Unit 6
Name _______________________________ # ________
Section 1: Bonding
1. What is the duet rule and what elements follow it? What is the octet rule
and what elements follow it?
2. How many valence electrons are in each of the following elements?
a.
b.
c.
d.
e.
C
H
Si
Kr
P
_______
_______
_______
_______
_______
f.
g.
h.
i.
j.
O
N
Cl
Ar
Cu
_______
_______
_______
_______
_______
Answer the following questions.
3. What types of elements form ionic bonds?
4. What happens to the electrons in an ionic bond?
5. What types of elements form covalent bonds?
6. What happens to the electrons in a covalent bond?
7. Explain the difference between polar covalent and nonpolar covalent
bonds.
8. Define electronegativity.
9. Express the periodic trend of electronegativity in words.
10. Record the electronegativities for the following elements.
a.
b.
c.
d.
C
H
Si
Kr
_______
_______
_______
_______
e.
f.
g.
h.
O
N
Cl
Ar
_______
_______
_______
_______
11. For each of the following pairs of elements, calculate the difference in their
electronegativities.
a.
b.
c.
d.
e.
bond between 2 hydrogen atoms in H2
bonds between H and O in water
bond between Na and Cl in sodium chloride
bond between C and Cl in carbon tetrachloride
bond between 2 chlorine atoms in Cl2
_______
_______
_______
_______
_______
12. Using the letters to represent the six pairs above, order the bonds from
those having the most polarity to those having the least.
___________________________________________________________
Section 2: Bond Types
Determine if the bond in the following pairs of atoms is either ionic, polar
covalent, or nonpolar covalent. If ionic show the charge (+ or -) and if polar
covalent show the partial charge(δ+ , δ- ) on the appropriate atoms.
1.
O–H
bond type: __________________________
2.
N–N
bond type: __________________________
3.
C–O
bond type: __________________________
4.
K – Cl
bond type: __________________________
5.
Mg – O
bond type: __________________________
6.
As – O
bond type: __________________________
7.
C–H
bond type: __________________________
8.
N -- F
bond type: __________________________
9.
P–H
bond type: __________________________
10.
C – Cl
bond type: __________________________
11.
P–O
bond type: __________________________
12.
Si – I
bond type: __________________________
13.
Fe – Cl
bond type: __________________________
14.
Sr – S
bond type: __________________________
Section 3: Binary Covalent Compounds
Correctly name the following compounds.
1. IF5
_________________________________________________
2. SeO
_________________________________________________
3. AsCl3 _________________________________________________
4. B2O3 _________________________________________________
5. P2O5
_________________________________________________
6. SiBr4 _________________________________________________
7. N2O5 _________________________________________________
8.
CO2
_________________________________________________
9. B2H6 _________________________________________________
10. H2O _________________________________________________
11. Cl2S _________________________________________________
Correctly write the formula for the following compounds.
12. phosphorus triiodide ________________
13. iodine monobromide ________________
14. nitrogen trichloride
________________
15. carbon monoxide
________________
16. sulfur hexafluoride
________________
17. dinitrogen tetroxide
________________
Section 4: Binary Covalent Compounds
1. Name the following compounds:
a. N2O5
__________________________
b. SO3
__________________________
c. P4O10
__________________________
d. N2O
__________________________
e. ClO2
__________________________
f. P2S5
__________________________
g. SO2
__________________________
h. CO2
__________________________
i. N2O4
__________________________
j. CO
__________________________
2. Write formulas for the following compounds:
a. boron trichloride
____________________
b. phosphorus pentachloride
____________________
c. nitrogen dioxide
____________________
d. carbon tetrachloride
____________________
e. dichlorine heptoxide
____________________
f. dichlorine monoxide
____________________
g. nitrogen trichloride
____________________
h. triphosphorus hexafluoride
____________________
Section 5: Lewis Dot Diagrams
I. Correctly draw the Lewis structure for the following elements in the spaces
provided.
a)
Na
b)
Cl
c)
B
d)
Ne
e)
Mg
f)
Si
g)
P
h)
Se
II. Correctly draw the Lewis structure for the following molecules in the spaces
provided.
a) H2
b)
MgF2
c)
BH3
d)
CH4
e)
H2Se
f)
AlCl3
g)
KBr
h)
SrI2
i)
NBr3
j)
LiH
k)
Na2S
l)
GeI4
Section 6: Lewis Dot Diagrams
Fill in the table below with the correct Lewis Dot Structure and the correct structural
formula (replace bonded electrons with lines).
CCl4
H2O
CO2
O2
PH3
SF2
Br2
N2
NCl3
SiS2
Section 7: Lewis Dot Diagrams
Correctly draw the Lewis Structure from each of the molecular formulas given in
the spaces provided.
1.
OH -1
2.
CHN
3.
SF6
4.
SCN -1
5.
XeF4
6.
CS2
7.
BrF3
8.
Br3-1
9.
BrO3-1
10.
SbI5
11.
NO3 -1
12.
SO2
13.
O -2
14.
PO4 -3
15.
CO3 -2
Section 8: VSEPR
For each of the following molecules, draw the Lewis dot diagram, determine the
electron group geometry, determine the molecular geometry, and draw the
molecular structure, 3-D shape, of the molecule.
1. CF4 Electron group geometry _____________________
Molecular geometry _________________________
Lewis Structure:
Molecular Structure:
2. NH3
Electron group geometry _____________________
Molecular geometry _________________________
Lewis Structure:
Molecular Structure:
3. SO2
Electron group geometry _____________________
Molecular geometry _________________________
Lewis Structure:
Molecular Structure:
4. CN-1 Electron group geometry _____________________
Molecular geometry _________________________
Lewis Structure:
Molecular Structure:
5. BrCl5 Electron group geometry _____________________
Molecular geometry _________________________
Lewis Structure:
Molecular Structure:
6. XeF4 Electron group geometry _____________________
Molecular geometry _________________________
Lewis Structure:
Molecular Structure:
7. CH2O Electron group geometry _____________________
Molecular geometry _________________________
Lewis Structure:
Molecular Structure:
8. IBr3
Electron group geometry _____________________
Molecular geometry _________________________
Lewis Structure:
Molecular Structure:
9. IBr
Electron group geometry _____________________
Molecular geometry _________________________
Lewis Structure:
Molecular Structure:
10. PF5
Electron group geometry _____________________
Molecular geometry _________________________
Lewis Structure:
Molecular Structure:
Section 9: Molecular Models
Lewis dot diagram Structural formula
(dots only)
(lines and dots)
Geometry
(Electron Group
and Molecular)
1. CH4
EG:
M:
2. O2
EG:
M:
Molecular
Structure
(3D)
3. H2
EG:
M:
4. NH3
EG:
M:
5. H2O
EG:
M:
6. SiCl4
EG:
M:
7. CO2
EG:
M:
8. PH3
EG:
M:
9. C2H6
EG:
M:
10. C2H4
EG:
M:
11. C2H2
EG:
M:
Section 10: Intermolecular Forces
1. What is the difference between a non-polar and a polar molecule?
2. Which of the following compounds are polar?
a. CO2
b. CH4
c. O2
d. H2O
e. NH3
3. What are the 2 main types of IMF’s?
4. What is a dispersion force?
5. What two factors affect the amount of dispersion forces in a non-polar
molecule? Describe the relationship between each factor and the
amount of force.
6. How are the three types of dipole-dipole interactions different from each
other?
Section 11: Substance Types
1. Explain why network covalent compounds are very hard and non-polar
covalent compounds are very soft.
2. You are given 4 compounds: diamond, methane (CH4), butane (C4H10),
and magnesium chloride. Rank them in the order of increasing
melting/boiling points. Justify your answer.
3. Why does surface tension increase with stronger IMF’s? What affect does
this have on the boiling point of the substance?
4. What is viscosity? What is the relationship between viscosity and IMF’s?