Chapter 5 Ionic Bonds

1
Chapter 5 Ionic Bonds

• Bonds attraction of two atoms for the same
• valence electrons
• Properties of Ionic Compounds (salts)
• A. Electronegativity difference
• 1.7 higher
• Calcium Chlorine

2

• B. formed by transfer of electrons
• -metals lose electrons
• -nonmetals gain electrons
• Oxidation lose electrons, get larger charge
• Ca ? Ca 2 2e-
• Reduction gain electrons, get smaller charge
• Cl 1e-? Cl1-

3

• C. Crystal lattice structure
• not separate molecules
• shape of crystals depends on
• size charges of ions
• D. Hard due to strong bonds
• E. Brittle ionic not lined up
• due to different sizes

4

• F. Very High Density, M.P. B.P.
• (more than metals)
• ions are packed close together
• have charges holding them together
• G. Good conductors if dissolved or melted
• -charges are on the atoms,
• so atoms must move to get charges to
  move
• H. Soluble in water
• water has and on its molecules
• that attracts and pulls on opposite
  charges
• in crystal lattice

5

• II. Writing Formulas Naming Ionic Cmpds
• A. Oxidation Numbers
• Group 1 1 Hydrogen 1
• Group 2 2
• Group 13 3
• Group 14 4 or 4- Negative ions
• Group 15 3- -ide ending
• Group 16 2-
• Group 17 1- Hydride 1-
• Group 18 0
• Silver 1 (Group 11)
• Zinc 2 (Group 12)

6

• B. Writing Formulas
• oxidation numbers must add to zero
• so electrons lost electrons gained
• Shortcut Criss-cross Reduce
• empirical formula simplest ratio
• subscript lowered , gives ratio
• Examples
• Silver sulfide
• Magnesium oxide
• Calcium phosphide

7

• C. Transition Metals ( others)
• d-electrons can give more than one
• oxidation number
• use Roman numeral to show charge
• Copper (II) Cu2
• Copper (I) Cu1
• Iron (II) Fe2
• Iron (III) Fe3
• Lead (II) Pb2
• Lead (IV) Pb4
• Example
• Copper (II) oxide

8

• D. Polyatomic Ions
• groups of atoms that act like one ion
• Ammonium NH41
• Hydroxide OH1-
• Acetate C2H3O21-
• Nitrate NO31-
• Chlorate ClO31-
• Carbonate CO32-
• Sulfate SO42-
• Phosphate PO43-
• Use parentheses if extra subscript is needed for
  polyatomic ions
• Example
• Ammonium sulfate

9

• Naming Ionic Compouds
• 1. ion first (metal or ammonium)
• use Roman numeral for Cu, Fe, Pb
• 2. ion last
• nonmetal elements -ide ending
• polyatomic ions no name change
• Examples
• CuSO4
• Fe2O3
• PbO2

10

• III. Measuring Ionic Compounds
• A. Molar Mass
• total of all atomic masses in a
• compound from periodic table
• grams in 1 mole of compound
• 6.02 x 1023 molecules

11

• Examples
• 1. Find the molar mass of calcium nitrate
• 2. How many moles are in 3.58 grams of
• sodium sulfide?

12

• 3. How many molecules are in 0.24 moles of
• potassium phosphate?
• 4. How many grams are in 1.91 x 1022
• molecules of aluminum sulfate?

13

• B. Percent Composition (by mass)
• part (element) x 100
• whole (compound)
• Remember s should add up to 100

14

• Examples
• 1. Find the composition of
• copper (I) sulfate
• 2. Find the percent of sulfur in
• ammonium sulfide.

15

• IV. Hydrated Salts
• ionic salts with water molecules trapped
• within their crystals
• Anhydrous salts have water removed from them
  (usually by heating)

16

• Prefixes in name tell of water molecules
• 1 mono 6 hexa
• 2 di 7 hepta
• 3 tri 8 octa
• 4 tetra 9 nona
• 5 penta 10 deca
• Example
• CaCl2 2H2O

17

• To Calculate Formulas
• 1. Find g water g anhydrous salt
• 2. Convert to moles water
• moles anhydrous salt
• 3. Get whole Mole Ratios
• -divide both by smallest moles

18

• Example
• A hydrated barium hydroxide salt has a mass of
  3.15 grams. The anhydrous salt has a mass of 1.71
  grams.
• 1. Find the formula of the hydrated salt give
  its name.

19

• 2. Find the percent of water in the hydrated
  salt.
• 3. Find the molar mass of the hydrated salt.