Electron Dots
and
VSEPR Theory
(mostly Chapter 9)
Metallic Bonding
• In metallic bonding the valence electrons are
shared between all the atoms in a positive
metal crystal.
 delocalized “sea” of electrons
 metallic bonded materials have
good thermal and electrical conduction.
Ionic Bonds
Occur when the nonmetal takes one or more
electrons away from a metal.
 The nonmetal becomes a negative ion
 metal becomes a positive ion.
The atoms are held together by their opposite
charges.
Ionic Bond Strength
Strength of crystal lattice depends on two
factors, size and charge transferred.
Smaller atoms have stronger ionic bonds.
Ex: NaF is stronger than NaCl
Atoms transferring more electrons are
stronger. Ex: MgCl2 is stronger than NaCl.
2 e- transferred
1 e- transferred
Covalent Bond
“the bonds of nature”
• Shared valence electrons
• Complete outer energy levels
• Molecule has 2 or more nonmetal atoms
covalently bond
– Carbohydrates, proteins, fats, DNA,
stupendous seven (H2, N2, O2, F2, Cl2, Br2, I2)
How do covalent bonds form?
distance is too
great
e- & e-  repulsive
p+ & e- attractive
p+ & p+  repulsive
repul. = attract
e- & e-  repulsive
Attractive forces balance the repulsive forces
Electronegativity
• Electronegativity is the
ability of atoms in a molecule
to attract electrons to
themselves.
• On the periodic table,
electronegativity increases as
you go…
– from left to right across a row.
– from the bottom to the top of a
column.
What types of bonds are they?
MgO, water, Calcium Carbide, Potassium Oxide,
Nitrogen trihydride
Electronegativity and Bond Type
Find the difference in 2 atoms’
electronegativies to predict bond type…
• Ionic Bonds: 1.7 or greater
• Polar Covalent Bonds: <1.7 and >0.2
• Pure or Nonpolar Covalent bonds: <0.2
Bonding Capacity
Atom
Carbon
Nitrogen
Oxygen
Halogens
Hydrogen
Number of
Valence
Electrons
Number of Bonding
Bonding
Capacity
Electrons
Electronegativity Table
Drawing Lewis Dot Structures
1. Count the valence electrons.
2. Predict the location of the atoms
a. Hydrogen is a terminal atom
b. The central atom has the smallest electronegativity.
3. Draw a pair of electrons between the central
atom and the surrounding atoms.
4. Use the remaining electrons to complete the
octets of each atom. If there are electrons left
over, place them on the central atom.
5. If the central atom does not have a complete
octet then try double or triple bonds.
a. If the atom has 1, 2, or 3 valence electrons, it
doesn’t require an octet.
STEP 1: count the total # of
valence e- for all atoms
involved in the bonding
Carbon: 1 carbon with
4 valence
electrons (1x4) = 4
CCl4
CCl4
4+28
=32
Chlorine: 4 chlorine
with 7 valence
electrons (4x7) = 28
STEP 2–place the single atom in the
center and other atoms around it
evenly spaced
CCl4
4+28
=32 e-
Cl
Cl C Cl
Cl
STEP 3: place the electrons in pairs
between the central atom
and each non-central atom
CCl4
4+28
=32
-8
=24
Cl
Cl C Cl
Cl
STEP 4: place the remaining electrons
around the non-central atom until each
has 8 electrons (H atoms have only 2e-)
CCl4
4+28
=32
-8
=24
-24
=0
Cl
Cl
C
Cl
Cl
Step 5: If you run out of electrons before
the central atom has an octet, form
multiple bonds until it does.
Example: HCN
Hydrogen- 1 electron
Carbon- 4 electrons
Nitrogen- 5 electrons
H:C:N
..
H:C:N:
..
TOTAL is 1+4+5 = 10 e-
H:C:::N:
Drawing Lewis Dot Structures
Draw Lewis Dot Structures for:
PH3
H2S
HCl
CCl4
SiH4
CH2Cl2
Draw Lewis Dot Structures
Cl2
NF3
CS2
BH3
CH4
SCl2
C2H6
BF3
(stop)
Covalent Bond Strength
• Based on proximity (closeness), also called
“bond length”
 Influenced by atom size and number of
shared electrons
 Smaller is stronger
F2 is stronger than Cl2 is stronger than Br2
F2:
O2
N2
1.43 x 10-10 m
1.21 x 10-10 m
1.10 x 10-10 m
single bond
double bond
triple bond
Bonding Orbitals
• When atoms bond together, their valence
shell electron orbitals overlap
• Overlapping electron orbitals create a
bonding orbital  an area with a high
probability of finding an electron
Types of Bonds
• When atoms form a molecule, their orbitals
can form different types of bonds:
–Sigma Bonds (σ)
•Orbitals overlap head-to-head
•Form first, there’s only 1
–Pi Bonds (π)
•Orbitals overlap side-to-side
•Form after sigma bonds
Every molecule has one sigma bond, but all subsequent
bonds between the same two atoms must have a different
way of connecting so they use pi bonds!
Multiple Covalent Bonds – Double
6 valence
electrons
6 valence
electrons
12 valence
electrons
1 sigma bond
Octet satisfied
More stable
and stronger
1 pi bond
(lines represent
bonded pairs of e-)
Multiple Covalent Bonds – Triple
5 valence
electrons
5 valence
electrons
10 valence
electrons
1 sigma bond
Octet satisfied
2 pi bonds
More stable
and stronger
Molecular Shapes
• The shape of a
molecule plays an
important role in its
reactivity.
• Look at bonding and
non-bonding electron
pairs
– You can predict the
shape of the molecule!
What Determines the Shape
of a Molecule?
• Electron pairs repel each other.
• Assuming electron pairs are placed as far as
possible from each other, we can predict the
shape of the molecule.
Valence Shell Electron Pair
Repulsion Theory (VSEPR)
“The best arrangement of a given number
of electron domains is the one that
minimizes the repulsions among them.”
Molecular Shape Chart
Formula
BeH2
BF3
CH4
NH3
H2O
Dot
Structure
Name of
Shape
Nonbonding
e- pairs
Bonding
electrons
Polarity
Hybridization
Bond Angle
Molecular Polarity
• Molecules can be polar and non-polar.
• Imagine you are turning over the 3D models on the table.
Are they still the same when you flip them over?
– If yes, then the molecule is non-polar (symmetrical)
Molecular Polarity
• Non-bonding electron
pair = polar
Nonbonding
electron pair
– The free pair pushes
the other atoms away
• Non-polar molecule
has equal pull from the
same atoms
(stop)
Bonding Orbital Hybridization
• Electron orbitals mix to
make a new set of bonding
orbitals (hybrids)
– These have different shapes
than regular atomic orbitals
– Requires energy but the
energy is returned during
bond formation
2s
Bonds can
form here
2p
This occurs to allow
more bonds!
2sp3
new hybridized
orbital
Hybrid Orbitals
L
Consider beryllium:
• In its ground state, it
would not be able to
form bonds because it
has no singly-occupied
orbitals.
Hybrid Orbitals
L
But by promoting an
electron from the 2s to
the 2p orbital, it can
now form two bonds.
J J
This new hybridized
orbital is called 2sp
2s
2sp orbitals
2s
2p
2sp
2s
2p
2sp2
Group 2A
elements
make sp
hybridized
orbitals
Group 3A
elements
make sp2
hybridized
orbitals
Group 4A elements have 4 valence electrons
- need 4 bonds to make an octet
- they will have sp3 hybridization.
2s2
2p2
2sp3
Endothermic and
Exothermic Reactions
• Endothermic Reactions – the energy
needed to break the bonds is greater than
the energy that is released, energy is used
– They feel cool
• Exothermic Reactions – the energy
needed to break the bonds is less than the
energy released, energy given off
– They feel warm
Questions
• Why do some solids dissolve in water
but others do not?
• Why are some substances gases at
room temperature, but others are liquid
or solid?
• What gives metals the ability to conduct
electricity, what makes non-metals
brittle?
• The answers have to do with …
Intermolecular forces
Intermolecular forces
(also called Van der Waal’s forces)
2 types of attraction in molecules:
Intramolecular bonds: (Covalent and ionic)
attraction between atoms in a molecule
Intermolecular forces (IMF): the attraction
between molecules
– 1) dipole-dipole
– 2) hydrogen bonding
– 3) London forces
Dipole - Dipole attractions
• Dipoles: a separation of charge
• This happens in both ionic and
polar covalent bonds
+
–
H
Cl
• Oppositely charged dipoles (+δ and –δ) are
attracted to each other in a molecule
+ –
+ –
Hydrogen Bonding
H-bonding is a special type of dipole - dipole
attraction that is very strong (5x stronger)
– Happens when N, O, or F are bonded to H
– Due to the high electronegativity difference
between the H and the other atom
– Compounds containing these bonds are
important in biological systems (special!)
London forces
• Named after Fritz London, sometimes called
dispersion forces
• London forces are due to small dipoles that exist
in non-polar molecules
• Random movement of electrons can sometimes
form temporary dipoles
• The resulting tiny dipoles cause attractions
between atoms/molecules
This is how non-polar molecules
can form solids and liquids!
London forces
Instantaneous dipole:
Induced dipole:
Sometimes the random
arrangement of electrons
forms tiny dipoles
A random dipole forms in one
atom or molecule, inducing a
dipole in the other
(stop)
IMF Strength and Molar Mass
• The size of a molecule (molar mass) affects
the strength of intermolecular forces (IMFs)
• Larger size = stronger forces
– Because the large molecule has more area and
electrons available for intermolecular attractions
such as London Forces
– (this is opposite of covalent bond strength)
Stronger IMFs
Weaker IMFs
IMF Strength and Molar Mass
• Consider the halogens (group 7A)
as an example
• F2 and Cl2 are gases, Br2 is liquid,
I2 is a solid
– Liquids and solids form when IMFs
are stronger
– Since they are further down the
group, the atoms are bigger
– Larger mass = stronger IMFs
Boiling Point and IMFs
• Boiling (liquid  gas) occurs when there is
enough energy to overcome
intermolecular attractions
• Boiling point tends to increase down a
group, as size of atoms in molecules
increases
Predicted and actual boiling points
What about these?
(such as H2O)
Boiling point
100
Group 4
Group 5
50
0
-50
Group 6
-100
-150
-200
Group 7
2
3
4
Period
5
This is because the
larger atoms/molecules
have stronger IMFs
so
it takes more energy to
break those attractions
 higher boiling point!
Hydrogen Bonds and Boiling Point
• H2O, HF, and NH3 have particularly high
boiling points
• This is because of hydrogen bonds!
• Because they are the strongest IMF, they
require more heat energy to break the
attraction  higher boiling points
Boiling point
Predicted and actual boiling points
100
Group 4
Group 5
50
0
-50
Group 6
-100
-150
-200
Group 7
2
3
4
Period
5
(end)
***Hints for IMF Lab***
• Activity 1, Question 3 asks you to draw the
Lewis Dot structures for acetone and
ethanol.
• Here are their shapes to help you…
ethanol
acetone