Transcript Lewis Dot (Electron Dot) Diagrams

2.1 The Formation of Ionic and
Covalent Bonds
Learning Goals …
… identify the intramolecular forces within a
compound as ionic, polar covalent or non
polar covalent
… draw Lewis structures for ionic and covalent
compounds
Lewis Theory of Bonding
Why do bonds form?
• atoms and ions are stable if they have a full valence shell
of electrons
• electrons are most stable when paired
• atoms form chemical bonds to achieve a full valence shell
• a full valence shell may be achieved by either transferring
(ionic bonds) or sharing electrons (covalent bonds)
A LEWIS STRUCTURE shows the arrangement of electrons and
bonds in a molecule or polyatomic ion
The type of intramolecular force can be determined by the
difference in electronegativity values (ΔEN)
If Δ EN > 1.7 then the bond is IONIC
If Δ EN < 1.7 and > 0.5 then the bond is POLAR COVALENT
If Δ EN < 0.5 then the bond is NON-POLAR COVALENT
Electronegativities
Δ EN
a) KF
Na = 0.82 F = 3.98
3.16
ionic
b) HCl
H = 2.10 Cl = 3.16
1.06
Polar
covalent
c) O2
O = 3.44 O = 3.44
0
bond type
Non polar
covalent
Ionic Bonds:
• ΔEN > 1.7
• The valence electrons are transferred.
• ions are formed and held together by electrostatic
attraction
Polar Covalent Bonds:
• ΔEN > 0.5 and < 1.7
• The valence electrons are shared unequally and
partial charges are formed (δ+, δ-)
Non-Polar Covalent Bonds:
• ΔEN < 0.5
• The valence electrons are shared equally
– no partial charges
Draw a Lewis Structure for each of the following compounds.
Draw the structural formula for the compound.
a) NCl3
Δ EN = 0.12
Cl
N
Cl
Mg
1 lone pair
3 single bonds
Cl
b) MgO
non-polar covalent – equal sharing
Δ EN = 2.13
O
Cl N
Cl
Cl
ionic – transfer eMg +2
O
–2
c) H2O
Δ EN = 1.34 Polar covalent – unequal sharing
2 lone pairs
δδ+
δ+
H
O
H O H
H
d) CH2O Δ ENCH = 0.45 Nonpolar covanent – equal sharing
Δ ENCO = 0.89 Polar covalent – unequal sharing
H
C
O
δ+
H C
δO
H
H
double bond
no lone pairs on C
Polyatomic ions can be a bit more complicated. We need to
add up the total number of electrons.
e) NO3-
Δ EN = 0.40 Non-Polar covalent – equal sharing
1 N: 5e3 O: 3(6e-) = 18e-1 charge = 1 eTotal: 24 eStart: 24 eUsed: 6
24e-e-
O
N
O
O
double bond
2 single bonds
O
N
O -
O
no lone pairs on N
Each O has a full octet
N needs 2 more e-, no remaining e-, move lone pair from O
into bonding position between N and O
f) NO2-
Δ EN = 0.40 Non-Polar covalent – equal sharing
1 N: 5e2 O: 2(6e-) = 12e-1 charge = 1 eTotal: 18 e-
O
N
O
-
O
N
O -
Start: 18 eUsed: 418
16e-eEach O has a full octet
N needs 4 more e-, place 2 remaining e- as lone pair on N,
move 2 e- from O to create a double bond
Try this …
ex) ClO33 O: 3(6e-) = 18e1 Cl: 7 e-1 charge = 1 eTotal: 26 e-
-
-
O
Cl
O
O
O
Cl
O
O
Central Atoms With an Expanded Valence Energy Level
• larger atoms can accommodate addition valence electrons
due to their size
H
Ex ) PH5
H
H
P
H
P
H
H
H
H
H
H
Each H has a full shell but P is overfilled. This is an exception
to the octet rule.
Can I …
… identify compounds as ionic, polar covalent or non
polar covalent
… draw Lewis structures for ionic and covalent
compounds
HOMEWORK
P63 #1-12, 14
WS “Lewis Dot Diagrams”