Ionic vs. Covalent
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IONIC
Electrons are transferred
Between metals and nonmetals (or cations (+)
and anions (-))
Usually solids at room
temperature
High melting and boiling
points
Good conductors
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COVALENT
Electrons are shared
Between non-metal and
non-metal
Usually gases at room
temperature
Low melting and boiling
points
Bad conductors
Ionic vs. Covalent Naming
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IONIC
Name the metal (or cation)
Name the non-metal by
dropping the ending and
changing it to –ide.
If a polyatomic ion is used, just
name it
If there is a transition metal, you
must include a roman numeral
in its name
Subscripts do not matter in
naming
Ex: Mg2Cl-magnesium chloride
Ex: Be2OH- beryllium hydroxide
Ex: Fe2O3- iron (III) oxide
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COVALENT
If the first element has
more than one atom, give
it a prefix.
The second element must
have a prefix and end in
–ide.
Subscripts matter in
naming!
Ex: C8O3- octacarbon
trioxide
Ex: NO2-nitrogen dioxide
REVIEW…
 All atoms bond together to get to ______ valence electrons
 Valence electrons are predicted by the group number
 Group 1 = 1
 Group 2 = 2
 Group 13 = 3
 Etc…………
 Lewis Dot Structures show the number of valence electrons
 Covalent bonds are ones where atoms share their electrons
THIS IS HOW WE DO IT!!!
 One bond represents two electrons
 Bonds are represented by lines
 Remember in a covalent bond, the
electrons are shared
Single Bonds
 Single covalent bond: one pair of
shared electrons (two electrons)
=
 Atoms may be the same
 H2,
F2, Cl2, Br2, I2, O2
 Atoms may be different
 HF,
HCl, HBr, HI
Formation of Single Covalent Bonds
Examples
H2
F2
HCl (chlorine)
Formation of Single Covalent Bonds
Your Turn!
1. Cl2
2. HBr
Lewis Structure Tips
1. Carbon is usually central.
2. Hydrogen is NEVER central.
3. Halogens and oxygen are rarely
central.
4. If you can’t decide what atom should
be central, pick the one with the
lowest electronegativity.
Formation of Multiple Single Bonds
Examples
NH3
SeCl2
Practice Time!
 Draw the
Lewis Dot
Structures for
the covalent
compounds to
the right 
Br2
2. HI
3. PH3
4. H2S
5. CCl4
1.
LDS and Multiple Covalent Bonds
Sometimes atoms must share
more than one pair of
electrons in order to become
stable.
 This results in double and triple
bonds.
Double and Triple Bonds
 Double covalent bond: two pairs
of shared electrons (four electrons)
=
 Triple covalent bond: three pairs
of shared electrons (six electrons)
=
Examples
O2
CO2
N2
COCl2