Nitrogen (N₂)
Introduction to Nitrogen
Nitrogen belongs to group 15 of the periodic table. In industrial processes, nitrogen is typically obtained by cooling air until it becomes a liquid. Liquid nitrogen is commonly used for rapid cooling purposes. In laboratory settings, nitrogen can be generated by slowly heating a solution of ammonium nitrite.
| Property | Value | Property | Value |
|---|---|---|---|
| Atomic number | 7 | Ionic radius | 171 pm |
| Relative atomic mass | 14.007 a.m.u | 1st Ionization energy | 14.02 kJ/mol |
| Physical appearance | Colourless Gas | Electronegativity | 3.0 |
Reactivity of Nitrogen
Nitrogen is a significant component of the air, known for its low reactivity due to its small size, symmetrical electronic cloud, and nonpolar triple bond. With an electronic configuration of 1s² 2s² 2p³, nitrogen requires three electrons to complete its octet, forming a triple bond by sharing three electrons with another nitrogen atom.
The bond between two nitrogen atoms has a bond enthalpy of +944 kJmol⁻¹. High energy is required to break this bond to form new bonds, making N₂ very unreactive. Another reason for its lack of reactivity is the non-polarity of its bond. Both atoms are the same having zero electronegativity difference.
Ammonia (NH₃)
Preparation and Properties
Ammonia (NH₃) is an important industrial compound of nitrogen, which is mainly used as a fertilizer. It is prepared industrially by Haber-Bosch process. In the laboratory, ammonia gas can be synthesized by heating an ammonium salt such as ammonium chloride (NH₄Cl) with a base like calcium hydroxide (Ca(OH)₂).
If a gas with a pungent smell is released and turns moistened red litmus paper blue, it indicates the presence of ammonium in the compound.
Structure and Basicity
Ammonia molecule has pyramidal shape due to lone pair of nitrogen. But when nitrogen atom in ammonia utilizes this lone pair of electrons to form ammonium, this ion adopts a tetrahedral shape in which all the bonds are of equal length and strength.
Ammonia behaves as a Lowry-Bronsted base by accepting a proton (H⁺) from an acid to form ammonium. It dissolves in water to form ammonium hydroxide (NH₄OH) and equilibrium is established between ammonia molecules and ammonium ions in the solution. Ammonia solution is a weak base due to the low basicity constant (Kb) and the equilibrium position being towards the far left side.
Oxides of Nitrogen
Common Oxides
Oxides of nitrogen are NO, N₂O, NO₂, N₂O₄ and N₂O₅ in which oxidation states range from I to V. N₂O₄ and N₂O₅ decay quickly to other oxides. NO and NO₂ are collectively called as NOₓ.
| Name and formula | Formal oxidation state | Properties | Uses |
|---|---|---|---|
| Nitrous oxide N₂O | 1+ | Colourless gas, water-soluble, neutral, sweet smelling | Dental anesthetic, propellant for whipped ice cream |
| Nitric oxide NO | 2+ | Colourless gas, slightly water-soluble, paramagnetic | Biochemical messenger, synthesis of nitrosyl carbonyls |
| Nitrogen dioxide NO₂ | 4+ | Reddish-brown gas, paramagnetic, reacts with water | Rocket propellant, HNO₃ formation by Ostwald process |
Sources and Environmental Impact
Natural sources include lightning, volcanoes, biological decay, forest fires, and denitrifying bacteria in soil. The main anthropogenic sources of NOₓ are the combustion of fossil fuels in vehicles and power plants. Other sources include chemical plants, biomass burning, welding, etc.
Photochemical smog (Los Angeles smog) forms in the atmosphere from NOₓ and volatile organic compounds (VOCs) in the sunlight. It contains photochemical oxidants, such as NO₂, ozone, and peroxyacetyl nitrates (PANs).
Sulfur
Physical Properties
Sulfur is a member of group 16 which is also called the Chalcogen family. Sulfur usually forms single bonds with other sulfur atoms instead of double bonds due to poor overlapping of the orbitals. As a result, it forms larger molecules and structures through a process called catenation. S₈ is a crown-like molecule.
| Property | Value | Property | Value |
|---|---|---|---|
| Atomic number | 16 | Ionic radius | 184 pm |
| Relative atomic mass | 32.06 a.m.u. | 1st Ionization energy | 1000 kJ/mol |
| Physical appearance | Solid Yellow | Electronegativity | 2.5 |
| Melting point | 113°C (honey-Yellow) | Common oxidation states | 4+ and 6+ |
Oxidation States and Reactions
Sulfur exhibits oxidation states of −2, 0, +2, +4, and +6. Under standard conditions, sulfur and oxygen react to produce sulfur dioxide (SO₂) in which sulphur has an oxidation state of +4. To form sulfur trioxide (SO₃) with an oxidation state of +6, high energy is required.
SO₂ + ½O₂ → SO₃ (Requires catalyst)
Sulfur can combine with many elements to form a wide variety of inorganic and organic compounds. It is unreactive to water under normal conditions, dilute non-oxidizing acids, and noble gases. Its ability to catenate allows it to form ring structures and linear chains.
Sulfuric Acid (H₂SO₄)
Contact Process
The major portion of sulfur, around 85% is used for the production of sulphuric acid (H₂SO₄). Sulphuric acid is produced by the contact process. The process starts with the combustion of molten sulphur or by heating pyrites such as iron pyrite (FeS) in excess of air to produce sulfur dioxide SO₂.
4FeS₂ + 11O₂ → 2Fe₂O₃ + 8SO₂
2SO₂ + O₂ ⇌ 2SO₃ (V₂O₅ catalyst, 450°C)
SO₃ + H₂SO₄ → H₂S₂O₇ (Oleum)
H₂S₂O₇ + H₂O → 2H₂SO₄
Purified SO₂ and air, preheated at 420°C-450°C, are fed to the first converter stage of the contact tower at 1-2 atm pressure. Here, these gases come in contact with vanadium pentoxide (V₂O₅) catalyst.
Properties and Uses
Sulphuric acid is soluble in water and hygroscopic in nature. It readily absorbs water vapour from the air. Anhydrous H₂SO₄ is a very polar liquid. It is highly corrosive to various materials. On contacting the skin, it causes chemical burns.
| Property | Value |
|---|---|
| Molar mass | 98.08 g/mol |
| Physical appearance | colourless viscous liquid |
| Melting point | 10°C |
| Boiling point | 290°C |
| Specific gravity (15°C) | 1.83 g/cm³ |
Sulphuric acid is considered a king of chemicals and its consumption is an indicator of the industrial progress of a country. Major uses include fertilizers, metal extraction, catalyst in oil refining, production of pesticides, explosives, and lead storage batteries.
Study Strategies for Chapter 12
Master Nitrogen Properties
Focus on understanding why nitrogen is relatively inert compared to other diatomic gases. Memorize the bond energy of N≡N bond (944 kJ/mol) and how this affects reactivity. Practice writing equations for nitrogen preparation from ammonium nitrite.
Ammonia Reactions
Learn the laboratory preparation of ammonia from ammonium salts. Understand the basicity of ammonia and how it forms ammonium ions. Practice drawing the structures of ammonia (pyramidal) and ammonium ion (tetrahedral).
Oxides of Nitrogen
Create a table comparing all oxides of nitrogen with their formulas, oxidation states, properties, and uses. Focus on NO and NO₂ as they are most important environmentally. Understand photochemical smog formation mechanisms.
Sulfur Chemistry
Memorize the common oxidation states of sulfur (-2, 0, +2, +4, +6). Understand the contact process for sulfuric acid production step by step. Learn the properties and uses of sulfuric acid as dehydrating and oxidizing agent.