Download PIB and HH - Unit 4 - Chemical Names and Formulas

1
Instructor Lesson Plan – (PIB Chemistry)
Unit 4: Chemical Names and Formulas
and Bonding
Instructor:
D. Ditkowsky (Thanks J. Galinski)
Introductory Resources:
Addison-Wesley v.5 - Chapter 6, 14, 15, 16
Addison-Wesley v.4 - Chapter 5, 12, 14, 15
Advanced Resources:
Brown v.4 - Chapter 5, 7, 8, 9
Zumdahl v.3 - Chapter 2, 7, 8, 9
Zumdahl v.5 - Chapter 2, 7, 8, 9
Main Idea Summary:
 Every substance is either an element or a compound.
 A compound consists of more than one kind of atom.
 A compound is either molecular or ionic in nature.
 Molecular compounds are composed of two or more nonmetals.
 A molecular formula shows the number and kinds of atoms present in a molecule of a
compound.
 Ionic compounds are composed of oppositely charged ions combined in electrically
neutral groupings.
 A formula unit gives the lowest whole-number ratio of ions in the compound.
 The charges of the ions of the representative elements can be determined by the
position of these elements in the periodic table.
 Most transition metals have more than one common ionic charge.
 A polyatomic ion is a group of atoms that behaves as a unit and has a charge.
 When a cation can have more than one ionic charge, a Roman numeral is used in the
name.
 There are simple rules for naming ionic and molecular compounds.
 Elements that have similar properties also have similar electron configurations and
are members of the same group.
 Regular changes in the electron configuration of the elements cause gradual changes
in both the physical and chemical properties within a group and within a period.
 Atomic radii decrease as you move from left to right in a given period.
 Ionization energy increases as you move from left to right in a given period.
 Atomic radii increase within a given group because the outer electrons are farther
from the nucleus as you go down the group.
 Ionization energy decreases as you move down through a group.
2
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Atoms in compounds are held together by chemical bonds. Chemical bonds result
from the sharing or transfer of valence electrons between pairs of atoms.
Bonded atoms attain the stable electron configuration of a noble gas. The noble gases
themselves exist as isolated atoms because that is their most stable condition.
For the representative elements, the number of valence electrons is equal to the
element’s group number in the periodic table.
The transfer of one or more valence electrons between atoms produces positively and
negatively charged ions, or cations and anions respectively.
The attraction between an anion and a cation is an ionic bond. A substance with ionic
bonds is an ionic compound. Nearly all ionic compounds are crystalline solids at
room temperature. They have high melting points. The total positive charge cancels
out the total negative charge, yielding a neutral compound.
When melted or in solution, ionic compounds can conduct electricity, because the
ions can move freely when a voltage is applied.
Metals bond by packing positive metal ions together, surrounded by a sea of
electrons. The sea of electrons in metallic bonding causes most of the properties of
metals (malleability, ductility, electrical conductivity).
Atoms form covalent bonds when they share electrons to form an octet.
A shared pair of valence electrons constitutes a single covalent bond. Additional
sharing of electron pairs constitute the formation of higher order bonds.
Resonance structures help to visualize the bonding in molecules when more than one
electron dot formula can be written.
Molecular orbital theory is a logical extension of the quantum mechanical description
of the electron structure of atoms. The Valence Shell Electron Pair Repulsion Theory
explains expected molecular geometry of a molecule.
3
4
Common Ion List
CATIONS:
Aluminum
Ammonium
Barium
Beryllium
Calcium
Cesium
Copper (I)
Copper (II)
Francium
Hydrogen
Iron (II)
Iron (III)
Lead (II)
Lead (III)
Lead (IV)
Lithium
Magnesium
Potassium
Radium
Rubidium
Silver
Sodium
Strontium
Zinc
ANIONS:
Al+3
NH4+1
Ba+2
Be+2
Ca+2
Cs+1
Cu+1
Cu+2
Fr+1
H+1
Fe+2
Fe+3
Pb+2
Pb+3
Pb+4
Li+1
Mg+2
K+1
Ra+2
Rb+1
Ag+1
Na+1
Sr+2
Zn+2
Acetate
Bromide
Carbonate
Chlorate
Chloride
Chlorite
Chromate
Cyanide
Dichromate
Fluoride
Hydride
Hydroxide
Hypochlorite
Iodate
Iodide
Nitrate
Nitrite
Oxalate
Oxide
Perchlorate
Permanganate
Phosphate
Silicate
Sulfate
Sulfide
Sulfite
Thiocyanate
Thiosulfate
C2H3O2-1
Br-1
CO3-2
ClO3-1
Cl-1
ClO2-1
CrO4-2
CN-1
Cr2O7-2
F-1
H-1
OH-1
ClO-1
IO3-1
I-1
NO3-1
NO2-1
C2O4-2
O-2
ClO4-1
MnO4-1
PO43SiO3-2
SO4-2
S-2
SO3-2
SCN-1
S2O3-2
5
Laboratory Activity 5A – Density is a Periodic Property
NAME: _____________________
The Periodic Table was developed by Dmitri Mendeleev in the mid-1800s. Mendeleev
put all of the elements in order based on mass, and noticed a periodic reoccurance of
chemical and physical properties. He arranged the elements in columns. Elements in
each column have similar properties. Occasionally, he would find a hole in the table. He
could use the surrounding information to predict the properties of a yet-to-be-discovered
element. As new elements were discovered, they neatly filled holes in the periodic table.
You will be given samples of Silicon, Tin, and Lead. All of these elements are in the
same column of the Periodic Table, and have similar properties. Using a balance, you
will find the mass of a sample of each element. Using a graduated cylinder and some
water, you will find the volume of each sample. Then, you can determine the density of
each element using calculations you have already mastered.
Your goal for this lab is to determine the density of Germanium, element #32, based on
what you learn about the other elements in the same column. Watch for patterns. Use
the data below.
Symbol
Element
Si
Silicon
Ge
Germanium
Sn
Tin
Pb
Lead
Qualitative
Observations
Mass
(g)
Volume
(cm3)
XXXXXXX
XXXXXXX
XXXXXXX
XXXXXX
XXXXXX
XXXXXX
XXXXXXX
XXXXXXX
XXXXXXX
Density
(g/cm3)
6
Periodic Trends
NAME: ____________________
1. Restate in one or two words: “The amount of energy required to remove one electron
from the valence shell of a neutral atom.”
2. Restate in one or two words: “The tendency of an atom to hold on to its valence
electrons while engaged in a chemical bond.”
3. Restate in one or two words: “The actions of the non-valence electrons, diluting the force
of the attraction between nucleus and valence electrons.”
4. Which has greater shielding, Au or Cu?
5. Which is larger, Au or Cu?
6. Which has greater ionization energy, Cu or Ag?
7. Which has greater shielding, Xe or Ar?
8. Which is larger, Ca or Cs?
9. Which has greater shielding, Se or Ra?
10. Which has greater nuclear charge, Zn or Se?
11. Which is larger, Mg or P?
12. Which has greater ionization energy, Fe or K?
13. Restate in one or two words: “The actions of the non-valence electrons, diluting the
force of the attraction between nucleus and valence electrons.”
14. Restate in one or two words: “Half the distance between the nuclei of two like atoms.”
15. Restate in one or two words: “The tendency of an atom to attract electrons when it is
chemically combined with another element.”
16. Which has greater ionization energy, Cl or Hf?
17. Which has greater shielding, P or Ar?
18. Which is larger, Os or Ta?
19. Which is a larger ion, sulfur ion or phosphorus ion?
20. Which has greater shielding, Ge or Ra?
21. Which has greater nuclear charge, Sb or Se?
7
Laboratory Activity 5B – Periodic Trends
Page One
NAME: ____________________
Part I: Periodic Trend Graphing Activity
Atomic
Number
1
2
3
4
5
6
7
8
9
10
11
12
13
14
15
16
17
18
19
20
21
22
23
24
25
26
27
Element
H
He
Li
Be
B
C
N
O
F
Ne
Na
Mg
Al
Si
P
S
Cl
Ar
K
Ca
Sc
Ti
V
Cr
Mn
Fe
Co
Ionization
Potential
(eV)
13.60
24.59
5.39
9.32
8.30
11.26
14.53
13.62
17.42
21.56
5.14
7.65
5.99
8.15
10.49
10.36
12.97
15.76
4.34
6.11
6.54
6.82
6.74
6.77
7.44
7.87
7.86
Atomic
Radius
(Å)
0.30
0.93
1.52
0.89
0.88
0.77
0.70
0.66
0.64
1.12
1.86
1.60
1.43
1.17
1.10
1.04
0.99
1.54
2.31
1.97
1.60
1.46
1.31
1.25
1.29
1.26
1.25
Atomic
Number
28
29
30
31
32
33
34
35
36
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
Element
Ni
Cu
Zn
Ca
Ge
As
Se
Br
Kr
Rb
Sr
Y
Zr
Nb
Mo
Tc
Ru
Rh
Pd
Ag
Cd
In
Sn
Sb
Te
I
Xe
Ionization
Potential
(eV)
7.64
7.73
9.39
6.00
7.90
9.81
9.75
11.81
14.00
4.18
5.70
6.38
6.84
6.88
7.10
7.28
7.37
7.46
8.34
7.58
8.99
5.79
7.34
8.64
9.01
10.45
12.13
Atomic
Radius
(Å)
1.24
1.28
1.33
1.22
1.22
1.21
1.17
1.14
1.69
2.44
2.15
1.80
1.57
1.41
1.36
1.30
1.33
1.34
1.38
1.44
1.49
1.62
1.40
1.41
1.37
1.33
1.90
1. On a sheet of graph paper, graph the ionization potential (y-coordinate) versus atomic
number (x-coordinate). Make sure to properly label the graph.
2. On a sheet of graph paper, graph the atomic radius versus the atomic number. Label
the graph.
3. What do the units “eV” and “Å” stand for? What do these measurements mean, in
your own words.
8
Part III: Periodic Table Identification Activity
1. On the blank Periodic Chart on Page Three, clearly locate the following, using the
following color code:
a.
b.
c.
d.
e.
f.
g.
h.
i.
Representative elements : Outline in HEAVY BLACK
Transition elements : (Sky Blue)
Metallic elements: Color these in shades of (Blue)
Nonmetallic elements: Color these in the shades of (Red)
Metalloids: Purple (Red and Blue mixed)
Alkali metals (Dark Blue)
Alkaline-earth metals (Greenish Blue)
Halogens (Pink)
Noble gases (Yellow)
2. On the same periodic chart, locate these elements and write in their atomic symbols:
a.
b.
c.
d.
e.
sodium
potassium
chlorine
nickel
bromine
f. phosphorus
g. carbon
h. magnesium
i. sulfur
j. calcium
k. barium
l. aluminum
m. silicon
n. zinc
o. lead
9
Laboratory Activity 5B – Periodic Trends
Page Three
NAME: ____________________
10
Laboratory Activity 5B – Periodic Trends
Page Four
NAME: ____________________
Part IV: Interpretation Questions
1. Notice that the graph of first ionization potential versus atomic number consists of
generally rising values followed by sharp drops. List the elements on the five major
peaks in this graph. What name is given to this group of elements?
2. List four elements located at the bottom of the sharp drops. What name is given to
this group of elements?
3. Assuming that the periodic trends indicated on the graph continue, what value do you
predict for the first ionization potential of cesium, Cs, atomic number 55?
4. What generalization can be made about the change in first ionization potential as the
atomic number increases in a period (such as Na to Ar)?
5. What generalization can be made about the change in first ionization potential as the
atomic number increases in a group (family)?
6. Looking at the atomic radius versus atomic number, what would you predict for the
atomic radius of Cs, atomic number 55? (Use Cl-Ar-K and Br-Kr-Rb as examples.)
11
Electronegativity and types of bonds
Carefully read text pages 405 – 406, 419, 422 for details about ionic bonds and
electronegativity. The difference between the electronegativity of the anion minus the
cation determines if a bond is ionic.
If the difference in electronegativity is greater than 1.7, the bond is considered ionic
because one atom possesses the electron more than 50% of the time. (Anything less than
0.3 is considered nonpolar covalent)
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 Mg
0.9 1.2
Al
1.5
Si
1.8
P
2.1
S
2.5
Cl
3.0
K
0.8
Ca
1.0
Sc
1.3
Ti
1.5
V
1.6
Cr Mn Fe
1.6 1.5 1.8
Rb
0.8
Sr
1.0
Y
1.2
Zr
1.4
Nb Mo
1.6 1.8
Tc
1.9
Cs
0.7
Ba
0.9
La
1.0
Hf
1.3
Ta
1.5
Re
1.9
W
1.7
Co
1.9
Ni
1.9
Cu
1.9
Zn
1.6
Ga
1.6
Ge
1.8
As
2.0
Se
2.4
Br
2.8
Ru
2.2
Rh
2.2
Pd
2.2
Ag
1.9
Cd
1.7
In
1.7
Sn
1.8
Sb
1.9
Te
2.1
I
2.5
Os
2.2
Ir
2.2
Pt
2.2
Au
2.4
Hg
1.9
Tl
1.8
Pb
1.9
Bi
1.9
Po
2.0
At
2.2
Percent Ionic Character
0%
5%
non(p.c.) polar covalent
0
0.3
50%
ionic
1.7
100%
3.3
1. What must always be true if a covalent bond is to be polar?
2. The bonds between the following pairs of elements are covalent. Arrange them
according to polarity, naming the most polar bond first.
H—Cl,
H—C,
H—F,
H—O,
H—H,
S—Cl
3. Arrange the following bonds in order of increasing ionic character.
Cl—F
N—N
K—O
C—H
S—O
Li—F
12
More Bonding Concepts
NAME: ______________________
1. Write electron configurations for the following and comment on the result:
a. N3− ________________________
b. O2− ________________________
c. F1− ________________________
d. Ne ________________________
2. Explain why molten MgCl2 does conduct an electric current although crystalline
MgCl2 does not.
3. Write the correct chemical formula (the formula unit) for the compounds formed from
each pair of ions.
a. K+, S2−
____________
b. Ca2+, O2−
____________
c. Na+, SO42−
____________
d. Al3+, PO43−
____________
4. Write formulas for each compound:
a. potassium nitrate
___________________________
b. barium chloride
___________________________
c. magnesium sulfate
___________________________
d. lithium oxide
___________________________
e. ammonium carbonate ___________________________
13
Worksheet 5.4 – Lewis Structures of Atoms
NAME: ______________________
Draw the appropriate number of dots around the following atoms to create correct Lewis
Dot Structures:
H
He
Li
Be
B
C
N
O
F
Ne
Na
Mg
Al
Si
P
S
Cl
Ar
K
Ca
Ga
Ge
As
Se
Br
Kr
Rb
Sr
In
Sn
Sb
Te
I
Xe
Cs
Ba
Tl
Pb
Bi
Po
At
Rn
Write lewis structures for the following:
CH4
CCl4
H
H
C
H
CF4
PCl3
OF2
Cl
H
Cl
C
Cl
CI4
Cl
Cl
P
Cl
F
O
Cl
PBr3
SF2
F
14
Lewis Structures of Ions
NAME: ______________________
Draw the appropriate number of dots around the following ions to create correct Lewis
Dot Structures:
[H]+
Heo
[Li]+
[Be]++
[B]3+
[C]4+
[N]3-
[O]2-
[F]1-
Neo
[Na]+
[Mg]++
[Al]3+
[Si]4+
[P]3-
[S]2-
[Cl]1-
Aro
[K]+
[Ca]++
[Ga]3+
[Ge]4+
[As]5+
[Se]4+
[Br]1-
Kro
[Rb]+
[Sr]++
[In]3+
[Sn]4+
[Sb]3+
[Te]4+
[I]1-
Xeo
[Cs]+
[Ba]++
[Tl]3+
[Pb]2+
[Bi]4+
[Po]4+
[At]4+
Rno
Write the Dot Diagrams for the following:
LiCl
Li3N
Ca3P2
Use the example on page 439 to show what happens to the electrons in these bonds.
Show the electrons of the cations in one color and the electrons of the anion a different
color.
Li __ __
Li __ __
Ca __ __ __ __ __ __ __ __ __ __
1s 2s
1s 2s
1s 2s
2p
3s
3p
4s
Cl __ ___ __ __ __ Li __ __
Ca __ __ __ __ __ __ __ __ __ __
1s 2s
2p
1s 2s
1s 2s
2p
3s
3p
4s
Li __ __
Ca __ __ __ __ __ __ __ __ __ __
1s 2s
1s 2s
2p
3s
3p
4s
P __ __ __ __ __ __ __ __ __
N __ ___ __ __ __
1s 2s
2p
3s
3p
1s 2s
2p
P __ __ __ __ __ __ __ __ __
1s 2s
2p
3s
3p
15
Group 1: How do I Identify general categories of elements? What are they?
Group 2: How are electron configurations related to the position of an element on
the periodic table?
Group 3: How is the atomic size related to the position of an element on the
periodic table?
Group 4: How is the ionic size related to the position of an element on the
periodic table? How is it related to the electron configuration.
Group 5: How is ionization energy related to the position of an element on the
periodic table, and its electron configuration?
Group 6: How is electronegativity related to the position in the periodic table?
Group 7: Make a powerpoint with animation to summarize all of the previous
presentations.
16
Cations
Ammonium Ammonium Ammonium Ammonium Ammonium Ammonium Ammonium
NH4+1 +
NH4+1 +
NH4+1 +
NH4+1 +
NH4+1 +
NH4+1 +
NH4+1 +
Potassium
K+1 +
Potassium
K+1 +
Potassium
K+1 +
Potassium
K+1 +
Potassium
K+1 +
Potassium
K+1 +
Potassium
K+1 +
Sodium
Na+1 +
Silver
Ag+1 +
+
Calcium
Ca+2
+
+
Magnesium
Mg+2
+
+
Iron (II)
Fe+2
+
+
Sodium
Na+1 +
Silver
Ag+1 +
+
Calcium
Ca+2
+
+
Magnesium
Mg+2
+
+
Iron (II)
Fe+2
+
+
Sodium
Na+1 +
Silver
Ag+1 +
+
Calcium
Ca+2
+
+
Magnesium
Mg+2
+
+
Iron (II)
Fe+2
+
+
Sodium
Na+1 +
Silver
Ag+1 +
+
Calcium
Ca+2
+
+
Magnesium
Mg+2
+
+
Iron (II)
Fe+2
+
+
Sodium
Na+1 +
Silver
Ag+1 +
+
Calcium
Ca+2
+
+
Magnesium
Mg+2
+
+
Iron (II)
Fe+2
+
+
Sodium
Na+1 +
Silver
Ag+1 +
+
Calcium
Ca+2
+
+
Magnesium
Mg+2
+
+
Iron (II)
Fe+2
+
+
Sodium
Na+1 +
Silver
Ag+1 +
+
Calcium
Ca+2
+
+
Magnesium
Mg+2
+
+
Iron (II)
Fe+2
+
+
Aluminum
Al+3
+
Aluminum
Al+3
+
Aluminum
Al+3
+
Aluminum
Al+3
+
Aluminum
Al+3
+
Aluminum
Al+3
+
Aluminum
Al+3
+
+
+
+
+
+
+
+
Ammonium Ammonium Ammonium Ammonium Ammonium Ammonium Ammonium
NH4+1 +
NH4+1 +
NH4+1 +
NH4+1 +
NH4+1 +
NH4+1 +
NH4+1 +
Potassium
Potassium
Potassium
Potassium
Potassium
Potassium
Potassium
+1
+1
+1
+1
+1
+1
K
+
K
+
K
+
K
+
K
+
K
+
K+1
+
Cut out the cations, store these in a plastic bag.
17
Anions
Dichromate Cr2O7-2
Nitrate
NO3-1
Acetate
C2H3O2-1 -
Dichromate Cr2O7-2
Nitrate
NO3-1
Acetate
C2H3O2-1
-
Dichromate Cr2O7-2
Nitrate
NO3-1
Acetate
C2H3O2-1
-
Dichromate Cr2O7-2
Nitrate
NO3-1
Acetate
C2H3O2-1
-
Dichromate Cr2O7-2
Nitrate
NO3-1
Acetate
C2H3O2-1
-
Dichromate Cr2O7-2
Nitrate
NO3-1
Acetate
C2H3O2-1
-
Dichromate Cr2O7-2
Nitrate
NO3-1
Acetate
C2H3O2-1 -
Dichromate Cr2O7-2
-
Dichromate Cr2O7-2
-
Dichromate Cr2O7-2
-
Dichromate Cr2O7-2
-
Dichromate Cr2O7-2
-
Dichromate Cr2O7-2
-
Dichromate Cr2O7-2
-
Sulfate
SO4-2
Sulfate
SO4-2
Sulfate
SO4-2
Sulfate
SO4-2
Sulfate
SO4-2
Sulfate
SO4-2
Sulfate
SO4-2
-
-
-
-
-
-
-
-
-
-
-
-
-
-
Carbonate CO3-2
-
Carbonate CO3-2
-
Carbonate CO3-2
-
Carbonate CO3-2
-
Carbonate CO3-2
-
Carbonate CO3-2
-
Carbonate CO3-2
-
Phosphate PO4-3
Phosphate PO4-3
Phosphate PO4-3
Phosphate PO4-3
Phosphate PO4-3
Phosphate PO4-3
Phosphate PO4-3
Bromide
Br-1
Chloride
Cl-1
-
Bromide
Br-1
Chloride
Cl-1
-
Bromide
Br-1
Chloride
Cl-1
-
Bromide
Br-1
Chloride
Cl-1
-
Bromide
Br-1
Chloride
Cl-1
-
Cut out the anions, store in a plastic bag.
Bromide
Br-1
Chloride
Cl-1
-
Bromide
Br-1
Chloride
Cl-1
-
18
Laboratory Activity 5C – Making Models of Compounds
Chemical Formulas
Match the appropriate cation to the appropriate anion so that the number of + charges = the number of anions.
Chloride
Bromide
Acetate
Sulfate
Carbonate
Nitrate
Cl-1
Br-1
C2H3O2-1
SO4-2
CO3-2
NO3-1
Ammonium ex:
NH4+1
NH4+
Cl↓
NH4Cl
Potassium
K+1
Sodium
Na+1
Silver
Ag+1
Calcium
Ca+2
Magnesium
Mg+2
Iron (II)
Fe+2
Aluminum
Al+3
Name ____________________
Phosphate
PO4-3
Dichromate
Cr2O7-2
19
Laboratory Activity 5C – Making Models of Compounds
Ammonium
NH4+1
Chloride
Cl-1
example:
ammonium
chloride
Potassium
K+1
Sodium
Na+1
Silver
Ag+1
Calcium
Ca+2
Magnesium
Mg+2
Iron (II)
Fe+2
Aluminum
Al+3
Bromide
Br-1
Acetate
C2H3O2-1
Names
Sulfate
SO4-2
Carbonate
CO3-2
Name ____________________
Nitrate
NO3-1
Phosphate
PO4-3
Dichromate
Cr2O7-2
20
Electron Dot Structures and VSEPR Shapes
NAME: ____________________
1. Classify the following compounds as ionic or covalent. Underneath the covalent
compounds, draw a plausible electron dot structure for the molecule.
a. MgCl2
b. OCl2
c. Na2S
d. SF2
2. Describe the difference between an ionic and a covalent bond.
3. Draw the electron dot structure of each of the following molecules.
a. F2
b. HCl
c. HCCH
d. NI3
4. Draw resonance structures for the nitrite ion, NO2−1. The oxygen atoms are attached
to the nitrogen.
5. Explain the difference between “paramagnetic” and “diamagnetic.”
6. Predict whether the following species are diamagnetic or paramagnetic.
a. BF3
b. O2−1
c. NO2
d. F2
21
More Work With VSEPR, Sigma and Pi Bonds
NAME: ____________________
1. Use VSEPR theory to predict the shapes of the following species. Draw the electron
dot structures at the right.
b. CO2
c. SiCl4
d. SO3
e. SCl2
f. CO
g. I3−1
2. Which of the following species would you predict to be diamagnetic? Paramagnetic?
a. NO3−1
b. OH−1
c. H2O
d. SO3
3. Suggest a reason why phosphorus and sulfur have more than an octet in many of their
compounds. Explain why compounds of nitrogen and oxygen never do.
22
Laboratory Activity 5D – Candy Dot Molecule Building
NAME: ______________________
Students will be required to use marshmallows (atoms) and toothpicks (bonds) to create
molecule models. The table provided below should be expanded to include all of the
molecules used in this activity. Students will complete a table for each of the following:
CCl4
PH4+1
AsCl3
CO2
PCl5
SiO2
GeCl4
SCl2
3-Dimensional
Drawing
BeI2
BF3
VSEPR
Shape
Bonds
Lewis Dot
Structure
H2S
SCl6
SeCl6
 Bonds
Molecular
Formula
NF3
SO3
BBr3
23
24
Polarity of Molecules
NAME: _____________________
1. Not every molecule with polar bonds is polar. Explain this statement, using CCl4 as
an example.
2. Draw the electron dot structure for each molecule. Identify polar covalent bonds by
assigning slightly positive (δ+) and slightly negative (δ−) symbols to the appropriate
atoms.
a. HOOH
c. HBr
b. BrCl
d. H2O
3. Based on the information you have amassed about molecular shapes, which of these
molecules would you expect to be polar?
a. SO2
c. CO2
b. H2S
d. BF3
4. Use VSEPR theory to predict the geometry of the following. Draw electron dot
structures of each compound at the right.
a. SiCl4
b. CO32−
c. CCl4
d. SCl2
25
Intermolecular Forces
NAME: ___________________
1. How does a network solid differ from most other covalent compounds?
2. Which of the following are characteristic of most covalent compounds?
a.
b.
c.
d.
e.
high melting point
shared bonding electrons
low water solubility
existence as molecules
composed of a metal and a nonmetal
3. Depict the hydrogen bonding between two ammonia molecules and between one
ammonia molecule and one water molecule.
4. Which compound in each pair exhibits the stronger intermolecular hydrogen
bonding?
a. H2S, H2O
c. HBr, HCl
b. HCl, HF
d. NH3, H2O
5. What is a hydrogen bond?
6. Why do compounds with strong intermolecular attractive forces have higher boiling
points than compounds with weak intermolecular attractive forces?
26
Practice Test – Page One
Naming, Formula Writing, Trends
NAME ______________________
1. Which is larger, an atom of Ni or an atom of Ag?
2. Which is smaller, an atom of Si or an atom of Ga?
3. Which has a higher electronegativity, Ti or Cr?
4. Which has more shielding, K or Zr?
_____
_____
_____
_____
5. Which is larger, an atom of Magnesium or a Magnesium ion?
6. Which is smaller, an ion of Sulfur or an ion of Sodium?
_______________
7. Which has greater nuclear charge, Sodium or Magnesium?
8. Which has a lower ionization energy, Ba or Rb?
9. Which holds electrons more tightly, N or B?
_______________
_______________
_____
_____
10. Which neutral atom has higher electronegativity, neutral atom “A,” which has an
electron configuration of 1s22s22p63s23p64s23d104p3 or atom “X,” which has an
electron configuration of 1s22s22p63s23p64s1? _____
11. Define electronegativity.
12. Compound: Fe2(CO3)3
Type of bond (Ionic or Covalent): ________________
Type of structure (Molecule or Lattice): ________________
Name of compound: ________________
13. Compound: N2O7
Type of bond (Ionic or Covalent): ________________
Type of structure (Lattice or Molecule): ________________
Name of compound: ________________
27
Practice Test – Page Two
Naming, Formula Writing, Trends
NAME ______________________
14. Compound: Manganese (II) Sulfate
Type of bond (Covalent or Ionic): ________________
Type of structure (Molecule or Lattice): ________________
Formula of compound: ________________
15. Compound: Oxygen Monochloride
Type of bond (Ionic or Covalent): ________________
Type of structure (Molecule or Lattice): ________________
Formula of compound: ________________
16. Compound: Al2S3
Type of bond (Ionic or Covalent): ________________
Action of electrons (Sharing or Transfer): ________________
Name of compound: ________________
17. Compound: Tricarbon Heptachloride
Type of bond (Ionic or Covalent): ________________
Action of electrons (Sharing or Transfer): ________________
Formula of compound: ________________
18. Define shielding.
19. At the right, put the following atoms or ions in order, from smallest to largest:
K, K+1, Ca, Mg+2, P, P-3, Cl, O
___ ___ ___ ___ ___ ___ ___ ___
28
PRACTICE TEST: Page Three
NAME: ____________________
I:
DRAW the Lewis Dot Diagrams for the following.
IDENTIFY the shape of the molecule just under each drawing.
(Phosphorus CAN but DOESN’T HAVE TO always violate the octet rule).
SCl2
OF2
AsCl3
CF4
SiO2
PI5
II:
Which of the above compounds are polar molecules?
III:
HCN, has a few bonds. Draw the structure of HCN and identify the number of
Sigma and Pi bonds it has.
Sigma:
Pi:
IV:
Draw the Lewis Dot structure for each of the resonance conformers of ozone, O3.
29
PRACTICE TEST Page Four
NAME: ____________________
V:
Draw Lewis structures for the following atoms and ions under their names.
Sulfur
Sulfide ion
Sodium
Sodium ion
VI:
Draw a Lewis structure of the compound formed when Sodium ionically bonds
with Sulfur.
VII:
Explain the differences between iron and steel.
VIII:
What will form stronger bonds, NH3 or NF3?
IX:
What is stronger, Ionic Bond or Hydrogen Bond?
X:
What is stronger, Dipole-Dipole Interaction or Covalent Bonds?
XI:
What type of intermolecular bonds will hold molecules of PCl3 together?
XII:
What type of intermolecular bonds will hold molecules of CO2 together?
XIII:
Which covalent bond is most polar? H-H, H-N, H-F, or H-O?
30