Study Guide Chapters 12 – 14

Key

1. Define: electronegativity, dipole, dipole moment, Van der Waals Forces.

• electronegativity: The tendency of a bonded atom to attract electrons towards itself. (example: when F bonds to O the F pulls the electrons closer to it because it’s electronegativity is higher.)

• Dipole: a polar molecule• Dipole moment: measurement of the amount of

polarity. (example: a molecule that is more polar would have a greater dipole moment).

• Van der Waals Forces: Attractive forces between adjacent molecules. (example: the bigger a molecule is and the more polar it is the better it is able to attract adjacent molecules).

2. State the differences and similarities between ionic, covalent, and metallic bonds (see the “Four Types of Bonding” table in your notebook).

Type of Bonding

Explanation of Bonding

Form and Properties.

EXAMPLES

1. Metallic Nondirectional bonds involving

positive metal ions with free electrons moving throughout

the metal.

Crystals with high melting points,

shiny, malleable, conductors,

copper, bronze, steel

2. Ionic The transfer of electron(s) from a metal atom to a nonmetal atom

which results in an attractive force

between a positive metal ion and a

negative nonmetal ion.

Crystals with high melting points

which are brittle and act as insulators.

NaCl, CuO, FeBr2, etc

3. Polar Covalent

Electrostatic attraction involving unequal sharing of

electron(s) in overlapping orbitals between different types of nonmetal

atoms.

Molecules with low melting points which if solid are brittle and act as

insulators.

H2O, CO2, sucrose

(C6H12O6), etc.

4. Nonpolar Covalent

Electrostatic attraction involving

equal sharing of electron(s) in

overlapping orbitals between the same type of nonmetal

atoms.

Molecules with low melting points which if solid are brittle and act as

insulators.

H2, N2, O2

3. Contrast the number of shared pairs, the number of electrons, the strength, and the length within

single, double, and triple bonds.

Single bonds have one shared pair of electrons (two shared electrons). They are the weakest and longest of the covalent bonds. Triple bonds have three shared pairs of electrons (6 shared electrons). They are the strongest and shortest of the covalent bonds. Double bonds have two shared pairs of electrons (four shared electrons). They have a strength and length between that of single and triple bonds.

4. What are the differences between shared pairs and unshared pairs?

• Shared pairs of electrons are represented in Lewis structures by –’s. They represent bonds and belong to both atoms which they connect.

• Unshared pairs (lone pairs) also called lone pairs are represented in Lewis structures by a pair of x’s, ’s or o’s. They belong only to the atom which they are placed on.

4. What are the differences between shared pairs and unshared pairs?

5. How does electronegativity vary within the groups and periods of the periodic table?

• The closer an atom is to “F” in the periodic table the higher the electronegativity. (Remember that the noble gases have no electronegativities).

6. How can we predict the type of bond formed between atoms by using (a) a periodic table and

(b) a table of electronegativities.

• A bond between a metal and a nonmetal is ionic. A bond between metals is metallic. A bond between different nonmetal atoms is polar covalent. A bond between the same nonmetal atoms is nonpolar covalent.

• The electronegativity difference can be used to determine bond type as well.

– If the electronegativity difference is greater than 1.67 between two atoms the bond between them is ionic.

– If the electronegativity difference is less than 1.67 the bond is covalent.

7. How do differences in electronegativities influence bond strength.

• The greater the electronegativity difference between two atoms the stronger the bond is between them.

8. Complete the table:

Number of atoms attached to the central atom

Number of unshared pairs attached to the central atom

Molecular Shape

4

0

2

1

Trigonal planer

2

2

Trigonal pyramidal

linear

9. Contrast the attractive forces within solids, liquids, and gases at room temperature.

• At any given temperature a solid has the greatest attractive forces and a gas has the least. The attractive forces in a liquid are somewhere in between.

10. How do the size and polarity of molecules affect their Van der Waals forces.

• The larger and more polar a molecule is the greater its Van der Waals forces.

11. PBr3

11. NO2-

11. ClF2+

11. FNO2

11. N3-

11. CF4