Chapter 8 Covalent Boding

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Chapter 8Covalent Boding
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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.
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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.
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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
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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.
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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.
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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.
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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 molecules structure. It does
not show either the arrangement of the various
atoms in space or which atoms are covalently
bonded to one another.
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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.
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End of Section 8.1
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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.
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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.
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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.
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Covalent Bonding
A pair of valence electrons that is not shared
between atoms is called an unshared pair. (or
lone pair or nonbonding pair)
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Covalent Bonding
Take a look at ammonia (NH3). The ammonia
molecules has one unshared pair of electrons.
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Covalent Bonding
Another example, methane (CH4)
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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
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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.
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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 carbon-oxygen
double bonds.
Carbon is an example of a triatomic molecule,
which is a molecule consisting of three atoms.
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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
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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-)
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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.
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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 -
HCO3- (C is central H is attached to O)
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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)
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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 gt double gt
single Bond dissociation energy Triple gt double
gt single
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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.
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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.
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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, double-headed
arrows are used to connect resonance structures.
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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.
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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.
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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.
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End of Section 8.2
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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.
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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.
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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.
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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.
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Molecular Orbitals
Atomic orbitals in pi bonding overlap less than
in sigma bonding. Therefore, pi bonds tend to be
weaker than sigma bonds.
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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.
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VSEPR Theory
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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.

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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)
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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.
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End of Section 8.3
End of Chapter 7
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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 e- When 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
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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.
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Electronegativity Values of Some Elements
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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
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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)
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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 d. Cl F e. Li O f. Br
Br a.moderately polar covalent b. ionic c.
moderately to very polar covalent d. moderately
to very polar covalent e. ionic
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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.
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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.
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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.

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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.
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Dipole Interactions
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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 molecules
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.
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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.
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Dispersions Forces
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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.
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Hydrogen Bonds
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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.
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End of Chapter 8