1. I. Introduction to Bonding (p. 161 – 163)
Ch. 6 & 7 - Chemical Bonding
I. Introduction to Bonding (p. 161 – 163)

2. A. Types of Bonds IONIC COVALENT
Bond Formation
e- are transferred from metal to nonmetal
e- are shared between two nonmetals
Type of Structure
crystal lattice
true molecules
Physical
State
solid
liquid or gas
Melting
Point
high
low
Solubility in
Water
yes
usually not
Electrical Conductivity
yes (solution or liquid) no
Other
Properties

3. A. Types of Bonds METALLIC e- are delocalized among metal atoms
Bond Formation
e- are delocalized among metal atoms
Type of Structure
“electron sea”
Physical
State
solid
Melting
Point
very high
Solubility in
Water no
yes (any form)
Electrical Conductivity
Other
Properties
malleable, ductile, lustrous

4. A. Types of Bonds
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5. A. Types of Bonds
RETURN

6. Ionic Bonding - Crystal Lattice
A. Types of Bonds
Ionic Bonding - Crystal Lattice
RETURN

7. Covalent Bonding - True Molecules
A. Types of Bonds
Covalent Bonding - True Molecules
Diatomic Molecule
RETURN

8. Metallic Bonding - “Electron Sea”
A. Types of Bonds
Metallic Bonding - “Electron Sea”
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9. B. Vocabulary Chemical Bond
electrical attraction between nuclei and valence e- of neighboring atoms that binds the atoms together
bonds form in order to…
decrease PE
increase stability

10. NaCl CO2 B. Vocabulary CHEMICAL FORMULA IONIC COVALENT Formula Unit
Molecular
Formula
NaCl
CO2

11. NaCl NaNO3 B. Vocabulary COMPOUND more than 2 elements 2 elements
Binary
Compound
Ternary
Compound
NaCl
NaNO3

12. Na+ NO3- B. Vocabulary ION 1 atom 2 or more atoms Monatomic Ion
Polyatomic
Ion
Na+
NO3-

13. C. Bond Polarity
Most bonds are a blend of ionic and covalent characteristics.

14. C. Bond Polarity Nonpolar Covalent Bond e- are shared equally
symmetrical e- density
usually identical atoms

15. + - C. Bond Polarity Polar Covalent Bond e- are shared unequally
asymmetrical e- density
results in partial charges (dipole) + -

16. C. Bond Polarity
Nonpolar
Polar
Ionic
View Bonding Animations.

17. C. Bond Polarity Electronegativity
Attraction an atom has for a shared pair of electrons.
higher e-neg atom  -
lower e-neg atom +

18. C. Bond Polarity Electronegativity Trend (p. 151)
Increases up and to the right.

19. C. Bond Polarity
Difference in the elements’ e-negs determines bond type

20. II. Molecular Compounds (p. 164 – 172, 211 – 213)
Ch. 6 & 7 - Chemical Bonding
II. Molecular Compounds (p. 164 – 172, 211 – 213)

21. A. Energy of Bond Formation
Potential Energy
based on position of an object
low PE = high stability

22. A. Energy of Bond Formation
Potential Energy Diagram
attraction vs. repulsion
no interaction
increased attraction

23. A. Energy of Bond Formation
Potential Energy Diagram
attraction vs. repulsion
increased repulsion
balanced attraction & repulsion

24. A. Energy of Bond Formation
Bond Energy
Energy required to break a bond
Bond Energy
Bond Length

25. A. Energy of Bond Formation
Bond Energy
Short bond = high bond energy

26. X O B. Lewis Structures 2s 2p Electron Dot Diagrams
show valence e- as dots
distribute dots like arrows in an orbital diagram
4 sides = 1 s-orbital, 3 p-orbitals
EX: oxygen X 2s 2p O

27. Ne B. Lewis Structures Octet Rule
Most atoms form bonds in order to obtain 8 valence e-
Full energy level stability ~ Noble Gases Ne

28. B. Lewis Structures - + Nonpolar Covalent - no charges
Polar Covalent - partial charges + -

29. C. Molecular Nomenclature
Prefix System (binary compounds)
1. Less e-neg atom comes first.
2. Add prefixes to indicate # of atoms. Omit mono- prefix on first element.
3. Change the ending of the second element to -ide.

30. C. Molecular Nomenclature
PREFIX
mono-
di-
tri-
tetra-
penta-
hexa-
hepta-
octa-
nona-
deca-
NUMBER 1 2 3 4 5 6 7 8 9 10

31. C. Molecular Nomenclature
CCl4
N2O
SF6
carbon tetrachloride
dinitrogen monoxide
sulfur hexafluoride

32. C. Molecular Nomenclature
arsenic trichloride
dinitrogen pentoxide
tetraphosphorus decoxide
AsCl3
N2O5
P4O10

33. C. Molecular Nomenclature
The Seven Diatomic Elements
Br2 I2 N2 Cl2 H2 O2 F2 H N O F Cl Br I

34. III. Ionic Compounds (p. 176 – 180, 203 – 211)
Ch. 6 & 7 - Chemical Bonding
III. Ionic Compounds (p. 176 – 180, 203 – 211)

35. A. Energy of Bond Formation
Lattice Energy
Energy released when one mole of an ionic crystalline compound is formed from gaseous ions

36. B. Lewis Structures Covalent – show sharing of e-
Ionic – show transfer of e-

37. B. Lewis Structures Covalent – show sharing of e-
Ionic – show transfer of e-

38. C. Ionic Nomenclature Ionic Formulas
Write each ion, cation first. Don’t show charges in the final formula.
Overall charge must equal zero.
If charges cancel, just write symbols.
If not, use subscripts to balance charges.
Use parentheses to show more than one polyatomic ion.
Stock System - Roman numerals indicate the ion’s charge.

39. C. Ionic Nomenclature Ionic Names
Write the names of both ions, cation first.
Change ending of monatomic ions to -ide.
Polyatomic ions have special names.
Stock System - Use Roman numerals to show the ion’s charge if more than one is possible. Overall charge must equal zero.

40. C. Ionic Nomenclature Consider the following:
Does it contain a polyatomic ion?
-ide, 2 elements  no
-ate, -ite, 3+ elements  yes
Does it contain a Roman numeral?
Check the table for metals not in Groups 1 or 2.
No prefixes!

41. C. Ionic Nomenclature
Common Ion Charges 1+ 2+ 3+ NA 3- 2- 1-

42. C. Ionic Nomenclature potassium chloride magnesium nitrate
copper(II) chloride
K+ Cl-
 KCl
Mg2+ NO3-
 Mg(NO3)2
Cu2+ Cl-
 CuCl2

43. C. Ionic Nomenclature NaBr Na2CO3 sodium bromide FeCl3
sodium carbonate
iron(III) chloride