p-Block Elements – Comprehensive Notes

The p-block elements are located in Groups 13 to 18 of the periodic table, occupying the right side. Their valence electrons are located in the outermost p-orbital. The general valence shell electronic configuration is ns2np1−6. These elements exhibit a wide variation in properties, ranging from highly metallic to distinctly non-metallic, with metalloids bridging the two.

1. General Characteristics of p-Block Elements

• Electronic Configuration: The outer electronic configuration is ns2np1−6. This unique configuration leads to a diverse range of chemical behaviors.
• Metallic Character: Metallic character increases as you go down a group and significantly decreases as you move across a period (from left to right). Non-metals are found predominantly on the top right side of the p-block (e.g., Nitrogen, Oxygen, Fluorine), metals on the bottom left (e.g., Aluminium, Lead, Bismuth), and metalloids (e.g., Boron, Silicon, Germanium, Arsenic, Antimony, Tellurium) form a diagonal line separating them.
• Oxidation States: These elements exhibit a variety of oxidation states, often showing both positive and negative values. A common trend is that the group oxidation state (equal to the group number, e.g., +3 for Group 13, +5 for Group 15) and an oxidation state of (Group number – 2) are frequently observed. The latter is due to the inert pair effect, where the ns2 electrons become increasingly reluctant to participate in bonding down the group, especially for heavier elements.
• Ionization Enthalpy: Generally high compared to s-block elements, but shows a decreasing trend down a group due to increasing atomic size and shielding effect. However, there are some irregularities. For example, the ionization enthalpy of Group 16 elements is lower than that of Group 15 elements in the same period due to the extra stability of the half-filled p-orbitals in Group 15 elements.
• Electronegativity: High electronegativity values, generally increasing across a period and decreasing down a group. Fluorine (Group 17) is the most electronegative element in the entire periodic table.
• Non-metallic Character: Non-metallic character increases across a period and decreases down a group. This is directly related to electronegativity and ionization enthalpy trends.
• Catenation: This is the unique ability of an atom to form bonds with other atoms of the same element to create long chains or rings. Carbon exhibits the strongest catenation, forming a vast number of organic compounds. Silicon and Nitrogen also show catenation to a lesser extent, while it is almost negligible for heavier elements in most p-block groups.
• Allotropy: Many p-block elements exist in more than one physical form, known as allotropes, which have different structures and properties. Examples include Carbon (diamond, graphite, fullerenes), Sulfur (rhombic, monoclinic), Phosphorus (white, red, black), and Oxygen (O2, O3).

2. Group 13 Elements: Boron Family (B, Al, Ga, In, Tl)

2.1. General Characteristics and Trends

• Electronic Configuration: ns2np1.
• Atomic and Ionic Radii: Generally increase down the group. Anomaly: Gallium (Ga) has a slightly smaller atomic radius than Aluminium (Al). This is due to the presence of 10 d-electrons in gallium, which provide poor shielding, leading to a stronger effective nuclear charge pulling the valence electrons closer.
• Ionization Enthalpy: Generally decreases down the group. Irregularities: IE of Ga > Al and IE of Tl > In. These anomalies are attributed to the ineffective shielding by the d and f electrons in the heavier elements, leading to increased effective nuclear charge.
• Electronegativity: Generally decreases down the group. Boron (B) has a relatively high electronegativity, accounting for its covalent nature.
• Oxidation States: All elements show a +3 oxidation state. However, due to the inert pair effect, heavier elements like Gallium, Indium, and especially Thallium also show a +1 oxidation state. The stability of the +1 oxidation state increases significantly down the group (e.g., Tl+ is much more stable than Tl3+, while B+ is highly unstable). This is because the ns2 electrons become more tightly bound and less available for bonding.
• Acidity of Oxides: The acidic character of oxides decreases down the group, and basic character increases.
  • Boron oxide (B2O3): Acidic
  • Aluminium oxide (Al2O3) and Gallium oxide (Ga2O3): Amphoteric
  • Indium oxide (In2O3) and Thallium oxide (Tl2O3): Basic

2.2. Important Compounds

• Borax (Sodium tetraborate decahydrate), Na2B4O7.10H2O:
  • Structure: Contains complex tetranuclear anion [B4O5(OH)4]2- units, with two triangular BO3 and two tetrahedral BO4 units linked.
  • Reactions:
    • On heating: Borax swells up (due to loss of water of crystallization) and then forms a transparent glassy bead called borax glass (a mixture of sodium metaborate and boric anhydride). Na2B4O7.10H2O (heat) -> Na2B4O7 (swells) (further heat) -> 2NaBO2 (sodium metaborate) + B2O3 (boric anhydride, glassy bead)
    • Borax Bead Test: The glassy bead can react with transition metal oxides to form colored metaborates, used for qualitative analysis. B2O3 + CoO (cobalt oxide) (heat) -> Co(BO2)2 (cobalt metaborate, gives a blue bead)
  • Uses: In borax bead test for identifying colored metal ions, in the manufacture of heat-resistant glasses (e.g., Pyrex), ceramics, enamels, and as a flux in soldering.
• Boric Acid, H3BO3 (or B(OH)3):
  • Structure: Planar BO3 units are linked together by extensive hydrogen bonding, forming a layered structure.
  • Nature: It is a weak monobasic Lewis acid, not a proton donor (Arrhenius acid). Instead, it accepts a hydroxide ion (OH-) from water. B(OH)3 + 2H2O -> [B(OH)4]- + H3O+
  • Preparation: From borax by reacting with sulfuric acid. Na2B4O7 + H2SO4 + 5H2O -> Na2SO4 + 4H3BO3
  • Uses: As a mild antiseptic, in eye washes (e.g., Boroline), as an insecticide, and in the production of glazes and enamels.
• Diborane, B2H6:
  • Structure: An electron-deficient molecule. It has a unique “banana bond” or “3-center-2-electron” bond structure. It contains two bridging hydrogen atoms and four terminal hydrogen atoms. The two boron atoms and the four terminal hydrogen atoms lie in one plane, while the two bridging hydrogen atoms are perpendicular to this plane.
  • Reactions: Highly flammable gas, readily hydrolyzed. B2H6 + 6H2O -> 2H3BO3 + 6H2
  • Uses: As a high-energy fuel for rockets, a reducing agent in organic synthesis, and in the preparation of other boron hydrides.
• Aluminium Chloride, AlCl3:
  • Structure: Anhydrous AlCl3 is a covalent compound. In the vapor phase below 473 K, it exists as a dimeric molecule (Al2Cl6) with two chlorine atoms forming bridges. In the solid state, it forms a polymeric sheet structure. Above 473 K, the dimer dissociates into monomeric AlCl3.
  • Nature: It is a strong Lewis acid due to the presence of an empty p-orbital on the aluminium atom, enabling it to accept a lone pair of electrons.
  • Uses: A very important catalyst in various organic reactions, particularly in Friedel-Crafts alkylation and acylation reactions.

3. Group 14 Elements: Carbon Family (C, Si, Ge, Sn, Pb)

3.1. General Characteristics and Trends

• Electronic Configuration: ns2np2.
• Atomic and Ionic Radii: Consistently increase down the group.
• Ionization Enthalpy: Decreases down the group. However, there are minor irregularities (e.g., IE of Pb is slightly higher than Sn) due to the poor shielding effect of d and f electrons in heavier elements.
• Electronegativity: Electronegativity values are relatively constant down the group (around 1.8-1.9), indicating a transition from non-metallic to metallic character.
• Oxidation States: Exhibit both +4 and +2 oxidation states.
  • The stability of the +4 oxidation state decreases down the group (C and Si mostly +4).
  • The stability of the +2 oxidation state increases down the group due to the inert pair effect (e.g., Pb2+ is much more stable than Pb4+, while C2+ is highly unstable).
• Catenation: The ability to form long chains and rings with similar atoms is strongest for Carbon, leading to the vast field of organic chemistry. Catenation decreases significantly for Silicon and Germanium, and is almost absent for Tin and Lead.
• Allotropy: Many elements in this group show allotropy:
  • Carbon: Diamond (hardest natural substance, insulator), Graphite (soft, good conductor, layered structure), Fullerenes (cage-like structures like C60).
  • Silicon and Germanium: Exist in amorphous and crystalline forms.
  • Tin: Has two main allotropic forms, grey tin (α-tin, brittle, non-metallic) and white tin (β-tin, metallic).
• Acidity of Oxides: The acidity of oxides generally decreases down the group, with metallic oxides being amphoteric.
  • Carbon monoxide (CO) and Carbon dioxide (CO2) are acidic.
  • Silicon dioxide (SiO2) is acidic.
  • Germanium dioxide (GeO2) is weakly acidic.
  • Tin dioxide (SnO2) and Lead dioxide (PbO2) are amphoteric.

3.2. Important Compounds

• Carbon Monoxide, CO:
  • Preparation: Incomplete combustion of carbon. Also by heating formic acid in the presence of concentrated sulfuric acid. C + 1/2 O2 -> CO HCOOH (conc. H2SO4, heat) -> CO + H2O
  • Nature: A neutral oxide. It is a highly poisonous gas because it forms a stable complex with hemoglobin (carboxyhemoglobin), preventing oxygen transport.
  • Uses: A powerful reducing agent in metallurgy (e.g., in the blast furnace for reduction of iron oxides), and as a fuel.
• Carbon Dioxide, CO2:
  • Preparation: Complete combustion of carbon or carbon-containing compounds. Also by heating carbonates (e.g., CaCO3).
  • Nature: An acidic oxide (dissolves in water to form carbonic acid).
  • Uses: Crucial for photosynthesis, used in fire extinguishers (as dry ice, solid CO2), as a refrigerant, and in carbonated beverages.
• Silicon Dioxide (Silica), SiO2:
  • Structure: A three-dimensional network solid where each silicon atom is tetrahedrally bonded to four oxygen atoms, and each oxygen atom is shared by two silicon atoms. This forms a giant covalent structure.
  • Nature: An acidic oxide.
  • Uses: Major component of glass, ceramics, and cement. Also used in semiconductors (pure silicon) and as a desiccant (silica gel).
• Silicones:
  • Organosilicon polymers containing repeated Si-O-Si linkages (siloxane linkages).
  • Properties: They are water repellent, have high thermal stability, are good electrical insulators, and are chemically inert.
  • Uses: Used as sealants, lubricants, water-proofing agents, high-temperature greases, and in cosmetics and surgical implants.

4. Group 15 Elements: Nitrogen Family (N, P, As, Sb, Bi)

4.1. General Characteristics and Trends

• Electronic Configuration: ns2np3. (Stable half-filled p-orbitals).
• Atomic and Ionic Radii: Increase down the group.
• Ionization Enthalpy: Very high due to stable half-filled p-orbitals. Generally decreases down the group.
• Electronegativity: High, but decreases down the group.
• Oxidation States: Can range from -3 to +5.
  • Stability of +5 oxidation state decreases down the group.
  • Stability of +3 oxidation state increases down the group due to the inert pair effect (e.g., Bi3+ is more stable than Bi5+).
  • Nitrogen exhibits a wide range of oxidation states (-3, -2, -1, +1, +2, +3, +4, +5) due to its small size and high electronegativity.
• Allotropy: Phosphorus (P) exists in several allotropic forms:
  • White Phosphorus (P4): Highly reactive, waxy solid, soluble in CS2, glows in the dark (chemiluminescence), highly poisonous. Tetrahedral P4 molecule.
  • Red Phosphorus: Polymeric structure of P4 units. Less reactive, insoluble in CS2, non-poisonous.
  • Black Phosphorus: Most stable form.
  • Arsenic (As) and Antimony (Sb) also show allotropy.
• Hydrides (MH3): Ammonia (NH3), Phosphine (PH3), Arsine (AsH3), Stibine (SbH3), Bismuthine (BiH3).
  • Basic Character: Decreases down the group (NH3 > PH3 > AsH3 > SbH3 > BiH3) due to decreasing availability of the lone pair on the central atom as size increases.
  • Reducing Character: Increases down the group (BiH3 is the strongest reducing agent) due to decreasing M-H bond strength.
  • Thermal Stability: Decreases down the group due to decreasing M-H bond strength.
• Oxides: Form E2O3 and E2O5 type oxides.
  • Acidity of E2O3: Decreases down the group: N2O3 (acidic), P4O6 (acidic), As4O6 (amphoteric), Sb4O6 (amphoteric), Bi2O3 (basic).
  • Acidity of E2O5: Decreases down the group: N2O5 (acidic), P4O10 (acidic), As4O10 (acidic), Sb4O10 (amphoteric), Bi2O5 (basic).

4.2. Important Compounds

• Ammonia, NH3:
  • Preparation (Haber’s Process): Industrial preparation for ammonia. N2 (g) + 3H2 (g) ⇌ 2NH3 (g) This is an exothermic and reversible reaction. Optimum conditions for good yield are high pressure (200 atm), moderate temperature (700 K), and the use of an iron catalyst with molybdenum as a promoter.
  • Nature: A colorless gas with a pungent smell. It is a weak base, forming ammonium salts with acids.
  • Uses: Primarily in the manufacture of fertilizers (e.g., urea, ammonium sulfate, ammonium nitrate), nitric acid, and as a refrigerant.
• Nitric Acid, HNO3:
  • Preparation (Ostwald Process): Industrial preparation of nitric acid involves three main steps:
    1. Catalytic oxidation of ammonia: 4NH3 (g) + 5O2 (g) Pt/Rhcatalyst,500K,9bar​ 4NO (g) + 6H2O (g)
    2. Oxidation of nitric oxide: 2NO (g) + O2 (g) -> 2NO2 (g)
    3. Absorption of nitrogen dioxide in water: 3NO2 (g) + H2O (l) -> 2HNO3 (aq) + NO (g) (NO is recycled)
  • Nature: A strong monobasic acid and a powerful oxidizing agent.
  • Uses: In the manufacture of fertilizers (e.g., ammonium nitrate), explosives (e.g., TNT, nitroglycerine), and in the nitration of organic compounds.
• Phosphorus Allotropes: (As described in General Characteristics). White phosphorus is very dangerous due to its reactivity and toxicity.
• Phosphine, PH3:
  • Preparation: By the reaction of calcium phosphide with water or dilute acids. Ca3P2 + 6H2O -> 3Ca(OH)2 + 2PH3
  • Nature: A colorless, highly poisonous gas with a rotten fish smell. It is weakly basic.
  • Uses: In Holme’s signals (containers with CaC2 and Ca3P2 are thrown into the sea, releasing acetylene and phosphine which ignites spontaneously, serving as a smoke screen).
• Phosphorus Halides (PCl3, PCl5): These are important chlorinating agents used in organic synthesis.
• Oxoacids of Phosphorus: Important oxoacids include H3PO2 (hypophosphorous acid), H3PO3 (phosphorous acid), and H3PO4 (phosphoric acid). Their acidity depends on the number of P-OH groups, and their reducing character depends on the number of P-H bonds.

5. Group 16 Elements: Oxygen Family (Chalcogens) (O, S, Se, Te, Po)

5.1. General Characteristics and Trends

• Electronic Configuration: ns2np4.
• Atomic and Ionic Radii: Increase down the group.
• Ionization Enthalpy: Decreases down the group. They generally have lower ionization enthalpies than Group 15 elements in the same period due to the less stable electron configuration (np4 vs np3).
• Electronegativity: High (Oxygen is the second most electronegative element, after Fluorine). Electronegativity decreases down the group.
• Electron Gain Enthalpy: Generally negative (exothermic). Sulfur has a more negative electron gain enthalpy than Oxygen, due to the very small size of Oxygen causing electron-electron repulsion.
• Oxidation States:
  • Oxygen: Predominantly -2 (most common). Also -1 (in peroxides), -1/2 (in superoxides), +1 (in O2F2), +2 (in OF2). Oxygen does not show +4 or +6 due to the absence of d-orbitals.
  • Other elements: -2, +2, +4, +6. The stability of the +6 oxidation state decreases down the group, while the stability of the +4 oxidation state increases (inert pair effect).
• Allotropy: Oxygen (O2 – dioxygen, O3 – ozone), Sulfur (rhombic, monoclinic, plastic, colloidal forms), Selenium, Tellurium all exhibit allotropy.
• Hydrides (H2E): Water (H2O), Hydrogen sulfide (H2S), Hydrogen selenide (H2Se), Hydrogen telluride (H2Te), Hydrogen polonide (H2Po).
  • Acidic Character: Increases down the group (H2O < H2S < H2Se < H2Te) due to decreasing bond dissociation enthalpy.
  • Reducing Character: Increases down the group (H2Te is the strongest reducing agent).
  • Thermal Stability: Decreases down the group.
  • Boiling Point: Water (H2O) has an abnormally high boiling point compared to other hydrides due to extensive intermolecular hydrogen bonding.
• Oxides: Form EO2 and EO3 types.
  • EO2 are generally acidic (SO2, SeO2, TeO2).
  • EO3 are generally acidic (SO3, SeO3).

5.2. Important Compounds

• Oxygen, O2 (Dioxygen):
  • Preparation: From liquid air by fractional distillation (separation of O2 and N2).
  • Uses: Essential for respiration, combustion, welding (oxy-acetylene flame), in steel manufacturing, and medical oxygen.
• Ozone, O3:
  • Preparation: By passing silent electric discharge through dry oxygen. (This prevents decomposition of ozone back to oxygen). 3O2(g)silentelectricdischarge​2O3(g)
  • Nature: A pale blue, pungent gas. It is a powerful oxidizing agent.
  • Uses: As a germicide, for sterilizing water, bleaching oils, starch, and fabrics. Crucially, it forms the ozone layer in the stratosphere, protecting Earth from harmful UV radiation.
• Sulfur Allotropes: Rhombic sulfur (alpha-sulfur) is the most stable and common allotrope, existing as S8 rings. Monoclinic sulfur (beta-sulfur) is stable above 369 K.
• Sulfuric Acid, H2SO4:
  • Preparation (Contact Process): The most important industrial method.
    1. Burning of sulfur or sulfide ores to produce sulfur dioxide: S+O2→SO2 or 2ZnS+3O2→2ZnO+2SO2
    2. Catalytic oxidation of sulfur dioxide to sulfur trioxide: 2SO2(g)+O2(g)V2O5catalyst​2SO3(g) (This is the key reversible step)
    3. Absorption of sulfur trioxide in concentrated sulfuric acid to form oleum (pyrosulfuric acid): SO3(g)+H2SO4(l)→H2S2O7(l) (oleum)
    4. Dilution of oleum with water to produce sulfuric acid: H2S2O7(l)+H2O(l)→2H2SO4(l)
  • Nature: A strong dibasic acid, a powerful dehydrating agent, and a strong oxidizing agent (especially when hot and concentrated).
  • Uses: Known as the “King of Chemicals” due to its widespread use in fertilizers (e.g., ammonium sulfate, superphosphate), petroleum refining, detergents, metallurgy, and in lead-acid storage batteries.

6. Group 17 Elements: Halogen Family (F, Cl, Br, I, At)

6.1. General Characteristics and Trends

• Electronic Configuration: ns2np5. (One electron short of a stable noble gas configuration, hence highly reactive).
• Atomic and Ionic Radii: Increase down the group.
• Ionization Enthalpy: High (due to strong effective nuclear charge), decreases down the group.
• Electronegativity: Extremely high. Fluorine is the most electronegative element in the entire periodic table. Electronegativity decreases down the group.
• Electron Gain Enthalpy: Highly negative values (highly exothermic process). Chlorine has the most negative electron gain enthalpy among all elements (more negative than Fluorine due to the very small size of F and significant electron-electron repulsion in its 2p orbital).
• Oxidation States: Exhibit -1 (most common due to high electronegativity). Other positive oxidation states (+1, +3, +5, +7) are also shown by Cl, Br, I due to the availability of d-orbitals for expansion of octet. Fluorine, being the most electronegative, only shows -1 oxidation state.
• Bond Dissociation Enthalpy: Generally decreases down the group. Exception: F2 has a lower bond dissociation enthalpy than Cl2 (due to strong lone pair-lone pair repulsions between the small fluorine atoms in F2, leading to weaker F-F bond). Order: Cl2 > Br2 > F2 > I2.
• Physical State: At room temperature: F2 (pale yellow gas), Cl2 (greenish yellow gas), Br2 (reddish brown liquid), I2 (violet solid, sublimes easily).
• Colour: They are colored due to the absorption of radiation in the visible region, causing excitation of outer electrons.
• Reducing Character of Halides (HX): Increases down the group (HI > HBr > HCl > HF) due to decreasing H-X bond strength.
• Oxidizing Nature: Strong oxidizing agents (tendency to accept electrons and get reduced). Oxidizing power decreases down the group. Order: F2 > Cl2 > Br2 > I2.

6.2. Important Compounds

• Hydrohalic Acids (HX): HF, HCl, HBr, HI.
  • Acidic Strength: HF < HCl < HBr < HI. This trend is explained by decreasing H-X bond dissociation enthalpy down the group. HF is a weak acid due to strong hydrogen bonding and high bond dissociation enthalpy.
  • Boiling Point: HF has the highest boiling point among HX due to extensive intermolecular hydrogen bonding.
• Chlorine, Cl2:
  • Preparation (Deacon’s Process): Catalytic oxidation of HCl. 4HCl (g) + O2 (g) CuCl2catalyst,723K​ 2Cl2 (g) + 2H2O (g)
  • Preparation (Electrolysis of Brine): Industrially, chlorine is obtained as a byproduct during the electrolysis of concentrated aqueous sodium chloride solution (brine), which produces NaOH.
  • Uses: A powerful bleaching agent (for paper pulp, textiles), used for sterilization of drinking water, and in the manufacture of PVC (polyvinyl chloride) and CFCs (chlorofluorocarbons).
• Hydrogen Chloride, HCl:
  • Nature: A colorless gas, which forms hydrochloric acid (a strong acid) when dissolved in water.
  • Uses: In the pickling of steel (removing oxide layers), production of glucose from corn starch, and in the manufacture of dyes and drugs.
• Oxoacids of Halogens: These acids contain halogen, oxygen, and hydrogen (e.g., Hypohalous acids (HOX), Halous acids (HXO2), Halic acids (HXO3), Perhalic acids (HXO4)).
  • Acidic Strength (for a given halogen): Increases with increasing oxidation state of the halogen (e.g., HOCl < HClO2 < HClO3 < HClO4). This is due to increasing stability of the conjugate base by dispersal of negative charge with more oxygen atoms.
  • Acidic Strength (for a given oxidation state): Decreases down the group (e.g., HOCl > HOBr > HOI) due to decreasing electronegativity of the halogen.
• Interhalogen Compounds: Formed between different halogens (e.g., ClF, BrF3, IF5, IF7).
  • Highly reactive (more reactive than halogens themselves, except F2).

7. Group 18 Elements: Noble Gases (He, Ne, Ar, Kr, Xe, Rn)

7.1. General Characteristics and Trends

• Electronic Configuration: ns2np6 (fully filled valence shells, except Helium: 1s2). This complete octet (or duplet for He) is responsible for their extreme stability.
• Reactivity: Traditionally known as “inert gases” because they are extremely unreactive due to their stable electronic configuration, very high ionization enthalpy, and positive (unfavorable) electron gain enthalpy.
• Atomic Radii: Largest in their respective periods (measured as van der Waals radii, which are larger than covalent or metallic radii). Atomic radii increase down the group.
• Ionization Enthalpy: Very high due to stable electron configuration. Decreases down the group as atomic size increases.
• Electron Gain Enthalpy: Positive values, meaning energy is absorbed when an electron is added, indicating no tendency to accept electrons.
• Melting and Boiling Points: Very low, as only weak London dispersion forces (van der Waals forces) exist between their monatomic molecules. Melting and boiling points increase down the group with increasing atomic mass and thus increasing London forces.
• Physical State: All are monatomic gases at room temperature.
• Abundance: Argon is the most abundant noble gas in the atmosphere (about 1% by volume). Helium is abundant in natural gas.

7.2. Chemistry of Noble Gases

• Discovery of Reactivity: Until 1962, they were considered completely unreactive. Neil Bartlett’s synthesis of the first noble gas compound, XePtF6 (Xenon hexafluoroplatinate(V)), revolutionized this understanding.
• Xenon Compounds: Xenon (Xe) is the most reactive noble gas (due to its relatively large size and lower ionization enthalpy) and forms compounds primarily with highly electronegative elements like Fluorine and Oxygen.
  • Xenon Fluorides: XeF2, XeF4, XeF6.
    • Preparation: Direct reaction of Xenon with Fluorine under specific conditions.
      • Xe + F2 (excess F2) (heat 673K, 1 bar) -> XeF2 (Linear structure)
      • Xe + 2F2 (1:5 molar ratio) (heat 873K, 7 bar) -> XeF4 (Square Planar structure)
      • Xe + 3F2 (1:20 molar ratio) (heat 573K, 60-70 bar) -> XeF6 (Distorted Octahedral structure, due to one lone pair)
  • Xenon Oxides: XeO3, XeO4.
    • Preparation: Hydrolysis of Xenon fluorides.
      • XeF6 + 3H2O -> XeO3 (Pyramidal structure) + 6HF
      • XeF8 (hypothetical) + 4H2O -> XeO4 (Tetrahedral structure) + 8HF
  • Xenon Oxofluorides: XeOF4, XeO2F2.
    • Preparation: Partial hydrolysis of XeF6.
      • XeF6 + H2O -> XeOF4 (Square Pyramidal structure) + 2HF
      • XeF6 + 2H2O -> XeO2F2 (Trigonal Bipyramidal with equatorial lone pair) + 4HF
• Uses:
  • Helium (He): Non-flammable and lighter than air. Used in meteorological balloons, cryogenics (low-temperature research), for inflating tires of aircraft, and as a diluent for oxygen in diving apparatus (to prevent bends).
  • Neon (Ne): Used in neon sign tubes (produces a distinct reddish-orange glow), discharge tubes, and fluorescent bulbs.
  • Argon (Ar): Used to provide an inert atmosphere for welding (e.g., arc welding) and in electric bulbs (to prevent oxidation of the filament), and in metallurgical processes.
  • Krypton (Kr) & Xenon (Xe): Used in high-intensity discharge lamps (e.g., headlights) and photographic flash tubes.
  • Radon (Rn): A radioactive noble gas. Used in cancer therapy.