Name______________________________
Period:________
Unit 4 Assignment Packet
A1: Ionic Bonding (Goals 1 – 5)
1. What are valence electrons?
2. How do valence electrons largely determine the chemical properties of an element?
3. Is the following sentence true or false? The group number of an element in the periodic table is related to the
number of valence electrons it has.
4. What is an electron dot structure?
5. Draw the electron dot structure of each of the following atoms
a. Argon
b. Calcium
c. Iodine
6. What is the octet rule?
7. Metallic atoms tend to lose their valence electrons to produce a(n) _______________ or a positively charged
ion. Most nonmetallic atoms achieve a complete octet by _______________ electrons.
8. Write the electron configuration for theses metals and circle the electrons lost when each metal forms a cation.
a. Mg
b. Al
c. K
9. Atoms of most nonmetallic elements achieve noble-gas electron configurations by gaining electrons to become
_______________, or negatively charged ions.
10. What property of nonmetallic elements makes them more likely to gain electrons than lose electrons?
11. Is the following sentence true or false? Elements of the halogen family lose one electron to become halide ions.
12. How many electrons will each element gain in forming an ion?
a. Nitrogen
c. Sulfur
b. Oxygen
d. Bromine
13. Write the symbol and electron configuration for each ion from question 12 and name the noble gas with the
same configuration.
a. Nitride
c. Sulfide
b. Oxide
d. Bromide
14. What is an ionic bond?
15. In an ionic compound, the charges of the _______________and _______________ must balance to produce an
electrically _______________ substance.
16. Why do beryllium and fluorine combine in a 1:2 ratio?
Complete the chart for each element.
Element
# Protons
# Electrons
# Valence
Electrons
Ion Charge
Sodium
Chlorine
Beryllium
Fluorine
Lithium
Oxygen
Phosphorus
For the each ionic bond do the following:
17. A. Use electron dot structures to show the transfer of electrons from one element to the other.
B. Write the electron configuration of each element before and after bonding.
Sodium
A.
B.
+
Fluorine

Sodium Fluoride
Magnesium
+
Phosphorus
Aluminum
+
Oxygen
Lithium +
Nitrogen

Magnesium Phosphide
A.
B.

Aluminum Oxide

Lithium Nitride
A.
B.
A.
B.
A2: Covalent Bonding (Goals 5-6)
Answer the following questions. Use complete sentences or pictures if necessary.
1. What is a covalent bond?
2. Is the following sentence true or false? A shared pair of electrons is represented by a double dash.
3. Structural formulas show the arrangement of _______________ in molecules and polyatomic ions.
4. What is the octet rule of covalent bonding?
5. Is the following sentence true or false? All diatomic molecules contain double bonds.
6. Complete the electron dot or structural formula for each molecule. Each molecule contains only single covalent
bonds.
H
Br
F
H
O
O
H
H
C
H
H
7. How many electrons are shared in a double bond?
8. Draw an electron dot or structural formula for two molecules that contain a double bond.
9. How many electrons are shared in a triple bond?
10. Draw an electron dot or structural formula for a molecule containing a triple bond.
Draw an electron dot or structural formula for the following Molecules and Polyatomic Ions.
1. SiF4
2. PCl3
3. N2O
4. OF2
5. SeO3
6. AsBr3
7. SO2
8. NH4+
9. NF2-
10. NO2-
11. NO3-
12. SO32-
13. SO42-
14. PO43-
15. OCN-
A3 VSEPR Geometry (Goals 7-9)
1. What is VSEPR theory?
2. When the central atom of a molecule has unshared electrons, the bond angles will be _______________ than
when all the central atom’s electrons are shared.
Molecular Formula
Structural Formula
Molecular Geometry
H
CH4
H
C
H
CO2
I2
H
tetrahedral
Sketch
Molecular formula
CCl4
NH3
H2O
SCl2
BF3
SO2
NO2-
SO32-
Structural Formula
Molecular Geometry
Sketch
A4: Bond Types and Intermolecular Forces (Goals 7-9)
18. Is the following statement true or false? Covalent bonds differ in the way electrons are shared by the bonded
atoms, depending on the kind and number of atoms joined together.
19. Describe how electrons are shared in each type of bond. Write equally or unequally
a. Nonpolar bond
b. Polar bond
20. Explain how you can use electronegativity values to classify a bond as non-polar, polar covalent, or ionic.
21. Why does the chlorine atom in hydrogen chloride acquire a slightly negative charge?
22. Circle the letter of each sentence that is true about polar molecules
a. Some regions of a polar molecule are slightly negative and some are slightly positive.
b. A molecule containing a polar bond is always polar.
c. A molecule that has two poles is called a dipolar molecule.
d. When polar molecules are placed in an electric field, they all line up with the same orientation in
relation to the charged plates.
23. Are the following molecules polar or nonpolar (hint: draw them)?
a. H₂O
b. CO₂
c. NH₃
d. CCl4
24. Describe the three types of weak intermolecular attractive forces that hold groups of molecules together. Rank
these forces form weakest to strongest.
25. Not every molecule with polar bonds is polar. Explain this statement using CCl4 as an example.
26. Depict the hydrogen bonding between two ammonia molecules and between one ammonia molecule and one
water molecule.
27. Circle which compound in each pair exhibits the stronger intermolecular hydrogen bonding?
a. H2S, H2O
b. HCl, HF
c. HBr, HCl
d. NH3, H2O
28. Why do compounds with stronger intermolecular attractive forces have higher boiling points than compounds
with weak intermolecular attractive forces?
29. What must always be true if a covalent bond is to be polar?
30. Explain the trend in the boiling points for the halogens using dispersion forces and polarizability.
A5: Chemical Names and Formulas (Goals 10-12)
1. What does a chemical formula represent?
2. There are two types of substances that represent the smallest unit of that substance.
i.
ii.
3. Define molecule –
4. Molecules are electrically ____________________.
5. What type of elements are typically bonded together to form a molecular compound?
6. Produce a list of physical properties that molecules and molecular compounds share.
7. There are also seven common elements that exist in nature as two atoms bonded together. They make a seven
on the periodic table, they are:
8. Define Ion –
a. Cations are _______________ because they have ________________ electrons. _____________
form these positive ions.
b. Anions are _______________ because they have ________________ electrons.
___________________ form these negative ions.
9. How do we name positive ions?
10. How do we name negative ions?
11. Create a list of physical properties that Ionic Compounds share.
12. ____________________ Formulas show the type and number of atoms in a molecule.
13. A _________________ written after the element symbol indicates the number of atoms of each element in the
molecular compound.
14. How many atoms of each element are in CH4?
15. How many atoms of each element are in 4H2O?
16. What information can be gathered from the molecular formula CO2?
17. Chemical formulas can also be written for ionic compounds. In this case, however, the formula does
not represent a molecule. There are no separate compounds, only a continuous array of __________.
18. Formula units are always made of _______________ and ______________________.
19. To represent ionic compounds, chemists use a ___________________________, which is the
____________ whole-number ratio of the _________ in the compound.
20. For Fe3P2:
a. What is the cation in the ionic compound?
b. What is the anion in the ionic compound?
c. What is the ratio of cations to anions in the ionic compound?
21. How many of each ion are present in AlCl3?
22. ____________________ions: Ions made of single atoms.
A6: Ionic Compounds (Goals 13-15)
1. Representative Elements: There is a pattern in predicting how many electrons are lost and gained for the
representative elements, can you guess it?
WRITE ON YOUR PERIODIC TABLES THE CHARGES OF THE REPRESENTATIVE ELEMENTS NOW.
2. The Transition metals have much more complicated patterns of valence electrons. There are two methods of
naming such cations. The preferred method is called the stock system. As part of this system, a roman numeral
in parentheses indicates the charge value of the cation.
Examples: Name the following transition metal ions:
a. tin (lost 2 electrons):
Tin (II)
b. tin (lost 4 electrons)
c. iron (lost 3 electrons)
d. iron (lost 2 electrons)
There are 3 exceptions to this rule:
1. DO NOT USE A ROMAN NUMERAL WHEN NAMING SILVER, ZINC AND CADMIUM IONS.
2. ALWAYS USE A ROMAN NUMERAL WHEN NAMING ANY OTHER TRANSITION METAL ION.
3. DO NOT USE A ROMAN NUMBERAL WHEN NAMING A REPRESENTATIVE ELEMENTS ION.
Write the symbol and charge of the following
elements.
a. sulfur
S2b. lead (4 electrons lost)
c. strontium
d. bromine
e. copper (1 electron lost)
f.
selenium
g. silver
h. cesium
i.
phosphorus
Ag+
Name the ion
Cation or Anion?
Sulfide
anion
3. Polyatomic Ions: Ions made of _________________________________.
4. What endings do most polyatomic ions receive when naming them?
5. There are 3 important exceptions, they are:
6. What are the “rules” for writing Binary Ionic Compounds?
a. Write the __________ ion first
b. Write the ______________ ion last
c. The net charge for the compound must add to _____ (positives + negatives = 0)
d. Use _____________ to indicate how many of each ion you need to “balance” the charge.
7. Write the formula for the ionic compound formed between potassium and chlorine.
8. Write the formula for the ionic compound formed between calcium and bromine.
Another approach to writing a balanced formula for a compound is to use the crisscross method. In this method, the
numerical charge of each ion is crossed over and used as the subscript for the other ion. The signs of the numbers are
dropped.
9. Use the crisscross method to write the formula for the ionic compound formed between iron (III) and oxygen.
10. Use the crisscross method to write the formula for the ionic compound formed between calcium and sulfur.
11. Write the formulas for the compounds formed between these pairs of ions.
a. Ba+2, S-2
c. Ca+2, N-3
b. Li+1, O-2
d. Cu+2, I-1
12. Write formulas for these compounds.
a. sodium iodide
b. potassium sulfide
c. tin (II) chloride (also called stannous chloride)
d. calcium iodide
13. Define ternary ionic compounds –
Remember, they are still just two ions, and all rules from before still apply!
14. Write the formula for lithium nitrate, a ternary compound:
15. Sometimes, we need to take more than one polyatomic ion to balance the charge to 0. If this happens, place
the polyatomic ion in parenthesis and the subscript outside of the parentheses.
a. Write the formula for lithium carbonate:
b. Write the formula for potassium sulfate:
c. Write the formula for ammonium phosphate:
16. Write the name & formulas for ionic compounds formed from these pairs of ions:
a. NH4+1, SO32b. Calcium ion, phosphate ion
c. Al 3+, NO3 -1
d. Potassium ion, chromate ion
17. Write formulas for these compounds
a. lithium hydrogen sulfate
b. chromium (III) nitrite
c. copper (II) bromide
d. ammonium dichromate
18. Name these compounds:
a. LiCN
e. KClO
b. (NH4)2CO3
f.
KMnO4
c. Fe(ClO3)3
g. Li2SO3
d. CaSO4
A7: Molecular Compounds (Goal 16)
23. What is a binary molecular compound?
24. We use prefixes when naming binary molecular compounds. Fill in the following:
Prefix
Number
1
2
3
4
5
Prefix
Number
6
7
8
9
10
25. Say the name of the first element, say the name of the second element, ending in –IDE, and put the appropriate prefix in to
indicate how many of each element there are in the formula:
If the prefix for the first element in a binary molecular compound is ________, it may be dropped. However, it must be said
if it’s for the second element. Don’t reduce the subscripts (like you did for binary ionic compounds)
26. Name these binary molecular compounds:
a. N2O
d. OF2
b. PCl3
e. Cl2O8
c. SF6
f.
SO3
27. Write formulas for the following binary molecular compounds:
a. nitrogen trifluoride
d. octoxygen dichloride
b. disulfur dichloride
e. trinitrogen pentoxide
c. dinitrogen tetroxide
Try these…
Name
Formula
1)
CCl4
11)
dihydrogen monoxide
2)
CO
12)
sulfur trifluoride
3)
P2O5
13)
iodine monochloride
4)
SF6
14)
nitrogen tribromide
5)
N2O3
15)
phosphorus pentachloride
6)
SO3
16)
xenon difluoride
7)
OCl2
17)
dichlorine octoxide
8)
CO2
18)
dinitrogen monoxide
9)
CS2
19)
trisilicon tetranitride
10)
SeCl2
20)
boron trichloride