1. Unit 5 – Lewis Structures of Covalent Molecules
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2. Lewis Dot Structure for Molecules
Nonmetals form covalent bonds to achieve a noble gas electron configuration:
8 valence electrons for all atoms, except
2 valence electrons for helium and hydrogen.
The pair of electrons in the bond is called the shared pair.
The pairs of electrons NOT in the bond are called unshared pairs.
Single Covalent Bonds + 

3. Lewis Structure of Molecules
Add up all the valence electrons. (skip 1 line for later)
This is the number of dots in the final structure.
Place least electronegative element in center. (Never H)
Place two electrons between central and outer atoms
Fill up outer atom octets.
Fill up central atom octet.
Skip 2 lines for later. Box these steps.
PCl3
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4. Lewis Structure of Molecules
Add up all the valence electrons.
If it is an anion, add one electron for each negative charge.
If it is a cation, subtract one electron for each positive charge.
PCl3
(7) = 26
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5. Lewis Structure of Molecules
The central atom is the least electronegative element (Never hydrogen!).
Connect the outer atoms to it by single bonds. (Each line represents 2 shared electrons.)
Keep track of the electrons:
= 20
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6. Lewis Structure of Molecules
4. Fill the octets of the outer atoms.
Keep track of the electrons:
= 20; = 2
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7. Lewis Structure of Molecules
5. Fill the octet of the inner atoms.
Keep track of the electrons:
= 20; = 2; = 0
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8. Lewis Structure Practice
H2S
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9. Lewis Structure Practice
NH3
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10. Lewis Structure Practice
Add up all the valence electrons. This is the number of dots in the final structure.
Place least electronegative element in center. (Never H)
Place two electrons between central and outer atoms
Fill up outer atom octets.
Fill up central atom octet.
CF4
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11. Lewis Structure Practice
Add up all the valence electrons.
CF4
(7) = 32
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12. Lewis Structure Practice
The central atom is the least electronegative element (Never hydrogen!).
Connect the outer atoms to it by single bonds. (Each line represents 2 shared electrons.)
Keep track of the electrons:
= 24
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13. Lewis Structure Practice
4. Fill the octets of the outer atoms.
Keep track of the electrons:
= 24; = 0
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14. Lewis Structure Practice
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15. Double and Triple Covalent Bonds
A bond that involves one shared pair of electrons (2 total) is a single covalent bond.
A bond that involves two shared pairs of electrons (4 total) is a double covalent bond.
A bond that involves three shared pairs of electrons (6 total) is a triple covalent bond.

16. Multiple Covalent Bonds
Bond length (distance between centers of bonded atoms) DEcreases with increased # of shared electrons.
Bond strength INcreases with increased # of shared electrons
Bond
# of shared electrons
Bond Length
Bond Strength
Single 2
Longest
Weakest
Double 4
Triple 6
Shortest
Strongest
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17. Lewis Structure of Molecules Double and Triple Bonds
Add up all the valence electrons. This is the number of dots in the final structure.
Place least electronegative element in center. (Never H)
Place two electrons between central and outer atoms
Fill up outer atom octets.
Fill up central atom octet.
Make double or triple bonds if necessary so all atoms meet the octet rule.
CO2
4 + 2(6) = 16

18. Lewis Structure of Molecules Double and Triple Bonds

19. Lewis Structure of Molecules Double and Triple Bonds
8 electrons in bonds,
Plus 8 electrons as lone pairs.
Total of 16 electrons.
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20. Carbon dioxide
To meet the octet rule for C and O requires two double bonds
Each atom can count all the electrons in the bond

21. Lewis Structure of Molecules Double and Triple Bonds
Add up all the valence electrons. This is the number of dots in the final structure. a) _____________
Place two electrons between central and outer atoms
Fill up outer atom octets.
Make double or triple bonds if necessary so all atoms meet the octet rule.
5 + 5 = 10
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22. Lewis Structure of Molecules Double and Triple Bonds
Final structure has 6 shared electrons (triple bond) and two unshared pairs of electrons.
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23. Polyatomic ions are MOLECULES with COVALENT bonds that have a charge.
A polyatomic ion, is a tightly bound group of atoms that has a positive or negative charge and behaves as a unit.
Polyatomic ions are MOLECULES with COVALENT bonds that have a charge.
Examples:
Carbonate: CO32-
Phosphate: PO43-
Nitrate: NO3-
Phosphite: PO33-
Nitrite: NO2-
Hydroxide: OH- and Cyanide: CN-
Sulfate: SO42-
Sulfite: SO32-
Ammonium NH4+
The polyatomic ammonium ion (NH4), present in ammonium sulfate, is an important component of fertilizer for field crops, home gardens, and potted plants.

24. Lewis Structure of Polyatomic Ions
Add up all the valence electrons. ADD 1 for every negative charge, SUBTRACT 1 for every positive charge. This is the number of dots in the final structure.
Place least electronegative element in center. (Never H)
Place two electrons between central and outer atoms
Fill up outer atom octets.
Fill up central atom octet.
Make double or triple bonds if necessary so all atoms meet the octet rule.
5 + 4x1 - 1 = 8
NH4+
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25. Lewis Structure of Polyatomic Ions
SO42–
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26. Lewis Structure of Polyatomic Ions
BrO3–
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27. Ammonium, NH4+ Total valence electrons = 8
Polyatomic Ions
Ammonium, NH4+ Total valence electrons = 8
Total number of electrons around each atom: 8 + 4x2 = 16
Total number of shared electrons: 16-8=8
The polyatomic ammonium ion (NH4), present in ammonium sulfate, is an important component of fertilizer for field crops, home gardens, and potted plants.

28. 8.4 Bond Polarity Polarity
The bonding pairs of electrons in covalent bonds are pulled by the nuclei.
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.

29. Nonpolar covalent bond
Polarity
Electronegativity values determine the charge distribution in a covalent bond.
> Electronegativity: a measure of an atom’s attraction for electrons
Nonpolar covalent bond
Bond between two nonmetal atoms with the same electronegativity
Bonding electrons are shared equally
Between identical atoms (Ex: Cl2, O2, N2)
8.4

30. Polarity
Polar covalent bond
Bond between 2 nonmetal atoms with the different electronegativity values
Bonding electrons are shared unequally
the atom with the higher electronegativity pulls more strongly and has a partial negative charge
the atom with the lower electronegativity pulls less strongly and has a partial positive charge
8.4

31. Polar Molecules
In a polar molecule, one end of the molecule is slightly negative and the other end is slightly positive.
A molecule that has two poles (opposite charges) is called a dipolar molecule, or dipole.

32. Electronegativity Values: H: 2.1 Cl: 3.0 8.4 Bond Polarity
The chlorine atom attracts the electron cloud more than the hydrogen atom does – forms a dipole
Electronegativity Values:
H: 2.1
Cl: 3.0
8.4
Bond Polarity
This electron-cloud picture of hydrogen chloride shows that the chlorine atom attracts the electron cloud more than the hydrogen atom does. Inferring Which atom is more electronegative, a chlorine atom or a hydrogen atom?

33. 8.4
Bond Polarity

34. 8.4 Bond Polarity Practice Problem
Identify the bonds between each pair as nonpolar covalent, moderately polar covalent, very polar covalent, or ionic:
Difference Type of Bond
a) H and Br
b) C and O
c) Br and Br
d) K and Cl
e) P and F
polar
polar
nonpolar
2.2 ionic
polar

35. valence-electron pairs stay as far apart as possible.
Shapes of Molecules
The repulsion between electron pairs causes molecular shapes to adjust so that
valence-electron pairs stay as far apart as possible.

36. 8.3
Shapes of Molecules
Methane is a tetrahedral molecule. The hydrogens in methane are at the four corners of a regular tetrahedron, and the bond angles are all 109.5°. Interpreting Diagrams How do the resulting H—C—H bond angles compare to the tetrahedral angle?

37. Shapes of Molecules The carbon dioxide molecule is linear.
Methane is a tetrahedral molecule. The hydrogens in methane are at the four corners of a regular tetrahedron, and the bond angles are all 109.5°. Interpreting Diagrams How do the resulting H—C—H bond angles compare to the tetrahedral angle?

38. Shapes of Molecules The shape of a water molecule is “bent triatomic."
Methane is a tetrahedral molecule. The hydrogens in methane are at the four corners of a regular tetrahedron, and the bond angles are all 109.5°. Interpreting Diagrams How do the resulting H—C—H bond angles compare to the tetrahedral angle?

39. Shapes of Molecules Methane has a tetrahedral shape.
Methane is a tetrahedral molecule. The hydrogens in methane are at the four corners of a regular tetrahedron, and the bond angles are all 109.5°. Interpreting Diagrams How do the resulting H—C—H bond angles compare to the tetrahedral angle?

40. Shapes of Molecules Ammonia (NH3) has a pyramidal shape.
Methane is a tetrahedral molecule. The hydrogens in methane are at the four corners of a regular tetrahedron, and the bond angles are all 109.5°. Interpreting Diagrams How do the resulting H—C—H bond angles compare to the tetrahedral angle?

41. Shapes of Molecules Nine Possible Molecular Shapes
Shown here are common molecular shapes.