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. Do Now
Draw Lewis structures for each molecule below. Then use VSEPR theory to predict their shapes.
SiCl4
SCl2 CO
CH2O
PH3
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3. Nonpolar: Electrons are shared equally.
Bond Polarity
Covalent Bonds
Nonpolar: Electrons are shared equally.
Polar: Electrons are shared unequally.
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4. Electronegativity Values for Selected Elements
Interpret Data
Electronegativity – ability of an atom to attract electrons when the atom is in a compound.
Electronegativity Values for Selected Elements H
2.1 Li
1.0 Be
1.5 B
2.0 C
2.5 N
3.0 O
3.5 F
4.0 Na
0.9 Mg
1.2 Al Si
1.8 P S Cl K
0.8 Ca Ga
1.6 Ge As Se
2.4 Br
2.8 Rb Sr In
1.7 Sn Sb
1.9 Te I Cs
0.7 Ba Tl Pb Bi
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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
Type of bond
Example
0.0–0.4
Nonpolar Covalent
H—H (0.0)
0.4–1.9
Polar Covalent
δ+ δ– δ+ δ–
H—Cl (0.9) 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. Al and Cl
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7. 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 has a slight negative charge.
H has a slight positive charge.
δ+ δ–
H—Cl
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8. 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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9. Dipole: Molecule that has two poles (electrically charged regions)
Bond Polarity
Polar Molecule: Molecules in which electrons are unequally distributed. One end 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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10. 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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11. Arrange the following covalent bonds in order of decreasing polarity.
Do Now
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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12.
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13. Attractions Between Molecules
Intermolecular Forces
Weaker than ionic or covalent bonds
Types
Hydrogen Bonds
Van der Waals Forces
Dipole Interactions
Dispersion Forces
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14. Attractions Between Molecules
Hydrogen bonds - a hydrogen weakly bonds to the unshared electron pair of another atom.
The positive region of a water molecule attracts the negative region of another water molecule.
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15. Attractions Between Molecules
Hydrogen Bonds
~5% strength of average covalent bond.
Strongest intermolecular force
Determines properties of water and biological molecules
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16. Attractions Between Molecules
Dipole Interactions
Polar molecules are attracted to each other
Similar to, but much weaker than, ionic bonds.
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17. 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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18. 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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19. END OF 8.4
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