Chem I: TYPES OF CHEMICAL BONDS #1
Name____________________
Period________
PART 1: Three Types of Bonds
1. In nature, most atoms are joined to other atoms by chemical bonds. The periodic table can help you
determine what type of bonds elements will make.
a. Bond between Metal and Nonmetal = ___________________________________________
b. Bond between Nonmetal and Nonmetal = ________________________________________
c. Bond between Metal and Metal = ______________________________________________
Using your periodic table, classify the following compounds as ionic, covalent or metallic:
a. CaCl2
g. CaF2
b. Cu (copper)
h. HCl
c. CO2
i. Cu-Zn (brass)
____________
d. H2O
j. Fe (iron)
____________
e. MgO
k. PCl2
____________
f. C6H12O6 (glucose)
l. NH4Cl
____________
PART 2: Electronegativity
2. Atoms form bonds to ___________________________________________
_____________________________________________________________
3. We can figure out for sure what type of bond will form between atoms by looking
at electronegativities.
a. An ionic bond will form between atoms who’s electronegativities differ by
(more/less than) ___________.
b. A covalent bond will form between atoms who’s electronegativities differ
by (more/less than) ___________.
Which type of bond will form between the following atoms?
Atom 1 –
Electronegativity
Fe =
Atom 2 –
electronegativity
O=
Fr =
Br =
Au =
As =
Li =
Cl =
Na =
Pd =
Mg =
N=
Electronegativity
Difference
Type of Bond
PART 3: Lewis Dot Diagrams
a) Write the shorthand electron configuration for the element.
b) Determine how many valence electrons. (circle the valence electrons)
c) Draw a Lewis Dot Diagram.
d) Based on the atoms valence electrons, write the charge that atom would have (it’s oxidation
number), when it gains/loses electrons to become stable.
EXAMPLE:
Valence Electrons: ___2___
[Ne]3s
2
..
Mg
Oxidation number: ____Mg2+___
1.
Valence Electrons: _______
Na
Oxidation Number: __________
2.
Valence Electrons: _______
F
Oxidation Number: __________
3.
Valence Electrons: _______
K
Oxidation Number: __________
4.
Valence Electrons: _______
S
Oxidation Number: __________
5.
Valence Electrons: _______
Al
Oxidation Number: __________
6.
Valence Electrons: _______
O
Oxidation Number: _________
Chem I: IONIC BONDING #2
Ionic Bonding occurs when a metal transfers one or more electrons to a nonmetal in an effort to attain a
stable octet of electrons. For example, the transfer of an electron from sodium to chlorine can be depicted
by a Lewis dot diagram.
Na +
Cl
→
Na+Cl-
Calcium would need two chlorine atoms to get rid of its two valence electrons.
Cl + Ca + Cl → Ca+2Cl −𝟐
Part I: Show the transfer of electrons in the following combinations.
1. K + F
2. Mg + I
3. Be + S
4. Na + O
5. Al + Br
Part II
Draw the Lewis structures for each atom, draw arrows to show the transfer of electrons, and then
write the chemical formula (include the charge for each ion). Remember that more than one
element may be needed in order for the ionic compound to be stable.
1. Mg + O
Mg + O
8. Al + Cl
Mg+2O-2
2. Na + S
9. Li + I
3. Mg + F
10. K + S
4. K + O
11. Ca + S
5. K + Cl
12. Ca + Br
6. Ba + F
13. Ca + N
7. Cs + N
14. Rb + F
Chemistry I: COVALENT BONDING #3
Drawing the structure formulas for the MOLECULES below:
1. Count the total number of valence e-.
2. Determine the central atom. The following are guides:
 Often the unique atom (only one of it) is the central atom.
 Or put the least electronegative element in the middle.
3. Arrange the other atoms around the central atom creating a skeleton.
4. Connect all bonded atoms in the skeleton with one bond.
5. Subtract the number of electrons already used for the single bonds; two for each bond.
6. Distribute the remaining electrons in pairs around the atoms, trying to satisfy the octet rule. Assign them to the most
electronegative atom first.
7. If you run out of electrons before all atoms have an octet of electrons, you need to form double or triple bonds.
8. If you have extra electrons and all of the atoms have an octet, put the extra electrons on the central atom in pairs.
1. Br2
Br Br
3. SiH4
Total # of valence electrons _ 14___
14
-2
12
-12
0
Total # of valence electrons_______
2. CH3Br
Total # of valence
electrons_______
4. H3P
Total # of valence
electrons_______
5. O2
Total # of valence electrons_______
6. CO
Total # of valence
electrons_______
7. N2
Total # of valence
electrons_______
9. SiO2
8. SeO3
Total # of valence
electrons_______
Total # of valence
10. SCl4
Total # of valence electrons_______
12. TeF6
Total # of valence electrons_______
electrons_______
11. ICl5
Total # of valence electrons_______
13. Only if we get to polyatomic ions
NO3 - 1
Total # of valence electrons_______
14. Only if we get to polyatomic ions
PO4-3
Total # of valence electrons_______
Chem I: SPECIAL COVALENT BONDS #4
Diatomic Molecules & Polyatomic Molecules
PART I:
Using your notes, write the definition of the following, include a statement on what force holds the
atoms together (sharing electrons or transferring electrons):
 Ionic Bond:

Covalent Bond:
Directions: Classify the following compounds as
 (I) ionic (metal + nonmetal)
 (C) covalent (nonmetal + nonmetal)
 (D) diatomic molecules (two of the same atom bound together)
 (P) polyatomic ion (groups containing more than two elements covalently bonded together that
carry an overall charge. )
1. CaCl2
11. N2
2. CO2
12. NH4Cl
3. H2O
13. HCl
4. BaSO4
14. KI
5. K2O
15. NaOH
6. NaF
16. NO2
7. Na2CO3
17. AlPO4
8. CH4
18. FeCl3
9. O2
19. P2O5
10. LiBr
20. N2O3
PART II:
Diatomic molecules are two of the same atom bound together. These atoms do not exist alone in nature
because they are extremely reactive. They always will come in pairs.
Elements that form diatomic molecules are: Br, I, N, Cl, H, O, F
They are called the BIG SEVEN.
*YOU MUST MEMORIZE THESE DIATOMIC MOLECULES.
1. On the periodic table below, color the diatomic molecules.
2. Draw the structure formula (using Lewis Dot drawings) for the following diatomic molecules.
2. Br2
Total # of valence electrons _ 14___
Br Br
3. I2
4. O2
4. Cl2
Total # of valence electrons_______
14
-2
12
-12
0
Total # of valence electrons_______
Total # of valence electrons_______
5. H2
6. N2
Total # of valence electrons_______
Total # of valence electrons_______
PART III:
Polyatomic ions are groups of covalently bonded atoms that carry an overall charge.
Drawing the structure formulas for the MOLECULES below:
Remember… Constructing Dot Diagrams for the polyatomic ions is the same as constructing Dot Diagrams for molecules,
except the difference in charge (+ or -) must be accounted for.
STEPS:
1. Count the total number of valence e-.
2. Determine the central atom. The following are guides:
 Often the unique atom (only one of it) is the central atom.
 Or put the least electronegative element in the middle.
3. Arrange the other atoms around the central atom creating a skeleton.
4. Connect all bonded atoms in the skeleton with one bond.
5. Subtract the number of electrons already used for the single bonds; two for each bond.
6. Distribute the remaining electrons in pairs around the atoms, trying to satisfy the octet rule. Assign them to the most
electronegative atom first.
7. If you run out of electrons before all atoms have an octet of electrons, you need to form double or triple bonds.
8. If you have extra electrons and all of the atoms have an octet, put the extra electrons on the central atom in pairs.
9. *Polyatomic Ions: Account for the charge (+ or -). Place structure in brackets with the charge indicated on outside.
1. NH3
Total # of valence electrons
# of electrons involved in single bonds
_______
_______
_________________________________________
2. NH4 +
Number of electrons to distribute
_______
*For Polyatomic Ions: what is the charge
_______
Total # of valence electrons
# of electrons involved in single bonds
_______
_______
_________________________________________
3. SO3 2-
Number of electrons to distribute
_______
*For Polyatomic Ions: what is the charge
_______
Total # of valence electrons
# of electrons involved in single bonds
_______
_______
_________________________________________
Number of electrons to distribute
_______
*For Polyatomic Ions: what is the charge
_______
4. HI
Total # of valence electrons_______
8. SiO2
Total # of valence electrons_______
5. CO2
Total # of valence electrons_______
9. PO4 -3
Total # of valence electrons_______
6. NO-3
Total # of valence electrons_______
10. SCl4
Total # of valence electrons_______
7. H3O+
Total # of valence electrons_______
11. N(CH3)3
Total # of valence electrons_______
Chemistry I: VSEPR-MOLECULAR SHAPES #5
Using the VSEPR Theory, name and sketch the shape of the following molecules.
STEPS:
1. Count the total number of valence e-.
2. Draw the Lewis Dot Structure
3. Determine the # of lone pairs on central atom
Molecule
Formula
1. N2
2. H2O
3. CO2
4. NH3
5. CH4
6. SO3
Valence
Electrons
*Show math*
Draw Lewis Dot
Structure
4. Write the VSEPR Formula (eg. AX2)
5. Draw the Ball & Stick Model
6. Write the name of the geometric shape
# of lone
pairs on
central
atom
VSEPR
Formula
Draw Geometric Shape
(Ball & Stick Method)
Name of
Geometric
Shape
Molecule
Formula
7. HF
8. CH3OH
9. H2S
10. I2
11. CHCl3
12. O2
Valence
Electrons
*Show math*
Draw Lewis Dot
Structure
# of lone
pairs on
central
atom
VSEPR
Formula
Draw Geometric Shape
(Ball & Stick Method)
Name of
Geometric
Shape