1. Chapter 8 Covalent Bonding 8.4 Polar Bonds and Molecules
8.1 Molecular Compounds
8.2 The Nature of Covalent Bonding
8.3 Bonding Theories
8.4 Polar Bonds and Molecules
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2. Nonpolar Covalent Bond Electrons are shared equally.
Bond Polarity
Nonpolar Covalent Bond
Electrons are shared equally.
Polar Covalent Bond
Electrons are shared unequally.
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3. Hydrogen Chloride (HCl)
Bond Polarity
Hydrogen Chloride (HCl)
Hydrogen has an electronegativity of 2.1 Chlorine has an electronegativity of 3.0.
This covalent bond is polar.
Cl acquires a slight negative charge.
H acquires a slightly positive charge.
δ+ δ–
H—Cl
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4. Hydrogen Chloride (HCl)
Bond Polarity
Hydrogen Chloride (HCl)
Partial charges can be shown as clouds of electron density.
They are also represented by an arrow pointing to the more electronegative atom.
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5. 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)
0.5–1.0
Moderately polar covalent
δ+ δ–
H—Cl (0.9)
1.0–1.9
Very polar covalent
H—F (1.9)
>2.0
Ionic
Na+Cl– (2.1)
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6. Sample Problem 8.3
Identifying Bond Type
Which type of bond will form between the following pairs of atoms?
a. N and H
b. F and F
c. Ca and Cl
d. Na and F
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7. Sample Problem 8.3
Solve
Based on the electronegativity difference, determine the bond type using Table 8.4.
a. N(3.0), H(2.1); 0.9; moderately polar covalent
b. F(4.0), F(4.0); 0.0; nonpolar covalent
c. Ca(1.0), Cl(3.0); 2.0; ionic
d. Na(0.9), F(4.0); 1.5; ionic
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8. Dipole: Molecule that has two poles (electrically charged regions)
Bond Polarity
Polar Molecule: Molecules in which the electrons are unequally distributed. One end of the molecule is slightly negative and the other end is slightly positive.
Dipole: Molecule that has two poles (electrically charged regions)
HCl is an example of a dipole
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9. Molecule polarity depends on the shape and orientation of the bonds.
Bond Polarity
Molecule polarity depends on the shape and orientation of the bonds.
Nonpolar Polar
O C O
bond polarities cancel bond polarities do not cancel
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10. Arrange the following covalent bonds in order of decreasing polarity.
H-Cl
H-C
H-F
H-O
H-H
S-Cl
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11. Arrange the following covalent bonds in order of decreasing polarity.
H-Cl (3.0 – 2.1 = 0.9)
H-C ( = 0.4)
H-F ( = 1.9)
H-O ( = 1.4)
H-H ( = 0)
S-Cl ( = 0.5)
c,d,a,f,b,e
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12. Attractions Between Molecules
Intermolecular Forces
Weaker than ionic or covalent bonds
Types
Van der Waals Forces
Dipole Interactions
Dispersion Forces
Hydrogen Bonds
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13. Attractions Between Molecules
Dipole Interactions
Polar molecules are attracted to each other
Similar to, but much weaker than, ionic bonds.
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14. Attractions Between Molecules
Dispersion Forces
Weakest of all molecular interactions
Caused by the motion of electrons.
When electrons happen to be more on one side of a molecule, they influence their neighbor’s electrons
Forces increase with number of electrons
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15. Attractions Between Molecules
Dispersion Forces
F and Cl have relatively few electrons and are gases at room temperature.
Bromine attracts other bromine molecules enough to make it a liquid at room temperature.
Iodine is a solid at room temperature.
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16. Attractions Between Molecules
Hydrogen bonds - a hydrogen covalently bonded to a very electronegative atom is also weakly bonded to an unshared electron pair of another atom.
The other atom may be in the same molecule or in a nearby molecule.
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17. Attractions Between Molecules
Hydrogen Bonds
The positive region of one water molecule attracts the negative region of another water molecule.
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18. Attractions Between Molecules
Hydrogen Bonds
A hydrogen bond has about 5 percent of the strength of the average covalent bond.
Hydrogen bonds are the strongest of the intermolecular forces.
They are extremely important in determining the properties of water and biological molecules.
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19. END OF 8.4
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