Chapter 6Covalent Bonds

1
Chapter 6-Covalent Bonds

• Properties of covalent compounds
• A. Formed by atoms sharing electrons
• (still obey octet rule, but both atoms
• can count shared electrons as part
• of octet)
• B. Formed by nonmetal atoms reacting
• together-both atoms want to take
• electrons

2

• C. Form separate molecules
• are amorphous-no regular shape
• D. Low density, M.P. B.P.
• molecules have spaces between them
• many are flammable
• E. Can have liquids, gases or solids at
• room temperature
• (depends on molecule size)

3

• F. Two Types of Covalent Bonds
• Polar Bonds
• electronegativity difference between
• 0.4 1.7
• b. electrons unequally shared
• pulled to higher electronegativity
• have small charges on molecules
• c. are poor conductors if dissolved or melted
• d. soluble with ionic or polar
• (like dissolves like-both have charges)

4

• 2. Nonpolar Bonds
• a. electronegativity difference
• 0.4 less
• b. electrons equally shared
• no charges on molecules
• c. Are nonconductors
• d. Soluble only with other nonpolar
• Weaker bonds have less charge
• (smaller electronegativity difference)

5

• II. Covalent Diagrams
• A. Element Diagrams
• 1. Orbital Notation Diagrams
• -draw with noble gas
• Examples
• Tin
• Argon

6

• 2. Lewis Dot Diagrams
• draw dots for valence electrons
• (s p electrons)
• keep paired single like
• orbital notation diagram
• Carbon has hybridization makes all electrons
  equal, bonds with all 4

7

• Use Periodic Table to draw Lewis dot
• Group Valence Electrons
• 1 1
• 2 2
• 13 3
• 14 4
• 15 5
• 16 6
• 17 7
• 18 8

8

• B. Compound Diagrams
• 1. Orbital Notation Diagrams
• circle shared electrons
• each pair 1 bond
• SF2 N2
• Forming bonds Exothermic
• Breaking bonds Endothermic

9

• 2. Lewis dot structures
• use different marks for each element
• most single electrons central atom
• replace bond electrons with lines
• All electrons should be paired when done
• All elements (except H) should have octet when
  done

10

• III. Covalent Molecules
• A. Molecule Shapes VSEPR theory
• depends on of atoms, of bonds,
• and of electron clouds
• 1. Linear shape
• 2 atoms or Double/Triple Bonds
• 2. Bent shape
• 3 atoms, single bonds
• 3. Triangular Pyramid shape
• 4 atoms, single bonds
• 4. Tetrahedral shape
• 5 atoms, single bonds

11

• B. Molecule Type
• depends on bond type shape
• Bond Type Molecule Type
• Ionic (EN 1.7 more) Ionic crystal
• Polar (EN 0.4-1.7)
• pulls remain Polar
• pulls cancel Nonpolar
• Nonpolar (EN 0.4 less) Nonpolar

12

• Examples
• (Be sure to draw Lewis structures with correct
  shape)
• SiCl4 PCl3

13

• IV. Naming Covalent Compounds
• Use prefixes to tell how many atoms of each
  element are present
• mono 1 - not on 1st element
• di 2
• tri 3
• tetra 4 penta 5
• -ide ending on 2nd element
• Do not Reduce Formulas!
• Examples
• Dinitrogen tetroxide
• NO

14

• V. Molecular Formulas
• empirical formula simplest ratio of elements
  in a compound (ionic formulas)
• molecular formula actual of atoms in each
  molecule (covalent formula)
• 1. Change grams to moles
• ( is mass, so can label grams)
• 2. Find Mole Ratios-divide by smallest
• (if any ratio is 1.5, 2.5, 3.5, etc.
• double all mole ratios to make whole s)

15

• 3. write empirical formula/get mass
• (mole ratios are subscripts)
• 4. Divide molar mass by
• empirical formula mass
• this is how many times larger you
• need to make molecular formula

16

• Example
• A compound is 69.9 iron and
• 30.1 oxygen.
• Its molar mass is 321 g/mole.
• What is its molecular formula?