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
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Atoms in contact will interact!
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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
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What about metals with other metals?
Metallic atoms share their valence electrons freely
in a “sea of electrons” to form alloys.
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Brass
White gold
14K gold
Steel
Cast iron
Bronze
Pewter
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Cu + Zn
Au + Ni or Pd
Au + Cu or Ag
Fe + C
Fe + C + Si
Cu + Sn
Sn + Cu or Sb or Pb
Pause for penny demo!
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Properties of other bonding:
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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
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Covalent
◦ Molecular arrangement
◦ Lower melting and boiling
temperatures (may even be
gases!)
◦ Ratio of atoms involved is
determined experimentally
◦ Generally non-conductive
Ionic Bond
Electrostatic force that holds oppositely
charged particles together in an ionic
compound
 Binary ionic compounds – contain only
two different elements
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◦ A metallic cation and a nonmetallic anion
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Electrolyte – ionic compound whose
aqueous solution conducts an electric
current
Ionic Bond
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# electrons lost must = # electrons
gained
◦ Calcium: 2+ charge
◦ Fluorine: 1- charge
◦ 1 Ca to every 2 F: CaF2
Example Ionic Bond
Sodium chloride
 Na+1 , Cl-1
 Methods: (p. 216)
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Electron configuration
Orbital notation
Electron-dot structures
Atomic models
Energy and Ionic Bonds
Endothermic – energy absorbed during a
chemical reaction
 Exothermic – energy released during a
chemical reaction
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◦ Ionic compounds always exothermic reaction
Energy and Ionic Bonds
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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
Crystal strength:
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Determined by ionic radius
◦ Smaller radii = higher lattice energy
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Determined by ionic charge
◦ Higher charge = higher lattice energy
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KI < KF < LiF < MgO
Predicting ionic ratios
Based on charge ratios (“formula units” –
simplest ratio of the ions)
 Cations first, anions second
 For example
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◦ 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”
Oxidation Number
Charge of a monatomic ion (one-atom
ion)
 Also known as oxidation state
 Group 1: +1
 Group 2: +2
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D-block cations
Have varying oxidation numbers
 Charges of these elements are indicated
with Roman numerals (Stock method)
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◦ Cu (I) or Cu (II)
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OR name changes (less common)
◦ “-ic” means higher option (cupric = 2+)
◦ “-ous” means lower option (cuprous = 1+)
Naming Binary Ionic Compounds
Name the cation (including charge if a dblock metal) and the anion with “-ide”
 Sodium chloride
Gold (III) iodide
 Beryllium oxide
Zinc nitride
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Polyatomic ions
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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
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If more than one needed – parenthesis
◦ Mg(ClO3)2
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Oxyanions- negatively charged polyatomic
ion containing oxygen
Make another ‘A’
Vocabulary
 Memorize polyatomic ions
 Read about alloys
 Read about properties of ionic
compounds
 Practice writing formulas and names
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Covalent bonding
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…not ‘til next chapter! ;0)
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The end!