Covalent Bonding
Chapter 9
Section - 9.1 Forming Compounds
1. Why do
Atoms Bond?
• Lower energy states make an atom more stable.
• Gaining or losing electrons makes atoms more
stable by forming ions with noble-gas electron
configurations.
• Atoms also gain stability when they share
electrons and form covalent bonds.
• Sharing valence electrons with other atoms also
results in noble-gas electron configurations.
• Atoms in non-ionic compounds share
electrons.
2. What is a
Covalent Bond? •
The chemical bond that results from sharing
electrons is a covalent bond.
• A molecule is formed when two or more
atoms bond.
What is a
Covalent
Bond? (cont.)
• Diatomic molecules (H2, F2 for example) exist
because two-atom molecules are more stable
than single atoms. For example:
• As two fluorine atoms approach each other two
forces become important.
• A repulsive force occurs between the like-charged
protons of the two atoms.
• An attractive force also occurs between the protons of
one fluorine atom and the electrons of the other atom.
• As the fluorine atoms move closer, the attraction of
both nuclei for the other atom’s electrons until the
maximum attraction is achieved.
• The most stable arrangement of atoms exists at
the point of maximum net attraction, where the
atoms bond covalently and form a molecule.
3. Single Covalent
Bonds
• When only one pair of electrons is
shared, the result is a single covalent
bond.
• The figure shows two hydrogen atoms
forming a hydrogen molecule with a single
covalent bond, resulting in an electron
configuration like helium.
• The shared electron pair, often
referred to as the bonding pair, is
represented by either a pair of dots or
a line.
• In a Lewis structure dots or a line are
used to symbolize a single covalent
bond.
Single Covalent
Bonds (cont.)
• The halogens—the group 17/7A elements—
have 7 valence electrons.
• To attain an octet, one more electron is
necessary.
• Therefore, they will form single covalent bonds
with atoms of other identical non-metals.
• Atoms in group 16 can share two electrons
and form two covalent bonds.
• Water is formed from one oxygen with two
hydrogen atoms covalently bonded to it .
• Atoms in group 15 form three single
covalent bonds, such as in ammonia.
• Atoms of group 14 elements form four
single covalent bonds, such as in methane.
Single Covalent
Bonds (cont.)
• Single covalent bonds are also called
Sigma bonds.
• Sigma bonds occur when the pair of
shared electrons is in an area centered
between the two atoms.
• When two atoms share electrons, the
valence atomic orbital of one atom
overlaps or merges with the valence
atomic orbital of the other atom.
• A sigma bond results if the atomic orbitals
overlap end to end, concentrating the
electrons in a bonding orbital between the
atoms
4. Multiple
Covalent Bonds
• Double bonds form when two pairs of
electrons are shared between two atoms.
• Triple bonds form when three pairs of
electrons are shared between two atoms.
Multiple
Covalent Bonds
(cont.)
• The pi bond is formed when parallel orbitals
overlap and share electrons.
• The shared electron pair of a pi bond
occupies the space above and below the
line that represents where the two atoms are
joined together.
• A multiple covalent bond consists of one
sigma bond and at least one pi bond.
5. Strength of
Covalent Bonds
• The strength depends on the distance
between the two nuclei, or bond length.
• Determined by the size of the atoms and how
many electrons pairs are shared.
• As the number of shared electron pairs
increases, bond length decreases.
• As length increases, strength decreases,
thus single bonds are weaker than triple
bonds.
• The amount of energy required to break a
bond is called the bond dissociation
energy.
• The shorter the bond length, the greater
the energy required to break it.
Strength of
Covalent Bonds
(cont.)
• Bond dissociation energy indicates the
strength of a chemical bond.
• A direct relationship exist between bond
energy and bond length.
• An endothermic reaction is one where a
greater amount of energy is required to
break a bond in reactants than is released
when the new bonds form in the products.
• An exothermic reaction is one where
more energy is released than is required
to break the bonds in the initial reactants.
Section 9.2 - Naming Molecules
6. Naming
Binary
Molecular
Compounds
1. The first element is always named first
using the entire element name.
2. The second element is named using its
root and adding the suffix –ide.
3. Prefixes are used to indicate the number of
atoms of each element in a compound.
•
One exception to using these prefixes is
that the first element in the formula never
uses the prefix mono-.
• Many compounds were discovered and
given common names long before the
present naming system was developed
(water, ammonia, hydrazine, nitrous oxide).
7. Naming
Acids
• If the compound produces hydrogen ions in
solution, it is an acid.
• Two common acids exist – binary acids and
oxyacids.
• Naming Binary Acids
• Binary acids contain hydrogen and one other
element.
• The first word has 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 plus the suffix –ic followed
by the word acid (hydrochloric acid is HCl in water).
Naming
Acids
(cont.)
• Naming Oxyacids
• An oxyacid is an acid that contains both a hydrogen
atom and an oxyanion.
• An oxyanion is a polyatomic ion that contains oxygen.
• Identify the oxyanion present.
• The name of an oxyacid consists of a form of the root
of the anion, a suffix, and the word acid.
• If the anion suffix is –ate, it is replaced with the suffix
–ic. When the anion ends in –ite, the suffix is
replaced with the suffix –ous.
• These hydrogen-containing compounds are named
as acids only when they are in water solution.
8. Writing Formulas
from Names
• The name of a molecular compound
reveals its composition and is important
in communicating the nature of the
compound.
• Subscripts are determined from the
prefixes used in the name because the
name indicates the exact number of
each atom present in the molecule
• The formula for an acid can be derived
from the name as well
9. Names and
Formulas for
Bases
• A base is an ionic compound that produces
hydroxide ions when dissolved in water.
• Bases are named in the same way as other
ionic compounds—the name of the cation is
followed by the name of the anion.
• For example, aluminum hydroxide consists of
the aluminum cation (Al3+) and the hydroxide
anion (OH–).
• The formula for aluminum hydroxide is Al(OH)3.
Section 9.3 - Molecular Structures
10. Structural
Formulas
• A structural formula uses letter
symbols and bonds to show
relative positions of atoms.
• The structural formula can be
predicted for many molecules by
drawing the Lewis structure.
• Drawing Lewis Structures
Structural
Formulas
(cont.)
1. Predict the location of certain atoms.
a. Hydrogen is always a terminal atom because it can share only one
pair of electrons with one other atom
b. The atom with the least attraction for shared electrons in the
molecule is the central atom (lowest electronegativity)
2. Find the number of electrons available for bonding.
3. Determine the number of bonding pairs.
4. Place one bonding pair between the central atom and each
terminal element.
5. Determine the number of valence electrons remaining.
6. Determine whether the central atom satisfies the octet rule.
• Atoms within a polyatomic ion are covalently bonded.
11. Resonance
Structures
• Resonance is a condition that occurs when
more than one valid Lewis structure can be
written for a molecule or ion.
• Resonance structures differ only in the position of
the electron pairs, never the atom positions.
• The molecule or ion behaves as if it has only one
structure.
• The bond lengths are identical to each other and
intermediate between single and double
covalent bonds.
• This figure shows three correct ways to draw
the structure for (NO3)-1.
12. Exceptions to
the Octet Rule
• Some molecules do not obey the octet
rule.
1. A small group of molecules might have
an odd number of valence electrons.
• Ex. NO2 has five valence electrons from
nitrogen and 12 from oxygen and cannot form
an exact number of electron pairs.
Exceptions to
the Octet Rule
(cont.)
2. A few compounds form stable
configurations with less than 8
electrons around the atom—a
suboctet.
• Ex. Boron forms three covalent bonds
with other nonmetallic atoms.
• Such compounds are reactive and can
share an entire pair of electrons donated
by another atom
• A coordinate covalent bond forms
when one atom donates both of the
electrons to be shared with an atom
or ion that needs two electrons.
Exceptions to
the Octet Rule
(cont.)
3. A third group of compounds has central
atoms with more than eight valence
electrons, called an expanded octet.
• Elements in period 3 or higher have an
empty or partially filled d-orbital and can form
more than four covalent bonds by expanding
into the empty or partially filled d-orbital.
Section 9.4 - Molecular Shape
13. VSEPR Model
• The shape of a molecule determines many
of its physical and chemical properties.
• Molecular shape, in turn, is determined by
the overlap of orbitals that share electrons.
• Once Lewis structure is drawn, you can
determine the molecular geometry
(shape).
• Molecular geometry can be determined
with the Valence Shell Electron Pair
Repulsion model, or VSEPR model .
• This model minimizes the repulsion of
shared and unshared atoms around the
central atom.
VSEPR Model
(cont.)
• Electron pairs repel each other and cause
molecules to be in fixed positions relative to each
other.
• The angle formed by any two terminal atoms and
the central atom is called bond angle
• Lone pairs also determine the shape of a
molecule.
• Because lone pairs are not shared between two
nuclei, they occupy a slightly larger orbital than
shared electrons
• Electron pairs are located in a molecule as far
apart as they can be, but are pushed together
slightly by lone pairs
14. Hybridization
• Hybridization is a process in which atomic
orbitals mix and form new, identical hybrid
orbitals.
• Carbon often undergoes hybridization,
which forms an sp3 orbital formed from one
s orbital and three p orbitals.
• The number of atomic orbitals mixed to
form the hybrid orbital equals the total
number of pairs of electrons.
• The number of hybrid orbitals formed
equals the number of atomic orbitals mixed
• Lone pairs also occupy hybrid orbitals.
Section 9.5 - Electronegativity and Polarity
15. Electronegativity
Difference and Bond
Character
• Electron affinity measures the
tendency of an atom to accept an
electron.
• Affinity increases with increasing
atomic number within a period, and
decreases with increasing atomic
number within a group.
• Noble gases are not listed because
they generally do not form
compounds.
• This table lists the character and type
of chemical bond that forms with
differences in electronegativity.
Electronegativity
Difference and Bond
Character
• A covalent bond formed between
atoms of different elements does
not have equal sharing of the
electron pair due to a difference in
electronegativity
• Unequal sharing of electrons results
in a polar covalent bond.
• Bonding is often not clearly ionic or
covalent.
• This graph summarizes the range of
chemical bonds between two atoms.
16. Polar
Covalent Bonds
• Polar covalent bonds form when atoms
pull on electrons in a molecule unequally.
• Electrons spend more time around one
atom than another resulting in partial
charges at the ends of the bond called a
dipole (two poles).
• The more electronegative atom is located
at the partially negative end, while the
less electronegative atom is located at
the partially positive end.
Polar Covalent
Bonds (cont.)
• Covalently bonded molecules are either polar
or non-polar, depending on the nature and
location of the covalent bonds they contain.
• Non-polar molecules are not attracted by an
electric field.
• Polar molecules align with an electric field.
• Compare water and CCl4.
• Both bonds are polar, but only water is a polar
molecule because of the shape of the molecule.
• The electric charge on a CCl4 molecule measured
at any distance from the center of the molecule is
identical to the charge measured at the same
distance on the opposite side.
Polar Covalent
Bonds (cont.)
• Solubility is the physical property of a
substance’s ability to dissolve in another
substance.
• Bond type and the shape of the molecules
present determine solubility
• Polar molecules and ionic substances are
usually soluble in polar substances.
• Non-polar molecules dissolve only in nonpolar substances.
• “Like dissolves like”
17. Properties
of Covalent
Compounds
• Differences in properties are a result of
differences in attractive forces.
• Covalent bonds between atoms (within a
molecule) are strong, but attraction forces
between molecules are weak.
• The weak attraction forces between molecules
are known as intermolecular forces, or van
der Waals forces.
• The forces vary in strength but are weaker than
the bonds that join atoms in a molecule or ions
in an ionic compound.
Properties of
Covalent
Compounds
(cont.)
• There are different types of intermolecular forces.
• Nonpolar molecules exhibit a weak dispersion
force, or induced dipole.
• The greater the number of electrons, the stronger the
dispersion force.
• The force between two oppositely charged ends
of two polar molecules stronger and is called a
dipole-dipole force.
• The more polar the molecule the stronger the dipoledipole force.
• A hydrogen bond is an especially strong dipoledipole force between a hydrogen end of one
dipole and a fluorine, oxygen, or nitrogen atom on
another dipole.
Properties of
Covalent
Compounds
(cont.)
• Many physical properties are due to intermolecular
forces.
• Weak forces result in the relatively low melting and
boiling points of molecular substances.
• Many covalent molecules are relatively soft solids.
• Molecules can align in a crystal lattice, similar to
ionic solids but with less attraction between
particles.
• Solids composed of only atoms interconnected by
a network of covalent bonds are called covalent
network solids.
• Quartz and diamonds are two common examples of
network solids.
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