__________________ = atoms tend to gain, lose or share
electrons so as to have 8 electrons
C would like to Gain 4 electrons
N would like to Gain 3 electrons
O would like to Gain 2 electrons
Electron Dot diagrams are…
• A way of showing & keeping
track of valence electrons.
• How to write them?
• Write the symbol - it represents
the nucleus and inner (core)
electrons
• Put one dot for each valence
electron (8 maximum)
• They don’t pair up until they
have to (Hund’s rule)
X
The Electron Dot diagram for
Nitrogen
Nitrogen has 5 valence
electrons to show.
 First we write the symbol.
Then add 1 electron at a
time to each side.
Now they are forced to pair up.
We have now written the electron dot
diagram for Nitrogen.

N
Electron Dot Structures
Symbols of atoms with dots to represent the valence-shell
electrons
Learning Check

A.
X would be the electron dot formula for
1) Na
B.

X

1) B
2) K
3) Al
would be the electron dot formula
2) N
3) P
The type of bond can usually be calculated by
finding the difference in electronegativity of the two
atoms that are going together.
Electronegativity
Difference
• If the difference in electronegativities is
between:
– 1.7 to 4.0: Ionic
– 0.3 to 1.7: Polar Covalent
– 0.0 to 0.3: Non-Polar Covalent
Example: NaCl
Na = 0.8, Cl = 3.0
Difference is 2.2, so
this is an ionic bond!
Formation of Ions from ________
 Ionic compounds result when metals react with
nonmetals
 Metals lose electrons to match the number of valence
electrons of their nearest noble gas
 Positive ions (cations) form when the number of
electrons are less than the number of protons
Group 1 metals  ion 1+
Group 2 metals  ion 2+
•
Group 13 metals  ion 3+
Formation of Sodium Ion
Sodium atom
Na 
–
e
Sodium ion

Na +
e- config:
e- config:
11 p+
11 e-
11 p+
10 e-
Formation of Magnesium Ion
Magnesium atom
Magnesium ion

Mg 
–
e- config:
12 p+
12 e-
2e 
Mg2+
e- config:
12 p+
10 e-
Electron Dots For Cations
• Let’s do Scandium, #21
• The electron configuration is:
1s22s22p63s23p64s23d1
• Thus, it can lose 2e- (making it
2+), or lose 3e- (making 3+)
Sc
2+
= Sc
Scandium (II) ion
Sc =
3+
Sc
Scandium (III) ion
Electron Dots For Cations
• Let’s do Silver, element #47
• Predicted configuration is:
1s22s22p63s23p64s23d104p65s24d9
• Actual configuration is:
1s22s22p63s23p64s23d104p65s14d10
Ag =
1+
Ag
(can’t lose any more, charges
of 3+ or greater are uncommon)
Learning Check
A. Number of valence electrons in aluminum
1) 1 e2) 2 e3) 3 eB. Change in electrons for octet (Al)
1) lose 3e2) gain 3 e3) gain 5 eC. Ionic charge of aluminum
1) 32) 53) 3+
D.
E.
F.
12 p+ and 10 e1) 0
2) 2+
50p+ and 46 e1) 2+
2) 4+
15 p+ and 18e2) 3+
2) 3-
3) 23) 43) 5-
Ions from __________ Ions
In ionic compounds, nonmetals in 15, 16, and 17
gain electrons from metals
Nonmetal add electrons to achieve the octet
arrangement (forming anions)
Nonmetal ionic charge:
3-, 2-, or 1-
Chloride Ion
unpaired electron
octet

:Cl 

e- config:
17 p+
17 e-

+
e

1-
: Cl:

e- config:
17 p+
18 eionic charge
________ Bond
• Between atoms of metals and nonmetals with
very different electronegativity
• Bond formed by transfer of electrons
• Produce charged ions in all states.
• Conductors and have high melting point.
• Examples; NaCl, CaCl2, K2O
Ionic bonding
• Its like taking candy from a baby…….
• The nonmetal (bullies) take electrons (candy)
from the metals (babies)
1). Ionic bond – electron from Na is transferred to Cl,
this causes a charge imbalance in each atom. The Na
becomes (Na+) and the Cl becomes (Cl-), charged
particles or ions.
Ionic Bonding
Lets do an example by combining
calcium and phosphorus:
Ca
P
• All the electrons must be accounted for, and
each atom will have a noble gas
configuration (which is stable).
Ionic Bonding
= Ca3P2
Formula Unit
This is a chemical formula, which
shows the kinds and numbers of atoms in
the smallest representative particle of the
substance.
For an ionic compound, the smallest
representative particle is called a:
Formula Unit
Properties of Ionic Compounds
1. ____________ solids - a regular repeating
arrangement of ions in the solid:
Ions are strongly bonded together.
– Structure is rigid.
2. High melting points
• Coordination number- number of ions of
opposite charge surrounding it
A. Energy of Bond Formation
• Lattice Energy
– Energy released when one mole of an
ionic crystalline compound is formed
from gaseous ions
– Greater lattice energy = greater ionic
bond
C. Johannesson
Do they Conduct?
•
Conducting electricity means allowing
charges to move.
• In a solid, the ions are locked in place.
• Ionic solids are insulators.
• When melted, the ions can move around.
3. Melted ionic compounds conduct.
–
–
NaCl: must get to about 800 ºC.
Dissolved in water, they also conduct (free to
move in aqueous solutions)
- Page 198
The ions are free to move when they are
molten (or in aqueous solution), and thus
they are able to conduct the electric current.
A. Oxidation Number
• The charge on an ion.
• Indicates the # of e- gained/lost to
become stable.
4-
1+
2+
3+ 4+ 3- 2- 1(1+ to +3)
0
B. Ionic _________
 Write the names of both elements,
cation first.
 Change the anion’s ending to -ide.
 Write the names of polyatomic ions.
 For ions with variable oxidation #’s, write
the ox. # in parentheses using Roman
numerals. Overall charge = 0.
Polyatomic Ions
• “Many atoms”
• A group of covalently-bonded atoms that have
a net charge.
• They act as a unit as an anion or cation.
• You should be able to recognize polyatomic
ions in a given chemical formula.
•
•
•
•
•
•
•
•
•
•
•
•
•
•
•
Symbol
CH3COO1–
NH41+
AsO43–
C6H5COO1–
HCO31–
BrO31–
CO32–
ClO31–
ClO21–
C6H5O73–
CN1–
Cr2O72–
OH1–
CrO42–
Poly atomic ions
Name
acetate ion
ammonium ion
arsenate ion
benzoate ion
bicarbonate ion
bromate ion
carbonate ion
chlorate ion
chlorite ion
citrate ion
cyanide ion
dichromate ion
hydroxide ion
chromate ion
Poly atomic ions
•
•
•
•
•
•
•
•
•
•
•
•
•
•
•
•
•
C6H5O73–
CN1–
Cr2O72–
OH1–
ClO1–
IO31–
PO31–
NO31–
NO21–
C2O42–
ClO41–
MnO41–
PO43–
SiO32–
SO42–
SO32–
S2O32–
citrate ion
cyanide ion
dichromate ion
hydroxide ion
hypochlorite ion
iodate ion
phosphite ion
nitrate ion
nitrite ion
oxalate ion
perchlorate ion
permanganate ion
phosphate ion
silicate ion
sulfate ion
sulfite ion
thiosulfate ion
Special Ions
NAME
Copper (I)
Copper (II)
Iron (II)
Iron (III)
Chromium (II)
Chromium (III)
Lead (II)
Lead (IV)
Scandium (II)
Scandium (III)
Ox. #
+1
+2
+2
+3
+2
+3
+2
+4
+2
+3
C. Ionic ___________
 Write each ion. Put the cation first.
 Overall charge must equal zero.
• If charges cancel, just write the symbols.
• If not, crisscross the charges to find
subscripts.
 Use parentheses when more than one
polyatomic ion is needed.
 Roman numerals indicate the oxidation #.
C. Ionic Formulas
 potassium chloride

• K+ Cl
 magnesium nitrate
• Mg2+ NO3

 copper(II) chloride
• Cu2+ Cl

C. Ionic Formulas
 calcium oxide
• Ca2+ O2

 aluminum chlorate
• Al3+ ClO3

 iron(III) oxide
• Fe3+ O2

B. Ionic Names
 Write the names of both elements,
cation first.
 Change the anion’s ending to -ide.
 Write the names of polyatomic ions.
 For ions with variable oxidation #’s, write
the ox. # in parentheses using Roman
numerals. Overall charge = 0.
B. Ionic Names
 NaBr
 Na2CO3
 FeCl3
Ionic hydrates
• Ionic Hydrates – ionic compounds that have
loosely held water molecules
• Example:
CuSO4●5H20(s)
Formula of
ionic compound
raised # of water
dot
molecules
• By heating an ionic hydrate, the water molecules are released
and the ionic compound becomes “anhydrous”
Example: CuSO4 ●5H20(s) + heat
CuSO4 (s)
(hydrated)
(anhydrous)
Name: ____________________________________