Study Notes on Hydrogen

4.1 Position of Hydrogen in Periodic Table

  • Elements in first period:
    • Hydrogen and Helium.
    • Hydrogen is reactive, Helium is inert.
  • Positioning of Hydrogen:
    • Not firmly correlated to main groups, often placed with:
    • Alkali metals (Group IA)
    • Halogens (Group VIIA or 17th).
    • Exhibits properties that differ from alkali metals and halogens.

Similarities with Alkali Metals

  1. Electronic Configuration:
    • Only one electron in the outer shell, similar to alkali metals.
    • Hydrogen: Configuration (1); Alkali metals (Li: 2, 8, 1) etc.
  2. Electropositive Character:
    • Tendency to lose one electron, forming unipositive ions (H⁺).
    • Similar reaction in electrolytic dissociation:
    • HCl → H⁺ + Cl⁻, NaCl → Na⁺ + Cl⁻
  3. Valency:
    • Forms unipositive cations, indicating a valency of +1.
  4. Oxidation State:
    • Often exhibits an oxidation state of +1 when combined with electronegative elements.
  5. Affinity for Non-Metals:
    • Forms compounds with halogens (halides), oxygen (oxides), and sulfur (sulfides).
  6. Reducing Nature:
    • Acts similarly as a reducing agent.
    • E.g.: CuO + H₂ → Cu + H₂O

Differences with Alkali Metals

  1. Ionisation Energy:
    • Hydrogen has a significantly higher ionisation energy (Hydrogen: 1312 kJ/mol, Lithium: 520 kJ/mol).
  2. Non-metallic Character:
    • Hydrogen is classified as a non-metal.
  3. Atomicity:
    • Exists as diatomic (H₂); alkali metals are monoatomic (individual atoms).
  4. Nature of Oxides:
    • Hydrogen's oxides are neutral (H₂O) whereas alkali metals form basic oxides.
  5. Nature of Compounds:
    • Hydrogen forms low-boiling covalent compounds (HF, HCl, etc.) vs. alkali metals forming high-melting ionic solids (LiF, NaCl).
  6. Ionic Radius:
    • Smaller ionic radius for H⁺ compared to alkali metal ions, limiting its existence as a free ion.

Similarities with Halogens

  1. Non-metallic Character:
    • Both hydrogen and halogens exhibit non-metallic characteristics.
  2. Electronic Configuration:
    • One electron less than the nearest inert gas (H: 1, Halogens: F - 2, 7; Cl - 2, 8, 7).
  3. Diatomic Nature:
    • Hydrogen forms diatomic molecules (H₂) similar to halogens (F₂, Cl₂).
  4. Electronegativity:
    • Can gain an electron to form a negative ion (H⁻).
  5. Ionisation Potential:
    • Ionisation potential of hydrogen (13.5 eV) is comparable to halogens.
  6. Covalent Compounds:
    • Similarity in covalent compound formations (e.g., CH₄, CCl₄).
  7. Oxidation State:
    • Shows -1 in reactions with metals (e.g., NaH, NaCl).
  8. Valency:
    • Exhibits both electrovalence and covalence in compounds.

Differences with Halogens

  1. Electronegativity:
    • Hydrogen is less electronegative; halogens readily form anions (X⁻).
  2. Nature of Oxides:
    • Halogen oxides are generally acidic while H₂O remains neutral.
  3. Absence of Unshared Pairs:
    • Hydrogen has no unshared pairs in the H₂ molecule; halogens have unshared electron pairs in their molecules.

Hydrogen’s Unique Role

  • Hydrogen is often referred to as a rogue element due to its unique properties.
  • Thomson suggested a separate position in the periodic table to avoid inconsistency.
  • It may also resemble carbon, both having a half-filled electronic shell.
  • Often treated as a group on its own.

4.2 Discovery and Occurrence

  • Discovery:
    • Prepared by Henry Cavendish in 1766 by reacting acids with metals.
    • Initially named 'inflammable air' and later termed hydrogen by Lavoisier (meaning water maker).
  • Occurrence in Nature:
    • Most abundant element in the universe, found in:
    • Free state: Sun's atmosphere (90% hydrogen), air, volcanic gases, natural gas.
    • Combined state: Water, organic compounds (hydrocarbons, carbohydrates, fats, proteins).

4.3 Preparation of Hydrogen

  • Main Sources:
    • Water:
    1. Cold Water with Metals:
      • Alkali metals: React vigorously with cold water.
    2. Hot Water/Steam with Metals:
      • React with metals like Zn, Fe, to liberate hydrogen.
      • Example reactions:
      • Zn + H₂O → ZnO + H₂
      • Fe + H₂O → Fe3O4 + H₂.
    3. Electrolysis of Water:
      • Requires addition of acid/alkali to conduct electricity.
      • Product: H₂ at cathode, O₂ at anode using electrodes.
      • Example: H₂O + SO₄²⁻ → H₂ + O₂ + H₂SO₄.
    • Acids:
    1. Reactive Metals in Aqueous Solutions:
      • Active metals (Zn, Mg, etc.) react with dilute acids to yield hydrogen gas.
      • Example: Zn + HCl → ZnCl₂ + H₂.
    2. Oxidation of Hydrides:
      • Ionic hydrides are hydrolyzed in water producing hydrogen.

4.4 Hydrogen from Alkalies

  • React with water producing hydrogen in varying methods.
  • Examples:
  1. Via Sodium Hydroxide:
    • Zn + 2NaOH → Na₂ZnO₂ + H₂.
  2. Using Electrode Reaction:
    • Water in electrolysis with NaOH yields hydrogen at cathode.

4.5 Properties of Hydrogen

Physical Properties

  1. State:
    • Colorless, odorless, tasteless gas.
  2. Density:
    • 0.08987 g/L, lighter than air.
  3. Solubility:
    • Slightly soluble in water (approx. 2 volumes in 100 volumes).
  4. Melting/Boiling Points:
    • Boiling point: -252.8°C; Melting point: -259.16°C.
  5. Other Constants:
    • Ionization potential: 13.6 eV; Electron affinity: ~72 kJ/mol.
    • Critical temperature: -236.8°C.
  6. Hydrogen Storage:
    • Metals (Pd, Pt, etc.) can adsorb large amounts of hydrogen, called occluded hydrogen.