (e.g., charcoal, hydrocarbons)
Examples: Combustion: C (s) + O2 (g) → CO2 (g)

Summary

• Oxygen exhibits distinct physical properties and is highly reactive.
• It forms oxides when combined with other elements.
• Oxygen's ability to support combustion makes it essential for various chemical reactions and processes.

D) Sulfur: Occurrence and Extraction

Natural Occurrence

• Found in compounds such as:

• Galena (PbS) (fool’s gold)

• Pyrites (FeS2)

• Cinnabar (HgS)

• Sphalerite (ZnS)

• Gypsum (CaSO4.2H2O)

• Barite (BaSO4)

• Hydrogen sulfide (H2S) in natural gas and crude oil

Frasch Process

• Industrial extraction method for underground sulfur deposits
• Involves three concentric pipes:

• Outermost pipe: Pumps superheated water (170°C)

• Innermost pipe: Compresses hot air

• middle tube:Air and water create a froth of molten sulfur, which rises to the surface
  • Molten sulfur is cooled and solidified

Summary

• Sulfur occurs naturally in various compounds.
• The Frasch process is an efficient method for extracting elemental sulfur from underground deposits.

Uses and Allotropes sulfur

Industrial Uses

• Nearly half of the sulfur used in chemical industries is obtained as a waste product from:

• Natural gas processing

• Crude oil purification

• Metal ore roasting

Environmental Benefits

• Using sulfur from waste products reduces the demand for natural sulfur and minimizes:

• Atmospheric air pollution

• Acid rain formation

Allotropes of Sulfur

• Rhombic sulfur (α-sulfur)
• Monoclinic sulfur (β-sulfur)

Rhombic Sulfur

• Most stable form of sulfur
• Consists of S_8 molecules
• Octahedral shape
• Yellow color
• Melting point: 385.8K
• Insoluble in water, soluble in organic solvents

Monoclinic Sulfur

• Less dense than rhombic sulfur
• Also contains S_8 molecules
• Monoclinic crystal structure
• Converts to rhombic sulfur below 95.3 °C

Summary

• Sulfur is used in various industrial processes, with a significant portion obtained as a waste product.
• Rhombic and monoclinic sulfur are the two main allotropes of sulfur, differing in their crystal structures and stability.

Physical and Chemical Properties Sulfur

Physical Properties

• Tasteless, odorless, brittle solid
• Pale yellow color
• Poor conductor of electricity
• Insoluble in water

Chemical Properties

• Relatively stable at room temperature
• Reacts with metals and non-metals when heated
• Reaction with Metals: Forms sulfides (e.g., 8Fe (s) + S8 (s) → 8FeS (s) • Reaction with Oxygen: Burns to form oxides (e.g., S8 (s) + 8O2 (g) → 8SO2 (g)

Industrial Significance

• Raw material for sulfuric acid production (Contact Process)

Contact Process

• Step 1: Oxidation of sulfur to sulfur dioxide (SO2)
• Step 2: Conversion of SO2 to sulfur trioxide (SO3) in the presence of a catalyst
• Step 3: Absorption of SO3 into concentrated sulfuric acid to produce oleum (H2S2O7)
• Step 4: Dilution of oleum with water to obtain desired sulfuric acid concentration

Summary

• Sulfur exhibits distinct physical properties and chemical reactivity.
• The Contact Process is a significant industrial method for producing sulfuric acid, using sulfur as the starting material.

E) Chlorine: Occurrence and Extraction

Occurrence

• Found in nature only in compound form
• Chiefly as chlorides of sodium, potassium, calcium, and magnesium
• Primary source: Sodium chloride (NaCl)

Industrial Production

• Electrolysis of concentrated aqueous NaCl solution
• Chlorine gas (Cl2) produced at anode
• Hydrogen gas (H2) and hydroxide ions (OH-) produced at cathode

Electrolysis Equations

• Anode reaction: 2Cl^- (aq) → Cl2 (g) + 2e^- • Cathode reaction: 2H2O (l) + 2e^- → H2 (g) + 2OH^- (aq) • Cell reaction: 2NaCl (aq) + 2H2O (l) → 2NaOH (aq) + Cl2 (g) + H2 (g)

Separation of Products

• NaOH and Cl2 must be kept separate to prevent reaction: NaOH (aq) + Cl2 (g) → NaOCl (aq) + HCl (aq)

Oxidation Number of Chlorine in NaOCl

• +1

Summary

• Chlorine is extracted from chlorides, primarily NaCl, through electrolysis.
• The electrolysis process produces Cl2 gas and NaOH, which must be separated to prevent a reaction that forms NaOCl.
• The oxidation number of chlorine in NaOCl is +1.
• Chlorine is a highly reactive non-metal that reacts with almost all elements except the noble gases.

Physical Properties

• Greenish-yellow gas at room temperature
• Melts at -102°C, boils at -34°C
• Fairly soluble in water
• Extremely poisonous

Chemical Properties

• Oxidizing agent
• Forms chloride salts with metals
• Reacts with hydrogen to form hydrogen chloride
• Displaces less reactive halogens from aqueous solutions
• Reacts with water to form hydrochloric acid and hypochlorous acid

Uses of chlorine

• Bleaching agent
• Disinfectant

Reaction of Chlorine with Bromine

• Bromine cannot displace chlorine in a chemical reaction because chlorine is more reactive than bromine.

Conclusion

• Chlorine is a versatile element with a wide range of applications, but it must be handled with care due to its toxicity.