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:
ionic (metals with nonmetals)
polar covalent (varying degrees)
nonpolar covalent (2 nonmetals)
See page 169
What about metals with other metals?

4. Metallic atoms share their valence electrons freely in a “sea of electrons” to form alloys.
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

5. Pause for penny demo!

6. Properties of other bonding:
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

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!