Chapter 8
Covalent Boding
Molecules & Molecular Compounds
In nature, matter takes many forms.
The noble gases exist as atoms. They are monatomic;
they consist of single atoms.
Hydrogen chloride (HCl) is a gas at room temperature.
Water is a liquid at room temperature.
Salts are crystalline solids with high melting points.
Compounds such as HCl and water are not ionic.
Their combing atoms do not give up electrons or accept
electrons as readily as sodium does in combing with
chlorine.
Molecules & Molecular Compounds
Instead, a “tug-of-war” for the electrons takes place
between the atoms, bonding the atoms together.
The atoms held together by sharing electrons are joined
by a covalent bond.
Many elements found in nature are in the form of
molecules. A molecule is a neutral group of atoms
joined together by covalent bonds.
Example is oxygen molecules. Each oxygen molecules
consists of two oxygen atoms joined by covalent bonds.
A diatomic molecule is a molecule consisting of two
atoms.
Molecules & Molecular Compounds
There are seven elements that are always found as
diatomic molecules
Iodine, Hydrogen, Nitrogen, Bromine, Oxygen, Chlorine,
Fluorine
I Have No Bright Or Clever Friends
CLIF H BrON
Molecules & Molecular Compounds
Atoms of different elements can combine chemically to
form compounds. In many compounds, atoms are
bonded to each other to form molecules. (H2O, CO)
A compound composed of molecules is called a
molecular compound.
The molecules of a given molecular compound are all the
same.
There is not such thing as a molecule of sodium chloride
(NaCl) or magnesium chloride. (MgCl2)
These ionic compounds exist as collections of + and charged ions arranges in repeating three dimensional
patterns.
Molecules & Molecular Compounds
Molecular compounds tend to have relatively lower
melting and boiling points than ionic compounds.
Many molecular compounds are gases or liquids at room
temperature.
Ionic compounds are formed from a metal combined with
a nonmetal.
Most molecular compounds are composed of atoms of
two or more nonmetals.
Molecular Formulas
A molecular formula is the chemical formula of a
molecular compound.
A molecular formulas shows how many atoms of each
element of a molecule contains.
A water molecule consists of two hydrogen atoms and
one oxygen atom. The molecular formula of water is
H2O.
The subscript written after the symbol indicates the
number of atoms of each element in the molecule. (If
there is only 1 atom, the subscript 1 is omitted.)
The molecular forumla of carbon dioxide is CO2.
Molecular Formulas
Ethane is C2H6. Ethane contains 2 carbon atoms and 6
hydrogen atoms.
A molecular formula reflects the actual number of atoms
in each molecule. The subscripts are not necessarily
lowest whole-number ratios.
Molecular formulas also describe molecules consisting of
one elements. Example O2
A molecular formula does not tell you about a molecule’s
structure.
It does not show either the arrangement of the various
atoms in space or which atoms are covalently bonded to
one another.
Questions
How are the melting points and boiling points of molecular
compounds usually different from ionic compounds?
Molecular compounds tend to have relatively lower melting
points and boiling points.
What information does a molecular formula provide?
Show how many atoms of each element one molecules of
a compound contains.
What are the only elements that exist in nature as
uncombined atoms & what term is used to describe them?
Noble gases, monatomic.
End of Section 8.1
Octet Rule for Covalent Bonding
Recall that in forming ionic compounds, electrons tend to
be transferred so that each ion acquires a noble gas
configuration.
In forming covalent bonds, electron sharing usually
occurs so that atoms attain the electron configurations of
noble gases.
Example: each H atom has one electron. But a pair of H
atoms share these two electrons when they form a
covalent bond in the H molecule.
Each H atom thus attains the electron configuration of
He, a noble gas with two electrons.
Covalent Bonding
Combinations of atoms of the nonmetallic elements in
Groups 4A, 5A, 6A, and 7A are likely to form covalent
bonds.
The atoms usually acquire a total of 8 electrons, or an
octet, by sharing electrons so that the octet rule applies.
The hydrogen atoms in a hydrogen molecule are held
together mainly by the attraction of the shared electrons
to the positive nuclei.
Two atoms held together by sharing
a pair of electrons are joined by a
single covalent bond.
Covalent Bonding
The halogens (7A) also form single covalent bonds in
their diatomic molecules.
Chlorine atom has 7 valence electrons (group 7A), it
needs one more to attain an octet.
By sharing electrons and forming a single covalent bond,
two chlorine atoms achieve an octet.
Covalent Bonding
A pair of valence electrons that is not shared between
atoms is called an unshared pair. (or lone pair or
nonbonding pair)
Covalent Bonding
Take a look at ammonia (NH3). The ammonia molecules
has one unshared pair of electrons.
Covalent Bonding
Another example, methane (CH4)
Practice Problems
Draw electron dot structures for each molecule: chlorine,
bromine, iodine
Draw an electron dot structure for each of the following
molecules that have single covalent bonds: H2O2, PCl3
Double and Triple Covalent Bonds
Atoms form double or triple covalent bonds if they can
attain a noble gas structure by sharing two pairs or three
pairs of electrons.
Double Covalent Bond – a bond that involves two
shared pairs of electrons
Triple Covalent Bond – a bond formed by sharing three
pairs of electrons.
Double and Triple Covalent Bonds
Single, double and triple covalent bonds can also exist
between unlike atoms.
CO2 - 2 oxygen atoms (6 valence e-), each share 2 e- with
carbon (4 valence e-) to form a total of two carbonoxygen double bonds.
Carbon is an
example of a
triatomic molecule,
which is a molecule
consisting of three
atoms.
Coordinate Covalent Bonds
Carbon monoxide (CO) is an example of a type of
covalent bonding different from that seen in water,
ammonia, methane and carbon dioxide.
A carbon atoms needs to gain four electrons to attain an
octet. An oxygen atom needs two electrons.
Coordinate covalent bond is a covalent bond in which
one atom contributes both bonding electrons.
CO has 2 covalent bonds and 1 coordinate covalent bond
Coordinate Covalent Bonds
If the structure is a molecular ion, add one valence
electron for each negative charge and remove one
valence electron for each positive charge.
Polyatomic ion – is a tightly bound group of atoms that
has a positive or negative charge and behaves as a unit.
(NH4+, SO42-, PO43-)
Polyatomic Ions
Most polyatomic cations and anions contain both
covalent and coordinate covalent bonds.
Therefore, compounds containing polyatomic ions
include both ionic and covalent bonding.
The electron dot structure for a neutral molecule contains
the same number of electrons as the total number of
valence electrons in the combining atoms.
The negative charge of a polyatomic ions shows the
number of electrons in addition to the valence electrons
of the atoms present.
Because a negatively charged polyatomic ion is part of
an ionic compound, the positive charge of the cation of
the compound balances these additional electrons.
Practice
Draw the electron dot structure of the following:
Hydroxide ion (OH-)
The polyatomic boron tetrafluoride anion (BF4-)
Sulfate - SO42-
(S is the central atom)
Carbonate - CO32-
(C is the central atom)
Hydrogen carbonate ion - HCO3attached to O)
(C is central & H is
Bond Dissociation Energies
A large quantity of heat is given off when hydrogen atoms
combine to form hydrogen molecules.
This suggests that the product is more stable than the
reactants.
The covalent bond in the hydrogen molecule (H2) is so
strong that it would take 435 kJ of energy to break apart
all of the bonds in 1 mole of H2.
The energy required to break the bond between two
covalently bonded atoms is the bond dissociation
energy.
Usually expressed as the energy needed to break on mol
of bonds. (kJ/mol)
Bond Dissociation Energies
A large bond dissociation energy corresponds to a strong
covalent bond.
A carbon-carbon single bond has a bond dissociation
energy of about 347 kJ/mol.
Strong carbon-carbon bonds help explain the stability of
carbon compounds.
Bond dissociation energy is a measure of bond strength
Bond strength - Triple > double > single
Bond dissociation energy – Triple > double > single
Resonance
Ozone in the upper atmosphere blocks harmful ultraviolet
radiation from the sun. At the lower elevations, it
contributes to smog.
Ozone molecule has two possible electron dot structures.
The electron pairs were thought to rapidly flip back and
forth, or resonate, between the different electron dot
structures.
Ozone consists of 1 single coordinate covalent bond, and
1 double covalent bond.
Resonance
Double covalent bonds are usually shorter than single
bonds, so it was once thought that the bond lengths in
ozone were unequal.
Experiments showed that was not the case, however.
The two bonds are the same length.
The actual bonding in the ozone molecule is the average
of the two electron dot structures.
The electron pairs do not actually resonate back and
forth.
The actual bonding of oxygen atoms in ozone is a hybrid,
or mixture of the extremes represented by the
resonance forms.
Resonance
Resonance structure is a structure that occurs when it
is possible to write two or more valid electron dot
formulas that have the same number of electron pairs for
a molecule or ion.
Resonance structures are a way to envision the bonding
in certain molecules.
Although no back-and-forth changes ocurr, doubleheaded arrows are used to connect resonance
structures.
Questions
What electron configurations do atoms usually achieve
by sharing electrons to form covalent bonds?
Noble gases configurations
How is an electron dot structure used to represent a
covalent bond?
Two dots represent each covalent bond
When are two atoms likely to form a double bond
between them? A triple bond?
When they can attain a noble gas structure by sharing
two pairs or three pairs of electrons.
Questions
How is a coordinate covalent bond different from other
covalent bonds?
The shared electron pair comes from one of the bonding
atoms. In other covalent bonds, each bonding atom
provides an electron.
How is the strength of a covalent bond related to its bond
dissociation energy?
A large bond dissociation energy corresponds to a strong
covalent bond.
Questions
List three ways in which the octet rule can sometimes fail
to be obeyed.
Molecules whose total number of valence electrons is an
odd number. Atom has fewer or more than a complete
octet of valence electrons.
Draw electron dot structures for the following molecules,
which have only single covalent bonds. H2S, PH3, CIF
Which bond is stronger? H2 has a dissociation energy of
435 kJ/mol, a carbon-carbon bond has a dissociation
energy of 347 kJ/mol
The H-H bond is stronger because it has a greater
dissociation energy.
End of Section 8.2
Molecular Orbitals
The model for covalent bonding we have been using
assumes that the orbitals are those of the individual
atoms.
There is a quantum mechanical model of bonding,
however, that describes the electrons in molecules using
orbitals that exist only for groupings of atoms.
When two atoms combine, the model assumes that their
atomic orbitals overlap to produce molecular orbitals that
apply to the entire molecule.
In some ways, atomic orbitals and molecular orbitals are
similar.
Molecular Orbitals
Just as an atomic orbital belongs to a particular atom, a
molecular orbital belongs to a molecule as a whole.
Each atomic orbital is filled if it contains two electrons.
Similarly, two electrons are required to fill a molecular
orbital.
A molecular orbital that can be occupied by two electrons
of a covalent bond is called a bonding orbital.
Molecular Orbital - Sigma Bond
When two atomic orbitals combine to form a molecular orbital
that is symmetrical around the axis connecting two atomic
nuclei, a sigma () bond is formed.
The distinguishing feature of a sigma bond is that the overlap
region lies directly between the two nuclei.
It does not matter what shapes the orbitals have or what types
they are. They can be s orbitals or p orbitals or hybrid orbitals.
Molecular Orbitals
The side by side overlap of atomic p orbitals produces pi
() molecular orbital.
When a pi molecular orbital is filled with 2 electrons, a
pi bond results.
The pi bond has orbital overlap off to the sides of the
line joining the two nuclei.
Molecular Orbitals
Atomic orbitals in pi bonding overlap less than in sigma
bonding.
Therefore, pi bonds tend to be weaker than sigma bonds.
VSEPR Theory
Electron dot structures fail to reflect the three
dimensional shapes of the molecules.
Valence Shell Electron Pair Repulsion Theory (VSEPR)
According to VSEPR theory, the repulsion between
electron pairs causes molecular shapes to adjust so that
the valence electron pairs stay as far apart as possible.
Unshared pairs of electrons – no bonding atom is vying
for these unshared electrons so they are held closer to
the central atom than are the bonding pairs.
The unshared pair strongly repels the bonding pairs
pushing them together.
VSEPR Theory
VSEPR Theory
Steric number = (number of lone pairs on central
atom) + (number of atoms bonded to central atom)
• The steric number is determined from the Lewis
structure.
Steric number determines the electron-pair
arrangement, the geometry that maximizes the
distances between the valence-shell electron pairs.
VSEPR
In predicting molecular shapes, it may be useful to start
with an electron dot structure.
The electron dot structure shows both the bonding and
nonbonding pairs of electrons around the central atom.
When using VSEPR theory to predict molecular shape,
double and triple bonds are viewed as single bonds.
When 4 pairs of electrons must be accommodated
around the central atom – tetrahedral (109.5º)
When 3 pairs – a trigonal planar maximizes space (120º)
When 2 pairs – a linear arrangement (180º)
Hybrid Orbitals
VSEPR works well when accounting for molecular shape,
but does not help much in describing the types of bonds
formed.
Orbital hybridization provides information about both
molecular bonding and molecular shape.
In hybridization, several atomic orbitals mix to form the
same total number of equivalent hybrid orbitals.
End of Section 8.3
End of Chapter 7
Bond Polarity
Covalent bonds involve electron sharing between atoms.
However, they differ in terms of how the bonded atoms
share the electrons.
The bonding pairs of electrons
are pulled between the nuclei
of the atoms sharing the eWhen the atoms in the bond pull
Equally (identical atoms are bonded), the bonding
electrons are shared equally and the bond is a nonpolar
covalent bond.
H2, O2, N2, and Cl2 all have nopolar covalent bonds
Bond Polarity
A polar covalent bond (polar bond), is a covalent bond
between atoms in which the electrons are shared
unequally.
The more electronegative atoms attracts electrons more
strongly and gains a lightly negative charge.
The less electronegative atoms has a slightly positive
charge.
Fluorine is the most electronegative and oxygen is the
second most electronegative element.
Electronegativity describes the attraction an atom has
for electrons when the atom is in a compound.
Electronegativity Values of Some Elements
Bond Polarity
In hydrogen chloride (HCl) molecule, the H has an
electronegativity of 2.1 and the chlorine has an
electronegativity of 3.0.
The values are significantly different, so the covalent
bond in HCl is polar.
The chlorine atom acquires a slightly negative chare and
the hydrogen atoms acquires a slightly positive charge.
Greek letter delta () indicated that the atoms involved in
the covalent bond acquire only partial charges, less than
1+ or 1-.
+
-
H – Cl
+
-
-
H–O–H
+ -
H - Br
Bond Polarity
The electronegativity difference between two atoms tells
you what kind of bond is likely to form.
As the electronegativity difference between two atoms
increases, the polarity of the bond increases.
If the electronegativity difference is greater than 2.0, it is
very likely that electrons will be pulled away completely
by one of the atoms. (an ionic bond will form)
EN Differences
Probable Bond Type
0.0 – 0.4
Nonpolar covalent
0.4 – 1.0
Moderately polar covalent
Example
H – H (0.0)
H – Cl (0.9)
1.0 – 2.0
> 2.0
H – F (1.9)
Na – Cl (2.1)
Very polar covalent
ionic
Problems
Place the following covalent bonds in order from least to
most polar. a. H – Cl b. H – Br c. H – S d. H - C
c & d (tie), b, a
Identify the bonds between atoms of each pair of
elements as nonpolar covalent, moderately polar
covalent, very covalent, or ionic.
a. H & Br b. K & Cl
c. C & O
e. Li & O
f. Br & Br
d. Cl & F
a.moderately polar covalent
b. ionic
c. moderately
to very polar covalent d. moderately to very polar
covalent e. ionic
Polar Molecules
The presence of a polar bond in a molecule often makes
the entire molecule polar.
In a polar molecule, one end of the molecule is slightly
negative and the other end is slightly positive.
In HCl, the partial charges on the hydrogen and chlorine
atoms are electrically charged regions or poles.
A molecule that has two poles is called a dipolar
molecule.
When polar molecules are placed between oppositely
charged plates, they tend to become oriented with
respect to the positive and negative plates.
Polar Molecules
The effect of polar bonds on the polarity of an entire
molecule depends on the shape of the molecule and the
orientation of the polar bonds.
CO2 has two polar bonds and is linear. The C and O lies
along the same axis, thus the bond polarities cancel
because they are in opposite directions.
CO2 is a nonpolar molecule despite the presence of two
polar bonds.
Water also has two polar bonds. The water molecule is
bent and the bond polarities do not
cancel and the water molecule is polar.
Attractions Between Molecules
Molecules can attract each other by a variety of forces.
Intermolecular attractions are weaker than either ionic or
covalent bonds.
Intermolecular forces are responsible for determining
whether a molecular compound is a gas, a liquid, or a
solid at a given temperature.
Van der Waals Forces – are the two weakest attractions
between molecules and include both:
• dipole interactions
• dispersion forces.
Dipole Interactions
Dipole interactions occur when polar molecules are
attracted to one another. The attraction involved occurs
between the oppositely charged regions of polar
molecules.
The slightly negative region is weakly attracted to the
slightly positive regions of another molecule.
Dipole interactions are similar to but much weaker than
ionic bonds.
Dipole Interactions
Dispersions Forces
Dispersion forces are the weakest of all molecular
interactions and are caused by the motion of electrons.
They occur between nopolar molecules.
When the moving electrons happen to be momentarily
more on the side of a molecule closest to a neighboring
molecule, their electric force influences the neighboring
molecule’s electrons to be momentarily more on the
opposite side.
This causes an attraction between the two molecules
similar but much weaker than the attraction between
polar molecules.
Dispersions Forces
The strength of dispersion forces generally increases as
the number of electrons in a molecule increases.
The halogen diatomic molecules attract each other
mainly by means of dispersion forces.
Fluorine and chlorine has relatively few electrons and are
gases at room temperature because of their weak
dispersion forces.
Bromine has a larger number of electrons and generates
larger dispersion forces. Bromine molecules attract each
other sufficiently to make it a liquid at room temperature.
Iodine, with a larger number of electrons, is a solid at
room temperature.
Dispersions Forces
Hydrogen Bonds
The dipole interactions in water produce an attraction
between water molecules.
Each O – H bond in the water molecule is highly polar
and the O acquires a slightly negative charge.
The positive region of one water molecules attracts the
negative region of another water molecule.
This attraction between the hydrogen of one water
molecule and the oxygen of another water molecule is
strong compared to other dipole interaction.
This strong attraction, (also found in other H containing
molecules) is called hydrogen bonding.
Hydrogen
Bonds
Hydrogen Bonds
Hydrogen bonds are attractive forces in which a
hydrogen covalently bonded to a very electronegative
atom is also weakly bonded to an unshared electron pair
of another electronegative atom.
The other atom may be in the same molecule or in a
nearby molecule.
Hydrogen bonding always involves hydrogen.
The bond between hydrogen and a very electronegative
atom is strongly polar.
Hydrogen bonds are the strongest of the intermolecular
forces.
End of Chapter 8