Valence electrons are the outer shell electrons of an
atom. The valence electrons are the electrons that
particpate in chemical bonding.
Group
e- configuration
# of valence e-
I
ns1
1
II
ns2
2
III
ns2np1
3
IV
ns2np2
4
V
ns2np3
5
VI
ns2np4
6
VII
ns2np5
7
9.1
Lewis Dot Symbols for the Representative Elements &
Noble Gases
9.1
Trends in the Periodic Table
 The elements have a regular (periodic) recurrence
of physical and chemical properties
 The horizontal rows are called periods
 The vertical columns are called groups
Atomic Size
 Covalent Atomic Radius – half of the distance between
the nuclei of two atoms in a homonuclear diatomic
molecule
 Group trend – atomic size increases as we move down a
group
 The farther away the e- from the nucleus, the less strongly
they are held
 Periodic trend – atomic size decreases as we move from
left to right
 Nuclear charge increases as we move left to right, pulling
the e- closer to the nucleus
Ionic Size
 Cations are always smaller than neutral atoms because
they have less e Loss of outer shell e- results in increased attraction by
the nucleus for the remaining e-
 Anions are always larger than neutral atoms because
they have more e Addition of outer shell e- results in less attraction to the
protons of the nucleus
Ionization Energy
 The energy required to remove an electron from a gaseous
atom
Na (g) + Energy  Na+(g) + e Group trend – Ionization energy decreases as we move
down a group because the size of the atom is increasing,
with more orbitals of e-, allowing the outermost e- to be
easily removed
 Periodic trend – Ionization energy increases from left to
right. Nuclear charge is increasing so more energy is
required to remove e-
Electron Affinity
 The energy change (release) that accompanies the
addition of an electron to a gaseous atom
F (g) + e-  F-(g) + Energy
 Group trend – Electron affinity generally decreases with
increasing atomic size (releases less energy as you go
down)
 Periodic trend – Generally increases from left to right
because atoms become smaller due to increased nuclear
charge (releases more energy as you go left to right)
Electronegativity
 The tendency for an atom to attract electrons to itself.
It is affected by the distance from the atom’s valence
electrons to the nucleus.
 Group trend – Electronegativity usually decreases as
you move down a group
 Periodic trend – Electronegativity usually increases as
you move across a period
 Do the Noble Gases have electronegativity?
 Why do substances bond?
 More stability
 Atoms want to achieve a lower energy state
Chemical bonds: an attempt to fill electron shells
1.
2.
3.
Ionic bonds –
Covalent bonds –
Metallic bonds
Ionic Bonding
 Between a metal and a non-metal with very different
electronegativity.
 Metals lose electrons becoming a cations, while nonmetals gain electrons becoming anions.
 An ionic bond is an electrostatic attraction between
the oppositely charged ions.
Ionic Bonds: One Big Greedy Thief
Dog!
. 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.
The Ionic Bond
Li + F
Li+ F -
1s22s1 1s22s22p5
[He]
1s1s2 [2Ne]
2s22p6
e- +
Li+ +
Li
Li+ + e-
F
F -
F -
Li+ F -
9.2
The Ionic Bond
Li + F
Li+ F -
1s22s1 1s22s22p5
[He]
1s1s2 [2Ne]
2s22p6
e- +
Li+ +
Li
Li+ + e-
F
F -
F -
Li+ F -
9.2
Ionic Structures
 In an ionic compound (solid), the ions are packed together
into a repeating array called a crystal lattice.
 The simplest arrangement is one in which the spheres in the
base are packed side by side. Opposite charges are attracted
to each other.
 Its called simple cubic packing (NaCl is an example)
 Ionic formulas are always Empirical Formulas (simplest)
Properties of Ionic Compounds
•All ionic compounds form crystals.
• Ionic
compounds tend to have high melting and
boiling points.
To break the positive and negative charges apart, it takes a huge
amount of energy.
• Ionic compounds are very hard and very brittle.
Again, this is because of the way that they're held together –
strong attraction of oppositely charged ions. These ions simply
don't move around - so they don't bend at all. This also explains
the brittleness of ionic compounds. If we give a big crystal a
strong enough whack with a hammer, we usually end up using
so much energy to break the crystal that the crystal doesn't
break in just one spot, but in a whole bunch of places. Instead of
a clean break, it shatters.
• Ionic
Compounds are poor conductors of
electricity in the solid state.
Ions are held tightly together and cannot move. Therefore ions
cannot conduct electricity
• Ionic compounds conduct electricity when
molten (melted).
Ions are mobile and can therefore conduct electricity.
• Ionic
compounds conduct electricity when
aqueous (dissolved in water).
If we take a salt and dissolve it in water, the water molecules
pull the positive and negative ions apart from each other.
Instead of the ions being right next to each other, they are able
to move around in the water and conduct electricity.
View here
Covalent Bonding
COVALENT BOND
bond formed by the
sharing of electrons
Covalent Bond
 Between nonmetallic elements of similar
electronegativity.
 Formed by sharing electron pairs
 Stable non-ionizing particles, they are not conductors
at any state
 Examples; O2, CO2, C2H6, H2O, SiC
Covalent bonds- Two atoms share one or more pairs of outer-shell
electrons.
Fluorine Atom
Fluorine Atom
Fluorine Molecule (F2)
A covalent bond is a chemical bond in which two or more
electrons are shared by two atoms.
Why should two atoms share electrons?
F
+
7e-
single covalent bond
lone pairs
F F
Dot diagram of F2
F
F F
7e-
8e- 8e-
lone pairs
F
F
lone pairs
single covalent bond
lone pairs
Lewis structure of F2
9.4
Lewis structure of water
H +
O +
H
single covalent bonds
H O H
or
H O
H
2e-8e-2eDouble bond – two atoms share two pairs of electrons
O C O
or
O
O
C
double bonds
- 8e8e- 8ebonds
double
Triple bond – two atoms share three pairs of electrons
N N
triple
bond
8e-8e
or
N
N
triple bond
9.4
Properties of Covalent Compounds
• Covalent compounds tend
to have low melting and
boiling points.
Many simple molecular substances are gases or liquids at room temperature.
• Most molecular compounds are poor conductors of
electricity in all states.
There are no free electrons available to move and conduct an electric
current.
• Covalent compounds are usually much softer than
ionic material.
• Covalent compounds tend to be flammable than ionic
compounds.
• Most molecular compounds are insoluble in water,
but will dissolve in nonpolar organic solvents.
9.4
Drawing Lewis Structures
1. Write the dot diagram for each atom present in the
compound.
2. Take the TWO atoms with the MOST unpaired dots
and join (bond) them together.
3. Add the remaining atoms where they are needed the
most. (starting with remaining atoms with most
unpaired dots)
Try the following:
HOF
N2H4
CH2O
H4CO
HCN
Bonds in all the
polyatomic
ions and
diatomics are
all covalent
bonds
In covalent bonding, one or
more pair of electrons are
shared. However, all
‘sharing’ is NOT the same.
Therefore, we can have
Nonpolar (Pure) Covalent
Bonds and Polar Covalent
Bonds.
NONPOLAR
COVALENT BONDS
when electrons are
shared equally
H2 or Cl2
Nonpolar Covalent bonds- Two atoms equally share one or more pairs of
outer-shell electrons.
Fluorine Atom
Fluorine Atom
Fluorine Molecule (F2)
POLAR COVALENT
BONDS
when electrons are
shared but shared
unequally.
H2O
Polar Covalent Bonds: Unevenly
matched, but willing to share.
- water is a polar molecule because oxygen is more electronegative than hydrogen,
and therefore electrons are pulled closer to oxygen.
Polar covalent bond or polar bond is a covalent bond
with greater electron density around one of the two
atoms (electrons are shared unequally)
electron poor
region
electron rich
region
e- poor
H
d+
H
F
e- rich
F
d-
To determine the type of Bonding
present in a compound, we compare
the electonegativities of the
bonding atoms. The difference in
electronegativity between these
atoms predicts the type of bond
present
The Electronegativities of Common Elements
Electronegativity is the ability of an atom to attract
toward itself the electrons in a chemical bond.
Electronegativity - relative, F is highest
electron poor
region
H
electron rich
region
F
Classification of bonds by difference in electronegativity
Difference
Bond Type
0  0.4
Nonpolar (Pure) Covalent
> 1.7
0.4 < and ≤1.7
Ionic
Polar Covalent
Creates a “Bonding Continuum”
Increasing difference in electronegativity
Nonpolar (Pure)Covalent
Polar Covalent
Ionic
e- equally
partial transfer of e-
transfer e-
share
Unequal sharing
Classify the following bonds as ionic, polar covalent,
or covalent: The bond in CsCl; the bond in H2S; and
the NN bond in H2NNH2.
Cs – 0.7
Cl – 3.0
3.0 – 0.7 = 2.3
H – 2.1
S – 2.5
2.5 – 2.1 = 0.4
N – 3.0
N – 3.0
3.0 – 3.0 = 0
Ionic
Nonpolar Covalent
NonPolar Covalent
9.5
Predicting Molecular Geometry
1. Draw Lewis structure for molecule.
2. Count number of lone pairs on the central
atom and number of atoms bonded to the
central atom.
3. Use VSEPR to predict the geometry of the
molecule.
Valence shell electron pair repulsion (VSEPR) model:
Predict the geometry of the molecule from the electrostatic
repulsions between the electron (bonding and nonbonding) pairs.
Class
# of atoms
bonded to
central atom
# lone
pairs on
central atom
Arrangement of
electron pairs
Molecular
Geometry
AB2
2
0
linear
linear
B
B
0 lone pairs on central atom
Cl
Be
Cl
2 atoms bonded to central atom
10.1
VSEPR
Class
# of atoms
bonded to
central atom
# lone
pairs on
central atom
Arrangement of
electron pairs
Molecular
Geometry
AB2
2
0
linear
linear
0
trigonal
planar
trigonal
planar
AB3
3
VSEPR
Class
# of atoms
bonded to
central atom
# lone
pairs on
central atom
Arrangement of
electron pairs
Molecular
Geometry
AB2
2
0
linear
linear
trigonal
planar
tetrahedral
AB3
3
0
trigonal
planar
AB4
4
0
tetrahedral
VSEPR
Class
# of atoms
bonded to
central atom
# lone
pairs on
central atom
AB3
3
0
AB2E
2
1
Arrangement of
electron pairs
Molecular
Geometry
trigonal
planar
trigonal
planar
trigonal
planar
bent
VSEPR
Class
# of atoms
bonded to
central atom
# lone
pairs on
central atom
AB4
4
0
tetrahedral
tetrahedral
AB3E
3
1
tetrahedral
trigonal
pyramidal
Arrangement of
electron pairs
Molecular
Geometry
VSEPR
Class
# of atoms
bonded to
central atom
# lone
pairs on
central atom
AB4
4
0
Arrangement of
electron pairs
Molecular
Geometry
tetrahedral
tetrahedral
AB3E
3
1
tetrahedral
trigonal
pyramidal
AB2E2
2
2
tetrahedral
bent
O
H
H
bonding-pair vs. bonding
pair repulsion
< lone-pair vs. bonding < lone-pair vs. lone pair
pair repulsion
repulsion
Polarity and shape
 The shape of the molecule directly influences the
overall polarity of the molecule.
 If there is symmetry the charges cancel each other out,
making the molecule non-polar
 If there is no symmetry, then its polar
 Polar bonds do not guarantee a polar molecule
 Ex: CCl4 and CO2 both have polar bonds, but
both are NON-POLAR molecules. They have a
dipole moment of zero (no net dipole or
‘resultant vector’.
 The greater the dipole moment, the more polar
the molecule
Why is molecular polarity
important?
 Polar molecules have higher melting and
boiling points (for example the BP of HF is
19.5° C, and the BP of F2 is –188° C).
 Polar solvents dissolve ionic and polar
molecules more efficiently than non-polar
solvents
Dipole Moments and Polar Molecules
electron poor
region
electron rich
region
H
F
d+
d-
10.2
‘polar molecule’
‘polar molecule’
Which of the following molecules have a dipole moment?
H2O, CO2, SO2, and CH4
O
S
dipole moment
polar molecule
dipole moment
polar molecule
H
O
C
O
no dipole moment
nonpolar molecule
H
C
H
H
no dipole moment
nonpolar molecule
Does BF3 have a
dipole moment?
Symmetrical Molecule – No Resultant Dipole
NONPOLAR MOLECULE
The bent shape creates an
overall positive end and negative end
of the molecule = POLAR
The symetry of the molecule
Cancels out the “charges”
Making this NON-POLAR
No overall DIPOLE
Examples to Try
 Determine whether the following molecules will
be polar or non-polar molecules.
 SI2
 CH3F
 AsI3
 H2O2
Summary of Polarity of
Molecules
 Linear:
 When two atoms attached to central atom are the
same, the molecule will be Non-Polar (CO2)
 When the two atoms are different the dipoles will
not cancel, and the molecule will be Polar (HCN)
 Bent:
 The dipoles created from this molecule will not
cancel creating a net dipole moment and the
molecule will be Polar (H2O)
Summary of Polarity of
Molecules
 Pyramidal:
 The dipoles created from this molecule will not
cancel creating a net dipole and the molecule will
be Polar (NH3)
 Trigonal Planar:
 When the three atoms attached to central atom are
the same, the molecule will be Non-Polar (BF3)
 When the three atoms are different the dipoles will
not cancel, resulting in a net dipole, and the
molecule will be Polar (CH2O)
Tetrahedral
 When the four atoms
attached to the central
atom are the same the
molecule will be NonPolar
 When three atoms are
the same, and one is
different, the dipoles
will not cancel, and
the molecule will be
Polar
METALLIC BOND
bond found in metals;
holds metal atoms
together very strongly
Metallic Bond
 Formed between atoms of metallic elements
 Electron cloud around atoms
 Good conductors at all states, lustrous, very high
melting points
 Examples; Na, Fe, Al, Au, Co
Metallic Bonds: Mellow dogs with
plenty of bones to go around.
Ionic Bond, A Sea of Electrons
This ‘sea of mobile electrons’ allows metals
to conduct an electricity current – electrons
are free to travel.
‘Sea of Electrons’ is also the reason why most
metals are malleable and ductile. As the
electrons are free to move, it allows the metal
to bend without breaking.