Covalent Compounds and
Intermolecular Forces
Bonds
• Chemical bonds are __________ forces
• They act between atoms ________ a
molecule
Why does bonding occur?
• Bonding occurs to maximize stability of
the atoms involved.
• More stable = LOWER potential energy
Bond types
• Dependent on the difference between the
electronegativities of the elements
involved in bond
▫ Electronegativity
 ______________________________
 Highest found in small non-metals
 ______________ are not ranked
Bond Types (cont)
• To Determine Type
▫ Subtract the electronegativities
▫ If difference is
 Zero – bond is ________________
 0.1-1.67- bond is ________________
 Greater than 1.67- __________
Bonding is all about the electrons!
• Bond type tells us what will happen to the
electron(s)
• Octet rule will give us an idea of how
many electrons will be involved
▫ Have to look at valance electrons
▫ Remember most atoms are stable with 8
▫ Common exceptions
 H, He, Li, and Be can be stable with _____
 B is stable with ______
 Elements with d orbitals (can have more
than 8)
Valence electrons
Covalent bonding
• Covalent bonding involves
the ___________ of
valence electrons.
• Atoms can share electrons
in order to have full
valence shell around them.
(usually 8 electrons)
Two Types of Covalent Bonds
• Non-polar
▫ _________ sharing of electrons
▫ Occurs between the same element
bonded together
▫ For example: ______
• Polar
▫ __________ sharing of electrons
▫ Occurs between two different
elements bonded together
▫ For example: _______
Polarity of Bonds
• Polarity indicates how much electrons are pulled
to one end of the molecule or the other.
• Creates two ends (or poles)
▫ One end is slightly negative and the other is
slightly positive
▫ Creates a ______________
This atom has
greater
electronegativity.
+
δ
δ
Labeling a Polar Bond
• Need to draw a vector to show polarity
• Vectors
▫ A vector is an arrow that represents the
strength of something (like a charge),
and the direction in which it is acting.
Vectors for Dipoles
• The vector points in the direction of
the partial ______________
• It has a + on the partial positive end.
Vector of dipole
+
δ
δ
Polarity
• Ionic compounds are very polar.
• They have one atom that’s so strong it can pull
one electron (or more) away from the other
atom. So the electrons are pulled to one end of
the molecule.
• This is a matter of degrees and there really isn’t a
distinct line between polar covalent and ionic.
Covalent Compounds
• _________ Covalent
▫ Form molecules
▫ Example: ________
• ________ Covalent
▫ Form large
interconnected
networks
▫ Example: _________
Naming Binary Covalent
Compounds
• Compound made from two nonmetals (covalent
bond)
▫ For example: H2O, CO, NH3, CH4
• No _________ Involved!!
• Form ___________
Creating Binary Covalent
Compound Formulas
• Can only be created from the name or a
description of how many atoms of each element
there are
Creating Binary Covalent
Compound Formulas (cont)
• For example,
▫ What is the formula for a compound made of 2
boron atoms and three oxygen atoms?
 _____________
Naming Binary Covalent
Compounds
•
•
•
•
Base name with _____________
First element has element name
Second element’s ending is changed to –ide
Prefix is put before each element to designate
how many atoms there are
▫ Mono- is never put in front of the first element
Binary Covalent Compound
Prefixes
•
•
•
•
•
•
•
•
•
•
Mono
Di
Tri
Tetra
Penta
Hexa
Hepta
Octa
Nona
Deca
1
2
3
4
5
6
7
8
9
10
Naming Binary Covalent
Compounds (cont)
• For example:
▫
▫
▫
▫
CO2 is _______________
BF3 is _______________
B2O3 is _______________
P2O5 is _________________
Representing bonds
• There are several ways we represent bonding so
that we can have a visual picture of how the
molecules look.
• The most common is called
▫
▫
▫
▫
Lewis structure
Lewis dot diagram
Electron dot diagram
(they all mean the same thing.)
Drawing Lewis Dots
• Write the symbol.
• Find out how many valence electrons.
• Draw that many dots
▫ 2 Dots max / side
▫ Max of 8 dots!
Examples of Lewis Dots
• Carbon – ______________
C
• Nitrogen – ____________
N
Lewis Structure for Compounds
• Ionic Compounds
▫ Show the Lewis Structure of each ion (including charge)
▫ Sit them next to each other
Lewis Structure for Compounds
• Covalent Compounds
▫ The central atom is usually the one with the lowest
electronegativity (but never ______)
▫ Determine total valence electrons
▫ Move electrons so that each terminal atom has an
octet (but not H who gets 2)
▫ Any extra electrons go on the _____________
Covalent bonding
•Let’s make carbon
tetrachloride, CCl4
•Start with the atoms
C
Cl
Cl
Cl
Cl
Covalent bonding
Double and Triple Bonds
• Atoms can form single, double or triple bonds
with other atoms.
• One example – carbon dioxide
Structural formula
• Electron dot diagrams show all valence electrons
as “dots”.
• We frequently represent BONDED electrons
(shared electrons) as lines. One line is a single
bond, two lines is a double bond, three lines is a
triple bond.
• This is called the structural formula.
• _______________ are still dots.
Examples
Lewis dot diagram
Structural formula
Cl
Cl
Cl C Cl
Cl
Shared pairs are dots
Cl - C - Cl
Cl
Shared pairs are lines, unshared
electrons are still dots.
Limitations of Lewis diagrams
• Work great on paper, but they are drawn in two
dimensions.
• Molecules exist in ________ dimensions.
Electrons repel each other
• In order to understand how molecules behave in
three dimensions, we need to realize that
electrons (or groups of electrons) repel each
other.
• When do we have groups of electrons on a Lewis
diagram?
▫ _______________
▫ ________________
VSEPR
• VSEPR (Valence Shell Electron Pair Repulsion)
Theory states that pairs of electrons repel each
other.
• This allows us to predict the shapes of molecules
in three dimensions.
Rules for VSEPR
• Used for covalently bonded molecules only (all
non-metals)
• “Electron groups” can be bonds or lone pairs.
• Double and triple bonds behave like single
bonds (so they’re really 2 or 3 pairs of electrons,
but they act like 1 group and we count them as 1
group).
Steps for using VSEPR to predict
molecular shape (geometry)
1. Draw the Lewis structure.
2. Count the number of electron groups (on the
central atom).
3. Look at the chart to determine electron group
geometry (first row).
4. Count the number of lone pairs on the central
atom (if none, you’re done) and move down the
chart to name the molecular geometry.
VSEPR Chart
No lone
pairs
1 lone pair
2 lone
pairs
3 lone
pairs
4 lone
pairs
2 E groups
3 E groups
4 E groups
5 E groups
6 E groups
LINEAR
TRIGONAL
PLANAR
TETRAHEDRAL TRIGONAL
OCTAHEDRAL
BIPYRAMIDAL
BENT
TRIGONAL
PYRAMIDAL
SEE-SAW
SQUARE
PYRAMIDAL
BENT
T-SHAPED
PLANAR
SQUARE
PLANAR
LINEAR
T-SHAPED
PLANAR
LINEAR
Linear
Trigonal Planer
Tetrahedral
Let’s try some examples:
 Draw the Lewis structure for methane, CH4.
 How many electron groups are on the central atom?
(electron groups = bonds + lone pairs)
 So the electron group geometry is TETRAHEDRAL.
 Since there are no lone pairs on the central atom,
the molecular geometry is the same as the electron
group geometry: TETRAHEDRAL.
 Now make a model of methane.
 In the Lewis structure, the terminal atoms are 90
degrees apart.
 In the 3-D model, they’re more like 109.5°.
Another example:
 Draw the Lewis structure for water.
 How many electron groups on the central atom?
 Electron group geometry is _______.
 How many lone pairs on the central atom?
 Find the 3D model that matches the electron
group geometry and pull of an atom for each
lone pair.
 Molecular geometry is ____________.
One more example:
 Draw the Lewis structure of COH2
 How many electron groups are on the central atom?
 Electron group geometry is _______.
 How many lone pairs are on the central atom?
 Molecular geometry is ________.
 Choose the best model to represent this.
 How does it compare to MY model of this?
 Does the double bonded atom repel the other
terminal atoms differently than a singly bonded
atom does?
Polar Bond vs Polar Molecule
• Polar Bond
▫ Between two different non-metals
▫ Electronegativity difference creates a pull in the
electrons to one side of bond
▫ Slightly negative and positive end of bond
• Polar Molecule
▫ Must contain at least one polar bond
▫ Polar bonds can cancel each other out
▫ Creates slightly negative and slightly positive end
of a molecule
Deciding if a Molecule is Polar
• Need to draw a vector to show polarity
▫ Vectors of equal strength and in opposite
directions cancel out!
Molecules with 2 Atoms
• Any 2-atom molecule with a polar bond
has a dipole moment.
• The molecule will be polar
Molecules with 3+ atoms.
• Vectors are very important for helping us
determine the polarity of a molecule that
has more than one bond.
• In some molecules, the polar bonds will
cancel each other out making the molecule
non-polar (even though it contains polar
bonds)
• This is bases on the molecules geometry
Bond Energy and Length
• The energy required to break a bond is
called bond energy.
• The distance between the nuclei in a
molecule is known as the bond length.
• Higher bond energy = lower bond length,
in general
▫ Short bonds are harder to break.
▫ Triple bonds are really short
▫ Double bonds are medium length
▫ Single bonds are longer
What is IMF?
• Force between different molecules
BONDS
IMF
IMF’s
• Also referred to as Van Der Waals forces
• ________ than bonds
• Two main types
▫ London dispersion forces
▫ Dipole-dipole
London Dispersion Forces
Very Weak
Occur in all molecule types
Electrons can become temporarily unbalanced
Can create a temporary polarity in a non-polar
molecule
 Opposite poles attract




London Dispersion Forces (cont)
 Forces increase
◦ With increasing molar
mass of compound
 For example: C2H6 has
a greater force than
CH4
◦ With elongated
geometries offering
more surface area
 For example: n-Butane
has a greater force than
isobutane
Dipole-Dipole
• Weak
• Attraction between opposite poles of polar
molecules
Hydrogen Bonding
• IMF not a bond
• Strong (but not as strong as a bond)
• Attraction between a H of one molecule and a N,
O, or F of another
Phases of Matter
• IMF’s explain the phases of matter
• Solid
▫ Particles move relatively slow (basically vibrating in
fixed position)
▫ IMF’s hold them together
• Liquid
▫ Particles move more (able to flow past each other)
▫ IMF’s still hold together but not as tightly
• Gas
▫ Particle move quickly (each in rapid, random motion)
▫ Little to no IMF’s (we usually assume none)
Intermolecular forces are Intermolecular forces
are very important.
of little significance
Intermolecular forces
must be considered.
Phase Changes
• Solid  Liquid
▫ Solid  Liquid = Melting
▫ Liquid  Solid = Freezing
• Liquid  Gas
▫ Liquid  Gas = Evaporation, Vaporization, or
Boiling
▫ Gas  Liquid = Condensation
.
• Solid  Gas
▫ Solid  Gas = Sublimation
▫ Gas  Solid = Deposition
Evaporation vs Vaporization
• Evaporation
▫ Occurs at temperatures below boiling point
▫ Some molecules have enough energy to escape
surface of liquid
• Vaporization
▫ Occurs at boiling point
▫ Change to gaseous phase occurs throughout
liquid
Affect of Forces Between Particles
• Stronger forces require more energy to break
• Substances with stronger forces will phase
change at higher temperatures
Network Covalent
• Very Strong Bonds
• Sublime or melt at very high temperatures
• Other properties
▫ Very Hard
▫ Do not conduct electricity
Another Bond Type
• Bond between metals
• Called alloys
• Have a special kind of bonding, called
metallic bonding.
Metallic bonding
• Electron sea
• The atoms in a metal share their valence
electrons in a “sea” of electrons.
• These electrons can flow from one atom to
another freely.
• This is why metals are good conductors of
electricity.
Metallic Bonds
• Depend on what metal(s) bond is between
• Range from weak to very strong
• Melting points range from low to very high
• Other properties
▫ Range from soft to hard
▫ Malleable and Ductile
▫ Good conductor of heat and electricity
Ionic Bonds
• Strong (created by attraction between cations
and anions)
• Moderate to high melting points
• Other Properties
▫ Conduct electricity in liquid or aqueous form
▫ Many soluble in polar substances like water
Hydrogen Bonding
• Strongest IMF’s
• Low to moderate melting points
• Other Properties
▫ Soluble in some hydrogen-bonded and some polar
solvents
Dipole-Dipole
• Strong IMF’s
• Low to moderate melting points
• Other properties
▫ Soluble in some polar and some non-polar
solvents
Dispersion Forces
• Weakest IMF’s
• Extremely low to moderate melting points
• Other properties
▫ Soft
▫ Soluble in non-polar solvents
Other Properties Affected by IMF’s
• Surface Tension
▫ “Skin” on surface of liquid
▫ With increasing IMF’s it is harder for
molecules to leave surface
 Strong IMF’s will have lower vapor
pressures
 Low IMF’s lead to high vapor pressures
(substance is considered to be volatile)
• Viscosity
▫ Resistance of flow in a liquid
▫ Substances are more viscous the stronger their
IMF’s