Ions and Ionic Compounds

Ions and Ionic Compounds

Ions and ionic compounds are important in chemistry, especially in acid-base and oxidation-reduction reactions in ionic solutions.

Stoichiometry and Ions

  • Goal: Identify oxidation states.
  • Importance: Determine electron equivalents, balance equations, and deduce chemical formulas from nomenclature.

Cations and Anions

  • Ionic compounds consist of:
    • Positively charged cations (usually metals).
    • Negatively charged anions (usually nonmetals).
  • Exception: Hydrogen can act as either a cation or an anion but is classified as a nonmetal.
  • Ionic bonds: Electrostatic attraction between oppositely charged particles.

Nomenclature of Ionic Compounds

  1. Elements with Multiple Positive Ions:

    • Charge indicated by Roman numeral in parenthesis.
    • Examples:
      • Fe2+Fe^{2+} Iron (II)
      • Cu+Cu^{+} Copper (I)
      • Fe3+Fe^{3+} Iron (III)
      • Cu2+Cu^{2+} Copper (II)

  2. Older Method (less common):

    • Use endings "-ous" or "-ic" to the Latin name root.
    • "-ous": lesser charge
    • "-ic": greater charge
    • Examples:
      • Fe2+Fe^{2+} Ferrous
      • Cu+Cu^{+} Cuprous
      • Fe3+Fe^{3+} Ferric
      • Cu2+Cu^{2+} Cupric

  3. Monoatomic Anions:

    • Drop ending of element name, add "-ide".
    • Examples:
      • H−H^{-} Hydride
      • F−F^{-} Fluoride
      • O2−O^{2-} Oxide
      • S2−S^{2-} Sulfide
      • N3−N^{3-} Nitride
      • P3−P^{3-} Phosphide

  4. Oxyanions (Polyatomic anions containing oxygen):

    • Two oxyanions: one with less oxygen ends in "-ite", one with more oxygen ends in "-ate".
    • Examples:
      • NO2−NO_2^{-} Nitrite
      • NO3−NO_3^{-} Nitrate
      • SO32−SO_3^{2-} Sulfite
      • SO42−SO_4^{2-} Sulfate

  5. Extended Series of Oxyanions:

    • Prefixes "hypo-" (less oxygen) and "per-" (more oxygen).
    • Examples:
      • ClO−ClO^{-} Hypochlorite
      • ClO2−ClO_2^{-} Chlorite
      • ClO3−ClO_3^{-} Chlorate
      • ClO4−ClO_4^{-} Perchlorate

  6. Polyatomic Anions with Hydrogen Ions:

    • Gain one or more H+H^{+} ions, reducing charge.
    • Named by adding "hydrogen" or "dihydrogen" prefix.
    • Older method: prefix "bi-" for a single hydrogen ion.
    • Examples:
      • HCO3−HCO_3^{-} Hydrogen carbonate or bicarbonate
      • HSO4−HSO_4^{-} Hydrogen sulfate or bisulfate
      • H<em>2PO</em>4−H<em>2PO</em>4^{-} Dihydrogen phosphate

Ion Charges

  • Cations: positive charge
  • Anions: negative charge
  • Some elements exist only in charged forms; others in charged or uncharged states.
  • Elements can have multiple oxidation states.
  • Active Metals:
    • Alkali metals (Group 1A or Group 1): +1 charge
    • Alkaline earth metals (Group 2A or Group 2): +2 charge
  • Nonmetals (right side of periodic table):
    • Generally form anions.
    • Halogens (Group 7 or Group 17): -1 charge (aim to fill octet).
  • Elements in the same group tend to form monatomic ions with the same charge.
  • Anionic species can contain metallic elements (e.g., MnO<em>4−MnO<em>4^{-}, CrO</em>42−CrO</em>4^{2-}, metals have positive oxidation states).
  • Oxyanions of halogens can have positive oxidation states (e.g., ClO−ClO^{-} , ClO2−ClO_2^{-} ).
  • Transition metals have numerous positively charged states.
  • Solution color can indicate oxidation state due to different electron transitions absorbing different light frequencies.

Electrolytes

  • Solid ionic compounds are poor conductors (ions fixed in crystal lattice).
  • Aqueous solutions: ions are free to move due to ion-dipole interactions with water, enabling electrical conductivity.
  • Electrolytes: enable solutions to carry current.
  • Electrical conductivity depends on ion presence and concentration.
  • Pure water is a poor conductor (few H+H^{+} and OH−OH^{-} ions from auto-dissociation).
  • Strong Electrolytes: Dissociate completely into ions.
    • Examples: NaCl, KI, HCl (in water).
  • Weak Electrolytes: Ionize/hydrolyze incompletely.
    • Examples: Hg<em>2I</em>2Hg<em>2I</em>2 (Ksp=4.5×10−29K_{sp} = 4.5 \times 10^{-29}), acetic acid, ammonia.
  • Non-electrolytes: Do not ionize, retain molecular structure.
    • Examples: O<em>2O<em>2(gas), CO</em>2CO</em>2(gas), glucose

Conclusion

  • Compounds consist of atoms of different elements in fixed ratios (empirical or molecular formula).
  • Each molecule has a defined mass (molecular weight).
  • Mass of one mole is determined by molar mass (grams per mole).
  • Basic Reaction Types: Combination, Decomposition, Combustion, Single Displacement, Double Displacement, Neutralization.
  • Balancing chemical reactions is essential.