Transcript Covalent Bonding - Biloxi Public Schools

Covalent Bonding
Chapter 9
Section 9.1
THE COVALENT BOND
Worldwide, scientists are studying ways to
increase food supplies, reduce pollution,
and prevent disease. Understanding the
chemistry of compounds that make up
fertilizers, pollutants, and materials that
carry genetic information is essential in
developing new technologies in these
areas. An understanding of the chemistry
of compounds requires an understanding
of their bonding.
Why do atoms bond?
You have learned that all noble gases have particularly
stable electron arrangements. This stable arrangement
consists of a full outer energy level. A full outer energy
level consists of two valence electrons for helium and
eight valence electrons for all other noble gases. Because
of this stability, noble gases, in general, don’t react with
other elements to form compounds.
You also learned that when metals and nonmetals react
to form binary ionic compounds, electrons are
transferred, and the resulting ions have noble-gas
electron configurations. But sometimes two atoms that
both need to gain valence electrons to become stable
have a similar attraction for electrons.
• Sharing of electrons is
another way that
atoms can acquire the
electron configuration
of noble gases.
• The octet rule states
that atoms lose, gain,
or share electrons to
achieve a stable
configuration of eight
electrons, or octet.
What is a
covalent bond?
You know that certain
atoms, such as magnesium
and chlorine, transfer
electrons from one atom
to another, forming an
ionic bond. However, the
number of ionic
compounds is quite small
compared with the total
number of known
compounds. What type of
bonding is found in all
these other compounds
that are not ionically
bonded?
• Atoms in these other compounds
share electrons.
• The chemical bond that results from
the sharing of valence electrons is
called a covalent bond.
• In a covalent bond, the shared
electrons are considered to be part
of the complete outer energy level
of both atoms involved.
• Covalent bonding generally happens
when elements are relatively close
to each other on the periodic table.
• The majority of covalent bonds form
between nonmetallic elements.
• A molecule is formed when two or
more atoms bond covalently.
Formation of a Covalent Bond
• Hydrogen, oxygen, nitrogen, chlorine,
bromine, iodine, and fluorine occur in
nature as diatomic molecules because
the molecules are more stable in pairs
than singly.
– HONClBrIF
• Why do two atoms that do not give up
electrons bond with each other?
Consider fluorine’s (F2) electron configuration.
Each fluorine molecule has seven valence
electrons and must have one additional electron
to form an octet. As two fluorine molecules
approach each other, two forces become
important. A repulsive force occurs between the
like-charged particles. An attractive force occurs
between oppositely-charged particles. As they
move closer, the nuclei become more attracted
to the other’s valence electron until maximum
attraction is achieved. At that point, the
attractive forces balance the repulsive forces. A
covalent bond now exists between the two
atoms.
Single Covalent Bonds
• When a single pair of
electrons is shared, a
single covalent bond
forms.
• Single covalent bonds are
also called sigma bonds.
• A sigma bond occurs
when the electron pair is
shared in an area
centered between the
two atoms.
Multiple Covalent Bonds
• Many molecules achieve stability by sharing more
than one pair of electrons between two atoms,
forming a multiple covalent bond.
• Carbon, nitrogen, oxygen, and sulfur most often
form multiple bonds.
• A double bond occurs when two pairs of
electrons are shared. (O2)
• A triple bond occurs when three pairs of
electrons are shared. (N2)
The Pi Bond
• A pi bond is formed when
parallel orbitals overlap to share
electrons.
• A multiple bond consists of one
sigma bond and at least one pi
bond.
• A pi bond always accompanies a
sigma bond when forming
double and triple bond.
Strength of Covalent Bonds
• Some covalent bonds are broken more easily than
others because they differ in strength.
• Several factors control the strength.
– How much distance separates the bonded nuclei (bond
length)
• The shorter the bond length, the stronger the bond.
– The amount of energy required to break a bond is called
bond dissociation energy.
• Endothermic reactions occur when a greater amount of energy
is required to break the bond(s) in the reactants than is
released when the new bonds form in the products.
• Exothermic reactions occur when more energy is released
forming new bonds than is required to break bonds in the
initial reactants.
Section 9.2
NAMING MOLECULES
You know that many atoms
covalently bond to form molecules
that behave as a single unit. These
units can be represented by
chemical formulas and names that
are used to identify them. When
naming molecules, the system of
rules is similar to the one you used
to name ionic compounds.
Naming Binary Molecular
Compounds
The anesthetic dinitrogen
oxide (N2O), commonly
known as nitrous oxide, is
a covalently bonded
compound. Binary
molecular compounds are
composed of two
different nonmetals and
do not contain metals or
ions. Although many of
these compounds have
common names, they also
have scientific names that
reveal their composition.
Use these rules to name
binary molecular
compounds.
1. The first element in the formula is
always named first, using the entire
element name.
2. The second element in the formula is
named using the root of the element
name and using the suffix –ide.
3. Prefixes are used to indicate the
number of atoms of each type that
are present in the compound.
Exceptions
One exception to using these prefixes is
that the first element in the formula never
uses the prefix mono-. Also, to avoid
awkward pronunciation, drop the final
letter in the prefix when the element name
begins with a vowel.
For example, CO is carbon monoxide not
monocarbon monooxide.
Practice Problems
Name the following binary covalent compounds:
1)
2)
3)
4)
5)
CCl4
As2O3
CO
SO2
NF3
Common Names of Some Molecular
Compounds
• Some compounds have “common names”
• Can you think of one for dihydrogen
monoxide?
Formula
Common Name
?
Molecular Name
Dihydrogen monoxide
Ammonia
Nitrogen trihydride
Hydrazine
Dinitrogen tetrahydride
Nitrous oxide (laughing gas)
Dinitrogen monoxide
Nitric oxide
Nitrogen monoxide
Naming Acids
Water solutions of some molecules are
acidic and are named as acids. Acids are
important compounds with specific
properties that will be discussed later. If the
compound produces hydrogen ions (H+) in
solution, it is an acid. For example, HCl
produces H+ in solution and is an acid. Two
common types of acids exist– binary acids
and oxyacids.
Naming Acids
Naming Binary Acids
• Contains a hydrogen
and one other element
• Use the prefix hydro• Root name of second
element plus suffix -ic
• Follow with word acid
• Example:
– HCl is hydrochloric acid
Naming Oxyacids
• Contains an oxyanion and
hydrogen
• Consists of a form of the root
of the anion, a suffix, and the
word acid
• -ate is replaced with –ic
• -ite is replaced with –ous
• Examples:
– HNO3 is nitric acid
– HNO2 is nitrous acid
Writing Formulas from Names
Section 9.3
MOLECULAR STRUCTURES
You can now identify atoms that bond covalently
and name the molecular compounds formed
through covalent bonding. In order to predict
the arrangement of atoms in each molecule, a
model, or representation is used. Several
different models can be used:
1)
Predict the location of certain atoms.
a)
b)
Structural Formulas
One of the most useful
molecular models is the
structural formula, which
uses letter symbols and
bonds to show relative
positions of atoms. The
structural formula can be
predicted for many
molecules by drawing
Lewis structures, but more
involved structures are
needed to help you
determine the shapes of
molecules.
2)
3)
4)
5)
6)
Hydrogen is always terminal and connects to
only one atom.
The atom with the least attraction for shared
electrons is usually central atom
Find the total number of electrons available for
bonding. This total is the number of valence
electrons in the atoms in the molecule.
Determine the number of bonding pairs by
dividing the number electrons available by 2.
Place one bonding pair between the central
atom and each of the terminal atoms.
Subtract the number of pairs you used in step 4
from the number of bonding pairs you
determined in step 3. The remaining electron
pairs include lone pairs as well as double and
triple bonds
If the central atom is not surrounded by four
electron pairs, it does not have an octet. You
must convert one or two of the lone pairs on
the terminal atoms to a double bond or triple
bond between the terminal and central atom.
Example Problem 9-3 (pg 253)
Ammonia is a raw material for the manufacture
of many materials, including fertilizers, cleaning
products, and explosives. Draw the Lewis
structure for ammonia (NH3).
Example Problem 9-4 (pg 254)
Carbon dioxide is a product of all cellular
respiration. Draw the Lewis structure for carbon
dioxide.
Resonance Structures
• Using the same sequence
of atoms, it is possible to
have more than one
correct Lewis structure
when a molecule has
both a double and single
bond.
• Resonance is a condition
that occurs when more
than one valid Lewis
structure can be written
for a molecule or ion.
Draw the Lewis resonance
structures for the following:
1)SO3
2)SO2
3)O3
4)NO2
Exceptions to the Octet Rule
Three Reasons:
1. Odd number of valence
electrons cannot form
octet around each atom
(NO2)
2. Form with fewer than 8
electrons present around
an atom (BH3)
3. Some have central atoms
that contain more than 8
electrons (PCl5)
Section 9.4
MOLECULAR SHAPE
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. Theories have
been develop to explain the overlap of
bonding orbitals and are used to
predict the shape of the molecule.
VSEPR Model
Valence
Shell
Electron
Pair
Repulsion
• Many chemical reactions,
especially those in living things,
depend on the ability of two
compounds to contact each other.
• The shape of the molecule
determines whether or not
molecules can get close enough to
react.
• Once a Lewis structure is drawn,
you can determine the molecular
geometry, or shape, of the
molecule.
• The model used to determine the
molecular shape is refereed to as
the VSEPR model.
• This model is based on an
arrangement that minimizes the
repulsion of shared and unshared
pairs of electrons around the
central atom.
Hybridization
A hybrid results from combining two
of the same type of object, and it
has characteristics of both. Atomic
orbitals undergo hybridization during
bonding. Let’s consider the bonding
involved in the methane molecule
(CH4). The carbon atom has 4
valence electrons and electron
configuration [He]2s22p2. You might
expect the two unpaired p electrons
to bond with other atoms and the 2s
electrons to remain in a loan pair.
However, carbon atoms undergo
hybridization, a process in which
atomic orbitals are mixed to form
new, identical hybrid orbitals. Each
hybrid orbital contains one electron
that it can share with another atom.
Determine the molecular geometry,
bond angle, and type of hybridization
for the following:
1) BF3
2) NH4+
3) OCl2
4) BeF2
5) CF4
Section 9.5
ELECTRONEGATIVITY AND
POLARITY
You now know that the type of bond
that forms when two electrons react
depends on which elements are
involved. What makes one type of
bond form when carbon burns and
another type form when iron
corrodes? The answer lies in how
much attraction each type of atom has
for electrons.
Electronegativity Difference and Bond
Character
• Electron affinity is a measure of the tendency of an atom
to accept an electron.
• Excluding noble gases, electron affinity increases as the
atomic number increases within a given period and
decreases with an increase in atomic number within a
group.
• The scale of electronegativity allows a chemist to evaluate
the electron affinity of specific atoms when they are
incorporated into a compound.
• The character and type of chemical bond can be
predicted using the electronegativity difference of the
elements that are bonded.
• For identical atoms which have an
electronegativity difference of 0, the electrons
are equally shared between the two atoms
and the bond is considered non-polar
covalent.
• Unequal sharing results in a polar covalent
bond.
• Electronegativity of about 1.7 or greater
typically means the bond is ionic and not
covalent.
• A polar molecule has a partial negative charge
on one side and a partial positive charge on
the other.
• The molecule is a dipole because to the two
partial charges.
• Symmetry also plays a role in polarity.
• Asymmetrical molecules tend to be polar as
long as their bonds are polar.
• Symmetrical molecules tend to be nonpolar
even if their bonds are polar.
Covalent Network Solids
• A number of solids are composed only of
atoms interconnected by a network of
covalent bonds.
• These solids are often called covalent network
solids.
• Quartz is a network solid, as is diamond.
• They are typically brittle, nonconductors of
heat or electricity, and extremely hard.