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Ionic Compounds Chapter 8.

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Presentation on theme: "Ionic Compounds Chapter 8."— Presentation transcript:

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!

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