Chemical Bonds: The Formation
of Compounds from Atoms
Dr. Bixler-Zalesinsky
11.1
PERIODIC TRENDS IN ATOMIC
PROPERTIES
Metals and Nonmetals (review)
Ionization energy is the energy required to remove an
electron; corresponds to their charge
Atomic Radii increase going
down a group and decrease
across a period
Valence Electrons
LEWIS DOT DIAGRAMS
Valence Electrons & Per. Table
Lewis Structures of an atom shows the
valence electrons (ones involved in bonding)
Octet Rule
• Every atom aspires to have eight electrons
in its outermost shell (2 s electrons and 6 p
electrons just like the noble gases)
• They must borrow (covalent molecules),
release or accept (ionic compounds)
electrons to get to the eight.
Types of bonding
Ions and
BONDING
Ionic bonding occurs between
metals (cations) and nonmetals
(anions)
The nonmetal accepts the
electron(s) and the metal
donates the electron(s)
ionic bond is the attractive
between oppositely charged
ions
Form large crystals; our formulas
are the smallest whole number
ratios not the true number of
atoms
The charges must cancel out and equal zero
to form stable compounds
If you have a +2 ion then you need either two
-1 ions or one -2 ion
Link to
Video clip on Ionic Bonding (1:39)
Bonding
COVALENT
V.
IONIC
11.5 Covalent Bonding:
Sharing Electrons
• Covalent bonding occurs between two
nonmetal atoms
• Electrons are shared between two atoms
11.7 Lewis Structures of Molecules
1.
2.
3.
4.
5.
Find the number of valence electrons for each element in the
structure
Multiply the number of valence electrons times the number of atoms
you have of that element
Determine which element can make the most bonds and put it in the
center and attach the other elements to it
Make each atom have 8 valence electrons around it.
Add up the number of electrons you used in the structure. This
number must match the total number of electrons you started with
Number of Valence Electrons:
Page 151 in textbook
1 2
H
2.1
Li
1.0
Na
0.9
K
08
Rb
0.8
Cs
0.7
Fr
0.7
Na
1.5
Mg
1.2
Ca
1.0
Sr
1.0
Ba
0.9
Ra
0.9
3 4 5 6 7
B
2.0
C
2.5
N
3.0
Al
1.5
Si
1.8
P
2.1
8
He
O
3.5
F
4.0
S
2.5
Cl
3.0
Ar
Br
2.8
Kr
Se
2.4
I
2.5
Ne
Xe
Rn
3
4 3 2 1 0 (number of bonds each can make)
S S/D/T S/T S/D S 0 (types of bonds
s=single, D= double,
T= triple)
Ex. Write the Lewis Dot Diagrams
for the following molecules
•
•
•
•
•
I2
H2O
FCl
CF4
NBr3
Molecule (Covalent) Nomenclature
• Naming: These binary inorganic molecules are
named using prefixes like mono-, di-, tri-, tetra-,
penta-, hexa-, hepta-, octa-, nona-, deca-,
• The first element only gets a prefix if there is
more than one of it; otherwise, the element name
remains the same.
• The second element ALWAYS gets a prefix and the
ending changes to –ide.
• Ex. CO is carbon monoxide (two words not
capitalized)
HW p. 164 # 2 a – f
Write the question and answer!
Multiple Bonds
Double and Triple Bonds
Knowing when NOT to use them is as
important as understanding when to use
them!
Multiple Bonds
• Some times using the correct number of
electrons will not give you a full octet.
When this happens:
• 1st double check your math and counting
• 2nd see if the atoms involved can make a
double or triple bond
Double Bonds
• O, S, and C can make double bonds with
each other than themselves but no others!
• Double bond is 4 electrons in a bond
• Symbolized by an = sign
• Take a look at CO2
Triple Bonds
• P, N, and C can make a triple bond with each
other or themselves
• A triple bond is 6 electrons in a bond
• The symbol for a triple bond is =
• Let’s try N2
Molecular Geometry
VSEPR Theory and Application
Structural Formula
• Shows how elements of a molecule are
connected to each other
VSEPR
•
•
•
•
•
•
V = valence
S = shell
E = electron
P = pair
R = repulsion
Electrons will arrange themselves as far
apart from one another as possible
• Unbonded pairs take up more room than
bonded ones
3-D
Hybridized orbitals, shapes, and
decision tree
Linear Shape
Bent Shape
Trigonal planar
Pyramidal
Tetrahedral
Video Review (3:21)
VSEPR
sp hybridization
Path to hybridization
Sp3 hybridization
Video Review (1:36)
HYBRIDIZATION
VSEPR Theory of Molecular
Geometry
# of
atoms
Central
Atom
Shape
Bond
angle
Example
VSEPR Theory of Molecular
Geometry
# of
atoms
2
Central
atom
None
Shape
Linear
180
HF
3
Any
Linear
180
CO2
3
S or O
Bent
105
H2O
4
B
Trig.
Planar
120
BCl3
4 (3-D)
P or N
Pyramidal
107
NH3
5 (3-D)
C or Si
Tetrahedral
109.5
CH4
Bond Angle
Example
(in degrees)
Shape Decision Tree
How many atoms?
2 = Linear
3
• No unshared pairs = Linear
• 2 unbonded pairs = bent
4
• No unshared pairs = Trigonal Planar
• 1 unbonded pair = Pyramidal
5 = Tetrahedral
Polar Covalent
v.
Nonpolar Covalent
Polar and Nonpolar Covalent
Bonds
• If they are shared equally they are said to be nonpolar bonds if they
are not equally shared they are said to be polar bonds
• Sharing of electrons has to do with the pull of one element compared
to the other element sharing the electron pair. This pulling is called
electronegativity (eneg) which increases across the period and up
the group
• The larger the electronegativity the greater the time the electrons
spend with the more electronegative atom giving it a slightly positive
charge and because of this imbalance we call this a polar molecule
• If the eneg difference lies between 0.5 to 1.6 it is a polar bond
Polar or Nonpolar Molecules
1.
2.
Determine the shape of the molecule
Determine how many polar bonds there are in
the molecule
3. If there are NO polar bonds the molecule must
be NONpolar.
4. If there is exactly one polar bond, the molecule
is polar.
5. If there is more than one polar bond, follow the
chart below.
Molecules with more than one
polar bond (assuming polarity is equal)
Shape
# of polar bonds
Molecular polarity
Linear
2
Nonpolar
Bent
2
Polar
Trig Planar
2
Polar
Trig Planar
3
Nonpolar
Pyramidal
2
Polar
Pyramidal
3
Polar
Tetrahedral
2
Nonpolar
Tetrahedral
3
Polar
Tetrahedral
4
Nonpolar
VSEPR Theory of Molecular
Geometry
# of
atoms
Central
Atom
Shape
Bond
angle
Example
Polarity of Molecules with more
than one polar bond
Shape
# of polar bonds Molecular
Polarity