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Ionic Compounds

Ionic Compounds. Chapter 8. Remember…. Chemical bond Electron-dot structure Ionization energy Electron affinity – how much attraction an atom has for electrons Electronegativity Octet rule Cation Anion. Atoms in contact will interact!. Based on electronegativity difference:

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Ionic Compounds

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  1. Ionic Compounds Chapter 8

  2. Remember…. • Chemical bond • Electron-dot structure • Ionization energy • Electron affinity – how much attraction an atom has for electrons • Electronegativity • Octet rule • Cation • Anion

  3. Atoms in contact will interact! • Based on electronegativity difference: • 1.8-3.3 ionic (metals with nonmetals) • 0.4-1.7 polar covalent (varying degrees) • 0.0-0.3 nonpolar covalent (2 nonmetals) • See page 169 • What about metals with other metals?

  4. Brass White gold 14K gold Steel Cast iron Bronze Pewter Cu + Zn Au + Ni or Pd Au + Cu or Ag Fe + C Fe + C + Si Cu + Sn Sn + Cu or Sb or Pb Metallic atoms share their valence electrons freely in a “sea of electrons” to form alloys.

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  6. Ionic Crystalline arrangement (brittle/will shatter) High melting and boiling temperatures Ratio of atoms involved is determined by charges Non-conductive unless molten, dissolved in water Covalent Molecular arrangement Lower melting and boiling temperatures (may even be gases!) Ratio of atoms involved is determined experimentally Generally non-conductive Properties of other bonding:

  7. Ionic Bond • Electrostatic force that holds oppositely charged particles together in an ionic compound • Binary ionic compounds – contain only two different elements • A metallic cation and a nonmetallic anion • Electrolyte – ionic compound whose aqueous solution conducts an electric current

  8. Ionic Bond • # electrons lost must = # electrons gained • Calcium: 2+ charge • Fluorine: 1- charge • 1 Ca to every 2 F: CaF2

  9. Example Ionic Bond • Sodium chloride • Na+1 , Cl-1 • Methods: (p. 216) • Electron configuration • Orbital notation • Electron-dot structures • Atomic models

  10. Energy and Ionic Bonds • Endothermic – energy absorbed during a chemical reaction • Exothermic – energy released during a chemical reaction • Ionic compounds always exothermic reaction

  11. Energy and Ionic Bonds • Lattice energy – energy required to separate one mole of ions of an ionic compound • Reflects strength of forces holding ions together • More negative lattice energy, stronger force of attraction

  12. Crystal strength: • Determined by ionic radius • Smaller radii = higher lattice energy • Determined by ionic charge • Higher charge = higher lattice energy • KI < KF < LiF < MgO

  13. Predicting ionic ratios • Based on charge ratios (“formula units” – simplest ratio of the ions) • Cations first, anions second • For example • Na 1+ and Cl 1- ; therefore, will combine 1:1 • NaCl “sodium chloride” • Na 1+ and S 2-; therefore, will combine 2:1 • Na2S “sodium sulfide” • Be 2+ and N 3-; therefore, will combine 3:2 • Be3N2 “beryllium nitride”

  14. Oxidation Number • Charge of a monatomic ion (one-atom ion) • Also known as oxidation state • Group 1: +1 • Group 2: +2

  15. D-block cations • Have varying oxidation numbers • Charges of these elements are indicated with Roman numerals (Stock method) • Cu (I) or Cu (II) • OR name changes (less common) • “-ic” means higher option (cupric = 2+) • “-ous” means lower option (cuprous = 1+)

  16. Naming Binary Ionic Compounds • Name the cation (including charge if a d-block metal) and the anion with “-ide” • Sodium chloride Gold (III) iodide • Beryllium oxide Zinc nitride

  17. Polyatomic ions • A group of atoms acting as one cation or anion • Memorize the chart on page 224 (Table 8.6) • Yes, all of it—test next Thursday • If more than one needed – parenthesis • Mg(ClO3)2 • Oxyanions- negatively charged polyatomic ion containing oxygen

  18. Make another ‘A’ • Vocabulary • Memorize polyatomic ions • Read about alloys • Read about properties of ionic compounds • Practice writing formulas and names

  19. Covalent bonding • …not ‘til next chapter! ;0) • The end!

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