Chemical Bonds
z Elements form bonds to be in a lower energy state
z 1. Ionic Bonds – transfer of electrons, between
metal and nonmetal
z 2. Covalent Bonds – sharing of electrons, between
two nonmetals
z 3. Metallic Bonds- neighboring atoms in solid
metals form bonds
z Octet rule: atoms tend to gain, lose, or share
electrons until they are surrounded by eight valence
electrons to achieve a stable octet (noble gas
configuration)
Electron Dot Symbols
z Valence electrons: reside in the highest occupied
energy level, reside in the outer s & p orbitals and
are the electrons involved in chemical bonding.
z Electron-dot symbols are convenient way of
showing the s & p electrons & tracking them in
bond formation
z -They consist of the chemical symbol for the
element plus a dot for each valence electron
z Consider sulfur whose electron configuration is
[Ne]3s23p4, thus there are six valence electrons:
S
Ch. 7 – Ionic and Metallic Bonding
I
II
III
IV
B. Ionic Bonds
• Ionic Bonds – atoms transfer electrons from a
cation (positive ion) to an anion (negative ion) to
achieve an octet.
• Ionic compounds are stable due to the
electrostatic forces between unlike charges
organizing the ions of ionic substances into
a rigid, organized three-dimensional
arrangement:
• The ions are drawn together
• Energy is released
• Ions form solid lattice structure
Ionic Rule 1: Metals with a Single
Oxidation Number Bound to Non-Metals
The metal will take its positive oxidation
number and the non-metal will have to
take its negative oxidation number.

Only one compound can be formed
Name of
Metal
(
Root of
Non-metal
)
-ide
Ionic Rule 1: Metals with a Single
Oxidation Number Bound to Non-Metals
Example 1: Sodium reacts with oxygen to
produce Na2O, what is the name of this
compound

Since there is only one possible compound,
we do not have to indicate the number of
elements
sodium oxide
Ionic Rule 1: Metals with a Single
Oxidation Number Bound to Non-Metals
Example 2: What is the chemical formula
for aluminum oxide



First write the symbols of the elements
Next write the oxidation number of each
element above that element
Switch the oxidation numbers and reduce
3
-2
2
Al O
Comprehension Check
What is the name of Mg3N2?

magnesium nitride
What is the name of Li2Se?

lithium selenide
What is the formula for indium chloride?

InCl3
What is the formula for potassium
phosphide?

K3P
Ionic Rule 2: Metals with Multiple
Oxidation Numbers Bound to Non-Metals
The metal will take one of its positive
oxidation numbers and the non-metal will
have to take its negative oxidation number.


Since the metal has more than one possible
oxidation number, multiple compounds can be
formed
We need a distinct name for each
Name of
Metal
(
Metal’s Oxidation State
as a Roman Numeral
)(
Root of
Non-metal
)
-ide
Ionic Rule 2: Metals with Multiple
Oxidation Numbers Bound to Non-Metals
Example 1: What is the name of IrBr6?

First we need to determine how many electrons
that iridium needs to lose in order to satisfy 6
bromine atoms.
Each bromine needs one electron
There is only one iridium in this compound
Therefore, the iridium atom will have to supply all six
electrons, giving it a +6 oxidation number.
iridium(VI) bromide
Ionic Rule 2: Metals with Multiple
Oxidation Numbers Bound to Non-Metals
Example 2: What is the formula for
mercury(II) nitride?



First write the symbols of the elements
Next write the oxidation number of each
element above that element
Switch the oxidation numbers and reduce
2
-3
3
Hg N
Comprehension Check
What is the name of RuN?

ruthenium(III) nitride
What is the name of MnO3

manganese(VI) oxide
What is the formula for paladium(IV)
bromide?

PdBr4
What is the formula for molybdenum(V)
sulfide?

Mo2S5
Ionic Rule 3: Metals with a single Oxidation
Number Bound to Polyatomic Ions
Polyatomic Ions – strongly bound group of
atoms that have either lost or gained
electrons and become charged.




List of common Polyatomic Ions are on the back
of your Periodic Table
Polyatomic ions act as a single atom, with a
single name
Subscripts within the ion cannot be changed
Since there is only one oxidation number for the
metals and Polyatomic Ion, only one compound
can be produced.
Ionic Rule 3 : Metals with a single Oxidation
Number Bound to Polyatomic Ions
Naming these compounds is just like rule
1, except we do not add –ide to the end of
the polyatomic ion
Name of
Metal
(
Name of
Polyatomic Ion
)
Ionic Rule 3 : Metals with a single Oxidation
Number Bound to Polyatomic Ions
What is the name of Mg(NO3)2



First, you should recognize that there are
more than two elements involved, which
means that a Polyatomic Ion is involved
Next, look up the Metal in the periodic table
and confirm that it has a single oxidation
number
Look up the name of the Polyatomic Ion
magnesium nitrate
Ionic Rule 3 : Metals with a single Oxidation
Number Bound to Polyatomic Ions
What is the formula for calcium iodite?




First, since the second name does not end in
–ide, a polyatomic ion is involved.
Write the symbol for calcium and formula for
iodite.
Write the oxidation numbers above the metal
and the polyatomic ion
Switch the numbers, and use parenthesis
around the polyatomic ion if necessary
2
-1
1
Ca (IO2 )
Comprehension Check
What is the name of KHSO4?


potassium hydrogen sulfate
potassium bisulfate
What is the name of In2(C2O4)3?

indium oxalate
What is the formula of strontium bromate?

Sr(BrO3)2
What is the formula for germanium
phosphate?

Ge3(PO4)4
Ionic Rule 4: Metals with Multiple Oxidation
Numbers Bound to Polyatomic Ions
When the metal has more than one
possible oxidation number, more than one
compound can be formed

We must use Roman Numerals to indicate
which oxidation number the metal is using
Name of
Metal
(
Metal’s Oxidation State
Name of
as a Roman Numeral Polyatomic Ion
)
Ionic Rule 4 : Metals with Multiple Oxidation
Numbers Bound to Polyatomic Ions
What is the name of RhSO4?


First, there are more than two elements involved
Look up the oxidation and name of SO4
Sulfate (-2)

Finally, figure out which oxidation number the
metal is using.
There is only one rhodium, so it must account for all of
the electrons & would have to take a +2 oxidation
number
rhodium (II) sulfate
Ionic Rule 4 : Metals with Multiple Oxidation
Numbers Bound to Polyatomic Ions
What is the formula for nickel(II) ferrocyanide?




First, since the second name does not end in -ide, a
polyatomic ion is involved
Write the symbol for nickel and formula for
ferrocyanide
Write the oxidation numbers above the metal and the
polyatomic ion
Switch the numbers, and use parenthesis around the
polyatomic ion if necessary and reduce
2
-4
4
Ni 2(Fe(CN)6)
Comprehension Check
What is the name of Cr(IO)3?

chromium(III) hypoiodite
What is the name of CuMnO4?

copper(II) manganate or copper(I) permanganate
What is the formula for palladium(IV)
ferricyanide?

Pd3(Fe(CN)6)4
What is the formula for molybdenum(VI)
dichormate?

Mo(Cr2O7)3
B. Properties of Ionic Compounds
• Most ionic compounds are crystalline
solids at room temperature
• Arranged in repeating threedimensional patterns
• Ionic compounds generally have high
melting points
• Large attractive forces result in very
stable structures
B. Properties of Ionic Compounds
• Ionic compounds can conduct an electric
current when melted or dissolved in
water
• When ionic compounds are dissolved
in water the crystalline structure breaks
down so ions are able to move freely
which results in conductivity
Ch. 7 – Ionic and Metallic Bonding
III. Bonding in Metals
(p. 201 – 203)
I
II
III
IV
A. Metallic Bonding
Metallic bonds: Consist entirely of metal
atoms.
• Bonding is due to valence electrons which
are delocalized throughout the entire solid
• The metal is held together by the strong
forces of attraction between the positive
nuclei and the delocalized electrons.
B. Metals
• Metals are good conductors of heat and
electricity because the valence electrons
are able to flow freely
• Valence electrons of metals can be thought
of as a ‘sea of electrons’, very mobile
C. Metallic Bond
Metallic Bonding - “Electron Sea”
D. Metallic Properties
• Have luster, are ductile and malleable
• Luster = shine
• Ductile = ability to be drawn into wires
• Malleable = ability to be shaped,
pounded, etc
D. Metallic Properties
• Properties can be explained by the mobility
of electrons in metals
• When subjected to pressure , the
cations easily slide past each other like
a ball bearing immersed in oil.
Covalent Bonding
What is a covalent bond?
• A chemical bond that results from the sharing
of electrons, to form a stable octet or duet
(Hydrogen only needs 2 to be stable)
• Molecule = two or more atoms
that are held together by
covalent bonds
H2O
• Majority of covalent bonds form between
nonmetals (CLOSE together on periodic table)
Covalent Rule 1: Nonmetals Bound
to Nonmetals
• Since nonmetals have more than one oxidation
number, there will always be more than one
compound produced
– Therefore we have to have a distinct name for each
compound
– To do this we use a prefix to indicate how many atoms of
each element are present
One – mono-
Five – penta-
Two – di-
Six – hexa-
Three – tri-
Seven – hepta-
Four – tetra-
Eight – octa-
Nine – nona
Ten – deca
Covalent Rule 1: Nonmetals Bound
to Nonmetals
Prefix + Nonmetal prefix + (root of nonmetal) -ide
• Using prefixes
– The prefix mono- is only used on the second
element
• Ex: PF3 is named phosphorus trifluoride
– If two vowels are adjacent, leave them
• Ex: NI3 is named nitrogen triiodide
– In the case of monoxide only, drop one “o”
– Cannot reduce subscripts on covalent Naming!
Covalent Rule 1: Nonmetals Bound
to Nonmetals
• Ex 1: What is the name of P2S3?
– diphosphorus trisulfide
• Ex 2: What is the name of As7I3?
– heptaarsenic triiodide
• Ex 3: What is the chemical formula of
dihydrogen monoxide?
– H2O
• Ex 4: What is the chemical formula of
dinitrogen pentaoxide?
– N2O5
Covalent Bonding Formation
• Diatomic molecule
– molecule containing the same two atoms
• Some elements always exist this way
because they are more stable than the
individual atoms
Cl2
B. Diatomic Elements
• The Seven Diatomic Elements
Br2 I2 N2 Cl2 H2 O2 F2
H
N O F
Cl
Br
I
Bonds in all the
polyatomic ions
and diatomics
are all covalent
bonds
Single Covalent Bonds
Two atoms held together by a sharing
of one pair of electrons are joined
together by a single covalent bond.
Single Covalent Bonds
An electron dot structure represents the
shared pair of electrons of the covalent bond
by two dots.
A structural formula represents the covalent
bonds by dashes and shows the
arrangement of covalently bonded atoms
Single Covalent Bonds
A pair of valence electrons that is not
shared between atoms is called an
unshared pair, also known as a lone pair
of a nonbonding pair.
Lone pair
Double and Triple Covalent Bonds
Atoms form double or triple covalent
bonds if they can attain a noble gas
structure by sharing two or three pairs of
electrons.
A double bond involves sharing two pairs
of electrons.
A triple bond involves sharing three pairs
of electrons.
Double and Triple Covalent Bonds
Molecular Structure
Lewis Diagrams
(p. 220 – 229)
I
II
III
Drawing Lewis Diagrams
1. Arrange atoms
•
•
•
Singular atom is usually in the center (often
Carbon)
If not Carbon, least e- neg atom is in center
Hydrogen is always terminal
2. Find total # of e- available to bond
(valence e- )
3. Place a pair of electrons between central
atom and each terminal atom
Drawing Lewis Diagrams
4. Place remaining electrons in pairs around
terminal atoms (except H) to satisfy octet
rule
• Any remaining pairs are assigned to
central atom
5. Determine whether or not central atom
satisfies octet
• If not, convert one or more lone pairs
from terminal atoms to double or triple
bonds
• Certain atoms can be exceptions to
octet rule – H, Be, B, S, P, Xe
Drawing Lewis Diagrams
CF4
1 C × 4e- = 4e4 F × 7e- = 28e32e- 8e24e-
F
F C F
F
Drawing Lewis Diagrams
CO2
1 C × 4e- = 4e2 O × 6e- = 12e16e- 4e12e-
O C O
Polyatomic Ions
To find total # of valence e-:
 Add 1e- for each negative charge
 Subtract 1e- for each positive
charge
Place brackets around the ion and
label the charge
Polyatomic Ions
ClO41 Cl × 7e- = 7e4 O × 6e- = 24e31e+ 1e32e- 8e24e-
O
O Cl O
O
Octet Rule
Exceptions:
F
F
F
B
F
 Boron & Beryllium get 6 & 4 valence e
H
O
H
respectively
F SF F
 Expanded octet  more than 8
valence
Fe (e.g. S, P,FXe)
 Hydrogen  2 valence e-
-
-
Resonance Structures



Molecules that can’t be correctly
represented by a single Lewis
diagram
Actual structure is an average of all
the possibilities
Show all possible structures
separated by double-headed arrows
C. Resonance Structures

SO3
O
O S O
O
O S O
O
O S O
Bond Polarity
• Most bonds are a
blend of ionic and
covalent
characteristics.
• Difference in
electronegativity
determines bond
type.
Bond Polarity
• Electronegativity
– Attraction an atom has for a shared pair of
electrons.
– higher e-neg atom  – lower e-neg atom +
Electronegativity
Difference
• If the difference in electronegativities is between:
– 1.7 to 4.0: Ionic
– Greater than 0.3 & less than 1.7: Polar Covalent
– 0.0 to 0.3: Non-Polar Covalent
The type of bond can usually be calculated by
finding the difference in electronegativity of
the two atoms that are bonded.
Compound
F2 or F-F
CF4
LiF or Li-F
Electronegativity
4.0 - 4.0 = 0 4.0 - 2.5 = 1.5 4.0 - 1.0 = 3.0
Difference
Non-polar (strong) Polar Ionic (nonType of Bond
covalent
covalent
covalent)
Bond Polarity
• Nonpolar Covalent Bond
– e- are shared equally
– symmetrical e- density
– usually identical atoms
• Ex: H2 or Cl2
Bond Polarity
• Polar Covalent Bond
– e- are shared unequally
– asymmetrical e- density
– results in partial charges (dipole)
• Ex: H2O
+


Ch. 8 – Molecular Structure
Molecular
Geometry
(p. 232 – 236)
I
II
III
VSEPR Theory
Valence Shell Electron Pair
Repulsion Theory
Electron pairs orient themselves in
order to minimize repulsive forces
VSEPR Theory
Types of e- Pairs
 Bonding pairs – form bonds
 Lone pairs – nonbonding e Total e- pairs– bonding + lone pairs
Lone pairs repel
more strongly than bonding
pairs!!!
A. VSEPR Theory
Lone pairs reduce the bond angle
between atoms
Bond Angle
Determining Molecular Shape
Draw the Lewis Diagram
Tally up e- pairs on central atom
(bonds + lone pairs)
 double/triple bonds = ONE pair
Shape is determined by the # of
bonding pairs and lone pairs
Common Molecular Shapes
2 total → Electronic Geometry = linear
2 bond
0 lone
BeH2
LINEAR
180°
Common Molecular Shapes
3 total → Electronic Geometry =
trigonal planar
3 bond
0 lone
BF3
TRIGONAL PLANAR
120°
Molecular Polarity
 Polar
molecule = one end slightly +
and one end slightly –
 Molecule with 2 poles = dipolar
molecule or dipole
Molecular Polarity
 Shape,
symmetry and bond
polarity determines molecular
polarity
 H – O bond is polar and water
is asymmetrical, so H2O is
polar
 C – Cl bond is polar, but CCl4
is symmetrical, so molecule is
nonpolar
Molecular Polarity
 Identify
each molecule as polar or
nonpolar
• O2
Nonpolar bonds → nonpolar
• CS2 Nonpolar bonds → Linear → nonpolar
Polar
bonds
→
Tetrahedral
→
nonpolar
• CF4
• H2O Polar bonds → Bent → polar
I. Intermolecular Forces
Intramolecular and
Intermolecular Forces
 Covalent
& Ionic bonds - the force that
holds atoms together making molecules.
• These are intramolecular forces.
 There are also forces that cause
molecules to attract each other. These
are called intermolecular forces.
Definition of IMF
 Attractive
forces between molecules.
 Much
weaker than
chemical bonds
within molecules.
 a.k.a.
van der Waals forces
Types of IMF