Chapter 6.2 and 6.5
Covalent Compounds
Covalent Bonds
• Sharing Electrons
• Covalent bonds form when atoms share
one or more pairs of electrons
• nucleus of each atom is attracted to
electron cloud of other atom
• neither atom removes an electron from
the other
• Usually formed by a nonmetal + a
nonmetal
Covalent Bonding
Covalent Bonds
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•
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•
Sharing Electrons
Covalent bonds
Molecules are formed
space where electrons move is called
molecular orbital
• made when atomic orbitals overlap
Molecules
Covalent Bonds
• Energy and Stability
• Noble gases are stable (full octet) (low
potential energy)
• Other elements are not stable (high
potential energy)
• covalent bonding decreases potential
energy because each atom achieves
an electron configuration like a
noble gas
Covalent Bonds
• Energy and Stability
• because potential decreases when
atoms bond, energy is released
• i.e., atoms lose potential energy when
they bond
• loss of potential energy implies
higher stability
Covalent Bonds
• Energy and Stability
• Potential energy determines bond length
• At minimum potential energy, the distance
between two bonded atoms is called bond
length
• The bonded atoms vibrate about the bond,
therefore, bond length is an average
length
Covalent Bonds
• Energy and Stability
• bonds vary in strength
• bond energy: the amount of energy
required to break the bonds in 1 mol of a
chemical compound
• bond energy predicts reactivity
• bond energy is equal to loss of potential
energy during formation
Bond Energies and
Lengths
Bonds and Energy
• Single bonds are the longest bonds
with the least bond energy
• Double bonds are shorter, stronger
and have intermediate bond energy
• Triple bonds are the shortest,
strongest, and have the highest bond
energy
Covalent Bonds
• Electronegativity
• Atoms share electrons equally or
unequally
• nonpolar covalent bond: bonding electrons
shared equally
• polar covalent bond: shared electrons
more likely to be found around more
electronegative atom (shared unequally)
Covalent Bonds
• Electronegativity
• The difference in electronegativity
can be used to predict the type of
bond (but boundaries are arbitrary)
• Electronegativities are listed on your
periodic table
Bond Types
0.3
1.7
Covalent Bonds
• Polar molecules have partially positive
and partially negative ends
• such molecules are called “dipoles”
• “δ” means “partial” in math and
science
• The positive end is designated as δ+
and the negative end as δ• example: Hδ+Fδ-
Practice: Calculate the
bond type
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N and H
F and F
Ca and Cl
C and O
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Polar covalent
Non-polar covalent
Ionic
Polar covalent
Covalent Bonds
• Polarity is related to bond
strength
• greater electronegativity
difference means greater
polarity, which means greater
bond strength
Electronegativity
Difference for Hydrogen
Halides
Covalent Bonds
• Bond type determines properties of
substances
• metallic bonds: electrons can move
from one atom to another—good
conductors
• ionic bonds: hard and difficult to break
apart
• covalent bonds: low melting/boiling
points
Properties of Substances with
Different Types of Bonds
Lewis Structures
• G. N. Lewis was an American
chemistry who developed the system
of showing valence electrons and the
resulting bonds
• Lewis structure: a structure in which
atomic symbols represent nuclei and
inner-shell electrons, and dots are
used to represent valence electrons
Drawing Lewis
Structures for Atoms
• Lewis structures represent valence
electrons with dots
• position of electrons is symbolic (not
literal)
• the symbol represents the nucleus
and core electrons
• dots around the atomic symbol
represent electrons
Lewis Structures of
Second-Period Elements
Vocabulary for Lewis
Structures
• Unshared pairs: AKA lone pairs, a
pair of valence electrons not involved
in bonding to another atom
• Single bond: a covalent bond in which
one pair of electrons is shared
between two atoms
• Multiple bond: double bond shares 2
•
pairs of electrons, triple bond
shares 3 pairs of electrons
Drawing Lewis
Structures for Molecules
• Lewis Electron-Dot Structures
• draw the Lewis structure for each atom in
the compound; place one electron on each
side before pairing
• Count and write down the total number of
valence electrons for all atoms involved
• Count and write down all the electrons in
lone pairs
Drawing Lewis
Structures for Molecules
• Subtract the number of electrons in
lone pairs from the total number of
valence electrons
• Divide this number by 2 and this is
usually the number of bonds you have
• Draw sticks for bonds (2 electrons
each), draw in the lone pairs, check
for octets
Drawing Lewis
Structures for Molecules
• The atom with the lowest
electronegativity is usually the
central atom
• If carbon is present, it is the central
atom
Lewis Structure Practice
• Draw Lewis structures for the
following molecules:
• SiF4
HBr
• PBr3
CO2
• SiS2
H2O
• CO
ammonia (NH3)
Drawing Ions
• You must consider the absence of an
electron or the extra electrons
• Subtract an absent electron from
the total
• Add any extra electrons to the total
(add on an atom that is not the
central atom)
Drawing Ions
• Place the entire structure in
brackets and put the charge in the
upper right corner outside the
brackets
• Coordinate-covalent bond: resulting
bond when a shared pair of electrons
originates from one atom
Drawing Ions
• Draw the Lewis structure for the
sulfate ion, SO42-
Resonance Structures
• Resonance Structures: bonding
in molecules or ions that cannot
be correctly represented by a
single Lewis structure
• Draw each possible structure
and the hybrid structure
Practice Resonance
Structure
• Draw the Lewis structure for SO3
• Possible structures:
Naming Covalent
Compounds
• First name: name of first element in
formula
• usually least electronegative
• requires a prefix if more than one of
them
• Second name: ends in –ide
• requires a prefix if more than one of them
• oxygen requires the prefix “mono” if only
one of them
Naming Covalent Compounds
Practice Naming
1.
2.
3.
4.
5.
6.
antimony tribromide
hexaboron silicide
chlorine dioxide
iodine pentafluoride
dinitrogen trioxide
ammonia
More Naming Practice
1.
2.
3.
4.
5.
6.
P4S5
PI3
NO
N2F4
CO2
H2O
Molecular Shapes
• Three-dimensional shape helps
determine physical and chemical
properties
• valence shell electron pair repulsion
(VSEPR) theory predicts molecular
shapes
• based on the idea that electrons
repel one another
Molecular Shapes
• Lewis structures show which atoms
are connected where, and by how
many bonds, but they don’t properly
show the 3-D shapes of molecules
• To find the actual shape, first draw
the Lewis structure, and then use
VSEPR theory
MOLECULAR GEOMETRY
VSEPR
• Valence Shell Electron Pair
Repulsion theory.
• Most important factor in determining
geometry is the relative repulsion
between electron pairs.
MOLECULAR
GEOMETRY
Molecule adopts
the shape
that minimizes
the electron
pair
repulsions.
VSEPR Rules
• To apply VSEPR theory:
• Draw the Lewis structure of the
molecule and identify the central
atom
• Count the number of electron charge
clouds (lone and bonding pairs)
surrounding the central atom.
• Predict molecular shape by assuming
that clouds orient so they are as far
away from one another as possible
Bond Angles
• Lone-pairs of electrons behave as if
they are slightly bigger than bonded
electron pairs and act to distort the
geometry about the atomic center so
that bond angles are slightly smaller
than expected
Polarity of Molecules
• Two atoms: bond polarity is the
molecular polarity
• More than 2 atoms: the geometry of
the molecule must be considered
• If the bonds are non-polar, the
molecular is non-polar
• Some molecules with polar bonds
•
can be non-polar
More
• Sometimes the partial charges cancel
each other out because they are
directly opposite each other
• Consider CO2 and CCl4
• The symmetrical distribution of the
bonds leads to cancellation of the
charges
Polar and Non-polar
Molecules
Linear with Only 2 Atoms
• Molecules made of only two atoms
are always linear
• Bond angle is 180°
• Examples: HCl, F2
• Bond polarity is the molecular
polarity
Linear with More than 2
Atoms
• Sometimes molecules made of more
than 2 atoms are shaped linearly
• Example: CO2
• The oxygens pull the electrons in
equal, but opposite, directions
• The molecule is sometimes nonpolar
and sometimes polar
Bent
• There are two shapes that are called
bent
• The first one has one lone pair on the
central atom and two bonded atoms
• Example: NOCl (nitrosyl chloride)
• Bond angle: 117°
• The molecule is polar
Bent
• The other bent shape has two lone
pairs on the central atom
• Example: H2O
• The bond angle is 105° (some books
report 104.5°)
• The molecule is polar
Trigonal Planar
• This shape resembles a triangle
• Example: CH2O and BF3
• Boron is an exception to the octet
rule
• The bond angle is 120°
• The molecule can be polar or nonpolar
depending on the nature of the atoms
bonded to the central atom
Tetrahedral
• This shape usually occurs when 4
atoms are bonded to the central
atom
• Examples: CF4, CH4 (methane)
• If the 4 bonded atoms are identical,
and the central atom obeys the octet
rule, the molecule is nonpolar
•
Bond angle is 109.5°
Trigonal Pyramidal
• There is one lone pair on the central
atom and three bonded atoms
• Examples: NH3, PCl3
• Bond angle is 107°
• The molecule is polar
Molecular Shapes
• Determine the molecular shapes,
bond polarity, molecular polarity,
name, and bond angles.
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CO
CO2
BF3
SO2
SiBr4
PCl3
Challenging one: C2H4 (ethene)