1. Basic Concepts of Chemical Bonding
Chapter 8
Basic Concepts of Chemical Bonding

2. 6.3 Describing Chemical Bonding
Page:

3. CHEMICAL BONDS IONIC COVALENT METALLIC
Electrostatic attraction between ions.
COVALENT
Sharing of electrons.
METALLIC
Metal atoms bonded to several other atoms.

4. The Ionic Bonding

5. Na(s) + ½Cl2(g)  NaCl(s)
8.2 The Ionic Bonding
Na(s) + ½Cl2(g)  NaCl(s)

6. transfer occurs readily?
The Ionic Bonding
Na: loss of an electron
Cl: gain of an electron
How Can you Know
Whether this electron
transfer occurs readily?

7. 4. Electronegativity Difference transfer occurs readily?
The Ionic Bonding
1. Ionization energy
2. Electron Affinity
3. Lattice Energy
4. Electronegativity Difference
How Can you Know
Whether this electron
transfer occurs readily?

8. One species must have very low ionization energy (Na)
The Ionic Bonding
Na: loss of an electron
Cl: gain of an electron
Ionization Energy
One species must have very low ionization energy (Na)

9. One species must have very high electron affinity (Cl)
The Ionic Bonding
Na: loss of an electron
Cl: gain of an electron
Ionization Energy
Electron Affinity
One species must have very high electron affinity (Cl)

10. ELECTROSTATIC ATTRACTIONS
The Ionic Bonding
Na(s) + ½Cl2(g)  NaCl(s)
Δ Hf˚= kJ
ELECTROSTATIC ATTRACTIONS

11. LE is a measure of a stability of ions arranged within an ionic solid
Lattice Energy
Attraction between oppositely
charged ions  release of energy
 formation of lattice
LE is a measure of a stability of ions arranged within an ionic solid

12. ΔEN > 1.7 An Ionic Bond would form between these species
Electronegativity difference
If the difference in electronegativity (ΔEN) between the species is:
ΔEN > 1.7
An Ionic Bond would form between these species

13. The Ionic Bonding SUMMARY

17. 6.3 THE COVALENT BONDING

18. Covalent Bonding

19. Covalent Bonding

20. ELECTRONEGATIVITY
is the ability of atoms in a molecule to attract electrons to themselves.

21. DIPOLE and ELECTRONEGATIVITY
DIPOLE: a partial separation of charge
One end of the molecule has slightly positive charge
The other end of the molecule has slightly negative charge

22. DIPOLE and ELECTRONEGATIVITY
Slight excess of
a negative charge
Slight excess of
a positive charge
Slight excess of
a positive charge

23. DIPOLES are caused by ELECTRONEGATIVITY
Hydrogen has lower electronegativity
The electrons spend LESS time around Hydrogen
Chlorine has higher electronegativity
The electrons spend MORE time around Chlorine

24. BOND POLARITY Homo nuclear diatomic compounds:
In different compounds, electrons are not shared equally
Homo nuclear diatomic compounds:
H2, Cl2, O2,N2 … share electrons equally =
NON POLAR COVALENT BOND

25. BOND POLARITY δ+ = slightly positive δ- = slightly negative
Hetero nuclear diatomic compounds: HF, HCl, NO, NaCl, LiO… DO NOT share electrons equally
POLAR COVALENT BOND(the attraction of one of the atoms for the bonding electrons is LARGE)
δ+ = slightly positive
δ- = slightly negative

26. BOND POLARITY
H – F δ+ δ-
Hetero nuclear diatomic compounds: HF, HCl, NO, NaCl, LiO… DO NOT share electrons equally
POLAR COVALENT BOND(the attraction of one of the atoms for the bonding electrons is LARGE)

27. BOND POLARITY
H – F δ+ δ-
Hetero nuclear diatomic compounds: HF, HCl, NO, NaCl, LiO… DO NOT share electrons equally
MOSTLY COVALENT BOND (the attraction of one of the atoms for the bonding electrons is SLIGHTLY GREATER)

28. POLAR, NONPOLAR, or ionic?
BOND POLARITY
Hetero nuclear diatomic compounds: HF, HCl, NO, NaCl, LiO… DO NOT share electrons equally
POLAR COVALENT BOND(the attraction of one of the atoms for the bonding electrons is LARGE)
IONIC BONDS (the attraction of one of the atoms for the bonding electrons is VERY LARGE)
HOW DO YOU KNOW
IF THE GIVEN BOND IS
POLAR, NONPOLAR, or ionic?

29. ELECTRONEGATIVITY
The greater the difference in electronegativity, the more polar the bond is

30. ΔE = 0 ΔE > 1.7 ELECTRONEGATIVITY 0 < ΔE < 0.4
NON POLAR COVALENT BOND
ΔE = 0
H2, Cl2, O2,N2
MOSTLY COVALENT BOND
0 < ΔE < 0.4
H2O, CO2, HF
POLAR COVALENT BOND
0.4 < ΔE < 1.7
H2O, CO2, HF
IONIC BOND
ΔE > 1.7
MgO, NaCl, LiF

31. NON POLAR COVALENT BOND
ELECTRONEGATIVITY
NON POLAR COVALENT BOND
H2, Cl2, O2,N2
POLAR COVALENT BOND
H2O, CO2, HF
IONIC BOND
MgO, NaCl, LiF
electronegativity
difference
> 2.1

32. In each case, which bond is more polar:
EXAMPLE
In each case, which bond is more polar:
a) B – Cl or C – Cl,
b) P – F or P – Cl? Indicate in each case which atom has the partial negative charge.
B – Cl
1.0
F – P
1.9 δ+ δ- δ- δ+
C – Cl
0.5
P – Cl
0.9 δ+ δ- δ- δ+

35. IONIC vs COVALENT Compounds
Bonding
COVALENT Bonding
Metal + nonmetal
High melting point
Lattice(crystal) structures
Strong electrolytes
Nonmetal + nonmetal
Low melting point
Low boiling point
Non - electrolytes

37. HOMEWORK
PAGE: PROBLEMS: all even

38. 6.4 LEWIS STRUCTURE DIAGRAMS

39. LEWIS DIAGRAMS - REVIEW
Show only an atom’s valence electrons and the chemical symbol.

40. LEWIS DIAGRAMS

41. Rule # 1
Dots representing valence electrons are placed around the element symbols

42. Rule # 2
Electron dots are placed singly until the fifth electron is reached then they are paired

43. Lewis Diagrams of IONS and IONIC BONDS
For positive ions, one electron dot is removed from the valence shell for each positive charge.
For negative ions, one electron dot is added to each valence shell for each negative charge.
Square brackets are placed around each ion to indicate transfer of electrons.

45. LEWIS STRUCTURE DIAGRAMS for MOLECULES
Lewis structures are representations of molecules showing all valence electrons:
bonding and
nonbonding

46. LEWIS STRUCTURE DIAGRAMS for MOLECULES
Lewis structures are representations of molecules showing all valence electrons:
bonding and
nonbonding
a lone pair

47. Lewis Structures and Multiple Bonds
When two electron pairs are shared, two lines are drawn, i.e. double bond
::O :: C :: O:: or
::O = C = O::

48. Lewis Structures and Multiple Bonds
When three electron pairs are shared, three lines are drawn, i.e. triple bond
:N ::: N: or
:N ≡ N:

49. N – N N = N N ≡ N BOND LENGTHS > > BOND STRENGTH?
The length of the bond between two atoms decreases as the number of shared electrons increases
N – N
>
N = N
>
N ≡ N
BOND STRENGTH?

50. If it is cation  subtract e-
STEPS TO FOLLOW…
Find the sum of valence electrons of all atoms in the polyatomic ion or molecule
If it is anion  add e-
If it is cation  subtract e-
PCl3
valence electrons
= 26 v.e.

51. PCl3 5 + 21 valence electrons = 26 v.e. STEPS TO FOLLOW…
2. Write the symbols for the atoms
Make one of the atoms a central atom (usually the least electronegative atom)
connect it with the other atoms by single bonds
PCl3
valence electrons
= 26 v.e.

52. 26 – 6 valence electrons = 20 v.e. left
STEPS TO FOLLOW…
2. Write the symbols for the atoms
Subtract those electrons from your total number of valence electrons
26 – 6 valence electrons
= 20 v.e. left

53. 3. Fill the octets of the outer atoms
STEPS TO FOLLOW…
3. Fill the octets of the outer atoms
20 valence electrons

54. 3. Fill the octets of the outer atoms
STEPS TO FOLLOW…
3. Fill the octets of the outer atoms
20 valence electrons

55. 20 valence electrons STEPS TO FOLLOW…
Subtract the added electrons from your total number of valence electrons
20 valence electrons

56. 20 – 18 valence electrons = 2 v.e. left
STEPS TO FOLLOW…
Subtract the added electrons from your total number of valence electrons
20 – 18 valence electrons
= 2 v.e. left

57. 4. Fill the octet of the central atom
STEPS TO FOLLOW…
4. Fill the octet of the central atom
2 valence electrons

58. 4. Fill the octet of the central atom
STEPS TO FOLLOW…
4. Fill the octet of the central atom
0 valence electrons

59. 5. If you run out of electrons before the central atom has an octet…
STEPS TO FOLLOW…
5. If you run out of electrons before the central atom has an octet…
…form multiple bonds until it does

60. Draw Lewis Structures for
WORKSHEET EXAMPLE
Draw Lewis Structures for
CH2Cl2
C2H4
BrO3-

61. Draw Lewis Structures for
WORKSHEET EXAMPLE
Draw Lewis Structures for NO
BF3
PF5

62. EXCEPTIONTS TO THE OCTET RULE
1. For molecules and polyatomic ions containing an odd number of valence electrons
2. For molecules and polyatomic ions in which an atom has fewer than an octet of valence electrons

63. EXCEPTIONTS TO THE OCTET RULE
3. For molecules and polyatomic ions in which an atom has more than an octet of valence electrons

64. 1. For molecules and polyatomic ions containing an odd number of valence electrons
Complete pairing of valence electrons is impossible due to the odd number of valence electrons
E.g.: ClO2, NO, NO2, O2-

65. Mostly Boron or Beryllium compounds
2. For molecules and polyatomic ions in which an atom has fewer than an octet of valence electrons
Not very common
Mostly Boron or Beryllium compounds

66. 3. For molecules and polyatomic ions in which an atom has more than an octet of valence electrons
Very common
Such molecules/ions are called HYPERVALENT
Only for atoms of 3rd period or higher
WHY?
1. They have available and unfilled d orbitals for bonding
2. Their central atom (P, S, I, Xe…) is large enough to be bonded to even five different atoms (Cl, F or O)

67. Draw a Lewis Structure for ion: ICl4-
EXAMPLE
Draw a Lewis Structure for ion: ICl4-

68. Draw a Lewis Structure for the thiocyanate ion: NCS-
EXAMPLE
Draw a Lewis Structure for the thiocyanate ion: NCS-

69. WHICH ONE IS THE MOST IMPORTANT?
WHICH ONE IS CORRECT ?
WHICH ONE IS THE MOST IMPORTANT?

70. Calculate THE FORMAL CHARGE of each ion to find out…

71. FORMAL CHARGE
The charge the atom would have if all the atoms in the molecule had the same electronegativity
ALL unshared electrons are assigned to the atom on which they are found
For a single/double/triple bond, half of the bonding electrons is assigned to to each atom in the bond

72. FORMAL CHARGE OF AN ATOM =
# of VALENCE ELECTRON – # OF ELECTRONS ASSIGNED
ALL unshared electrons are assigned to the atom on which they are found
For a single/double/triple bond, half of the bonding electrons is assigned to to each atom in the bond

73. EXAMPLE
What are the formal charges of C and N in the cyanide ion: CN-?
ALL unshared electrons are assigned to the atom on which they are found
2. For a single/double/triple bond, half of the bonding electrons is assigned to to each atom in the bond
[:C ≡ N:]-
# of VALENCE ELECTRON – # OF ELECTRONS ASSIGNED

74. EXAMPLE
What are the formal charges of C and N in the cyanide ion: CN-?
ALL unshared electrons are assigned to the atom on which they are found
2. For a single/double/triple bond, half of the bonding electrons is assigned to to each atom in the bond
[:C ≡ N:]- C N
Valence e- 4 5
Assigned e-
Formal Charge -1 -1
# of VALENCE ELECTRON – # OF ELECTRONS ASSIGNED

75. What are the formal charges on the thiocyanate ions?
WORKSHEET EXAMPLE
What are the formal charges on the thiocyanate ions?
1. The most important (most dominant) Lewis structure is the one which has its value closest to 0 (the one with the fewest charges)
2. The most important (most dominant) Lewis structure is the one which has any negative charges reside on the more electronegative atoms
WHICH STRUCTURE IS THE MOST IMPORTANT (DOMINANT) ONE?

76. What are the formal charges the thiocyanate ions?
WORKSHEET EXAMPLE
What are the formal charges the thiocyanate ions?
1. The most important (most dominant) Lewis structure is the one which has its value closest to 0 (the one with the fewest charges)
2. The most important (most dominant) Lewis structure is the one which has any negative charges reside on the more electronegative atoms
WHICH STRUCTURE IS THE MOST IMPORTANT (DOMINANT) ONE?

77. The cyanate ion, NCO-, has three possible Lewis structures.
WORKSHEET EXAMPLE
The cyanate ion, NCO-, has three possible Lewis structures.
Draw these three structures assign formal charges in each.
Which Lewis structure is dominant?

78. Draw a Lewis Structure of OZONE, O3
8.6 RESONANCE STRUCTURES
Draw a Lewis Structure of OZONE, O3

79. Not a single resonance structure
RESONANCE STRUCTURES
Not a single color
OZONE:
A mix of different resonance structures
GREEN PAINT: A mix of different colors
Not a single resonance structure

80. Somewhere between single bond and double bond
RESONANCE STRUCTURES
Same bond lengths
Somewhere between single bond and double bond

81. Is the use of two or more Lewis structures to represent a molecule
RESONANCE STRUCTURES
Is the use of two or more Lewis structures to represent a molecule
this is because that molecule can not be represented by only a single Lewis structure

82. RESONANCE STRUCTURES
In truth, the electrons that form the second C—O bond in the double bonds below do not always sit between that C and that O, but rather can move among the two oxygen atoms and the carbon.
They are not localized; they are delocalized.

83. RESONANCE STRUCTURES
The organic compound benzene, C6H6, has two resonance structures.
It is a hexagon with a circle inside to signify the delocalized electrons in the ring.

84. What are the resonance structures of nitrate ion: NO3-
EXAMPLE
What are the resonance structures of nitrate ion: NO3-

85. WORKSHEET EXAMPLE
Using Formal Charges and the concept of Resonance, show why the correct Lewis structure of BF3 is the one in which B has an incomplete octet and only single bonds are present.

86. 8.8 STRENGTHS OF COVALENT BONDS
Learn on your own

87. HOMEWORK
PAGE: PROBLEMS: 49, 51, 53, 55, 57, 61, 67, 69, 71, 73,

88. ANSWERS TO WORKSHEET EXAMPLES

89. Copy this table and fill out the missing information.
WORKSHEET EXAMPLE
Copy this table and fill out the missing information.

90. WORKSHEET EXAMPLE
Which substance would you expect to have the greatest lattice energy, MgF2, CaF2, and ZrO2 ?
ZrO2

91. WORKSHEET EXAMPLE
Ionizing an H2 molecule to H2+ changes the strength of the bond. Based on the description of covalent bonding given previously, do you expect the bond H—H in H2+ to be weaker or stronger than the bond in H2?
Weaker. In both H2 and H2+ the two H atoms are principally held
together by the electrostatic attractions between the nuclei and the electron(s) concentrated between them. H2+ has only one
electron between the nuclei whereas H2 has two and this results in the H—H bond in being stronger.

92. WORKSHEET EXAMPLE
How does the ELECTRONEGATIVITY of an element differ from its ELECTRON AFFINITY?
Electron affinity measures:
the energy released when an isolated atom gains an electron to form a 1- ion.
The electronegativity measures:
the ability of the atom to hold on to its own electrons
the ability of the atom to attract electrons from other atoms in compounds.

93. In each case, which bond is more polar:
WORKSHEET EXAMPLE
In each case, which bond is more polar:
a) B – Cl or C – Cl,
b) P – F or ? Indicate in each case which atom has the partial negative charge.

94. The cyanate ion, NCO-, has three possible Lewis structures.
WORKSHEET EXAMPLE
The cyanate ion, NCO-, has three possible Lewis structures.
Draw these three structures assign formal charges in each.
Which lewis structure is dominant?

95. WORKSHEET EXAMPLE
Using Formal Charges and the concept of Resonance, show why the correct Lewis structure of BF3 is the one in which B has an incomplete octet and there are only single bonds present.
Giving boron a filled octet places a negative charge on the boron and a positive charge on fluorine.
This would not be an accurate picture of the distribution of electrons in BF3.
Therefore, structures that put a double bond between boron and fluorine are much less important than the one that leaves boron with only 6 valence electrons.