Ionic Bonding
Valence Electrons,
Lewis Dot Structures,
and
Electronegativity
Valence Electrons
• valence electrons – the outermost electrons.
6 protons = C = carbon
Outer electrons
= 4 e–available
for bonding
Inner electrons =
2 e– not available
for bonding
4 electrons in valence shell
Lewis Dot Structures
• Lewis dot structures are a convenient way to
show how many valence electrons an atom
has.
• Example: Draw the Lewis dot structure for
hydrogen.
H
Lewis Dot Structures
dots = number of valence electrons.
The maximum number of dots is 8.
Look for the number at the top of the
column (e.g. 5A).
Exception: Helium only has 2 dots.
Ne
He
More Lewis Structure Practice
Draw the Lewis structure for oxygen.
O
Draw the Lewis structure for
magnesium.
Mg
Draw the Lewis structure for chlorine.
Cl
Even More Lewis Structure Practice
Draw the Lewis structure for carbon.
C
Draw the Lewis structure for
potassium.
K
Draw the Lewis structure for
phosphorus.
P
Electronegativity
• electronegativity – how much an atom wants
to keep hold of its electrons.
• ionization energy – the energy required to
remove an electron from an atom.
Lower
electronegativity
Greater
electronegativity
Metals
Nonmetals
Role Models: The Noble Gases
• An atom’s electrons are at their most stable
when they reorganize their electrons to more
closely resemble the electron configuration of
a noble gas.
• All atoms want to have stable electron
configurations.
Role Models: The Noble Gases
• All atoms wish their electrons were like the
noble gases’ electrons.
Example:
Beryllium and Oxygen
The 2 valence e–
Now
beryllium’s
matches that
of helium
4 protons =
Be = beryllium
Now oxygen’s
matches that
of neon
The 6 valence e–
8 protons =
O = oxygen
Same Thing, but in Lewis Dot Structure
Be
O
Semi-metals
or metalloids
Nonmetals
Metals
Three General Bonding Types
• Metal with Nonmetal
- form ionic compounds
• Metal with Metal
- form metallic compounds
• Nonmetal with Nonmetal
- form covalent compounds
Metal with Nonmetal Bonding
• Ionic compounds – the metal gives all of its
valence e– to the nonmetal.
• Known as – Salts, ions
Ions
• ion – an atom that gained or lost electrons to
become more like a noble gas.
+ Metals lose electrons to become positively
charged ions. We call them cations (cat-ions)
the “t” looks like a “+”. [e.g. 2A  +2]
– Nonmetals gain electrons to become
negatively charged ions. We call them anions
(an-ions) “n” for negative “–”.
[e.g. (8 – 6A) × -1 -2]
Writing the Charges
• Write out the ion that sodium forms.
Na+
• Write out the ion that chlorine forms.
Cl–
• Write out the ion that magnesium forms.
Mg2+
• Write out the ion that oxygen forms.
O2–
So where do the electrons go?
• Usually atoms that become cations give their
electrons to anions.
2+
Be
2–
O
• Now both the cations and anions resemble
noble gases, however now both have net
charges.
Basic Electrical Charge Laws
+ and – : Attract
(pull together)
Naming (aka nomenclature)
• Metals keep their names unchanged.
(e.g. sodium, aluminum, calcium)
• Transition metals have their charge shown as roman
numerals in parenthesis after the name.
Fe2+  iron (II)
Cu1+  copper (I)
Fe3+  iron (III)
Cu2+  copper (II)
• Nonmetals have the last one or two syllables
of their names altered with an –ide ending.
fluorine  fluoride
chlorine  chloride
nitrogen  nitride
oxygen  oxide
Naming (aka nomenclature)
• Metals keep their names unchanged.
(e.g. sodium, aluminum, calcium)
• If there are more than one possible charge for
a metal (the transition metals), the charge will
be specified in roman numerals after the
name.
Fe2+  iron (II)
Cu1+  copper (I)
Fe3+  iron (III)
Cu2+  copper (II)
Naming Continued
• Nonmetals have the last one or two syllables
of their names altered with an –ide ending.
• Examples:
carbon  carbide
fluorine  fluoride
nitrogen  nitride chlorine  chloride
oxygen  oxide
bromine  bromide
sulfur  sulfide
iodine  iodide
phosphorus  phosphide
Naming Continued
• Now put the metal and nonmetal ion names
together and you get the name for the ionic
compound.
• Examples:
LiF  lithium fluoride
NaCl  sodium chloride
KBr  potassium bromide
MgS  magnesium sulfide
CuI  copper (I) iodide
CuO  copper (II) oxide
FeN  iron (III) nitride
Sodium Chloride – NaCl
Na+
Cl–
Na+
Cl–
Cl–
Na+
Cl–
Na+
Na+
Cl–
Na+
Cl–
+6
–6
0
Magnesium Chloride
Mg2+
Cl–
Cl–
Cl–
Cl–
Mg2+
Cl–
Cl–
Cl–
Mg2+
Cl–
Cl–
Cl–
Mg2+
Mg2+
Cl–
Mg2+
Cl–
+ 12
– 12
6
+ 60
Magnesium Chloride
6x
Mg2+
12 x
6 = 1
12 2
Cl–
Mg6Cl12
MgCl2
Magnesium Chloride
Empirical Formula
Cl–
Cl–
Mg2+
MgCl2
Formula Unit (f.u.) – the
smallest amount of an ionic
compound that still has the
same ratio of ions as in the
formula.
Iron (III) Oxide
Fe3+ O2–
criss-cross
Fe2O3
2 x (+3) = +6
3 x (–2) = –6
0
Sodium Cloride
+
Na
–
Cl
criss-cross
Na1Cl1
NaCl
Magnesium Sulfide
Mg2+ S2–
criss-cross
Mg2S2
MgS
2= 1
2 1
Sodium Oxide
Na+ O2–
criss-cross
Na2O
REMEMBER
Fe3+ O2–
Fe2O3
Top right corner:
Charge
Bottom right corner:
How many atoms/ions
Polyatomic Ions
• Poly – many
• Atomic – having to do with atoms
• Polyatomic ions – ions made from multiple
atoms
• List on p.257
Polyatomic Ions (List on p.257)
Ion name
Formula
Acetate
CH3COO–
Ammonium
NH4+
Carbonate
CO32–
Chromate
CrO42–
Cyanide
CN–
Dichromate
Cr2O72–
Hydroxide
OH–
Ion name
Formula
Nitrate
NO3–
Nitrite
NO2–
Permanganate MnO4–
Peroxide
O22–
Phosphate
PO43–
Sulfate
SO42–
Sulfite
SO32–
Thiosulfate
S2O32–
Magnesium Nitrate
Mg2+ NO3–
criss-cross
Mg (NO3)2
1 x Mg
2xN
6xO
REMEMBER
Mg2+
NO3–
Mg (NO3)2
Top right corner:
Charge
Bottom right corner:
How many atoms/ions
Sodium Chloride – NaCl
Na+
Cl–
Na+
Cl–
Cl–
Na+
Cl–
Na+
Na+
Cl–
Na+
Cl–
Sodium Chloride – NaCl
Sodium Chloride – NaCl
Sodium Chloride – NaCl
Sodium Chloride – NaCl
Sodium Chloride – NaCl
Magnesium Chloride (MgCl2)
Mg2+
Cl–
Cl–
Cl–
Cl–
Mg2+
Cl–
Cl–
Cl–
Mg2+
Cl–
Cl–
Cl–
Mg2+
Mg2+
Cl–
Mg2+
Cl–
Magnesium Chloride (MgCl2)
Magnesium Chloride (MgCl2)
Calcium Fluoride (CaF2)
Calcium Fluoride (CaF2)
Calcium Fluoride (CaF2)
Calcium Fluoride (CaF2)
Calcium Fluoride (CaF2)
Calcium Fluoride (CaF2)
H
He
Li Be B
C
N
O
F
Na Mg Al
Si
P
S
Cl Ar
K
Ca
Ne
Br Kr
I
Xe
4 e– in valence shell
Ions
• ion – an atom that gained or lost electrons.
• metal ions lose electrons to become more
positively charged. (e.g. 2A  +2)
• Nonmetal ions gain electrons to become more
negatively charged. (e.g. 8 – 6A –2)
Ions
• ion – an atom that gained or lost electrons.
• metal ions lose electrons to become more
positively charged. (e.g. 2A  +2)
• Nonmetal ions gain electrons to become more
negatively charged. (e.g. 8 – 6A –2)
Ions - again
• ion – an atom that gained or lost electrons to
become more like a noble gas.
+ Metal ions = cations (cat-ions)
the “t” looks like a “+”.
– Nonmetal ions = anions (an-ions). “–”