1. Academic Chemistry Mrs. Teates Newport High School
Chapter 6 – Chemical Bonding
Academic Chemistry Mrs. Teates Newport High School

2. Lesson 1 – Introduction to Chemical Bonding
Lesson Essential Questions:
Why do atoms form chemical bonds?
How is the type of chemical bond determined?
Vocabulary: chemical bond, ionic bonding, covalent bonding, nonpolar-covalent bonding, polar, polar-covalent bonding

3. Vocabulary Chemical Bond
attractive force between atoms or ions that binds them together as a unit
bonds form in order to…
decrease potential energy (PE)
increase stability

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

5. 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
Solid, liquid, or gas
Melting
Point
high
low
Solubility in
Water
yes
usually not
Electrical Conductivity
yes (solution or liquid) no
Other
Properties
odorous

6. 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

7. Ionic Bonding - Crystal Lattice
Types of Bonds
Ionic Bonding - Crystal Lattice
RETURN

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

9. Metallic Bonding - “Electron Sea”
Types of Bonds
Metallic Bonding - “Electron Sea”
RETURN

10. Bond Polarity
Most bonds are a blend of ionic and covalent characteristics.
Difference in electronegativity determines bond type.

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

12. Bond Polarity
Electronegativity Trend
Increases up and to the right.

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

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

15. Bond Polarity
Nonpolar
Polar
Ionic
View Bonding Animations.

16. Bond Polarity 3.0-3.0=0.0 Nonpolar 3.0-2.1=0.9 Polar 3.0-0.9=2.1 Ionic
Examples:
Cl2
HCl
NaCl
=0.0
Nonpolar
=0.9
Polar
=2.1
Ionic

17. More Bond Polarity Practice
What type of bonding would be expected between the following atoms?
Li and Cl
Ca and Ga
I and Cl
K and Na

18. Lesson 2 – Covalent Bonding and Molecular Compounds
Lesson Essential Questions:
How is a molecular compound formed?
What are some of the characteristics of a covalent bond?
Vocabulary: molecule, chemical formula, molecular formula, bond energy, electron-dot, Lewis structure, structural formula, single bond, multiple bonds, resonance

19. Vocabulary
Covalent bond – bond that is created by the sharing of electrons
Molecule – neutral group of atoms held together by covalent bonds
Molecular compound – chemical compound made of molecules

20. NaCl CO2 Vocabulary CHEMICAL FORMULA IONIC COVALENT Formula Unit
Molecular
Formula
NaCl
CO2

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

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

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

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

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

26. Lewis Structures Electron Dot Diagrams Pick the central atom
Count the valence electrons (they are what electron dot diagrams show)
Place electrons around the atom

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

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

29. Practice Drawing Lewis Structures
On page 186 in your text book do practice problems #1-4
Draw the Lewis structure of ammonia, NH3
Draw the Lewis structure for hydrogen sulfide, H2S
Draw the Lewis structure for silane, SiH4
Draw the Lewis structure for phosphorus trifluoride, PF3

30. Multiple Covalent Bonds
Some elements can share more than one electron pair.
Double bond (two pairs of electrons are shared)
Triple bond (three pairs of electrons are shared)

31. Practice of Lewis Structures for multiple bonds
Draw Lewis structures for each of the following molecules: O2
CO2 N3 N2

32. Resonance
Occurs when more than one valid Lewis structure can be written for a particular molecule (due to position of double bond)
These are resonance structures of benzene.
The actual structure is an average (or hybrid) of these structures.

33. Resonance in Ozone Note the different location of the double bond
Neither structure is correct, it is actually a hybrid of the two. To show it, draw all varieties possible, and join them with a double-headed arrow.

34. Polyatomic ions – note the different positions of the double bond.
Resonance in a carbonate ion (CO32-):
Resonance in an acetate ion (C2H3O21-):

35. 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.

36. Molecular Nomenclature
PREFIX mono- di- tri- tetra- penta- hexa- hepta- octa- nona- deca-
NUMBER

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

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

39. Lesson 3 – Ionic Bonding and Ionic Compounds
Lesson Essential Questions:
How is an ionic bond formed?
What are some of the characteristics of an ionic bond?
Vocabulary: ionic compound, formula unit, lattice energy, polyatomic ion

40. Vocabulary:
Ionic compound – composed of positive and negative ions that are combined so that the charges are equal.

41. NaCl CO2 Vocabulary CHEMICAL FORMULA IONIC COVALENT Formula Unit
Molecular
Formula
NaCl
CO2

42. Forming Ionic Compounds
Electron dot notation is used to note changes.
Form to create an atmosphere of stability

43. Lewis Structures and Ionic Compounds
Covalent – show sharing of e-
Ionic – show transfer of e-

44. Characteristics of Ionic Bonding
Ions minimize potential energy in crystals by forming a crystal lattice.
Distance between all ions represent a balance of attraction between oppositely charged particles and repulsion between like charged particles

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

46. Ionic vs. Covalent Ionic High melting temperature High boiling point
Hard
Brittle, because slight shift of crystal can cause it to break
Conduct electricity when dissolved in water
Covalent
Low melting temperature
Low boiling point
Do not conduct electricity
Not as brittle

47. 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.

48. 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.

49. 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!

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

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

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

53. Lesson 4 – Metallic Bonding
Lesson Essential Questions:
How is a metallic bond formed?
What are some of the characteristics of a metallic bond?
Vocabulary: metallic bond, alloy

54. Characteristics of Metallic Bonds
Metal ions held together by attraction to free floating electrons. (Sea of electrons)
Good conductors of electricity – Why?

55. Characteristics of Metallic Bonds Cont.
Malleable
Ductile
Bond strength – related to enthalpy of vaporization
The more energy required to vaporize, the stronger the bond.
See table on page 196.

56. Alloys
A mixture of two or more substances, one of which must be a metal.
Common alloys include steel, 14K gold, 18K gold, cast iron, sterling silver, and bronze.
Within different alloys, there can be different types of mixtures – ex. Steel
Where do we find alloys?

57. Compare/Constrast
Use the 3 circle Venn diagram to compare and contrast ionic, metallic, and covalent bonding.

58. Lesson 5 – Molecular Geometry
Lesson Essential Questions:
How is the VSEPR Theory useful?
What are the different forces present in bonding?
Vocabulary:
VSEPR theory, hybridization, dipole, hydrogen bonding, London dispersion forces

59. VSEPR Theory Valence Shell Electron Pair Repulsion Theory
Electron pairs orient themselves in order to minimize repulsive forces.

60. Lone pairs repel more strongly than bonding pairs!!!
VSEPR Theory
Types of e- Pairs
Bonding pairs - form bonds
Lone pairs - nonbonding e-
Lone pairs repel more strongly than bonding pairs!!!

61. VSEPR Theory Lone pairs reduce the bond angle between atoms.

62. Determining Molecular Shape
Draw the Lewis Diagram.
Tally up e- pairs on central atom.
double/triple bonds = ONE pair
Shape is determined by the # of bonding pairs and lone pairs.
Know the 8 common shapes & their bond angles!

63. VSEPR Table

64. Common Molecular Shapes
2 total
2 bond
0 lone
BeH2
LINEAR
180°

65. Common Molecular Shapes
3 total
3 bond
0 lone
BF3
TRIGONAL PLANAR
120°

66. Common Molecular Shapes
3 total
2 bond
1 lone
SO2
BENT
<120°

67. Common Molecular Shapes
4 total
4 bond
0 lone
CH4
TETRAHEDRAL
109.5°

68. Common Molecular Shapes
4 total
3 bond
1 lone
NH3
TRIGONAL PYRAMIDAL
107°

69. Common Molecular Shapes
4 total
2 bond
2 lone
H2O
BENT
104.5°

70. Common Molecular Shapes
5 total
5 bond
0 lone
PCl5
TRIGONAL BIPYRAMIDAL120°/90°

71. Common Molecular Shapes
6 total
6 bond
0 lone
SF6
OCTAHEDRAL
90°

72. Examples
PF3
4 total
3 bond
1 lone
TRIGONAL PYRAMIDAL
107°

73. Examples
CO2
2 total
2 bond
0 lone
LINEAR
180°

74. Practice Problems
Identify the molecular geometry for the following molecules: HI
CBr4
CH2Cl2

75. Intermolecular Forces
Intermolecular forces = forces between molecules.
The boiling point of a liquid is a good measure of the intermolecular forces between its molecules: the higher the boiling point, the stronger the forces between the molecules.
Types of intermolecular forces
Dipole-dipole forces
Hydrogen bonding
London dispersion forces

76. Dipole – Dipole Forces
Dipole – created by equal but opposite charges that are separated by a short distance.
A dipole is represented by an arrow with its head pointing toward the negative pole and a crossed tail at the positive pole. The dipole created by a hydrogen chloride molecule is indicated as follows:
Dipole-dipole forces are the forces of attraction between polar molecules.

77. Dipole-dipole forces cont.
The negative region in one polar molecule attracts the positive region in adjacent molecules. So the molecules all attract each other from opposite sides.
Dipole-dipole forces act at short range, only between nearby molecules.

78. Hydrogen Bonding
Hydrogen bonding = intermolecular force in which a hydrogen atom that is bonded to a highly electronegative atom is attracted to an unshared pair of electrons in a nearby molecule.

79. London Dispersion Forces
London Dispersion Forces = intermolecular attractions resulting from the constant motion of electrons and the creation of instantaneous dipoles.

80. Works Cited Modern Chemistry Textbook www.nclark.net