Chemistry 30
Unit 2
Covalent Bonding
 Is the sharing of electrons
 It occurs between a:
 nonmetal and nonmetal
 semimetal and nonmetal
 semimetal and semimetal
 We already know how to name covalent compounds
(By adding prefixes and ending the last element in
“ide”.)
How do Covalent Bonds Form?
 Atoms share electrons with other atoms in order to
complete their shell.
 Octet Rule: Atoms can have a maximum of 8 valence
electrons in their outer shell.
 Bond – The attraction between atoms
 Lone Pair - a pair of electrons that are left on their own
around a central atom.
 Central atom – the one that has the most atoms attached to
it (usually the one with the lowest electronegativity;
exception is hydrogen, it is never the central atom).
How to Draw Lewis Dot Structures with
Molecules
 Count the total number of valence electrons for the molecule
 Choose a central atom (usually the one with the lowest electronegativity)
 Place electrons around atoms so that the octet rule is satisfied
 If the molecule is an ion, place brackets around the entire structure and
write the charge on the outside of the bracket.
Example: CH4
H
H
H
C
H
H
H
H C H
H
C
H
H
H
A bond is
indicated by a
dash/line
Example:
OH
First, count up the total number of valence electrons of
all the atoms.
O
Oxygen has 6
valence
electrons
H
Hydrogen has
1 valence
electron
That equals 7 valence electrons. However, we must take a look at
the charge. In this case, it is -1, which means that there is one
extra electrons, giving us a grand total of 8 valence electrons.
O H
Now put brackets around
the molecule and add the
charge on the outside.
Double and Triple Bonds
 Double bonds occur when there are 4 shared electrons in one spot. Ex:
CO2
 Triple bonds occur when there are 6 shared electrons in one spot. Ex: CO
Ionic vs. Covalent Bonding
 Ionic: Transfer of
electron(s) from one
atom to another.
Example is Sodium
Chloride, NaCl
Na
+
Na
H H C
Cl
Cl
Covalent: The atoms share
electrons. Example is
Methane, CH4
-
H
H C H
H
H
H
H
H C H
H
MOLECULAR GEOMETRY
VSEPR Theory
 Valence Shell Electron
Pair Repulsion theory.
 Most important factor
in determining
geometry is relative
repulsion between
electron pairs.
 (Show video)
A Molecule adopts
the shape that
minimizes the
electron pair
repulsions.
VSEPR charts
 Use the Lewis structure to determine the geometry




of the molecule
Electron arrangement establishes the bond angles
Geometry of the molecule can depend on either
the regions of electrons (Electron Pair Geometry)
or on the number of atoms (Molecular Geometry).
Charts look at the CENTRAL atom for all data!
Think REGIONS OF ELECTRON DENSITY rather
than bonds (for instance, a double bond would
only be 1 region)
Electron Pair Geometry
 Count up the total number of regions of bonds and #
of lone pairs around the central atom
 (If double bond or triple bond, it counts as 1)
# of Regions of
Electrons
Electron Pair Geometry
Bond Angle
2
Linear
180
3
Trigonal Planar
120
4
Tetrahedral
109.5, 107, 104.5
Electron Pair Geometry Examples
 BeH2 has two regions of electrons, therefore it is
linear.
 CO32- has three regions of electrons, therefore it is
trigonal planar.
 H2O has four regions of electrons, therefore it is
tetrahedral.
Molecular Geometry
 Depends on Electron Pair Geometry as well as the
number of atoms around the central atom
 Count up the number of atoms that are connected to
the central atom
Electron Pair
Geometry
Bond Angle
# of Atoms around
Central
Molecular
Geometry
Linear
180
2
Linear
Trigonal Planar
120
2
Bent
120
3
Trigonal Planar
104.5
2
Bent
107
3
Trigonal Pyramidal
109.5
4
Tetrahedral
Tetrahedral
Molecular Geometry Examples
 CO2 has two atoms around C, therefore it is linear.
 NO2- has two atoms around N, therefore it is bent.
 H2O has two atoms around O, therefore it is bent.
Bond Angles
 The angle between atoms
 Depends on Electron Pair Geometry and Molecular Geometry
Linear Electron Pair Geometry
180° Bond Angle
Trigonal Planar Electron Pair Geometry
120°
Tetrahedral Electron Pair Geometry
Tetrahedral
Trigonal Pyramidal
Molecular Geometries
Bent
Exceptions to the Octet Rule
 Sometimes there are exceptions, and an atom doesn’t need to satisfy the
octet rule (there aren’t enough electrons)….
 Sometimes an atom exceeds the octet rule…
But don’t worry
about these!
Bonding between
molecules or atoms
in solids or liquids
 Recall that molecules are
farthest apart in gases, but
closest together in solids.
Physical and chemical properties depend on
the type of bonds involved
 Ionic compounds typically have higher boiling points and melting points
than molecular compounds, due to the strength of the ionic attraction.
 Recall that sodium chloride is a solid at room temperature, while carbon
dioxide is a gas.
•
vs
Properties of Ionic Compounds
 High melting/boiling point
 Dissolve in water
 Form crystals when solid
 Conduct electrical current
Why do Ionic Compounds have High Melting
Points?
 Recall that ionic compounds form from oppositely charged ions.
 This creates strong bonds!
 Thus a lot of energy is
needed to separate the
atoms.
Properties of Covalent (Molecular) Compounds
 Due to weak intermolecular forces, are generally liquids and gases.
 Conduct little to no electricity
 Generally have low melting points and boiling points
Properties of Molecular Compounds Vary
 C0valent bonds differ in terms of how the bonded atoms share the
electrons. The number and type of atoms joined together determine the
molecular properties.
 The electrons which make up the
covalent bond are being pulled, like
a tug-of-war, toward each nucleus.
Nonpolar Covalent Bonds
 Recall that a magnet has a north and south pole. When the atoms in the
bond pull equally, the bonding electrons are shared equally and there are
no ‘north or south poles’ formed in the bond. We call this bond
NONPOLAR.
 Polarity increases intermolecular forces and therefore increases boiling
and melting points.
 The diatomic elements are nonpolar covalently bonded (e.g. hydrogen)
and thus are gases at room temperature (except bromine).
Polar Covalent Bonds
 Formed when the electrons are shared unequally between atoms
 Is a result of electronegativity
 Electronegativity: the ability of an atom to attract electrons when the
atom is in a compound (aka how hard it pulls in the tug-of-war)
 The more electronegative atom attracts electrons more strongly and gains
a slightly negative charge. The less electronegative atom has a slightly
positive charge
Describing Polar Bonds
 In hydrochloric acid (HCl), hydrogen has an electronegativity of 2.1 and
chlorine has 3.0. These values are significantly different, so the covalent
bond is polar.
 Chlorine pulls the electrons closer towards itself and becomes slightly
negative, leaving hydrogen slightly positive as shown:
Describing Polar Bonds
 Water is also a polar molecule (elecronegativities H: 2.1, O: 3.5)
 This explains why most ionic compounds are soluble (can dissolve) in
water:
Determining Bond Type
 Using the electronegativity chart, we can determine which bond type will
occur:
Attractions Between Molecules
 How do the strengths of intermolecular attractions compare with ionic
and covelent bonds?
 Intermolecular (attraction between molecules) are weaker than either
ionic or covalent bonds. However, these interactions still impact physical
properties.
 They include:
 Van der Waals forces:
 Dipole interactions
 Dispersion forces
 Hydrogen bonding
van der Waals Forces
 The two weakest interactions between molecules
 Named after Dutch chemist Johannes van der Waals
 Includes:
 Dispersion forces:



Weakest of all forces; occurs in all molecules
Caused by the motion of electrons
Very weak, very temporary attraction between slightly charged regions of a molecule
and its neighbours
 Dipole interaction:

Attraction between the slightly charged regions of polar molecules:
Hydrogen Bonding
 Attractive forces in which a hydrogen atom covalently bonded to a very
electronegative atom is also strongly attracted to an unshared electron pair of
another electronegative atom
 In other words, it is a dipole interaction that involves hydrogen and an
electronegative atom (N, O, F, Cl)
 This is a relatively strong attraction which serves to increase the melting and
boiling point of the substances affected by it
Hydrogen Bonding is Responsible for:
 Surface Tension
 Ice floating
 Helical structure of DNA
How many Drops of Water can a Penny hold?
Intermolecular Attractions and Molecular
Properties
 Recall that the physical properties of a compound depend on the type of
bonding it displays – in particular, whether it is ionic or covalent.
 A great range of physical properties occurs among covalent compounds.
 The diversity of physical properties among covalent compounds is mainly
because of widely varying intermolecular attractions.
 A few solids that consist of molecules break our rules – they will not melt
unless at extremely high temperatures or will not melt at all
Network Solids
 Aka network crystals
 Solids in which all of the atoms are covalently bonded to each other
 Melting a network solid would require breaking covalent bonds
throughout the solid
 E.g. diamond
Physical Properties
 The greater the strength within the bonds (INTRAMOLECULAR FORCES)
and between the molecules (INTERMOLECULAR FORCES) of a substance,
the more energy you need to break those bonds (i.e. to change state by
melting or vaporization)
That explains why you see such a variety in physical properties:
Lesson Check!
 Intramolecular Forces Worksheet
Melting and Boiling Points
 Dependent on intermolecular
forces
 The stronger the intermolecular
force is, the more energy is
required to melt or boil a solid
or liquid
 Therefore, intermolecular forces
raise the melting and boiling
points
Properties of Covalent (Network) Compounds
 Network = connected in many ways to molecules around it
 Have high melting and boiling points
 Cannot conduct electricity
Properties of Metal Compounds
 High melting points
 Very good electrical conductors
 Crystal arrangement
 Malleable (ability to
keep shape without
breaking)
 Dense
 Shiny