1. CHEMICAL
BONDS

2. INDEX INTRODUCTION ELECTRONEGATIVITY
CHEMICAL BONDS AND THEIR FORMATION
HYBRIDATION
COVALENT BONDING CAPACITIY OF THE SECOND ROW ELEMENTS
DOUBLE AND TRIPLE BONDS
RESONANCE STRUCTURES
EXCEPTIONS TO THE OCTET RULE

3. Intramolecular Intermolecular Ionic bond Covalent Bond CHEMICAL BONDS
Between
same
atoms
Between
molecules
Ionic bond
Covalent
Bond
Hydrogen
Dipole-Dipole
Polar
Van Der Waals
Nonpolar
Metallic Bond
Coordinate
Crystal structure
Van Der Waals

4. INTRODUCTION
The force of attraction that holds atoms or ions together is called a chemical bond.

5. During the process of forming a chemical bond, energy is given out.
This energy is equal to that required to break the same chemical bond.
To gain a better understanding of chemical bonds we need to study electronegativity.
Electronegativity plays an important role in bond formation.

6. ELECTRONEGATIVITY
First proposed in 1934 by American
physicist R.S. Mulliken.
Electronegativity is the tendency of an atom to attract the bonding electrons within a compound to itself.
Depends upon the nuclear charge (proton
number) and the atomic radius of the atom.

7. Electronegativity decreases from top to bottom in a group and increases from left to right across a period.

8. CHEMICAL BONDS
AND THEIR FORMATION
The noble gases (He, Ne, Ar, Kr, Xe and Rn) which form group 8A in the periodic table are the most stable elements.
They all have the ns2 np6 electron configuration (except He which has the 1s2 configuration).

9. The tendency of atoms to make the number of their valence electrons eight, like
the nobel gases, is known as the octet rule.
There are two ways for the elements
to gain their octet and obtain a noble gas electron configuration.
1. Electron transfer (creates an ionic bond)
2. Electron sharing (creates an covalent bond)

10. e e e e e e e Mg Na K O Cl Ne Ar He Kr

11. The electron configuration of fluorine is
Orbital Representation of Chemical Bonds:
The electron configuration of fluorine is
1s22s22p5

12. Fluorine molecule, F2:

13. Electron Dot Representation (Lewis Symbol)
of Chemical Bonds:
Electron dot representation of fluorine atom is
When two fluorine atoms combine with each other a F2 molecule forms.

16. Gilbert Newton Lewis ( ): Lewis was an American scientist born in 1875, in Massachusetts USA.
He started this academic career in 1912 and proposed the theory of electron sharing in 1916 which as we have seen is of great importance to chemists. Because of this theory “electron dot representation” is also named “Lewis dot structure”.

18. The Line Representation of Chemical Bonds:
Two electrons (:) are shown by a line (–).
The line representation of fluorine molecule is;

19. IONIC BONDS
Formed by the transfer of electrons
Formed between atoms that have an electronegativity difference greater than about 1.7

21. To complete their octets, sodium gives one electron to chlorine.
and
Na+ cation and Cl- anion are formed.
An ionic bond is formed between the sodium and chlorine ions.

26. COVALENT BONDS
Formed as a result of electron sharing
Generally formed between two nonmetals
Covalent bonds can be classified into three groups nonpolar, polar and coordinate covalent bonds.

30. Nonpolar Covalent Bonds
Formed between two atoms with the same electronegativity values.
For example; H2, N2 , Cl2 and O2 molecules
The electronegativity difference between the atoms which form the bond is zero
The charge distribution within the bond is equal.

32. Polar Covalent Bonds
Formed between two atoms with different electronegativity values.
For example; HCl, HI, HF…
Electronegativity value of hydrogen is 2.1
Electronegativity value chloride is 3.0
Electronegativity difference is 0.9 which is smaller than 1.7

37. Coordinate Covalent Bonds
Both of the shared electrons come from only one of the atoms

38. HYBRIDIZATION
The mixing of orbitals from different energy levels to form new orbitals all with the same energy is called hybridization.
The new orbitals formed at this new energy level are called hybrid orbitals.
There are three types of hybridization between `s` and `p`orbitals which are sp,sp2,sp3.

39. sp hybridization 4Be : 1s2 2s2 2p0 4Be : 1s2 2s1 2p1
sp hybrid orbitals are formed as a result of mixing one s orbital with one p orbital.
4Be : 1s2 2s p0
Unhybridized
ground state
4Be : 1s2 2s p1
hybridized
excited state

40. BeH2 molecule
4Be : 1s2 2s p1
1H : 1s1

43. sp2 hybridization 5B : 1s2 2s2 2p1 5B : 1s2 2s1 2p2
sp2 hybrid orbitals are formed as a result of mixing one s orbital with two p orbitals.
5B : 1s2 2s p1
Unhybridized
ground state
5B : 1s2 2s p2
hybridized
excited state

44. BH3 molecule
5B : 1s2 2s p2
1H : 1s1

45. All angles are 120 degrees.

47. sp3 hybridization 6C : 1s2 2s2 2p2 6C : 1s2 2s1 2p3 Unhybridized
ground state
6C : 1s2 2s p3
hybridized
excited state

48. CH4 molecule
6C : 1s2 2s p3
1H : 1s1

49. All angles are degrees

52. COVALENT BONDING CAPACITY OF THE SECOND ROW ELEMENTS
The number of unpaired valence electrons of an element shows the covalent bonds of that element.
The number of half-filled orbitals indicates the number of bonds that the atoms can form.

53. Bonding Capacity of Lithium:
Lithium is a metal so it tends to form bonds with nonmetals. The compound lithium hydride, LiH, is made up of crystals with cubic lattice structure.
3Li : 1s2 2s1
Bonds : Polar
Molecule: Polar
Shape : Linear
1H : 1s1

55. Bonding Capacity of Beryllium :
The electron configuration of Be is 1s22s2 it has two valence electrons in its ground state. It shouldn't be able to form a covalent bond as the electrons are paired.
4Be : 1s2 2s p0
Unhybridized
ground state
4Be : 1s2 2s p1
hybridized
excited state

56. 4Be : 1s2 2s1 2p1 1H : 1s1 Bonds : Polar Molecule: Nonpolar
Shape : Linear

57. The direction of orbitals is linear.
The shape of the BeH2 molecule.

58. Bonding Capacity of Boron:
Has three valence electrons, only one of them is unpaired in the ground state.
5B : 1s2 2s p1
Unhybridized
ground state
5B : 1s2 2s p2
hybridized
excited state

59. BH3 molecule 5B : 1s2 2s1 2p2 1H : 1s1 Bonds : Polar
Molecule: Nonpolar
Shape : trigonal planar

60. The shape of the BH3 molecule.
Orientation of orbitals is trigonal planar.
The shape of the BH3 molecule.

61. Bonding Capacity of Carbon:
Has four valence electrons, of which only two are unpaired in the ground state.
6C : 1s2 2s p2
Unhybridized
ground state
6C : 1s2 2s p3
hybridized
excited state

62. CH4 molecule 6C : 1s2 2s1 2p3 1H : 1s1 Bonds : Polar
Molecule: Nonpolar
Shape : Tetrahedral

63. The shape of the CH4 molecule.

64. Bonding Capacity of Nitrogen:
Has five valence electrons.
7N : 1s2 2s p3
1H : 1s1
Bonds : Polar
Molecule: Polar
Shape : Trigonal pyramidal

65. The shape of the NH3molecule is trigonal pyramidal.
The orientation of the orbitals in the NH3 molecule is trigonal
pyramidal.

67. Bonding Capacity of Oxygen:
Has six valence electrons.
8O : 1s2 2s p4
1H : 1s1
Bonds : Polar
Molecule: Polar
Shape : Angular,bent

68. The orientation of the orbitals in the H2O molecule.
The shape of the H2O molecule

69. Bonding Capacity of Fluorine:
Has seven valence electrons, only one of them is unpaired, so the fluorine atom can form one bond.
9F : 1s2 2s p5
1H : 1s1
Bonds : Polar
Molecule: Polar
Shape : Linear
HF,HCl …

70. Distribution of bonding electrons in the HF molecule.

71. Fluorine can form molecules with eachother
9F : 1s2 2s p5
9F : 1s2 2s p5
Bonds : Nonpolar
Molecule: Nonpolar
Shape : Linear
F2 , Cl2 …

74. Bonding Capacity of Neon:
Has eight valence electrons and all of them are paired.
Neon is very unreactive and does not bond with any other element.
Helium and argon are very unreactive.
Krypton and xenon may form bonds under certain conditions.

75. double bond contains one σ and one π bond.
DOUBLE AND TRIPLE COVALENT BONDS
Some atoms(C,O,N…) can form double or triple bonds as well as single bonds.
Two types of bonds may be formed when orbitals overlap which are sigma (σ) and pi (π) bonds.
single bond is σ bond.
double bond contains one σ and one π bond.
triple bond contains one σ and two π bonds.

76. Japanese scientist Prof. Dr
Japanese scientist Prof. Dr. Masaru Emoto in his book “Water crystals” where there are 70 pictures of crystals: Water is not a lifeless; it is composed of living and sensible crystals.
Three years ago Emoto in his researches by microscope discovered that water crystals give different reactions to external influences. According to these researches, water crystals give reactions to music, word and concepts as well as influences of external environment.

78. SIGMA (σ) BONDS
Formed by the end to end overlap of two orbitals.
Overlap can take place between s orbitals, p orbitals or hybrid orbitals.
In methane there are 4 C-H sigma bonds whereas in ethane there are 6 C-H and 1 C-C, sigma bonds.

81. Pi (π) BONDS
Formed by the side by side overlap of two parallel p orbitals.
The electron cloud lies above and below the plane formed by sigma(σ) bonds.
Pi (π) bonds are weaker than σ bonds.

83. In the C2H4 molecule, unhybridized p orbitals overlap in side by side and form a p bond.

85. In the C2H2 molecule, unhybridized p orbitals form two π bonds by overlapping side by side.

86. RESONANCE STRUCTURES
If the valence electrons in a molecule are capable of several alternative arrengements which differ only a small amount in energy each arrangement is called a resonance.
The bonds in the ozone molecule, O3, are identical and have a length of 128 pm.

87. The structure of the ozone molecule

88. EXCEPTIONS TO THE OCTET RULE
Most atoms complete their valence shell with eight (octet) electrons to become stable.
However, some exceptions occur.
ELECTRON DEFICIENCY
EXPANDED OCTETS
FREE RADICALS

89. ELECTRON DEFICIENCY
Beryllium and boron do not complete their octet in their covalent compounds because these atoms have less than four valence electrons.

90. EXPANDED OCTETS
Some atoms in 3rd period may have more than eight electrons in their valence orbitals.
Expanded octet in the PF5 molecule

91. Orbital orientation in PF5 the molecule is trigonal bipyramidal
Molecular model of PF5

92. FREE RADICALS
Compounds that have unpaired electrons in their structures are called free radicals.
NO and NO2 are two examples of free radicals.
Free radicals are chemically active substances.