1. Hydrogen and Chlorine:
Covalent Bonding: H Cl
or H-Cl H + Cl
Sharing electrons
2) Hydrogen & hydrogen.
(Single covalent bond)
3) Fluorine & fluorine.
4) Nitrogen & hydrogen.
5) Carbon & hydrogen.

2. Covalent Bonding: The Octet Rule in Covalent Bonding
In covalent bonds, electron sharing
usually occurs so that atoms attain
the electron configurations of noble
gases.
The pair of shared electrons forming the
covalent bond is also often represented
as a dash, as in H—H for hydrogen
(structural formula )

3. Two pairs of electrons shared (Double covalent bond)
Covalent Bonding:
Represent the formation of the chemical bond between the following atoms (Lewis dot structure):
1)Oxygen and oxygen: O O
or O=O O + O
Two pairs of electrons shared
(Double covalent bond)
2) Nitrogen & nitrogen.

4. Three pairs of electrons shared
2) Nitrogen & nitrogen: N + N
or N N N N
Three pairs of electrons shared
(Triple covalent bond)
A double covalent bond is a bond that
involves two shared pairs of electrons.
Similarly, a bond formed by sharing three
pairs of electrons is a triple covalent bond.

5. Using electron-dots structure represent the following molecules:
CO2
2) C2H4
3) C2H2

6. Bond Polarity:
How do electronegativity values determine the charge distribution in a polar bond?
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7. Bond Polarity
The bonding pairs of electrons in covalent bonds are pulled between the nuclei of the atoms sharing the electrons.
The nuclei of atoms pull on the shared electrons, much as the knot in the rope is pulled toward opposing sides in a tug-of-war.

8. Bond Polarity
A polar covalent bond, known also as a polar bond, is a covalent bond between atoms in which the electrons are shared unequally.
The more electronegative atom attracts more strongly and gains a slightly negative charge.
The less electronegative atom has a slightly positive charge.
δ+ δ–
H—Cl
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9. Bond Polarity H-H Br-Br O=O
When the atoms in the bond pull equally (as occurs when identical atoms are bonded), the bonding electrons are shared equally, and each bond formed is a nonpolar covalent bond.
H-H
Br-Br
O=O
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10. Bond Polarity
The electronegativity difference between two atoms tells you what kind of bond is likely to form.
Electronegativity Differences and Bond Types
Electronegativity difference range
Most probable type of bond
Example
0.0 – 0.4
Nonpolar covalent
H—H (0.0)
0.4 – 1.0
Moderately polar covalent
δ+ δ–
H—Cl (0.9)
1.0 – 1.5
>1.5
Very polar covalent
Ionic (if a metal is present)
H—F (1.9)
≥ 1.5
Ionic
Na+Cl– (2.1)
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11. Electronegativity Chart
Sample Problem
Identifying Bond Type
Which type of bond (nonpolar covalent, moderately polar covalent, very polar covalent, or ionic) will form between each of the following pairs of atoms?
a. N and H
b. F and F
c. Ca and Cl
d. Al and Cl
Electronegativity Chart
Page 181

12. Calculate the electronegativity difference between the two atoms & determine the bond type.
a. 3.0(N) - 2.1(H) =
0.9
Moderately polar covalent.
Nonpolar covalent.
b. 4.0(F) - 4.0(F) =
0.0
c. 3.0(Cl) - 1.0(Ca) =
2.0
Ionic.
d. 3.0(Cl) - 1.5(Al) =
1.5
Very polar covalent.
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13. A molecule that has two poles is called a dipolar molecule, or dipole.
Describing Polar Covalent Molecules:
In a polar molecule, one end of the molecule is slightly negative, and the other end is slightly positive. δ+ δ-
A molecule that has two poles is called a dipolar molecule, or dipole.
Example:
δ+ δ–
H—Cl
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14. A carbon dioxide molecule has two polar bonds, it is linear. O C O
The effect of polar bonds on the polarity of an entire molecule depends on the shape of the molecule and the orientation of the polar bonds.
A carbon dioxide molecule has two polar bonds, it is linear.
O C O
Therefore, the molecule is nonpolar.
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15. The water molecule also has two polar bonds.
However, the water molecule is bent rather than linear.
Therefore, the bond polarities do not cancel and a water molecule is polar.
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16. Intermolecular attractions:
Intermolecular attractions are weaker than either ionic or covalent bonds.
Molecules can be attracted to each other by a variety of different forces:
The weakest attractions between
molecules are called Van der Waals forces (London dispersion forces).
Can be found between nonpolar molecules as
O2(g), H2(g), Cl2(g), CO2(g).
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17. In hydrogen's case the attractions are so weak that the molecules have to be cooled to 21 K
(-252°C) before the attractions are enough to condense the hydrogen as a liquid.
Helium's molecules won't stick together to form a liquid until the temperature drops to 4 K (-269°C).

18. How molecular size affects the strength of the Van der Waals Forces?
The boiling points of the noble gases are
Helium oC
Neon oC
Argon oC
Krypton oC
Xenon oC
Radon oC

19. How molecular shape affects the strength of the Van der Waals Forces?
Long thin molecules can also lie closer together - these attractions are at their most effective if the molecules are really close (stronger attraction):

20. 2) Dipole interactions occur when polar
molecules are attracted to one another.
Ex. HCl(g), HBr(g). _ _ + +
++=
++=
These interactions are stronger than Van der Waals, but much weaker than ionic bonds.

21. 3) Hydrogen Bonds:
Are attractive forces in which a hydrogen covalently bonded to a very electronegative atom (F, O, N) is also weakly bonded to an unshared electron pair of another electronegative atom.
Oδ O δ-
Hδ Hδ Hδ Hδ+
O δ-
Hδ Hδ+
Hydrogen bonding always involves hydrogen.

22. Hydrogen bonds are the strongest of the intermolecular forces.
They are extremely important in determining the properties of water and biological molecules.
Ex: Water is liquid at room temperature because the hydrogen bond between its molecules.
Boiling Points
H2O oC
H2S oC
H2Se oC

23. Intermolecular attraction forces (weaker than ionic & covalent bonds)
Van der Waals
(the weakest)
Nonpolar molecules
Intermolecular attraction forces
(weaker than ionic & covalent bonds)
Dipole-dipole
(stronger than )
Polar molecules
Hydrogen bonds
(the strongest) Molecules containing F-H; O-H or N-H

24. Examples: Cl2(g), Br2(l), CCl4(l).
Properties of Molecular Compounds:
1) Weak attraction forces between molecules.
2) Low melting and boiling points.
3) Gases, liquids or solids at room temperature.
4) The solids are very soft.
5) Poor conductors of electricity.
6) Polar covalent: Show solubility in water.
Ex. HCl(g), NH3(g)
No polar covalent: Show solubility in non polar
solvents (CCl4, oil, gasoline, acetone, etc.)
Examples: Cl2(g), Br2(l), CCl4(l).

25. Covalent Network Solids:
A network solid or covalent network solid is a chemical compound in which the atoms are bonded by covalent bonds in a continuous network extending throughout the material
Examples:
Quartz or Silica
(SiO2 silicon dioxide),
diamond & graphite
(made by carbon).
diamond graphite

26. Covalent Network Solids:
Properties:
1) Brittle & extremely hard solids.
2) Very high melting points.
3) Non conductors (heat or electricity),
except graphite (good conductor)