Summary of Group 15 Elements PDF

Summary of Group 15 Elements Group 15 elements, known as pnictogens, include nitrogen (N), phosphorus (P), arsenic (As), antimony (Sb), and bismuth (Bi). Their properties are influenced by trends seen in groups 13 and 14, such as increasing metallic character and the stability of lower oxidation st...

Summary of Group 15 Elements Group 15 elements, known as pnictogens, include nitrogen (N), phosphorus (P), arsenic (As), antimony (Sb), and bismuth (Bi). Their properties are influenced by trends seen in groups 13 and 14, such as increasing metallic character and the stability of lower oxidation states as you move down the group. Unlike typical ionic chemistry, these elements primarily form covalent compounds rather than simple ions. While metal nitrides and phosphides may be considered to contain N³⁻ and P³⁻ ions, their ionic nature is uncertain. The only well-defined cationic species is Bi³⁺. Nitrogen and phosphorus have extensive chemistry, forming various oxoacids, oxoanions, and compounds in oxidation states ranging from +5 to -3. However, the chemistry of arsenic, antimony, and bismuth is less extensive. Their electronic configuration follows the general pattern ns²np³, with a half-filled p orbital that contributes to their stability and lower reactivity. The penultimate shell electron count varies among elements, affecting their properties. Electronic Configurations: - N (7) → [He] 2s²2p³ - P (15) → [Ne] 3s²3p³ - As (33) → [Ar] 3d¹⁰4s²4p³ - Sb (51) → [Kr] 4d¹⁰5s²5p³ - Bi (83) → [Xe] 4f¹⁴5d¹⁰6s²6p³ This variation in electronic structure influences their reactivity and bonding characteristics. General Properties of Group 15 Elements 1. Physical State: - Nitrogen is a gas, phosphorus is a volatility solid, while arsenic, antimony, and bismuth are solids with varying melting points. - Nitrogen liquefies only at extremely low temperatures. 2. Atomic and Ionic Size: - Atomic size is smaller than Group 14 elements and increases irregularly down the group. - Poor shielding by d- and f-electrons causes irregularities in size. - Ionic size increases as expected down the group. 3. Ionization Energy (IE): - IE decreases down the group due to increasing atomic size. - Higher IE compared to Group 14 due to stable half-filled orbitals. 4. Atomic Volume and Density:  - Both increase down the group. 5. Metallic Character, Melting & Boiling Points: - Metallic character increases down the group: N, P (nonmetals); As, Sb (metalloids); Bi (metal). - Melting & boiling points increase, but Bi has an exceptionally low melting point. 6. Electronegativity: - Decreases down the group, supporting the gradual transition from nonmetallic to metallic character. 7. Oxidation States: - Expected oxidation states: +3 and +5 due to ns²np³ configuration. - Nitrogen prefers -3 oxidation state due to high electronegativity and small size. - Nitrogen does not form covalency of five due to lack of empty d-orbitals. - Stability of +5 oxidation state decreases down the group due to the inert-pair effect, while +3 oxidation state becomes more stable. - N and P can donate lone pairs to form adducts, showing +4 oxidation state in some cases. 8. Nature of Bonds: - Elements prefer covalent bonding over ionic due to high ionization energy. - N and P form nitrides and phosphides, but Sb and Bi form trivalent ions that hydrolyze in water. 9. Catenation (Self-Linking Ability): - Much weaker than Group 14 due to low bond energies. - Nitrogen catenates up to three atoms (e.g., N₂H₄, N₃⁻), while phosphorus shows more catenation. Chemical Properties of Group 15 Elements 1. Thermal Stability of Trihydrides (MH₃) - Stability decreases down the group as M-H bond length increases and bond strength weakens. - BiH₃ is highly unstable, while NH₃ is the most stable. - Reducing character increases down the group, making PH₃, AsH₃, SbH₃, and BiH₃ good reducing agents. 2. Melting & Boiling Points of Hydrides - NH₃ has the highest melting & boiling points due to strong hydrogen bonding. - From PH₃ to BiH₃, melting and boiling points increase due to stronger Van der Waals forces. 3. Basic Character of Trihydrides - All act as Lewis bases due to a lone pair on the central atom.  - NH₃ is the strongest base, but basicity decreases down the group as electron donation ability reduces. 4. Bond Angles of Trihydrides - All trihydrides are sp³ hybridized and pyramidal. - NH₃ has a bond angle of 107.3° due to lone pair-bond pair repulsions. - Bond angle decreases from NH₃ to BiH₃ due to decreasing electronegativity and weaker lone pair repulsions. 5. Hydrolysis of Trihalides - NF₃ does not hydrolyze, while trichlorides of P, As, Sb, and Bi hydrolyze. - PCl₃ & AsCl₃ hydrolyze completely, forming HCl and oxoacids/oxides. - SbCl₃ & BiCl₃ hydrolyze partially, forming oxychlorides (SbOCl, BiOCl) and HCl. 6. Formation of Pentachlorides - Only P, As, and Sb form pentahalides (sp³d hybridization) due to vacant d-orbitals. - N does not form pentahalides due to the lack of d-orbitals. - Bi does not form stable pentahalides as the +5 oxidation state is unstable due to the inert-pair effect. Analogous Behavior of Nitrogen Nitrogen differs from other Group 15 elements due to its **small size, high electronegativity, absence of vacant d-orbitals, and tendency to form multiple bonds**. Key differences include: 1. Physical State: - Nitrogen is a gas, while other elements exist as solids in allotropic forms. 2. Molecular Structure: - Nitrogen is diatomic (N≡N) with a strong triple bond. - Other elements form tetratomic (P₄, As₄, Sb₄) or layered (Bi) structures. 3. Hydrogen Bonding: - NH₃ exhibits strong H-bonding, making it more stable. - Other hydrides lack H-bonding due to lower electronegativity. 4. Oxides Formation: - Nitrogen forms multiple oxides with oxidation states from +5 to -3**. - Other elements form fewer oxides, which are mostly dimeric**. 5. Stability of Trihalides: - NCl₃ and NI₃ are unstable and explosive, while NF₃ is stable. - Other trihalides (PCl₃, AsCl₃, etc.) are fairly stable. 6. Hydrides Stability:  - NH₃ is highly stable, while hydrides of P, As, Sb, and Bi are less stable. 7. Complex Formation: - Nitrogen cannot expand its octet, so it does not form complexes. - Other elements form complexes due to vacant d-orbitals. Occurrence & Extraction of Group 15 Elements Occurrence - Nitrogen (N₂): - Makes up 78% of Earth’s atmosphere. - Found in ammonium compounds, nitrates, proteins, and fertile soil. - Phosphorus (P): - Found in plant and animal tissues, bones, teeth, and DNA (phosphate esters). - Occurs in apatite minerals (Ca₅X(PO₄)₃), including fluorapatite, chlorapatite, and hydroxyapatite. - Major deposits in North Africa, North America, Asia, and the Middle East. - Arsenic (As): - Occurs in elemental form and ores like arsenopyrite (FeAsS), realgar (As₄S₄), and orpiment (As₂S₃). - Antimony (Sb): - Rare in elemental form; mainly found in stibnite (Sb₂S₃). - Bismuth (Bi): - Found as a free element and in ores like bismuthinite (Bi₂S₃) and bismite (Bi₂O₃). Extraction - Nitrogen: - Lab method: Heating sodium nitride (NaN₃). - Commercial method: Fractional evaporation of liquid air, as N₂ (b.p. 77.2 K) evaporates first. - Phosphorus: - Extracted from phosphate rock (Ca₃(PO₄)₂) by heating with sand (SiO₂) and coke(C) in an electric furnace, producing white phosphorus (P₄). - Arsenic: - Extracted from FeAsS by heating or by air oxidation of arsenic sulfides (As₂S₃) to As₂O₃, followed by reduction with carbon. - Antimony: - Extracted from Sb₂S₃ (stibnite) by reduction with iron (Fe) or conversion to Sb₂O₃, then reduction with carbon. - Bismuth: - Extracted from Bi₂S₃ or Bi₂O₃ by reduction with carbon. - Also obtained as a by-product of Pb, Cu, Sn, Ag, and Au refining. Chemistry of Nitrogen Physical Properties - Colourless, tasteless, and odourless gas. - Slightly soluble in water (2.3 volumes dissolve in 100 volumes of water at 273 K). - Liquefies at 77.2 K and solidifies at 62.5 K under sudden pressure release. - Non-poisonous and does not support combustion. Chemical Properties - Highly inert due to the strong triple bond (N≡N), requiring 945 kJ/mol to break. - Reacts with: 1. Hydrogen (Haber’s process) under high pressure with iron and molybdenum to form ammonia (NH₃). 2. Oxygen under lightning or electric arc to form nitric oxide (NO). 3. Calcium carbide (CaC₂) at 1000°C to form calcium cyanamide (CaNCN), a fertilizer. 4. Electropositive metals (e.g., Li, Na, Ca) to form ionic nitrides (N³⁻ compounds). 5. Less electropositive metals (e.g., Group 2, 13, 14, transition metals) forming various types of nitrides. Active Nitrogen - Highly reactive form of nitrogen produced by passing a discharge through nitrogen at low pressure. - Shows brilliant luminescence, influenced by oxygen, CO₂, or CO. - Reacts with metals (Na, Hg, Zn, Cd) and nonmetals (S, P, I) but not with H₂ or O₂. - Converts NO into N₂O₃ and decomposes organic compounds to cyanogen (C₂N₂). Uses of Nitrogen and its Compounds - Inert atmosphere in industries (e.g., electronics, transistors) and laboratories. - Liquid nitrogen (b.p. 77 K) used as a coolant in freezing processes. - Important nitrogen-based chemicals: - Fertilizers (e.g., ammonium nitrates). - Nitric acid (HNO₃) & nitrates (used in explosives like TNT & nitroglycerine). - Nitrites (used in food preservation). - Cyanides & azides (e.g., airbags releasing N₂). Compounds of Nitrogen 1. Hydrides of Nitrogen - Ammonia (NH₃): Most important nitrogen compound, mainly used in fertilizers and nitric acid production.  - Produced via Haber process: N₂ + 3H₂ → 2NH₃. - Highly soluble in water, forming NH₄⁺ and OH⁻ ions. - Hydrazine (N₂H₄): A strong reducing agent. - Produced by heating ammonia with sodium hypochlorite. - Reacts with O₂, O₃, and halogens, and reduces metal salts to metals. - Hydrazoic acid (HN₃): Explosive and toxic. - Covalent azides like Pb(N₃)₂ and AgN₃ are shock-sensitive explosives, while ionic azides are more stable. 2. Oxides of Nitrogen - Nitrous Oxide (N₂O): Colorless, relatively unreactive. - Nitric Oxide (NO): Paramagnetic, unstable at 25°C. Formed by nitric acid reduction. - Dinitrogen Trioxide (N₂O₃): Exists as a blue solid, dissociates into NO and NO₂. - Nitrogen Dioxide (NO₂) & Dinitrogen Tetroxide (N₂O₄): Exist in equilibrium, NO₂ is a brown gas, while N₂O₄ is colorless. - Dinitrogen Pentoxide (N₂O₅): An oxidizing agent, decomposes into NO₂ and O₂. 3. Oxoacids and Anions of Nitrogen - Includes hyponitrous acid (H₂N₂O₂), nitrous acid (HNO₂), nitric acid (HNO₃), and peroxonitric acid (HNO₄). - Nitrous acid (HNO₂): Unstable, acts as an oxidizing and reducing agent. - Nitric acid (HNO₃): Strong acid, highly corrosive, prepared from NaNO₃ and H₂SO₄. 4. Nitrides - Metals react with nitrogen at high temperaturesbto form nitrides. - Example: 3Mg + N₂ → Mg₃N₂. - Classified as ionic, covalent, or interstitial based on bonding nature. 5. Nitrogen Halides - General formula: NX₃ (X = halogen); pentavalent compounds (NX₅) are unstable. - Fluorine analogs of hydrazine (N₂F₄) and diimine (N₂F₂) exist. 6. Oxyhalides of Nitrogen - Nitrosonium (XNO) and nitryl halides (XNO₂) (X = F, Cl, Br)**. - Prepared by reaction of halogens with NO: 2NO + X₂ → 2XNO. - Nitryl chloride (ClNO₂) formed from ClSO₃H + HNO₃ → ClNO₂ + H₂SO₄. Chemistry of Phosphorus Allotropic Forms of Phosphorus Phosphorus exists in several allotropes, with the most common being white, red, and black phosphorus. 1. White Phosphorus - Appearance: White or yellowish, waxy solid, highly poisonous. - Properties: - Soft, can be cut with a knife. - Highly reactive: Spontaneously ignites in air (low ignition temp: 308 K) and glows in the dark (phosphorescence). - Solubility: Insoluble in water but dissolves in turpentine, ether, and CS₂. - Structure: Exists as P₄ tetrahedral molecules. - Reactivity: - Burns in air to form P₄O₁₀. - Reduces acids & metal salts: With HNO₃: P₄ + 20HNO₃ → 4H₃PO₄ + 20NO₂ + 2H₂O With CuSO₄: P₄ + 8CuSO₄ + 14H₂O → 8Cu + 8H₂SO₄ + 3H₃PO₃ + 2H₃PO₄ - Forms phosphides with metals and halides with halogens. - Reacts explosively with oxidizers and sulfur. - Health Hazard: Causes phossy jaw, a severe bone disease. 2. Red Phosphorus - Preparation: Heated white phosphorus at 253°C in an inert atmosphere (e.g., CO₂, coal gas). - Properties: - Dark red, powdery, crystalline. - Non-toxic, odorless, and more stable than white phosphorus. - Higher ignition temperature (533 K), does not ignite at room temp. - Reactivity: - Burns in air above ignition temp to form P₄O₁₀. - Reacts with metals, halogens, and sulfur only on heating. - Can convert back to white phosphorus on boiling. - Structure: Polymeric chains of P₄ tetrahedra linked together. 3. Black Phosphorus - Preparation: Heated white phosphorus at 474 K under 4000 atm, catalyzed by mercury. - Properties: - Most stable form, crystalline, flaky, black. - Does not ignite even at 400°C. - Melts at 587°C. - Structure: Layered, corrugated sheets, similar to graphite, with strong in-layer P-P bonds and weak interlayer forces. Uses of Phosphorus - Fertilizers: Major component in phosphate fertilizers. - Industrial Applications:  a- Phosphoric acid (H₃PO₄): Used in fertilizers, detergents, food additives (gives soft drinks their sharp taste), and metal cleaning. b- PCl₃, POCl₃, P₄O₁₀, P₄S₁₀: Used in chemical synthesis, including flame retardants, insecticides, and nerve agents. c- Steel and phosphor bronze** production. - Consumer Products: Red phosphorus: Used in safety matches, fireworks, and smoke bombs. Chemistry of Arsenic Allotropic Forms of Arsenic Arsenic exists in three allotropes: 1. Gray Arsenic: Stable form: Metallic, brittle due to a hexagonal rhombohedral crystal structure. -Properties: Good heat conductor, fair electricity conductor. Sublimes at 790 K without melting, melts at 773 K under pressure. Vapor density studies suggest As₂ at 1975 K and As₄ at 917 K. 2. Yellow Arsenic: Produced by rapidly cooling the vapor of arsenic. 3. Black Arsenic: Formed when arsenic vapor is condensed in the presence of mercury, in an orthorhombic structure. Uses of Arsenic - Toxicity: Arsenic compounds, including salts and arsines, are highly toxic and were historically used in weedkillers, cattle dips, and poisons. Their use has become less common. - Semiconductors: Arsenic is used as a doping agent in semiconductors, particularly in Gallium Arsenide (GaAs), which is widely used in solid-state devices. - Alloys: Arsenic is used to strengthen alloys, for example, increasing the strength of lead (Pb). - Batteries: Arsenic also has applications in batteries. Chemistry of Antimony Allotropic Forms of Antimony 1. Metallic Form: Most common: Silver-white luster, brittle with a layered structure**. Melting point: 903.7 K (in the presence of CO), boiling point: 1653 K. Poor conductor** of heat and electricity. 2. Yellow Form (α-antimony): Produced by passing ozonized oxygen on liquid stibine at 183 K. Yellowish color, slowly converts to metallic antimony. Soluble in carbon disulfide. 3. Black Form (β-antimony): Amorphous and explosive. Obtained by slow electrolysis of antimony trichloride in hydrochloric acid. Considered an intermediate form between yellow and metallic antimony. Reactivity - Antimony is inert at ordinary temperatures, unaffected by air or water. - It oxidizes slowly in moist air and burns with a blue flame in air, forming antimony oxides: 4Sb + 3O₂ → Sb₄O₆ 4Sb + 4O₂ → Sb₄O₈ - Concentrated acids affect antimony, but dilute acids have no impact. Uses of Antimony - Toxicity: Antimony compounds are less toxic than others, but large doses can cause liver damage. - Medicinal use: Potassium antimony tartrate was historically used as an emetic and expectorant but has been replaced by less toxic alternatives. - Industrial uses: ​ Sb₂O₃ is used in paints, adhesives, plastics, and as a flame retardant. ​ It is also used in photoelectric devices and electrophotographic materials. Chemistry of Bismuth Form: Bismuth exists in a single form with a layered structure. Melting point: 544 K, boiling point: 1837 K. It is a white crystalline metal with a reddish tinge and metallic luster. When molten, it cools to form crystals with slight expansion. It is a fair conductor of heat and electricity and is strongly diamagnetic. - Reactivity: Bismuth is not attacked by dry air or cold, air-free water. It burns in air and decomposes steam at red heat: 4Bi + 3O₂ → 2Bi₂O₃ 2Bi + 3H₂O → Bi₂O₃ + 3H₂ It dissolves readily in both dilute and concentrated acids, including aqua regia. Uses of Bismuth - Bismuth is one of the **less toxic heavy metals**. - **Bismuth subcarbonate** (**BiO₂CO₃**) is used in **stomach remedies**, particularly for treating **peptic ulcers. Compounds of Group 15 Elements** 1. Hydrides - Group 15 elements form binary compounds with hydrogen, known as hydrides (e.g., PH3, AsH3, SbH3, BiH3). - Basicity and Stability: Hydrides of heavier elements (like PH3) are less basic and less stable than NH3. For example, PH3 is a weaker Brønsted base than NH3. - Phosphonium Salts: Phosphines can form stable salts with strong acids like HI, producing compounds like PH4I. Reactivity: Phosphine burns easily and is toxic. The hydrides of arsenic, antimony, and bismuth are unstable at elevated temperatures and are prepared by hydrolyzing metal compounds. Examples: Phosphine can be made by reacting phosphorus with sodium hydroxide or water, while arsenic hydride (arsine) is obtained by reacting sodium arsenide with water. Toxicity: The hydrazine analogs and compounds like P2H4 are highly toxic and unstable. 2. Oxides ​ Phosphorus Oxides: P4O6 is the primary oxide formed by phosphorus, which can further decompose to form complex oxides. P4O10 (phosphorus pentoxide) is significant in manufacturing phosphoric acid. ​ Arsenic & Antimony Oxides: As4O6 and Sb4O6 are similar to P4O6 and are produced by burning arsenic or antimony in air. Arsenic and antimony also form +5 oxides (As4O10, Sb4O10) when reacting with nitric acid. ​ Bismuth Oxide: Bismuth forms Bi2O3, a less important oxide compared to the others. 3. Oxoacids ​ Phosphoric Acid (H3PO4): A major industrial acid produced in large quantities. ​ Arsenic Acid (H3AsO4): Similar to phosphoric acid, but less common. ​ Bismuth Acid (H3BiO3): A very weak acid, unlike phosphoric or arsenic acid. ​ Phosphorous Acid (H3PO3): Formed when P4O6 reacts with water or when PCl3 is hydrolyzed. 4. Halides ​ Trihalides: The group VA elements form both +3 and +5 halogen compounds. These compounds, such as phosphorus trihalides (e.g., PCl3), are important as starting materials for other chemical compounds. ​ Bismuth: Does not usually form trihalides with bromine and iodine. Fertilizer Production Fertilizer production is crucial for feeding a growing global population. As arable land shrinks, effective fertilizers, especially those containing phosphates and nitrogen, are essential for increasing food production. Phosphate Fertilizers - Source: The primary resource for phosphate fertilizers is naturally occurring calcium phosphate (Ca3(PO4)2), which is abundant but insoluble. - Process: To make it soluble, calcium phosphate reacts with sulfuric acid, producing calcium dihydrogen phosphate (Ca(H2PO4)2), which is more soluble and a key component of fertilizers. This process also produces gypsum (CaSO4·2H2O). - Sulfuric Acid: Sulfuric acid, produced in large quantities, is used in the reaction to convert calcium phosphate into a soluble form, and it plays a significant role in fertilizer production. Nitrogen Fertilizers - Ammonium Nitrate: Nitrogen fertilizers like ammonium nitrate (NH4NO3) are produced by reacting nitric acid (HNO3) with ammonia (NH3). These fertilizers are used both for agricultural purposes and in explosives. Other Nitrogen Fertilizers: - Ammonium sulfate (NH4)2SO4 is made by reacting ammonia with sulfuric acid. - Ammonium phosphate (NH4)3PO4 is produced by reacting ammonia with phosphoric acid. - Urea ((NH2)2CO) is made by reacting ammonia with carbon dioxide. Dual Fertilizers: Fertilizers containing both phosphorus and nitrogen can be made by reacting calcium phosphate with nitric acid. Inorganic Chemistry and Importance - Scale: The production of these fertilizers is done on a massive scale, with sulfuric acid and ammonium nitrate being produced in billions of pounds annually. - Global Need: As the world population is projected to reach 12 billion by 2030, the importance of these fertilizers in food production will only increase.