Unit 9: Chemistry of Non-Metals

9.1 Hydrogen, Oxygen and Ozone

Hydrogen

Hydrogen is the lightest and most abundant element in the universe, typically found as a diatomic molecule, H2. It exhibits unique chemical properties due to its simple atomic structure.

Atomic and Nascent Hydrogen

• Atomic Hydrogen: Refers to a single hydrogen atom, [H]. It is highly unstable and reactive, formed by dissociating molecular hydrogen at high temperatures or by electrical discharge.
• Nascent Hydrogen: This is hydrogen generated in situ (at the moment of reaction) by the reaction of a metal with an acid, e.g., Zn + H2SO4 → ZnSO4 + 2[H]. Nascent hydrogen is much more reactive than molecular hydrogen (H2) because it exists momentarily as individual atoms or highly reactive species before combining to form H2. Its enhanced reactivity allows it to reduce substances that molecular hydrogen cannot.

Isotopes of Hydrogen

Hydrogen has three main isotopes, differing in the number of neutrons in their nuclei:

• Protium (1H): Contains one proton and no neutrons. It is the most common isotope, making up over 99.98% of natural hydrogen. Its properties are those typically associated with hydrogen.
• Deuterium (2H or D): Contains one proton and one neutron. It is also known as "heavy hydrogen."
  • Properties: Slightly higher boiling and melting points than protium, lower vapor pressure, and slower reaction rates (kinetic isotope effect).
  • Uses: Used in nuclear magnetic resonance (NMR) spectroscopy as a solvent, as a tracer in chemical and biological studies, and in the production of heavy water.
• Tritium (3H or T): Contains one proton and two neutrons. It is a radioactive isotope with a half-life of 12.32 years.
  • Properties: Radioactive, emitting low-energy beta particles.
  • Uses: Used in self-powered lighting devices (e.g., exit signs), as a radioactive tracer in research, and as a fuel in experimental nuclear fusion reactors.

Application as Fuel

Hydrogen is considered a promising clean fuel for the future:

• Clean Burning: When hydrogen burns, it combines with oxygen to produce only water, 2H2(g) + O2(g) → 2H2O(g). This means no greenhouse gases (like CO2) or harmful pollutants are released, unlike fossil fuels.
• High Calorific Value: Hydrogen has the highest calorific value per unit mass of any known fuel (approximately 142 kJ/g), meaning it releases a large amount of energy for its weight.
• Water as Product: The primary product of hydrogen combustion is water, which can be recycled to produce more hydrogen, creating a sustainable energy cycle.

Heavy Water (D2O)

Heavy water is deuterium oxide, where the hydrogen atoms in water are replaced by deuterium.

• Preparation: Primarily prepared by the prolonged electrolysis of ordinary water. Due to the kinetic isotope effect, H2O molecules are electrolyzed slightly faster than D2O, leading to an enrichment of D2O in the residual water. Fractional distillation can also be used.
• Properties:
  • Higher melting point (3.82 °C) and boiling point (101.42 °C) than ordinary water.
  • Higher density (1.104 g/mL at 25 °C).
  • Slightly lower dielectric constant.
  • Slower reaction rates in chemical and biological processes compared to H2O.
• Uses as Moderator in Nuclear Reactors: Heavy water is an excellent neutron moderator in nuclear fission reactors. It slows down the fast neutrons produced during fission without absorbing them, allowing for a sustained chain reaction. This is crucial for reactors that use natural uranium as fuel, as it does not require enriched uranium.

Oxygen

Oxygen is a highly reactive non-metal, essential for life, and the most abundant element in the Earth's crust.

Allotropes of Oxygen

• Dioxygen (O2): The most common allotrope, a colorless, odorless gas vital for respiration and combustion. It consists of two oxygen atoms bonded together.
• Ozone (O3): A triatomic allotrope of oxygen, discussed in detail below.

Types of Oxides

Oxides are compounds formed when oxygen reacts with another element. They can be classified based on their acid-base character:

• Acidic Oxides: Non-metallic oxides that react with water to form acids or react with bases to form salt and water. E.g., Sulfur dioxide (SO2), SO2(g) + H2O(l) → H2SO3(aq).
• Basic Oxides: Metallic oxides that react with water to form bases or react with acids to form salt and water. E.g., Sodium oxide (Na2O), Na2O(s) + H2O(l) → 2NaOH(aq).
• Neutral Oxides: Oxides that show neither acidic nor basic properties and do not react with acids or bases. E.g., Carbon monoxide (CO), Nitrous oxide (N2O).
• Amphoteric Oxides: Oxides that can react as both an acid and a base, depending on the reactant. E.g., Aluminium oxide (Al2O3), Al2O3(s) + 6HCl(aq) → 2AlCl3(aq) + 3H2O(l) (basic), and Al2O3(s) + 2NaOH(aq) + 3H2O(l) → 2Na[Al(OH)4](aq) (acidic).
• Peroxides: Oxides containing the peroxide ion (O2^2-), where oxygen has an oxidation state of -1. They contain an O-O single bond. E.g., Sodium peroxide (Na2O2).
• Mixed Oxides: Oxides that are composed of two different simple oxides of the same metal. E.g., Trilead tetroxide (Pb3O4) which is 2PbO.PbO2, or Ferric ferro-oxide (Fe3O4) which is FeO.Fe2O3.

Hydrogen Peroxide (H2O2)

Hydrogen peroxide is a pale blue liquid, slightly more viscous than water, and a strong oxidizing agent.

• Preparation:
  • From Barium Peroxide: BaO2(s) + H2SO4(aq) → BaSO4(s) + H2O2(aq). Barium sulfate is insoluble and easily filtered off.
  • Electrolytic Method: Electrolysis of 50% sulfuric acid or ammonium hydrogen sulfate solution produces peroxodisulfuric acid, which upon hydrolysis yields hydrogen peroxide. 2H2SO4 → H2S2O8 + H2 (electrolysis) H2S2O8 + 2H2O → 2H2SO4 + H2O2 (hydrolysis)
  • Auto-oxidation of 2-ethylanthraquinol: This industrial process involves the catalytic oxidation of 2-ethylanthraquinol to 2-ethylanthraquinone and hydrogen peroxide, followed by reduction of the anthraquinone to regenerate the starting material.
• Decomposition: Hydrogen peroxide is thermodynamically unstable and decomposes into water and oxygen, a reaction catalyzed by light, heat, heavy metals, or rough surfaces. 2H2O2(l) → 2H2O(l) + O2(g)
• Disproportionation: In this reaction, hydrogen peroxide acts as both an oxidizing and a reducing agent.
  • As an oxidizer: H2O2 + 2HI → I2 + 2H2O
  • As a reducer: H2O2 + Ag2O → 2Ag + H2O + O2
• Uses:
  • Oxidizer: Used in chemical synthesis and waste water treatment.
  • Bleach: Effective bleaching agent for textiles, paper, and hair (hair bleach).
  • Antiseptic: Dilute solutions are used as a mild antiseptic for minor cuts and wounds, often marketed as a topical disinfectant.

Ozone

Ozone (O3) is an allotrope of oxygen with distinct properties and crucial environmental roles.

• Occurrence:
  • Stratosphere (Ozone Layer): Naturally occurring in the stratosphere (10-50 km above Earth's surface), where it forms the ozone layer. This layer absorbs harmful ultraviolet (UV) radiation from the sun, protecting life on Earth.
  • Troposphere (Pollutant): In the troposphere (ground level), ozone is a harmful air pollutant. It forms from reactions involving nitrogen oxides and volatile organic compounds in the presence of sunlight, contributing to smog and respiratory problems.
• Preparation: Ozone is prepared by passing a silent electrical discharge through dry, cold oxygen. This process prevents the decomposition of ozone back into oxygen. 3O2(g) → 2O3(g) ΔH = +142 kJ/mol (endothermic reaction)
• Structure: Ozone is an angular or bent molecule with a bond angle of approximately 117°. It is a resonance hybrid of two equivalent structures, where the central oxygen atom is bonded to two other oxygen atoms, one by a single bond and the other by a double bond, with the electron density delocalized over both bonds.
• Test: Ozone turns starch-iodide paper blue. This occurs because ozone oxidizes iodide ions (I-) to iodine (I2), which then forms a blue complex with starch. O3 + 2I- + H2O → O2 + I2 + 2OH-
• Ozone Depletion:
  • CFCs as Cause: Chlorofluorocarbons (CFCs), once widely used in refrigerants and aerosols, are the primary cause of stratospheric ozone depletion. In the stratosphere, UV radiation breaks down CFCs, releasing chlorine atoms (Cl·). These chlorine atoms act as catalysts, repeatedly reacting with and destroying ozone molecules. Cl· + O3 → ClO· + O2 ClO· + O· → Cl· + O2 (Net: O3 + O· → 2O2)
  • UV Radiation Increase: Depletion of the ozone layer leads to an increase in the amount of harmful UV-B radiation reaching the Earth's surface.
  • Effects: Increased UV radiation has severe consequences, including higher rates of skin cancer, cataracts in humans, damage to marine ecosystems (phytoplankton), and reduced crop yields.
  • Montreal Protocol Control Measures: In response to ozone depletion, the Montreal Protocol on Substances that Deplete the Ozone Layer (1987) was signed. It is an international treaty designed to phase out the production and consumption of ozone-depleting substances, including CFCs. This protocol has been highly successful in reducing the release of these chemicals, leading to a gradual recovery of the ozone layer.

9.2 Nitrogen

Nitrogen is a non-metal belonging to Group 15 of the periodic table. It is the most abundant gas in Earth's atmosphere (approx. 78% by volume).

Inertness

Molecular nitrogen (N2) is remarkably inert under normal conditions. This inertness is attributed to the presence of a strong triple bond between the two nitrogen atoms (N≡N). Breaking this triple bond requires a very high amount of energy (bond dissociation enthalpy of 945 kJ/mol), making N2 unreactive at room temperature.

Active Nitrogen

Active nitrogen is a highly reactive form of nitrogen produced by passing an electrical discharge through nitrogen gas at low pressure. It consists of nitrogen atoms and excited nitrogen molecules. This active form is very reactive and can combine with metals like mercury to form nitrides, and reacts with substances like phosphorus and sulfur at room temperature, which molecular nitrogen does not.

Ammonia (NH3)

Ammonia is a colorless gas with a pungent odor, highly soluble in water, and a fundamental compound in industrial chemistry.

Properties

• Action with Copper(II) Sulfate (CuSO4): Ammonia reacts with aqueous CuSO4 to form a deep blue solution due to the formation of the tetraamminecopper(II) complex ion. CuSO4(aq) + 2NH3(aq) + 2H2O(l) → Cu(OH)2(s) + (NH4)2SO4(aq) (initial pale blue precipitate) Cu(OH)2(s) + 4NH3(aq) → [Cu(NH3)4](OH)2(aq) (deep blue soluble complex)
• Action with Water: Ammonia dissolves in water to form ammonium hydroxide, a weak base. NH3(g) + H2O(l) → NH4OH(aq) ⇌ NH4+(aq) + OH-(aq)
• Action with Iron(III) Chloride (FeCl3): Ammonia water precipitates brown hydrated iron(III) oxide from FeCl3 solution. FeCl3(aq) + 3NH4OH(aq) → Fe(OH)3(s) + 3NH4Cl(aq) (brown precipitate)
• Action with Concentrated HCl: Ammonia gas reacts with concentrated hydrochloric acid vapor to produce dense white fumes of ammonium chloride. NH3(g) + HCl(g) → NH4Cl(s)
• Action with Mercurous Nitrate Paper: Ammonia reacts with mercurous nitrate (Hg2(NO3)2) to form a black precipitate, a mixture of finely divided mercury and basic mercury(II) amido-nitrate. This is a characteristic test for ammonia. 2Hg2(NO3)2 + 4NH3 + H2O → Hg2O·NH2NO3(s) (white) + Hg(l) (black) + 3NH4NO3 (Simplified: Hg2^2+ + 2NH3 → Hg(NH2)NO3 + Hg + NH4^+)
• Combustion in Oxygen: Ammonia burns in oxygen with a yellowish flame to produce nitrogen gas and water. 4NH3(g) + 3O2(g) → 2N2(g) + 6H2O(g) In the presence of a platinum catalyst, it oxidizes to nitric oxide (Ostwald process). 4NH3(g) + 5O2(g) → 4NO(g) + 6H2O(g)

Applications

• Refrigerant: Liquid ammonia is used as a refrigerant in large industrial cooling systems due to its high latent heat of vaporization.
• Fertilizer Production: The vast majority of ammonia produced is used to manufacture nitrogenous fertilizers (e.g., urea, ammonium nitrate, diammonium phosphate).
• Cleaning Agent: Aqueous ammonia (ammonia water or household ammonia) is a common cleaning agent for glass, tiles, and other surfaces due to its ability to emulsify grease and oil.

Harmful Effects

• Toxic: Ammonia gas is toxic and corrosive. Exposure to high concentrations can cause severe irritation to the respiratory tract, eyes, and skin.
• Environmental Pollution: Excess ammonia from agricultural runoff (fertilizers) can lead to eutrophication of water bodies, harming aquatic life. Atmospheric ammonia contributes to particulate matter formation and acid rain.

Oxy-acids of Nitrogen

Nitrogen forms several oxy-acids, where nitrogen is bonded to oxygen and hydrogen (as hydroxyl groups).

• Nitrous acid (HNO2): Unstable, typically formed in solution.
• Nitric acid (HNO3): Strong acid, highly important industrially.
• Hyponitrous acid (H2N2O2)
• Nitroxyl (H2NO2)
• Pernitrous acid (HNO3, isomer of nitric acid)
• Peroxonitrous acid (HNO3, isomer of nitric acid)
• Peroxynitric acid (HNO4)
• Orthonitric acid (H3NO4)

Nitric acid (HNO3)

Nitric acid is a strong mineral acid, a powerful oxidizing agent, and a key industrial chemical.

As Acid

• Forms Nitrates: Nitric acid reacts with metals, metal oxides, hydroxides, and carbonates to form nitrate salts. CuO(s) + 2HNO3(aq) → Cu(NO3)2(aq) + H2O(l)
• Neutralizes Bases: It is a strong acid that readily neutralizes bases. NaOH(aq) + HNO3(aq) → NaNO3(aq) + H2O(l)

As Oxidizing Agent

Nitric acid is a strong oxidizing agent, with its reducing power depending on its concentration, temperature, and the nature of the reducing agent. It rarely produces hydrogen gas when reacting with metals.

• Action with Zinc (Zn):
  • Very dilute HNO3: Reduces to ammonium nitrate. 4Zn(s) + 10HNO3(very dilute) → 4Zn(NO3)2(aq) + NH4NO3(aq) + 3H2O(l)
  • Dilute HNO3: Reduces to nitrous oxide (N2O). 4Zn(s) + 10HNO3(dilute) → 4Zn(NO3)2(aq) + N2O(g) + 5H2O(l)
  • Concentrated HNO3: Reduces to nitrogen dioxide (NO2). Zn(s) + 4HNO3(conc.) → Zn(NO3)2(aq) + 2NO2(g) + 2H2O(l)
• Action with Magnesium (Mg): With very dilute nitric acid, magnesium can produce hydrogen gas. Mg(s) + 2HNO3(very dilute) → Mg(NO3)2(aq) + H2(g)
• Action with Iron (Fe): Iron reacts with dilute nitric acid to form ferrous nitrate and nitrous oxide. With concentrated nitric acid, iron becomes passive due to the formation of a protective oxide layer. 4Fe(s) + 10HNO3(dilute) → 4Fe(NO3)2(aq) + N2O(g) + 5H2O(l)
• Action with Copper (Cu):
  • Dilute HNO3: Forms nitric oxide (NO). 3Cu(s) + 8HNO3(dilute) → 3Cu(NO3)2(aq) + 2NO(g) + 4H2O(l)
  • Concentrated HNO3: Forms nitrogen dioxide (NO2). Cu(s) + 4HNO3(conc.) → Cu(NO3)2(aq) + 2NO2(g) + 2H2O(l)
• Action with Sulfur (S): Oxidizes sulfur to sulfuric acid. S(s) + 6HNO3(conc.) → H2SO4(aq) + 6NO2(g) + 2H2O(l)
• Action with Carbon (C): Oxidizes carbon to carbon dioxide. C(s) + 4HNO3(conc.) → CO2(g) + 4NO2(g) + 2H2O(l)
• Action with Sulfur dioxide (SO2): Oxidizes SO2 to sulfuric acid. SO2(g) + 2HNO3(aq) + 2H2O(l) → H2SO4(aq) + 2NO(g)
• Action with Hydrogen sulfide (H2S): Oxidizes H2S to sulfur. 3H2S(g) + 2HNO3(aq) → 3S(s) + 2NO(g) + 4H2O(l)

Ring Test for Nitrate

The brown ring test is a common qualitative test for the presence of nitrate ions (NO3-).

Procedure: To an aqueous solution containing nitrate ions, add freshly prepared ferrous sulfate (FeSO4) solution. Then, carefully add concentrated sulfuric acid (H2SO4) down the side of the test tube, allowing it to form a layer below the aqueous solution. A dark brown ring forms at the junction of the two layers, indicating the presence of nitrate ions.

Explanation: The nitrate is reduced by Fe2+ ions to nitric oxide (NO) in the presence of concentrated sulfuric acid. The NO then reacts with excess Fe2+ ions to form a brown-colored complex, pentaaquanitrosyliron(II) ion ([Fe(H2O)5NO]2+).

NO3-(aq) + 3Fe2+(aq) + 4H+(aq) → 3Fe3+(aq) + NO(g) + 2H2O(l)

[Fe(H2O)6]2+(aq) + NO(g) → [Fe(H2O)5NO]2+(aq) + H2O(l) (Brown ring complex)

9.3 Halogens

The halogens are a group of five non-metallic elements in Group 17 of the periodic table: Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I), and Astatine (At). They are highly reactive and exist as diatomic molecules in their elemental form.

General Characteristics

• Group 17 Elements: They are located in Group 17, having seven valence electrons.
• High Electronegativity: They are among the most electronegative elements, with fluorine being the most electronegative of all. This tendency to gain electrons makes them strong oxidizing agents.
• Diatomic Molecules: In their elemental state, they exist as diatomic molecules (F2, Cl2, Br2, I2) due to the formation of a single covalent bond.
• Physical State: Fluorine and chlorine are gases at room temperature, bromine is a liquid, and iodine is a solid.

Chemical Properties of Cl2, Br2, I2

Action with Water

• Chlorine (Cl2): Reacts with water to form hydrochloric acid and hypochlorous acid (chlorine water). Hypochlorous acid is unstable and decomposes to release nascent oxygen, responsible for its bleaching and germicidal properties. Cl2(g) + H2O(l) → HCl(aq) + HOCl(aq) HOCl(aq) → HCl(aq) + [O]
• Bromine (Br2): Reacts similarly to chlorine, forming hydrobromic acid and hypobromous acid, but the reaction is less extensive. Br2(l) + H2O(l) → HBr(aq) + HOBr(aq)
• Iodine (I2): Reacts very little with water; the equilibrium lies far to the left. I2(s) + H2O(l) → HI(aq) + HOI(aq)

Action with Alkali

Halogens react with alkalis (bases), with the products depending on temperature and concentration.

• Cold, Dilute Alkali (e.g., NaOH): Forms a halide and a hypohalite. Cl2(g) + 2NaOH(aq) → NaCl(aq) + NaOCl(aq) + H2O(l) (Sodium hypochlorite, used as bleach)
• Hot, Concentrated Alkali: Forms a halide and a halate. 3Cl2(g) + 6NaOH(aq) → 5NaCl(aq) + NaClO3(aq) + 3H2O(l) (Sodium chlorate)

Action with Ammonia

• Excess Ammonia: Halogens react with excess ammonia to produce ammonium halide and nitrogen gas. 3Cl2(g) + 8NH3(g) → 6NH4Cl(s) + N2(g)
• Excess Halogen: If the halogen is in excess, highly explosive nitrogen halides are formed. NH3(g) + 3Cl2(g) → NCl3(l) + 3HCl(g) (Nitrogen trichloride)

Oxidizing Character

The oxidizing power of halogens decreases down the group (F2 > Cl2 > Br2 > I2). A halogen can oxidize halide ions of less electronegative halogens.

• Cl2(aq) + 2NaBr(aq) → 2NaCl(aq) + Br2(aq) (Chlorine oxidizes bromide to bromine)
• Br2(aq) + 2KI(aq) → 2KBr(aq) + I2(aq) (Bromine oxidizes iodide to iodine)

Bleaching Action

Chlorine and bromine water exhibit bleaching action due to the formation of hypohalous acids, which release nascent oxygen, an oxidizing agent. Iodine has very weak bleaching properties.

• HOCl → HCl + [O] (Nascent oxygen)
• Coloring matter + [O] → Colorless matter

Tests for Halogens

• Chlorine (Cl2): Turns moist starch-iodide paper blue-black. Chlorine oxidizes iodide (I-) to iodine (I2), which then reacts with starch. It also has a pungent, suffocating smell.
• Bromine (Br2): Appears as orange-brown vapor (or red-brown liquid). It also turns moist starch-iodide paper blue-black (less intensely than chlorine).
• Iodine (I2): Appears as violet vapor (or shiny black solid). It turns starch solution deep blue-black (this is the most characteristic test for iodine).

Haloacids (HCl, HBr, HI)

These are binary acids formed between hydrogen and a halogen (HX).

Preparation

• Hydrogen Chloride (HCl):
  • Laboratory: Reaction of concentrated sulfuric acid with sodium chloride. NaCl(s) + H2SO4(conc.) → NaHSO4(s) + HCl(g)
  • Industrial: Direct synthesis from hydrogen and chlorine. H2(g) + Cl2(g) → 2HCl(g)
• Hydrogen Bromide (HBr) and Hydrogen Iodide (HI): Cannot be prepared by reacting metal halides with concentrated sulfuric acid because H2SO4 is a strong enough oxidizing agent to oxidize HBr and HI to Br2 and I2, respectively. Instead, they are prepared by the hydrolysis of phosphorus trihalides. PX3 + 3H2O → H3PO3 + 3HX (where X = Br, I) e.g., PBr3 + 3H2O → H3PO3 + 3HBr

Properties

• Reducing Strength: Increases down the group (HCl < HBr < HI). This is because the bond strength decreases down the group (H-I is weakest), making it easier for HI to donate its hydrogen and get oxidized.
• Acidic Nature: Increases down the group (HCl < HBr < HI). This is due to the decreasing bond dissociation enthalpy of the H-X bond, making proton release easier.
• Solubility: All haloacids are highly soluble in water, forming strong acids.
• Thermal Stability: Decreases down the group (HCl > HBr > HI). HI decomposes most easily upon heating.

Uses

• Hydrochloric Acid (HCl): Used in the production of chlorides, cleaning metal surfaces (pickling), in leather industries, and in the production of aqua regia.
• Hydrobromic Acid (HBr): Used in the synthesis of organic bromo compounds, pharmaceuticals, and dyes.
• Hydriodic Acid (HI): Used as a strong reducing agent in organic synthesis, and in the preparation of alkyl iodides.

9.4 Carbon and Phosphorus

Carbon

Carbon is a non-metal with unique ability to form stable bonds with itself (catenation) and with other elements, leading to a vast array of compounds.

Allotropes

Carbon exists in several allotropic forms with distinct physical and chemical properties.

• Diamond:
  • Properties: Hardest known natural substance, extremely high melting point, excellent thermal conductivity, electrical insulator, transparent, high refractive index. Each carbon atom is sp3 hybridized and covalently bonded to four other carbon atoms in a tetrahedral arrangement, forming a rigid 3D network.
  • Uses: Abrasives (cutting, grinding), drilling tools, jewelry, and specialized optical components.
• Graphite:
  • Properties: Soft, slippery, good electrical and thermal conductor, opaque, high melting point. Each carbon atom is sp2 hybridized and bonded to three other carbon atoms in hexagonal rings, forming planar layers. These layers are held together by weak van der Waals forces, allowing them to slide past each other.
  • Uses: Lubricants, electrodes, pencil leads, moderator in nuclear reactors, and in the production of graphene.
• Fullerene (C60 Buckminsterfullerene):
  • Properties: Spherical molecule resembling a soccer ball, composed of 60 carbon atoms arranged in 12 pentagons and 20 hexagons. It is a semiconductor, soluble in organic solvents, and has unique electronic and mechanical properties.
  • Uses: Potential applications in nanotechnology, medicine (drug delivery), superconductors, and catalysts.

Carbon Monoxide (CO)

Carbon monoxide is a colorless, odorless, and highly toxic gas.

• Reducing Action: Carbon monoxide is a strong reducing agent, especially at high temperatures. It is used in metallurgy to reduce metal oxides. Fe2O3(s) + 3CO(g) → 2Fe(s) + 3CO2(g)
• Reaction with Metals/Nonmetals:
  • Forms metal carbonyls with transition metals (e.g., Ni(CO)4), which are important in catalysis.
  • Burns in air with a blue flame to form carbon dioxide. 2CO(g) + O2(g) → 2CO2(g)
• Toxicity: Carbon monoxide is highly toxic because it binds irreversibly to hemoglobin in red blood cells about 200-250 times more strongly than oxygen, forming carboxyhemoglobin. This prevents hemoglobin from transporting oxygen, leading to oxygen deprivation and potentially death.
• Uses: Fuel gas (component of producer gas and water gas), reducing agent in metallurgical processes, and as a raw material in the synthesis of various organic compounds (e.g., methanol).

Phosphorus

Phosphorus is a reactive non-metal found in Group 15, existing in several allotropic forms.

Allotropes

• White Phosphorus (P4):
  • Properties: Waxy, translucent solid, highly reactive, pyrophoric (ignites spontaneously in air), glows in the dark (chemiluminescence), soluble in carbon disulfide, highly toxic. Consists of discrete P4 tetrahedral molecules with highly strained P-P bonds.
  • Uses: Production of red phosphorus, in matches, and for smoke screens.
• Red Phosphorus:
  • Properties: Amorphous or crystalline solid, less reactive and more stable than white phosphorus, non-toxic, insoluble in carbon disulfide. It is a polymeric structure formed by linking P4 tetrahedra.
  • Uses: In the striking surface of safety matches, in some fireworks, and as a flame retardant.
• Black Phosphorus:
  • Properties: Most stable allotrope, has a layered structure similar to graphite, good electrical conductor.
  • Uses: Potential applications in electronics and as a semiconductor.

Phosphine (PH3)

Phosphine is a colorless, highly toxic gas with a characteristic "rotten fish" smell.

• Preparation:
  • Laboratory: By the reaction of white phosphorus with concentrated sodium hydroxide solution in an inert atmosphere. P4(s) + 3NaOH(aq) + 3H2O(l) → PH3(g) + 3NaH2PO2(aq) (Sodium hypophosphite)
  • From Metal Phosphides: By the hydrolysis of metal phosphides (e.g., calcium phosphide). Ca3P2(s) + 6HCl(aq) → 3CaCl2(aq) + 2PH3(g)
• Properties:
  • Basic Nature: Phosphine is a weak base, reacting with acids to form phosphonium salts (e.g., PH4Cl with HCl). PH3(g) + HCl(g) → PH4Cl(s)
  • Reducing Nature: Phosphine is a strong reducing agent. It reduces metal salts (e.g., silver nitrate) to the corresponding metals. 6AgNO3(aq) + PH3(g) + 3H2O(l) → 6Ag(s) + H3PO3(aq) + 6HNO3(aq)
  • Action with Halogens: Reacts vigorously with halogens to form phosphorus trihalides or pentahalides. PH3(g) + 3Cl2(g) → PCl3(l) + 3HCl(g)
  • Action with Oxygen: Burns in air (spontaneously if impure due to P2H4) to form phosphorus pentoxide. 2PH3(g) + 4O2(g) → P2O5(s) + 3H2O(g)
• Uses: Used in "Holme's signals" (containers with calcium carbide and calcium phosphide which, when thrown into the sea, produce phosphine and acetylene that ignite spontaneously), as a fumigant for grains, and in the semiconductor industry.

9.5 Sulphur

Sulphur is a non-metal in Group 16, known for its ability to form various allotropes and a wide range of compounds.

Allotropes

Sulphur exhibits several allotropic forms, primarily:

• Rhombic Sulphur (α-sulphur):
  • Properties: Yellow, transparent, octahedral crystals, stable below 95.6 °C, most stable allotrope at room temperature, insoluble in water but soluble in carbon disulfide. Composed of S8 puckered ring molecules.
  • Uses: Starting material for other sulfur forms, used in vulcanization of rubber.
• Monoclinic Sulphur (β-sulphur):
  • Properties: Pale yellow, needle-shaped crystals, stable above 95.6 °C, transforms to rhombic sulphur below this temperature, insoluble in water, soluble in carbon disulfide. Also composed of S8 rings but arranged in a different crystal lattice.
• Plastic Sulphur (μ-sulphur):
  • Properties: Amorphous, rubbery, dark brown mass formed when molten sulfur is poured into cold water. Consists of long, helical chains of sulfur atoms. It is unstable and slowly reverts to rhombic sulfur.

Hydrogen Sulphide (H2S)

Hydrogen sulphide is a colorless, highly toxic gas with the characteristic odor of rotten eggs.

• Preparation from Kipp's Apparatus: In the laboratory, H2S is typically prepared in a Kipp's apparatus by the reaction of dilute hydrochloric acid with iron(II) sulfide. FeS(s) + 2HCl(aq) → FeCl2(aq) + H2S(g)
• Properties:
  • Acidic Nature: H2S is a weak diprotic acid in aqueous solution, dissociating to produce H+ and HS-, and then S2- ions. H2S(aq) ⇌ H+(aq) + HS-(aq)
  • Reducing Nature: H2S is a strong reducing agent because the sulfur atom in H2S is in its lowest oxidation state (-2) and can easily be oxidized to elemental sulfur (0) or higher oxidation states. 2H2S(g) + O2(g) → 2S(s) + 2H2O(l) (Incomplete combustion)
  • Analytical Reagent for Qualitative Analysis: H2S is widely used in qualitative inorganic analysis to precipitate metal sulfides from solutions. The solubility products of these sulfides vary, allowing for selective precipitation at different pH values (e.g., Group II metal sulfides precipitate in acidic medium, Group IV in basic medium).
• Uses: Used as a reagent in analytical chemistry, in the production of elemental sulfur (Claus process), and in the purification of natural gas.

Sulphur Dioxide (SO2)

Sulphur dioxide is a colorless gas with a pungent, suffocating odor.

• Properties:
  • Acidic Nature: SO2 is an acidic oxide. It dissolves in water to form sulfurous acid (H2SO3), a weak acid. SO2(g) + H2O(l) → H2SO3(aq) It reacts with bases to form sulfites. SO2(g) + 2NaOH(aq) → Na2SO3(aq) + H2O(l)
  • Reducing Nature: SO2 can act as a reducing agent, especially in the presence of strong oxidizing agents, where sulfur is oxidized from +4 to +6. SO2(g) + 2HNO3(aq) + 2H2O(l) → H2SO4(aq) + 2NO(g)
  • Oxidizing Nature: SO2 can also act as an oxidizing agent, particularly with strong reducing agents, where sulfur is reduced from +4 to 0 or -2. SO2(g) + 2H2S(g) → 3S(s) + 2H2O(l)
  • Bleaching Action: SO2 acts as a temporary bleaching agent by reducing the coloring matter. It is often used for delicate materials like wool and silk, as the bleaching effect can be reversed by oxidation. Coloring matter + SO2 + H2O → Colorless addition product
• Uses: Used in the manufacture of sulfuric acid (Contact Process), as a bleaching agent for wool and silk, as a preservative for food and wine, and as a disinfectant.

Sulphuric Acid (H2SO4)

Sulphuric acid, also known as oil of vitriol, is a strong mineral acid, highly corrosive, and one of the most important industrial chemicals.

• Contact Process: The industrial manufacture of sulfuric acid involves three main steps:
  1. Production of SO2: Burning sulfur or roasting sulfide ores (e.g., pyrite). S(s) + O2(g) → SO2(g)
  2. Catalytic Oxidation of SO2 to SO3: SO2 is oxidized to sulfur trioxide (SO3) using a vanadium(V) oxide (V2O5) catalyst at 400-450 °C and 1-2 atm pressure. 2SO2(g) + O2(g) → 2SO3(g) (Reversible, exothermic)
  3. Absorption of SO3 and Dilution: SO3 is absorbed in concentrated sulfuric acid to form oleum (fuming sulfuric acid, H2S2O7), which is then diluted with water to produce sulfuric acid of desired concentration. Directly dissolving SO3 in water forms a highly corrosive mist. SO3(g) + H2SO4(conc.) → H2S2O7(l) H2S2O7(l) + H2O(l) → 2H2SO4(aq)
• Properties:
  • Acidic Nature: A strong diprotic acid, it completely dissociates in water to produce hydrogen ions. H2SO4(aq) + H2O(l) → H3O+(aq) + HSO4-(aq) HSO4-(aq) + H2O(l) → H3O+(aq) + SO4^2-(aq)
  • Oxidizing Nature: Concentrated sulfuric acid is a strong oxidizing agent, especially when hot. It oxidizes metals and non-metals, often being reduced to SO2. Cu(s) + 2H2SO4(conc.) → CuSO4(aq) + SO2(g) + 2H2O(l) C(s) + 2H2SO4(conc.) → CO2(g) + 2SO2(g) + 2H2O(l)
  • Dehydrating Nature: Concentrated sulfuric acid is a powerful dehydrating agent, meaning it removes water from compounds. It can char organic compounds by removing hydrogen and oxygen in the ratio of water. C12H22O11(s) + H2SO4(conc.) → 12C(s) + H2SO4·11H2O(l) (Dehydration of sugar)
• Uses:
  • Lead Storage Battery: Used as the electrolyte in lead-acid batteries.
  • Fertilizer Production: Essential in the manufacture of phosphate fertilizers (e.g., superphosphate) and ammonium sulfate.
  • Detergent Production: Used in the sulfonation of organic compounds to produce detergents.
  • Other Uses: In petroleum refining, pickling of metals, manufacture of paints, pigments, and explosives.

Sodium Thiosulphate (Na2S2O3)

Sodium thiosulphate is an inorganic compound, commonly known as "hypo."

• Formula: Na2S2O3·5H2O (pentahydrate is the most common form).
• Uses:
  • Photography: Used as a fixing agent in photographic processing, where it dissolves unexposed silver halides from photographic film and paper.
  • Medicine: Used as an antidote for cyanide poisoning and as a treatment for certain skin conditions (e.g., ringworm, tinea versicolor).
  • Iodometric Titration: Widely used in analytical chemistry as a standard reducing agent in iodometric titrations to determine the concentration of oxidizing agents. It reacts quantitatively with iodine: 2Na2S2O3(aq) + I2(aq) → Na2S4O6(aq) + 2NaI(aq) (Sodium tetrathionate)