1. Chemical Bonding
Chapter 6

2. Chemical bond
A mutual electrical attraction between the nuclei and valence electrons of different atoms that binds them together
Chemistry chapter 6

3. Types of bonds
Ionic
electrons are transferred from one atom to another, forming ions
Electrical attraction of large numbers of cations and anions holds the compound together
Covalent – electron pairs are shared between the atoms
Chemistry chapter 6

4. Ionic Bonds
Ions get very close together and are attracted to each other.
Example: NaCl
Cl-
Na+
Chemistry chapter 6

5. Covalent Bonds
The atoms share electrons, and their electron clouds overlap.
Example: carbon monoxide O C
Chemistry chapter 6

6. Electronegativity difference
Can be used to predict how two atoms will react with each other.
Has no units, because it is a comparison.
When it is high, the bond is ionic.
When it is low, the bond is covalent.
Chemistry chapter 6

7. Ionic – covalent continuum
Real bonds are usually sort of ionic and sort of covalent.
See figure 6-2 on page 162.
Difference of 1.7 or less is considered covalent
50% ionic character or less
Difference of more than 1.7 is considered ionic
More than 50% ionic character
Chemistry chapter 6

8. Nonpolar-covalent bond
Bonding electrons are shared equally
Balanced distribution of electrical charge
0 to 5% ionic character
0 to 0.3 electronegativity difference
Chemistry chapter 6

9. Polar
Having uneven charge distribution
Chemistry chapter 6

10. Polar-covalent bond
The bonded atoms have unequal attraction for the shared electrons
5 to 50% ionic character
0.3 to 1.7 electronegativity difference
Chemistry chapter 6

11. Polar and nonpolar bonds
Electron density is greater around Cl than H because it has greater electronegativity
d+ and d- indicate partial charge
Electron cloud drawings
Chemistry chapter 6

12. Comparing bond types Comparison between ionic and covalent bonding
Chemistry chapter 6

13. Discuss Page 163 Sample problem Practice problem
Section review 1 and 2
Chemistry chapter 6

14. Molecule
A neutral group of atoms that are held together by covalent bonds
A single molecule is capable of existing on its own
May consist of atoms of the same element or different elements
Chemistry chapter 6

15. Molecular compound Compound whose simplest units are molecules
Chemistry chapter 6

16. Chemical formula
Uses atomic symbols and subscripts to indicate the relative numbers of atoms of each kind in a chemical compound.
Chemistry chapter 6

17. Molecular formula
Shows the types and numbers of atoms in a single molecule of a molecular compound
H2O O2
C12H22O11
Chemistry chapter 6

18. Diatomic molecule
Contains only two atoms
Chemistry chapter 6

19. Bonding
Bonding makes atoms be at lower potential energy levels, which is favored by nature.
Attractive forces (nucleus-electron) are balanced by repulsive forces (nucleus-nucleus and electron-electron).
Chemistry chapter 6

20. Bond length
Average distance between bonded nuclei at their lowest potential energy
Chemistry chapter 6

21. Bond energy
The amount of energy that must be added to break a chemical bond and form neutral isolated atoms.
The same amount of energy that the atoms had to release when they bonded.
Chemistry chapter 6

22. Octet rule
Chemical compounds tend to form so that each atom, by gaining, losing or sharing electrons, has an octet of electrons in its highest occupied energy level
The outermost s and p sublevels want to be full
Chemistry chapter 6

23. Exceptions to octet rule
Hydrogen only needs two electrons, not eight
Boron tends to only be surrounded by six electrons.
Some atoms can have more than eight when combining with fluorine, oxygen, and chlorine. This expanded valence also includes the d orbitals.
Chemistry chapter 6

24. Electron dot diagrams
Usually only the valence electrons are involved in chemical reactions.
Electron dot diagrams allow us to draw these electrons around the symbol for an element.
Chemistry chapter 6

25. Drawing electron dot diagrams
The element’s symbol represents the nucleus and all the electrons in the inner energy levels.
Write the noble gas configuration of the element. Add the superscripts of the s and p orbitals to get the number of valence electrons.
Chemistry chapter 6

26. Draw dots on the sides to represent electrons.
Chemistry chapter 6

27. Examples
Examples
Chemistry chapter 6

28. Molecules Can be represented by dot diagrams Example: HCl
Shared pairs can be replaced by long dashes
Chemistry chapter 6

29. Lone pair Unshared pair of electrons Not involved in bonding
Chemistry chapter 6

30. Discuss
Read sample problem 6-2 on page 170
Chemistry chapter 6

31. Lewis diagram
Formulas in which atomic symbols represent the nuclei and inner-shell electrons
Dot diagrams with dots or dashes for shared electrons
Chemistry chapter 6

32. Structural diagrams Leaves out the unshared pairs
Indicates the kind, number, arrangement, and bonds of atoms
Chemistry chapter 6

33. Single bond
Covalent bond produced when one pair of electrons is shared
Chemistry chapter 6

34. Discuss
Sample problem 6-3 on pages
Chemistry chapter 6

35. Double bond
Covalent bond produced when two pairs of electrons are shared
Example: CO2
Chemistry chapter 6

36. Triple bond
Covalent bond formed when three pairs of electrons are shared
Example: N2
Chemistry chapter 6

37. Multiple bonds Double or triple bonds Have higher bond energies
Have shorter bond lengths
Chemistry chapter 6

38. Resonance structures
Bonding in molecules or ions that cannot be correctly represented by a Lewis structure
A double-headed arrow is placed between the two possibilities
Example: ozone
Chemistry chapter 6

39. Discuss Covalent Bonds More Covalent Bonds
Sample problem 6-4 on page 174
Chemistry chapter 6

40. Examples
Draw a Lewis Structure and the structural formula for each of the following:
H2O
CH4
CH4O O2
C2H2
C2H4
Chemistry chapter 6

41. Discuss
Section Review on page 175
Chemistry chapter 6

42. Ionic compound
Composed of positive and negative ions that are combined so that the numbers of positive and negative charges are equal
Usually crystalline solids
Not composed of independent units like molecules
Ionic bonding
More ionic bonding
Chemistry chapter 6

43. Chemical formula For ionic compounds
Shows the simplest ratio of the ions that gives no net charge
Chemistry chapter 6

44. Formula unit
The simplest collection of atoms from which an ionic compound’s formula can be established.
For sodium chloride, NaCl
The simplest way to show a one to one ratio
Chemistry chapter 6

45. Ratio of ions Depends on the charges of the ions combined.
Na+ and Cl- form NaCl
Mg2+ and F- form MgF2
Chemistry chapter 6

46. Forming ionic compounds
Examples:
NaCl
MgF2
Chemistry chapter 6

47. Crystal lattice
Ions minimize their potential energy by forming an orderly arrangement
Chemistry chapter 6

48. Lattice energy
The energy released when one mole of an ionic crystalline compound is formed from gaseous ions.
Used to compare bond strengths
Chemistry chapter 6

49. Molecular vs. ionic compounds
Forces between molecules are weaker than ionic bonding forces
Molecular compounds have lower melting points
Ionic compounds are hard but brittle.
Ionic compounds conduct electricity when melted or dissolved in solution
Chemistry chapter 6

50. Polyatomic ions A charged group of covalently bonded atoms.
Form ionic bonds just like other ions.
Examples
Ammonium ion NH4+
Sulfate ion SO42-
Chemistry chapter 6

51. Discuss
Use electron-dot notation to demonstrate the formation of ionic compounds involving
Li and Cl
Ca and I
What basic unit are molecular compounds composed of? Ionic compounds?
Chemistry chapter 6

52. Metals High electrical conductivity.
Metals form crystals and their orbitals overlap.
All the atoms have the same attraction for their electrons, so the electrons can move easily from one atom to another.
Chemistry chapter 6

53. Valence electrons Metals have very few Their p sublevels are empty
Transition metals also have many vacant d sublevels
Chemistry chapter 6

54. Metallic bonding
Delocalized electrons – don’t belong to one specific atom
Sea of electrons formed around the metal atoms
Metallic bonding results from the attraction between metal atoms and the surrounding sea of electrons
Chemistry chapter 6

55. Metallic properties High electrical and thermal conductivity
Ability to absorb a wide range of light frequencies
Electrons enter excited states and then return to ground states
This makes metals look shiny
Chemistry chapter 6

56. Metallic properties Metals bond the same way in all directions
Their layers can slide past each other without breaking
Malleabiltiy - the ability to be hammered or beaten into thin sheets
Ductility – the ability to be drawn, pulled or extruded into a thin wire
Chemistry chapter 6

57. Metallic bond strength
When there is a bigger nucleus and more electrons, the bond is stronger.
Stronger bonds have higher heats of vaporization
Chemistry chapter 6

58. Discuss Describe the electron-sea model of metallic bonding.
What is the relationship between metallic bond strength and heat of vaporization?
Explain why most metals are malleable and ductile but ionic crystals are not.
Chemistry chapter 6

59. Molecular polarity Uneven distribution of molecular charge
Determined by the polarity of each bond and the geometry of the molecule
Chemistry chapter 6

60. Molecular geometry There are two equally successful theories
One accounts for molecular bond angles
The other describes the orbitals that contain the valence electrons
Chemistry chapter 6

61. VSEPR theory Valence-shell, electron-pair repulsion
The repulsion between the sets of valence-level electrons surrounding an atom causes these sets to be oriented as far apart as possible
To predict geometry, one must consider the location of all electron pairs around the bonded atoms.
Chemistry chapter 6

62. Linear molecules A is the central atom B is an atom bonded to it
Formulas like AB2 form linear molecules, with the B atoms on opposite sides of the A atom
Chemistry chapter 6

63. Trigonal Planar AB3 A is in the center.
The 3 Bs make a triangle around it.
The molecule is flat (planar)
Chemistry chapter 6

64. Tetrahedral AB4 A pyramid with A at the center and a B at each corner
Chemistry chapter 6

65. When there are lone pairs
The repel the other electrons just like shared pairs do
Use an E to represent the unshared pairs
Chemistry chapter 6

66. Bent or angular
AB2E
Like trigonal planar, but without the top of the triangle
Chemistry chapter 6

67. Trigonal pyramidal AB3E
Like tetrahedral, but without the top of the pyramid
A is at the center, the Bs are the base, and E is at the top
Chemistry chapter 6

68. Bent or angular
AB2E2
Like a tetrahedral, but only the atoms (not the unshared pairs) determine the shape
Chemistry chapter 6

69. Trigonal bipyramidal AB5 Like two pyramids on top of each other
Chemistry chapter 6

70. Octahedral
AB6
8 sided shape
Chemistry chapter 6

71. VSEPR
Double and triple bonds are treated the same way as single bonds.
Use Lewis structures and table 6-5 on page 186 to predict molecule shapes
Chemistry chapter 6

72. Hybridization
Explains how orbitals get rearranged when atoms form covalent bonds
Describes how atoms are bonded
Chemistry chapter 6

73. Hybrid orbitals
Orbitals of equal energy produced by the combining of two or more orbitals on the same atom
The number of hybrid orbitals equals the number of orbitals they were made from
Chemistry chapter 6

74. Example: Methane
CH4
To get 4 equal orbitals, the s and p orbitals combine.
They form 4 sp3 orbitals
Chemistry chapter 6

75. Other hybrid orbitals
Chemistry chapter 6

76. Intermolecular forces
The forces of attraction between molecules
Usually weaker than covalent, ionic, and metallic bonds
Can be measured by boiling point
Higher boiling point means stronger intermolecular forces
Chemistry chapter 6

77. Dipole Equal but opposite charges separated by a short distance
Direction is from the positive pole to the negative pole
Chemistry chapter 6

78. Dipole-dipole forces Forces of attraction between polar molecules
Strongest intermolecular forces
Only act over short distances
Chemistry chapter 6

79. More than two atoms
Polarity of the molecule depends on the polarity and orientation of each bond
Chemistry chapter 6

80. Chemistry chapter 6

81. Induced dipoles
A polar molecule can induce a dipole in a nonpolar molecule by attracting its electrons.
It is temporary
Chemistry chapter 6

82. Hydrogen bonding
A hydrogen atom in a compound with a very electronegative atom is attracted to an unshared pair of electrons in a nearby molecule.
Examples:
Water (H2O)
Hydrogen fluoride (HF)
Ammonia (NH3)
Chemistry chapter 6

83. Hydrogen bonding
Causes higher boiling points
Chemistry chapter 6

84. Polarity and Hydrogen Bonding
Chemistry chapter 6

85. London dispersion forces
The intermolecular attractions caused by the constant motion of electrons and the creation of instantaneous dipoles
Act between all atoms and molecules
The only intermolecular forces among noble gas atoms and nonpolar molecules
Low boiling points
Chemistry chapter 6

86. discuss
What is the difference between intramolecular forces and intermolecular forces?
What is a dipole? In what kind of molecule are dipoles common?
What is hydrogen bonding?
What are London dispersion forces?
Molecular Geometry review
Chemistry chapter 6