COVALENT COMPOUNDS – ACIDS –
MOLECULAR GEOMETRY –
INTERMOLECULAR FORCES
Why do Atoms Share Electrons?
• In the last unit we learned that some
metals and nonmetals react to form binary
ionic compounds.
• Electrons are transferred and the resulting
ions have noble gas electron
configurations.
• Compounds are then formed because the
ions are attracted to one another.
Why do Atoms Share Electrons?
• Sometimes two atoms that both need to
gain valence electrons have a similar
attraction for the electrons.
• Sharing electrons is one way these atoms
can acquire the electron configuration of a
noble gas, even though it will be on a parttime basis.
Why do Atoms Share Electrons?
• In a covalent bond, atoms do not lose or
gain electrons. Instead they share pairs of
electrons to achieve stability by filling their
outer energy levels so they can have a
stable octet.
• A molecule is formed when two or more
atoms bond covalently. They are often
called molecular compounds.
• They are also called covalent
compounds.
Why do Atoms Share Electrons?
Covalent Bond
• An attractive force exists between the
outer electrons of one atom and the
nucleus of a nearby atom.
• The force of attraction brings the two
atoms together until the force of repulsion
between the two nuclei and the force of
repulsion between the electron clouds
forces the atoms apart.
Covalent Bond
• If the forces of attraction are greater than
the forces of repulsion, then a covalent
bond forms between the atoms.
• Besides the comparative strengths of the
attractive and repulsive forces, another
reason the attractive forces can be
stronger is that a pair of electrons shared
between atoms in a stable covalent bond
have opposite spins and occupy less
space than a pair of electrons in an orbital
of only one atom.
Covalent Bond
• The bond is not rigid. It is much like a
spring where the atoms vibrate back and
forth at some average distance where the
attractive force and the repulsive force are
balanced.
Covalent Bond – Sharing More Than Two
Electrons
• Covalent bonds between atoms can involve
sharing more than two electrons.
• When a single pair of electrons (2 electrons)
is shared, this is known as a:
• single bond
• When two pairs of electrons (4 electrons)
are shared, this is known as a:
• double bond
• When three pairs of electrons (6 electrons)
are shared, this is known as a:
• triple bond
Bond Length and Bond Energy
• The average distance that separates the
atoms in a bond is known as the bond
length.
• Bond lengths are never really fixed
distances because the atoms vibrate.
They can also vary depending on the other
bonds present in a molecule.
• Bond energy is the energy required to
break a chemical bond to produce
individual atoms, each keeping its own
electrons.
Bond Length and Bond Energy
• Bond length and bond energy are
inversely related.
• A short bond length requires higher bond
energy to break it while a long bond length
requires less energy to break it.
Bond Properties
• Few chemical bonds are either totally
molecular or totally ionic.
• The bonds in most compounds have
characteristics of both.
• The electrons in a bond are not
necessarily shared equally. To determine
whether this uneven sharing will be very
small or very large, one compares the
ability of each atom to pull electrons
toward itself.
Bond Properties
• This property is called electronegativity.
• The electronegativity table is used to
provide numbers for comparison.
• The greater the difference in
electronegativity values between two
atoms, the more unequal the sharing and
the more ionic character the bond will
have.
Bond Properties
• A covalent bond formed between two
atoms with equally shared bonding
electrons is said to be a:
• nonpolar covalent bond
• Examples are: H – H, O – O, F – F
• When atoms of different elements bond,
the sharing of electrons can never be truly
equal.
Bond Properties
• A covalent bond formed between two
atoms in which the bonding electrons are
more strongly attracted to one atom over
the other is said to be a:
• polar covalent bond
• Examples are: Rb – O, Al – N, C – O
Determining Bond Polarity
• We make use of Pauling’s
electronegativity values:
Determining Bond Polarity
• Look up the electronegativity value of each
atom in the bond, and then subtract the
smaller value from the larger value.
• The difference is always positive.
• Use this table to determine the bond
polarity:
Determining Bond Polarity
• The uneven sharing causes the more
electronegative atom to have a partial
negative charge while the less
electronegative atom will have a partial
positive charge.
• EXAMPLE: Determine the bond polarity of
the following bonds:
C–H
2.55 – 2.20 = .35 nonpolar
N–H
3.04 – 2.20 = .84 polar
C–O
3.44 – 2.55 = .89 polar
Naming Covalent Compounds
• Naming covalent compounds is similar to
naming ionic compounds.
• One can use either the Stock naming
system or one that makes use of prefixes,
roots, and suffixes.
• The latter is known simply as the prefix
naming system.
Naming Covalent Compounds
• The root comes from the name of the
element and then the prefixes and suffixes
are added.
• Examples are:
• CO
• CO2
carbon monoxide
carbon dioxide
Naming Covalent Compounds
• The first element named is usually the one
with the lower electronegativity value.
• If there is only one of the first element,
then NO prefix is used.
• The ending –ide is used as it was in
naming ionic compounds.
Naming Covalent Compounds
• The common
prefixes are:
1 mono
2 di
3 tri
4 tetra
5 penta
6 hexa
7 hepta
8 octa
9 nona
10 deca
Naming Covalent Compounds
• Name the following covalent compounds:
P2O5
diphosphorus pentoxide
CCl4
carbon tetrachloride
As2O3
diarsenic trioxide
SO2
sulfur dioxide
NF3
nitrogen trifluoride
Naming Covalent Compounds
• The Stock system can be used to name
covalent compounds.
P2O5
phosphorus (V) oxide
CCl4
carbon (IV) chloride
As2O3
arsenic (III) oxide
SO2
sulfur (IV) oxide
NF3
nitrogen (III) fluoride
Writing Formulas for Covalent Compounds
• If the compound is named using the prefix
system, simply translate the prefixes as
written.
• dinitrogen tetroxide
• N2O4
• phosphorus trichloride
• PCl3
• disulfur trioxide
• S2O3
Naming Acids
• Water solutions of some molecules are
acidic and are therefore named as acids.
• A binary acid contains:
• hydrogen
• an anion
• but, NO OXYGEN
Naming Acids – Binary Acids
• When naming a binary acid, use the prefix
hydro to name the hydrogen part of the
compound.
• The rest of the name consists of a form of
the root of the second element, or
polyatomic ion, plus the suffix –ic.
• Then add the word acid.
• HBr in a water solution (aqueous) is known
as: hydrobromic acid.
Naming Acids - Examples
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Name the following acids:
HF
hydrofluoric acid
HCN
hydrocyanic acid
HI
hydroiodic acid
H2S
hydrosulfuric acid
Naming Acids – Oxyacids
• Any acid that contains hydrogen and an
oxyanion is known as an oxyacid.
• To name it, first identify the anion present.
• The name of the oxyacid consists of the
root of the anion, a suffix, and the word
acid.
Naming Acids – Oxyacids
• If the anion suffix is ATE, change it to IC.
• If the anion suffix is ITE, change it to OUS.
• ATE – IC, ITE – OUS
Naming Acids – Oxyacids - Examples
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Name the following oxyacids:
HNO2
nitrous acid
HClO3
chloric acid
HClO4
perchloric acid
HC2H3O2
acetic acid
• Notice there is NO use of the hydro.
Naming Acids – Oxyacids - Exceptions
• The following oxyacids were named before
the rules went into effect, so they must be
memorized:
• H2SO3
• sulfurous acid
• H2SO4
• sulfuric acid
• H3PO4
• phosphoric acid
Molecular Geometry – Lewis Structures
• In order to predict the arrangement of
atoms in a molecule, a model is used.
• The nuclei and inner-shell electrons are
represented by the element’s symbol. The
valence electrons are represented as dots
placed around each side of the symbol, up
to two per side.
Molecular Geometry – Lewis Structures
• Bonds between atoms are represented
either as pairs of dots or lines between the
atoms involved in the bond.
• Unshared pairs are represented as pairs of
dots placed around the appropriate atoms.
Molecular Geometry – Drawing
Lewis Structures
(a) determine the total number of valence
electrons in the compound by adding up
all the valence electrons of the atoms in
the compound.
(b) arrange all the element symbols
according to which element can form
more than one bond and those that can
only form one bond
Molecular Geometry – Drawing
Lewis Structures
• Atoms that can form only one bond are: H
and F. Cl, Br, and I will normally form only
one bond unless outnumbered by O or F.
• Atoms that “love” being in the middle of
things are: B, C, N, O, Si, P, S, As, Se, Sb.
(c) draw single lines between all the atoms
that are bonded together.
Molecular Geometry – Drawing
Lewis Structures
(d) count each line, multiply by two, and
subtract that number from the total
number of valence electrons.
(e) this gives you the electrons left to
distribute to all the elements still needing
electrons so each can have an octet.
(f) if there are not enough electrons available
to distribute, then one or more of the
single bonds may have to be made
double or triple bonds.
Molecular Geometry – Drawing
Lewis Structures - Examples
• Draw the Lewis structure for iodomethane
(CH3I)
C = 4e3H = 3eI = 7eval e- = 14
bond e- =
unshared e- =
used e- =
remaining e- =
Molecular Geometry – Drawing
Lewis Structures - Examples
• Draw the Lewis structure for: methanol
(CH3OH)
Molecular Geometry – Drawing
Lewis Structures - Examples
• Draw the Lewis structure for: dinitrogen
difluoride (N2F2)
Molecular Geometry – Drawing
Lewis Structures - Examples
• Draw the Lewis structure for:
formaldehyde (H2CO)
Molecular Geometry – Drawing
Lewis Structures - Examples
• Draw the Lewis structure for the hydroxide
ion (OH-)
Molecular Geometry – Drawing
Lewis Structures - Examples
• Draw the Lewis structure for the
ammonium ion (NH4+)
Molecular Geometry – Drawing
Lewis Structures - Resonance
• Sometimes a single Lewis structure is not
enough to accurately depict a molecule.
• When that happens, more than one
equivalent structure is used to represent
the molecule.
• When more than one Lewis structure can
be drawn for a molecule, the molecule is
said to be a resonance hybrid.
Molecular Geometry – Drawing
Lewis Structures - Resonance
• An example is:
All of these structures are equivalent and only
the double bond moves to different oxygens
around the sulfur.
Molecular Geometry – Drawing
Lewis Structures - Resonance
• Draw the three resonance structures for:
dinitrogen monoxide (N2O).
Molecular Geometry – Drawing Lewis
Structures – Exceptions to the Octet Rule
• There are three types of ions or molecules
that do not follow the octet rule:
• Ions or molecules with an odd number of
electrons (NO2)
Molecular Geometry – Drawing Lewis
Structures – Exceptions to the Octet Rule
• Ions or molecules with less than an octet.
(BF3)
Molecular Geometry – Drawing Lewis
Structures – Exceptions to the Octet Rule
• Ions or molecules with more than eight
valence electrons (an expanded octet).
Molecular Geometry – Drawing Lewis
Structures – Exceptions to the Octet Rule
• It is thought that the extra electrons go into
empty “d” orbitals, thus permitting the
central atom to exceed the octet rule.
• When it is necessary to exceed the octet
rule for one of several third row (or higher)
elements, assume the extra electrons
should be placed on the central atom.
Molecular Geometry – Determining the Shape
of the Molecule
• Shape is an important factor in
determining the chemical properties of a
molecule.
• An example is the difference between
normal red blood cells and one in a person
suffering from sickle cell anemia.
Molecular Geometry – Valence Shell Electron
Pair Repulsion Theory
• The theory used to predict shapes is called
the Valence Shell Electron Pair
Repulsion Theory or VSEPR.
• It is based on the idea that electron pairs
surrounding the central atom in a bond will
arrange themselves to be as far apart as
possible.
Molecular Geometry – Valence Shell Electron
Pair Repulsion Theory
• Shapes cannot be predicted from the
molecular formula.
• One must know:
• the Lewis structure
• the bonding (shared) pairs of electrons
attached to the central atom
• the nonbonding (unshared) pairs of
electrons attached to the central atom
Molecular Geometry – Valence Shell Electron
Pair Repulsion Theory
• Once the bonding pairs and nonbonding
pairs attached to the central atom have
been determined, use this chart to
determine the molecular geometry.
Molecular Geometry – Valence Shell Electron
Pair Repulsion Theory
Molecular Geometry – Valence Shell Electron
Pair Repulsion Theory - Example
• Determine the shape of ammonia (NH3)
(1) draw the Lewis structure
(2) number of shared (bonding) pairs =
(3) number of unshared (nonbonding) pairs =
(4) shape:
Molecular Geometry – Valence Shell Electron
Pair Repulsion Theory – More Examples
• Determine the shape of each of the
following molecules:
• water (H2O)
• dichlorodifluoromethane (CF2Cl2)
• arsenic pentafluoride (AsF5)
• selenium hexabromide (SeBr6)
Molecular Dipole
• Remember that each bond within a
molecule can be either nonpolar or polar.
• If the bond is polar, then one end of the
bond appears to have a slight positive
charge while the other appears to have a
slight negative charge.
• This creates a dipole.
Molecular Dipole
• If the molecule has polar bonds and the
shape of the molecule causes the
polarities to cancel, then the molecule is
nonpolar.
• If the molecule has polar bonds and the
shape of the molecule does not cause the
polarities to cancel, then the molecule will
be polar.
• If the molecule has all nonpolar bonds,
then no matter what the shape is, the
molecule will be nonpolar.
Molecular Dipole
• Determine the polarity of each of the
following molecules:
• ammonia (NH3)
• water (H2O)
• dichlorodifluoromethane (CF2Cl2)
• arsenic pentafluoride (AsF5)
• selenium hexabromide (SeBr6)
Intermolecular Forces
• Atoms and molecules are attracted to one
another. But, these forces of attraction are
not as strong as the forces of attraction
between atoms in the bonding process
(ionic and covalent bonds).
• Intermolecular forces are forces that
cause attractions between molecules.
• Intermolecular forces of attraction between
molecules, or between atoms and
molecules, do not involve the transferring
or sharing of electrons.
Intermolecular Forces
• The weakest intermolecular forces are
London dispersion forces. These are
forces of attraction between nonpolar
substances.
London dispersion forces
Intermolecular Forces
• Next in increasing strength are dipoledipole forces. These intermolecular
forces occur between polar molecules
because there is a dipole in the molecule.
The positive dipole end of one molecule
attracts the negative dipole end of a
nearby molecule.
Intermolecular Forces
• The last type of intermolecular force is
known as hydrogen bonding. This
especially strong force of attraction occurs
between molecules containing hydrogen
bonded directly to a highly electronegative
atom such as F, O, or N.