CP Chemistry
Mrs. Klingaman
Chapter 6:
Chemical
Bonding
Ionic Bonding:
Electron Transfer
Covalent Bonding:
Electron Sharing
Name: _______________________________________________
Mods: _____________________________
Chapter 6: Chemical Bonding
Reading Guide
6.1 – Introduction to Chemical Bonding (pgs. 165-167)
1. What is a chemical bond?
2. Distinguish between an ionic bond and a covalent/molecular bond:

Ionic bond -

Covalent bond -
3. Bonding between atoms of different elements is rarely purely ionic or purely covalent; it usually falls
somewhere between these two extremes. The degree to which bonding between atoms of two
elements is ionic or covalent can be estimated by calculating the difference in the elements’
__________________________.
4. Use Figure 1.2 to fill in the following table.
Difference in Electronegativity
Bond Type
0.0 - 0.3
0.4 - 1.7
1.7 - 3.3
5. How do polar covalent bonds and nonpolar covalent bonds differ?
 Nonpolar covalent -
 Polar covalent -
6.2 – Covalent Bonding and Molecular Compounds (pgs. 168-179)
6. Define the following words:

Molecule-

Molecular formula-
7. Why do atoms form chemical bonds? Nature favors chemical bonding because atoms have
_____________________________________________ when they are bonded to other atoms than
they have when they are independent particles.
8. How do atoms form chemical bonds? The attractive and repulsive forces between two atoms play
an important role in the formation of a chemical bond. (refer to Figure 2.3)
 Attractive Forces –
 Repulsive Forces –
 The relative strength of attraction and repulsion between the charged particles (nucleus &
electrons) depends on the ________________________ separating the atoms.
9. Define the following terms:

Bond length-

Bond energy-
The Nature of Covalent Bonding
10. Atoms tend to form bonds to follow the octet rule. What is the octet rule?
11. Covalent bonding involves only the atom’s outermost, or _____________________ electrons.
To keep track of these electrons, it is helpful to use __________________________ notation.
12. Using Lewis Dot Structures to Represent Compounds:
 Using Lewis Dot Structures, show how two fluorine atoms form a single covalent bond to fulfill the
octet rule (pg. 175):
 The pair of electrons in between the two fluorine atoms represents the ________________
pair of electrons involved in the covalent bond. For this reason, they are also known as the
bonding electrons. This pair of electron-dots is often replaced by a long _____________, which
represents the covalent bond.
 An ___________________ pair, or ____________ pair, of electrons are those that are not
involved in the bonding and belong exclusively to one atom. For this reason, they are also known
as nonbonding electrons.
13. Lewis structures are drawn to show the composition of a molecule, which atoms are bonded to each
other, and what type of bond holds them together. There are three types of covalent bonds:
 Single Bond –
 Double Bond –
►
Example: ethene, C2H4 (pg. 176)
 Triple Bond –
►
Example: nitrogen gas, N2 (pg. 177)
14. Using Figure 2.10 on pg. 177, fill in the table below
Bond
Bond Length (pm)
Bond Energy (kJ/mol)
C–C
C=C
C
C
 When a multiple bond forms, bond length ______________and bond energy ______________.
15. Define resonance and give an example (ozone, O3 - pg. 179) of structures that exhibit resonance:
6.3 – Ionic Bonding and Ionic Compounds (pgs. 180-184)
The Nature of Ionic Bonding
16. Define the following terms:

Ionic compound-

Formula unit-
17. Atoms of metals, such as sodium, readily ________ electrons to form ______________, while
atoms of nonmetals, such as chlorine, readily _________ electrons to form ______________.
18. Show the formation of sodium chloride (NaCl) from sodium and chlorine ions (pg. 181)
19. All ionic compounds have an orderly arrangement known as a _______________ lattice structure.
This structure arranges the cations and anions so that they surround one another giving the
compound its lowest potential energy, thus making it very stable.
20. Define lattice energy. How is this different from bond energy?
21. The strong attraction of (+)/(-) ions is described by the lattice energy (kJ/mol) in Figure 3.6. All
lattice energies are negative, which means that energy is released when ionic bonds are formed.
This is a good thing, because all atoms and compounds want to be at their lowest possible energy
state. In fact, the higher the lattice energy (the greater the energy released) the stronger the
ionic bond and the more stable the compound.
22. The forces of attraction between molecules are much _______________________ than the forces
of attraction among formula units in ionic bonding. Thus, when comparing the strength of ionic and
covalent bonds, _______________________ compounds have a much stronger attraction which
gives rise to different properties in the two types of compounds which we will discuss further in class.
23. Define polyatomic ion and give one example:
6.4 – Metallic Bonding (pgs. 185-186)
24. Define the following terms:

Metallic bonding-

Malleability-

Ductility-
25. The amount of energy as heat required to vaporize a metal is a measure of the strength of the
metallic bonds that hold it together. This is known as the metal’s _______________________
___________________________________. Generally, the greater the metal’s nuclear charge and
the more electrons the metal has, the stronger the metallic bonds (refer to Figure 4.3).
Summary of Chemical Bonds and Their Properties

Ionic Bond: a nonmetal ____________ electrons from a metal due to a __________________
difference in electronegativity, leading to an attraction between the resulting cation and anion.
Example:

Polar Covalent Bond: an __________________ difference in electronegativity leads to an unequal
sharing of electrons between two ___________________.
Example:

Nonpolar Covalent Bond: two ____________________ with ________________ electronegativity
or only a ______________ difference in electronegativity equally share electrons.
Example:

Metallic Bond: results from the attraction between ___________ atoms and the surrounding sea of
__________________.
Example:
Compound
Melting Point
Conductivity
of
Solid
Solubility in
Water
(polar solvent)
Conductivity
in
Water
Solubility in
Hexane
(nonpolar solvent)
Hardness
Brittleness
Ionic Bond
Polar Covalent
Bond
Nonpolar Covalent
Bond
Metallic Bond
Sodium Chloride
Sucrose
Stearic Acid
(NaCl-table salt)
(C12H22O11-table sugar)
(CH3(CH2)16COOH-wax)
Aluminum (Al)
and
Mercury (Hg)
Determining the Type of Chemical Bond
Bond types based on differences in electronegativity:

Ionic > 1.7

Polar Covalent = 0.4 – 1.7

Nonpolar Covalent < 0.4
Directions: Determine the electronegativity difference between the atoms listed in the table below
and then determine the corresponding type of chemical bond.
Atoms that are bonding
1
Hydrogen and oxygen
2
Magnesium and sulfur
3
Strontium and fluorine
4
Carbon and oxygen
5
Magnesium and nitrogen
6
Nitrogen and oxygen
7
Chlorine and bromine
8
Hydrogen and nitrogen
9
Lithium and oxygen
10
Oxygen and fluorine
11
Bromine and oxygen
12
Hydrogen and iodine
13
Magnesium and hydrogen
14
Hydrogen and hydrogen
Electronegativity
Difference
Bond Type
Electron Dot Structure & Ions
1)
Draw electron dot structures for the following atoms:
a. Magnesium (Mg)
e. Oxygen (O)
b. Bromine (Br)
f.
c. Krypton (Kr)
g. Lithium (Li)
d. Silicon (Si)
h. Phosphorous (P)
Aluminum (Al)
2) Fill in the table below by stating if the atom will lose or gain electrons to form an ion, how
many electrons it will lose/gain, what the resulting ion symbol will be, and if it is a cation or
an anion.
Atom
Lose or Gain?
How Many
Electrons?
Ex) Nitrogen
gain
3
Magnesium
Bromine
Krypton
Silicon
Oxygen
Aluminum
Lithium
Phosphorous
Ion Symbol
Cation or
Anion?
anion
The Formation of Ions from Atoms
Answer the following questions about CALCIUM:

Is calcium a metal or a nonmetal? _____________________

How many total electrons does a calcium atom have? ____________

How many valence electrons does a calcium atom have? _____________

Draw the electron dot structure for a calcium atom in the space below:

Calcium will ___________ ____ electrons to become an ion.

Draw the symbol for a calcium ion in the space below:

Is the calcium ion a cation or an anion? ______________________

The calcium ion has the same electron configuration as ____________________.
Answer the following questions about SULFUR:

Is sulfur a metal or a nonmetal? _____________________

How many total electrons does a sulfur atom have? ______________

How many valence electrons does a sulfur atom have? _____________

Draw the electron dot structure for a sulfur atom in the space below:

Sulfur will ___________ ____ electrons to become an ion.

Draw the symbol for a sulfide ion in the space below:

Is the sulfide ion a cation or an anion? ______________________

The sulfide ion has the same electron configuration as ____________________.
Valence Electrons Review
1) Define valence electrons-
2) Fill in the following table:
Element
Barium
Sulfur
Aluminum
Bromine
Sodium
Carbon
Lithium
Neon
Group #
Block
(s, p, d, f)
Period #
# of Valence
Electrons
Ion Symbol
Lewis Structures – Ionic Compounds
Directions: Draw Lewis structures for the following ionic compounds showing the transfer of electrons and
the ions formed.
1) SrCl2
2) AlF3
3) BaI2
4) LiBr
5) Al2S3
6) CaO
Lewis Structures – Ionic Compounds Continued
Directions: Draw Lewis structures for the following ionic compounds showing the transfer of electrons and
the ions formed.
7) K3P
8) Mg3N2
9) Na2O
10) FrF
11) Be3As2
12) MgSe
Rules for Drawing Lewis Dot Structures
1) Determine the number of atoms of each element present in the molecule
2) Sum the valence electrons from all atoms

For an ANION  add one electron to the total for each negative charge

For a CATION  subtract one electron from the total for each positive charge
3) Determine the central atom in the molecule and attach all other atoms to it

The central atom is the least electronegative element in the compound excluding
hydrogen – hydrogen can NEVER be central! (Note: typically, the central atom is the
one written first in the molecular formula)

Write the symbols for all the other atoms around the central atom. Connect the atoms
to the central with a single bond (a dash). Keep track of the electrons being used.
Each single bond made uses 2 electrons.

Chemical formulas are often written in the order in which the atoms are connected in
the molecule (ex: HCN  carbon is the central atom)
4) Complete the octets around all the outer atoms bonded to the central

Note: Hydrogen atoms may only have a single electron pair around them

Keep track of the electrons being used to complete the octets
5) Place any leftover electrons on the central atom, even if doing so results in more than
an octet of electrons around the atom.

Note: Expanded octets occur when more than 8 electrons surround an atom. This
exception to the octet rule is allowed for any atom in the 3 rd row of the periodic table
and after!

Keep track of the electrons – make sure the total number of valance electrons
available were used in the Lewis structure
6) If there are not enough electrons to give the central atom an octet, multiple bonds are
needed.

Remove one or more of the nonbonding pairs of electrons on one of the outer atoms
and draw a double bond (2nd dash) connecting the outer atom to the central. If need
be, a triple bond (3rd dash) may be formed by removing another nonbonding pair from
the same outer atom.

Keep track of the electrons – make sure the total number of valance electrons
available were used in the Lewis structure
Drawing Lewis Structures – Examples
Lewis Dot Structures:
1) Methane: CH4
4) Sulfite ion: SO32-
2) Hydrochloric Acid: HCl
5) Hydrocyanic Acid: HCN
3) Ammonia: NH3
6) Methanal (aka: formaldehyde): H2CO
Resonance Structures: multiple Lewis structures are used to describe molecules which have
double or triples bonds that can be moved to different sides of the central atom and still stay
bonded to an identical outer atom.
7) Nitrite ion: NO2 –
Exceptions to the Octet Rule:
Less than an Octet  central
atoms which have FEWER than
8 electrons (seen with B & Be)
8) BF3
Expanded Octet  central atoms which have MORE
than 8 electrons (this is allowed for central
atoms found in the 3rd period and after)
9) SBr6
10) XeI4
Lewis Structures – Covalent Compounds
Directions: Draw Lewis structures for the following molecules and polyatomic ions. Remember to consider
exceptions to the octet rule, double/triple bonds, ion charges, and resonance structures.
1) CBr4
4) NBr3
2) SiS2
5) CO32–
3) SF4
6) PCl5
7) N2
10) H2S
8) CN–
11) CCl2O
9) NH4 +
12) BCl3
Chapter 6 Review
1)
Identify and define the three major types of chemical bonding.
___________________________________________________________________________
___________________________________________________________________________
___________________________________________________________________________
___________________________________________________________________________
___________________________________________________________________________
___________________________________________________________________________
2)
In a covalent bond, how is the octet rule fulfilled so that every atom in a molecule has 8 valence
electrons?
___________________________________________________________________________
___________________________________________________________________________
___________________________________________________________________________
3)
In an ionic bond, how is the octet rule fulfilled?
___________________________________________________________________________
___________________________________________________________________________
___________________________________________________________________________
4)
In writing Lewis structures, how do you know when a multiple bond is needed?
___________________________________________________________________________
___________________________________________________________________________
___________________________________________________________________________
5)
In general, how do ionic and covalent compounds compare in terms of melting points and boiling
points? What does this tell you about the relative strength of ionic vs. covalent bonds?
___________________________________________________________________________
___________________________________________________________________________
___________________________________________________________________________
___________________________________________________________________________
6)
How do the properties of metals differ from those of both ionic and molecular compounds?
___________________________________________________________________________
___________________________________________________________________________
___________________________________________________________________________
___________________________________________________________________________
7)
Complete the table below by determining what type of bond (ionic, polar covalent, or nonpolar covalent)
will form between each set of atoms.
Bonded Atoms
Electronegativity
Difference
Bond Type
More Electronegative
Atom
H and I
S and O
K and Br
Si and Cl
K and Cl
Se and S
C and H
8)
List the bonding pairs from Question 8 in order of increasing ionic character.
__________________________________________________________________________________
__________________________________________________________________________________
__________________________________________________________________________________
9)
The lattice energy of sodium chloride, NaCl is -787.5 kJ/mol. The lattice energy for potassium chloride,
KCl, is -715 kJ/mol. In which compound is the bonding between ions stronger? Why?
__________________________________________________________________________________
__________________________________________________________________________________
__________________________________________________________________________________
__________________________________________________________________________________
10) Use electron dot notation to illustrate the number of valence electrons in each of the atoms below.
a. Iodine
b. Strontium
c. Carbon
d. Nitrogen
Drawing Lewis Diagrams: Covalent vs. Ionic
1. K2O
Type:___________________
2. OF2
Type:___________________
3. PO43-
Type:__________________
4. BeS
Type:___________________
5. AlBr3
Type:__________________
6. CH2O
Type:__________________
7. IBr5
Type:____________________
8. CaF2
Type:__________________
9. CH2F2
Type:__________________
10. H2S
Type:____________________
11. Ba3N2
Type:__________________
12. SiO2
Type:___________________
**Challenge** -- NO31-
Type: _________________________