SCH3U: Chemical
Bonding
Q: what did the bartender say
when oxygen, hydrogen, sulphur,
sodium and phosphorous walked
into the room?
A: OH SNaP!
Agenda
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Questions from Last Class?
Ionic, Covalent and Metallic Bonding
Lewis Diagram
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Single Atoms; hints for drawing Lewis diagrams
Compounds
 Total Valence Electrons
 Steps to draw Lewis Diagram for compounds
Practice Questions
Reminder (Homework; Quiz; Gizmo)
Ionic bonding (Li + Cl)

Ionic bonding (stealing/transfer of electrons)
Li + Cl  [Li]+[Cl]–
Chemical Formula
Rutherford Diagram
1e3p+ 2e-1e4n0
Li
7e- 8e- 2e-
Cl
17p+
18n0
3p+ 2e- 8e-8e-2e
4n0
[Li]+ [ Cl ]–
Lewis Diagram
17p+
18n0
Ionic bonding (Mg + O)
Mg + O  [Mg]2+[O]2–
Chemical Formula
1e12p+ 2e- 8e- 2e12n0
6e- 2e-
8p+
8n0
12p+ 2e- 8e- 8e- 2e- 8p+
8n0
12n0
1eRutherford Diagram
Mg
O
[Mg]2+ [
O
]2–
Lewis Diagram
Ionic bonding: Li + O
2Li + O [Li]2+[O]2– or 4Li + O2 2[Li]2+[ O]2–
3p+ 2e-1e4n0
1e6e- 2e-
3p+ 2e-1e4n0
Li
Li
1e-
O
3p+ 2e4n0
8e- 2e-
8p+
8n0
8p+
8n0
3p+ 2e4n0
[Li]2+ [
O
]2–
Ionic vs. Covalent
Types of Covalent Bonds

a. non Polar Covalent Bond: When EN
difference is zero. Equal sharing of
electrons.

b. Polar Covalent Bond: When EN
difference is >0 and <1.7. Results in unequal
sharing of electrons, one atom has a higher
affinity for the electrons.
Polar Covalent Bonds
Metallic Bond (neither ionic or
covalent)
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Occurs between two metal atoms.
Valence electrons are released into a sea of
electrons.
Atoms release their electrons in a shared
pool leads to stronger metallic bonds.
Electrons are free to move and are not held
together by a lattice formation (like that of
ionic bonds
That is why pure metals such as copper,
sodium and gold are soft elements
Metallic Bonds
Chemical Bonding
Lewis Structures
Lewis Diagram
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Single Atom: the chemical symbol and dots indicating
the number of electrons in the valence energy level
Compound: representation of covalent bonding based on
Lewis symbols; Shared electrons are shown as lines and
lone pairs as dots
To Write Lewis Diagrams
1. Write the element symbol. Around this, draw dots one for each
valence electron.
2. The dots should be spread over four sides. Dots are not
paired until all sides have at least one dot.
3. It does not matter on which side dots are placed. For example,
hydrogen can be drawn four ways:
4. The number of valence electrons is equal to the group
number. For example, hydrogen is in group IA (group 1) and it
has one valence electron. Neon is in 0 (group 8) and it has 8
valence electrons. The only exception is He which is in group
8 but has 2 valence electrons.
Example: AsBr3
STEP 1:place the single atom
in the center and other atoms
around it evenly spaced
Br As Br
Br
STEP 2: count the total #
of valence e- for all atoms
involved in the bonding
Arsenic: 1 Arsenic with
5 valence
AsBr3
electrons (1x5) = 5
5+21
=26
Bromine: 3 bromine
with 7 valence
electrons (3x7) = 21
STEP 3: place the electrons
in pairs between the central
atom and each non-central
atom
Br
As
Br
Br
STEP 4: place the remaining
electrons around the non-central
atom until each has 8 electrons
(H atoms have only 2e-)
Ex: AsBr3
5 + (7x3) = 26
26 - 6 = 20
Br
As
Br
Br
STEP 5: if electrons remain
they are placed in pairs around
the central atom
Ex: AsBr3
Br
As
Br
5 + (7x3) = 26
20 - 18 = 2
Br
Br As
Br
Structural
Br
Exception: if the central atom is in
group 14, 15, 16, 17 or 18, the
octet rule must be satisfied by moving
electron pairs from non-central
atoms, creating multiple bonds.
Ex: SO2
6 +(6x2)
=18
O S O
O- S = O
Structural Formula
Resonance Structure
When several structures with different electron distributions among
the bonds are possible, all the structures contribute to the electronic
structure of the molecule. These structures are called resonance
structures
EX) CO3
When
considering the resonance structure…
Formal charge closest to zero is favoured
Negative formal charges on the more negative elements are favoured
Summary
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STEP 1:place the single atom in the center and other atoms around
it evenly spaced
STEP 2: count the total # of valence e- for all atoms involved in the
bonding
STEP 3: place the electrons in pairs between the central atom and
each non-central atom
STEP 4: place the remaining electrons around the non-central atom
until each has 8 electrons (H atoms have only 2e-)
STEP 5: if electrons remain they are placed in pairs around the
central atom
Exception: if the central atom is in group 14, 15, 16, 17 or 18, the
octet rule must be satisfied by moving electron pairs from noncentral atoms, creating multiple bonds
Summary / Hints
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When drawing Lewis Diagrams…
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Symmetrical arrangements are favoured
Hydrogen is never a central atom
The least electronegative element is usually the
central atom
When considering the resonance structure…
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Formal charge closest to zero is favoured
Negative formal charges on the more negative
elements are favoured
Practice Questions
Draw Lewis Structure/ Structural Formula
for the following compounds

1.
2.
3.
4.
5.
6.
CCl4
HCl
NH3
H2O
OH
H2
Reminder
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Do the practice problems and assigned
homework
Gizmo activity is due this Thursday!