We begin with our understanding of the
relationship between chemical behavior and atomic structure. That
is, we assume the Periodic Law that the chemical and physical
properties of the elements are periodic functions of atomic number.
We further assume the structure of the atom as a massive,
positively charged nucleus, whose size is much smaller than that of
the atom as a whole, surrounded by a vast open space in which move
negatively charged electrons. These electrons can be effectively
partitioned into a core and a valence shell, and it is only the
electrons in the valence shell which are significant to the
chemical properties of the atom. The number of valence electrons in
each atom is equal to the group number of that element in the
Periodic Table.
Goals
The atomic molecular theory is extremely
useful in explaining what it means to form a compound its component
elements. That is, a compound consists of identical molecules, each
comprised of the atoms of the component elements in a simple whole
number ratio. However, the atomic molecular theory also opens up a
wide range of new questions. We would like to know what atomic
properties determine the number of atoms of each type which combine
to form stable compounds. Why are some combinations observed and
other combinations not observed? Some elements with very dissimilar
atomic masses (for example, iodine and chlorine) form very similar
chemical compounds, but other elements with very similar atomic
masses (for example, oxygen and nitrogen) form very dissimilar
compounds. What factors are responsible for the bonding properties
of the elements in a similar group? In general, we need to know
what forces hold atoms together in forming a molecule.
We have developed a detail understanding of
the structure of the atom. Our task now is to apply this
understanding to develop a similar level of detail about how atoms
bond together to form molecules.
Observation 1: Valence and the Periodic
Table
To begin our analysis of chemical bonding, we
define the
valence of an atom by its tendencies to form molecules.
The inert gases do not tend to combine with any other atoms. We
thus assign their
valence as 0, meaning that these atoms tend to form 0
bonds. Each halogen prefers to form molecules by combining with a
single hydrogen atom (e.g. HFHF,
HClHCl).
We thus assign their valence as 1, also taking hydrogen to also
have a valence of 1. What we mean by a valence of 1 is that these
atoms prefer to bind to only one other atom. The valence of oxygen,
sulfur, etc. is assigned as 2, since two hydrogens are required to
satisfy bonding needs of these atoms. Nitrogen, phosphorus, etc.
have a valence of 3, and carbon and silicon have a valence of 4.
This concept also applies to elements just following the inert
gases. Lithium, sodium, potassium, and rubidium bind with a single
halogen atom. Therefore, they also have a valence of 1.
Correspondingly, it is not surprising to find that, for example,
the combination of two potassium atoms with a single oxygen atom
forms a stable molecule, since oxygen's valence of 2 is be
satisfied by the two alkali atoms, each with valence 1. We can
proceed in this manner to assign a valence to each element, by
simply determining the number of atoms to which this
element's atoms prefer to bind.
In doing so, we discover that the periodic
table is a representation of the valences of the elements: elements
in the same group all share a common valence. The inert gases with
a valence of 0 sit to one side of the table. Each inert gas is
immediately preceded in the table by one of the halogens: fluorine
precedes neon, chlorine precedes argon, bromine precedes krypton,
and iodine precedes xenon. And each halogen has a valence of one.
This "one step away, valence of one" pattern can be
extended. The elements just prior to the halogens (oxygen, sulfur,
selenium, tellurium) are each two steps away from the inert gases
in the table, and each of these elements has a valence of two (e.g.
H2OH2O,
H2
SH2
S). The
elements just preceding these (nitrogen, phosphorus, antimony,
arsenic) have valences of three (e.g.
NH3NH3,
PH3PH3),
and the elements before that (carbon and silicon most notably) have
valences of four
(CH4CH4,
SiH4SiH4).
The two groups of elements immediately after the inert gases, the
alkali metals and the alkaline earths, have valences of one and
two, respectively. Hence, for many elements in the periodic table,
the valence of its atoms can be predicted from the number of steps
the element is away from the nearest inert gas in the table. This
systemization is quite remarkable and is very useful for
remembering what molecules may be easily formed by a particular
element.
Next we discover that there is a pattern to
the valences: for elements in groups 4 through 8 (e.g. carbon through neon), the valence of each
atom
plus the number of electrons in the valence
shell in that atom always equals
eight. For examples, carbon has a valence of 4
and has 4 valence electrons, nitrogen has a valence of 3 and has 5
valence electrons, and oxygen has a valence of 2 and has 6 valence
electrons. Hydrogen is an important special case with a single
valence electron and a valence of 1. Interestingly, for each of
these atoms, the number of bonds the atom forms is equal to the
number of vacancies in its valence shell.
To account for this pattern, we develop a
model assuming that each atom attempts to bond to other atoms so as
to completely fill its valence shell with electrons. For elements
in groups 4 through 8, this means that each atom attempts to
complete an "octet" of valence shell electrons. (Why
atoms should behave this way is a question unanswered by this
model.) Consider, for example, the combination of hydrogen and
chlorine to form hydrogen chloride, HClHCl. The chlorine atom has
seven valence electrons and seeks to add a single electron to
complete an octet. Hence, chlorine has a valence of 1. Either
hydrogen or chlorine could satisfy its valence by
"taking" an electron from the other atom, but this
would leave the second atom now needing two electrons to complete
its valence shell. The only way for both atoms to complete their
valence shells simultaneously is to
share two electrons. Each atom donates a single
electron to the electron pair which is shared. It is this sharing
of electrons that we refer to as a chemical bond, or more
specifically, as a
covalent bond, so named because the bond acts to
satisfy the valence of both atoms. The two atoms are thus held
together by the need to share the electron pair.
Observation 2: Compounds of Carbon and
Hydrogen
Many of the most important chemical fuels are
compounds composed entirely of carbon and hydrogen,
i.e. hydrocarbons. The smallest of these is
methane
CH4CH4,
a primary component of household natural gas. Other simple common
fuels include ethane
C2H6C2H6,
propane
C3
H8C3
H8,
butane
C4
H10C4
H10,
pentane
C5H12C5H12,
hexane
C6H14C6H14,
heptane
C7
H16C7
H16,
and octane
C8
H18C8
H18.
It is interesting to note that there is a consistency in these
molecular formulae: in each case, the number of hydrogen atoms is
two more than twice the number of carbon atoms, so that each
compound has a molecular formula like
CnH2n+2CnH2n+2.
This suggests that there are strong similarities in the valences of
the atoms involved which should be understandable in terms of our
valence shell electron pair sharing model. In each molecule, the
carbon atoms must be directly bonded together, since they cannot be
joined together with a hydrogen atom. In the easiest example of
ethane, the two carbon atoms are bonded together, and each carbon
atom is in turn bonded to three hydrogen atoms. Thus, in this case,
it is relatively apparent that the valence of each carbon atom is
4, just as in methane, since each is bonded to four other atoms.
Therefore, by sharing an electron pair with each of the four atoms
to which it is bonded, each carbon atom has a valence shell of
eight electrons.
In most other cases, it is not so trivial to
determine which atoms are bonded to which, as there may be multiple
possibilities which satisfy all atomic valences. Nor is it trivial,
as the number of atoms and electrons increases, to determine
whether each atom has an octet of electrons in its valence shell.
We need a system of electron accounting which permits us to see
these features more clearly. To this end, we adopt a standard
notation for each atom which displays the number of valence
electrons in the unbonded atom explicitly. In this notation, carbon
and hydrogen look like Figure 1, representing the single valence electron in
hydrogen and the four valence electrons in carbon.
Figure 1
Using this notation, it is now relatively easy
to represent the shared electron pairs and the carbon atom valence
shell octets in methane and ethane. Linking bonded atoms together
and pairing the valence shell electrons from each gives Figure 2.
Figure 2
Recall that each shared pair of electrons
represents a chemical bond. These are examples of what are called
Lewis structures, after G.N. Lewis who first invented
this notation. These structures reveal, at a glance, which atoms
are bonded to which, i.e. the structural formula of the molecule. We
can also easily count the number of valence shell electrons around
each atom in the bonded molecule. Consistent with our model of the
octet rule, each carbon atom has eight valence electrons and each
hydrogen has two in the molecule.
In a larger hydrocarbon, the structural
formula of the molecule is generally not predictable from the
number of carbon atoms and the number of hydrogen atoms, so the
molecular structure must be given to deduce the Lewis structure and
thus the arrangement of the electrons in the molecule. However,
once given this information, it is straightforward to create a
Lewis structure for molecules with the general molecular formula
CnH2n+2CnH2n+2
such as propane, butane, etc. For example, the Lewis structure for
"normal" butane (with all carbons linked one after
another) is found here.
Figure 3
It is important to note that there exist no
hydrocarbons where the number of hydrogens exceeds two more than
twice the number of carbons. For example,
CH5CH5
does not exist, nor does
C2H8C2H8.
We correspondingly find that all attempts to draw Lewis structures
which are consistent with the octet rule will fail for these
molecules. Similarly,
CH3CH3
and
C2
H5C2
H5
are observed to be so extremely reactive that it is impossible to
prepare stable quantities of either compound. Again we find that it
is not possible to draw Lewis structures for these molecules which
obey the octet rule.
We conclude from these examples that, when it
is possible to draw a Lewis structure in which each carbon has a
complete octet of electrons in its valence shell, the corresponding
molecule will be stable and the hydrocarbon compound will exist
under ordinary conditions. After working a few examples, it is
apparent that this always holds for compounds with molecular
formula
CnH2n+2CnH2n+2.
On the other hand, there are many stable
hydrocarbon compounds with molecular formulae which do not fit the
form
CnH2n+2CnH2n+2,
particularly where the number of hydrogens is less than
2n+22n2.
In these compounds, the valences of the carbon atoms are not quite
so obviously satisfied by electron pair sharing. For example, in
ethene
C2
H4C2
H4
and acetylene
C2H2C2H2
there are not enough hydrogen atoms to permit each carbon atom to
be bonded to four atoms each. In each molecule, the two carbon
atoms must be bonded to one another. By simply arranging the
electrons so that the carbon atoms share a single pair of
electrons, we wind up with rather unsatisfying Lewis structures for
ethene and acetylene, shown here.
Figure 4
Note that, in these structures, neither carbon
atom has a complete octet of valence shell electrons. Moreover,
these structures indicate that the carbon-carbon bonds in ethane,
ethene, and acetylene should be very similar, since in each case a
single pair of electrons is shared by the two carbons. However,
these bonds are observed to be chemically and physically very
different. First, we can compare the energy required to break each
bond (the bond energy or bond strength). We find that the carbon-carbon bond
energy is 347 kJ in
C2H6C2H6,
589 kJ in
C2H4C2H4,
and 962 kJ in
C2H2C2H2.
Second, it is possible to observe the distance between the two
carbon atoms, which is referred to as the
bond length. It is found that carbon-carbon bond
length is 154 pm in
C2H6C2H6,
134 pm in
C2H4C2H4,
and 120 pm in
C2H2C2H2.
(1picometer=1pm=10-12m1picometer1pm
10-12m).
These observations reveal clearly that the bonding between the
carbon atoms in these three molecules must be very
different.
Note that the bond in ethene is about one and
a half times as strong as the bond in ethane; this suggests that
the two unpaired and unshared electrons in the ethene structure
above are also paired and shared as a second bond between the two
carbon atoms. Similarly, since the bond in acetylene is about two
and a half times stronger than the bond in ethane, we can imagine
that this results from the sharing of three pairs of electrons
between the two carbon atoms. These assumptions produce the Lewis
structures here.
Figure 5
These structures appear sensible from two
regards. First, the trend in carbon-carbon bond strengths can be
understood as arising from the increasing number of shared pairs of
electrons. Second, each carbon atom has a complete octet of
electrons. We refer to the two pairs of shared electrons in ethene
as a
double bond and the three shared pairs in acetylene as
a
triple bond.
We thus extend our model of valence shell
electron pair sharing to conclude that carbon atoms can bond by
sharing one, two, or three pairs of electrons as needed to complete
an octet of electrons, and that the strength of the bond is greater
when more pairs of electrons are shared. Moreover, the data above
tell us that the carbon-carbon bond in acetylene is shorter than
that in ethene, which is shorter than that in ethane. We conclude
that triple bonds are shorter than double bonds which are shorter
than single bonds.
Observation 3: Compounds of Nitrogen, Oxygen, and the
Halogens
Many compounds composed primarily of carbon
and hydrogen also contain some oxygen or nitrogen, or one or more
of the halogens. We thus seek to extend our understanding of
bonding and stability by developing Lewis structures involving
these atoms. Recall that a nitrogen atom has a valence of 3 and has
five valence electrons. In our notation, we could draw a structure
in which each of the five electrons appears separately in a ring,
similar to what we drew for C. However, this would imply that a
nitrogen atom would generally form five bonds to pair its five
valence electrons. Since the valence is actually 3, our notation
should reflect this. One possibility looks like this.
Figure 6
Note that this structure leaves three of the
valence electrons "unpaired" and thus ready to join in
a shared electron pair. The remaining two valence electrons are
"paired," and this notation implies that they therefore
are not generally available for sharing in a covalent bond. This
notation is consistent with the available data,
i.e. five valence electrons and a valence of 3.
Pairing the two non-bonding electrons seems reasonable in analogy
to the fact that electrons are paired in forming covalent
bonds.
Analogous structures can be drawn for oxygen,
as well as for fluorine and the other halogens, as shown here.
Figure 7
With this notation in hand, we can now analyze
structures for molecules including nitrogen, oxygen, and the
halogens. The hydrides are the easiest, shown here.
Figure 8
Note that the octet rule is clearly obeyed for
oxygen, nitrogen, and the halogens.
At this point, it becomes very helpful to
adopt one new convention: a pair of bonded electrons will now be
more easily represented in our Lewis structures by a straight line,
rather than two dots. Double bonds and triple bonds are represented
by double and triple straight lines between atoms. We will continue
to show non-bonded electron pairs explicitly.
As before, when analyzing Lewis structures for
larger molecules, we must already know which atoms are bonded to
which. For example, two very different compounds, ethanol and
dimethyl ether, both have molecular formula
C2
H6OC2
H6O.
In ethanol, the two carbon atoms are bonded together and the oxygen
atom is attached to one of the two carbons; the hydrogens are
arranged to complete the valences of the carbons and the oxygen
shown here.
Figure 9
This Lewis structure reveals not only that
each carbon and oxygen atom has a completed octet of valence shell
electrons but also that, in the stable molecule, there are four
non-bonded electrons on the oxygen atom. Ethanol is an example of
an
alcohol. Alcohols can be easily recognized in Lewis
structures by the C-O-H group. The Lewis structures of all alcohols
obey the octet rule.
In dimethyl ether, the two carbons are each
bonded to the oxygen, in the middle, shown here.
Figure 10
Ethers can be recognized in Lewis structures by the
C-O-C arrangement. Note that, in both ethanol and dimethyl ether,
the octet rule is obeyed for all carbon and oxygen atoms.
Therefore, it is not usually possible to predict the structural
formula of a molecule from Lewis structures. We must know the
molecular structure prior to determining the Lewis
structure.
Ethanol and dimethyl ether are examples of
isomers, molecules with the same molecular formula but
different structural formulae. In general, isomers have rather
different chemical and physical properties arising from their
differences in molecular structures.
A group of compounds called
amines contain hydrogen, carbon, and nitrogen. The
simplest amine is methyl amine, whose Lewis structure is here.
Figure 11
"Halogenated" hydrocarbons have
been used extensively as refrigerants in air conditioning systems
and refrigerators. These are the notorious
"chlorofluorocarbons" or "CFCs" which have
been implicated in the destruction of stratospheric ozone. Two of
the more important CFCs include Freon 11,
CFCl3CFCl3,
and Freon 114,
C2
F4Cl2C2
F4Cl2,
for which we can easily construct appropriate Lewis structures,
shown here.
Figure 12
Finally, Lewis structures account for the
stability of the diatomic form of the elemental
halogens, F2F2,
Cl2Cl2,
Br2Br2,
and
I2I2.
The single example of
F2F2
is sufficient, shown here.
Figure 13
We can conclude from these examples that
molecules containing oxygen, nitrogen, and the halogens are
expected to be stable when these atoms all have octets of electrons
in their valence shells. The Lewis structure of each molecule
reveals this character explicitly.
On the other hand, there are many examples of
common molecules with apparently unusual valences, including:
carbon dioxide
CO2CO2,
in which the carbon is bonded to only two atoms and each oxygen is
only bonded to one; formaldehyde
H2COH2CO;
and hydrogen cyanide
HCNHCN.
Perhaps most conspicuously, we have yet to understand the bonding
in two very important elemental diatomic molecules,
O2O2
and
N2N2,
each of which has fewer atoms than the valence of either
atom.
We first analyze
CO2CO2,
noting that the bond strength of one of the
COCO
bonds in carbon dioxide is 532 kJ, which is significantly greater
than the bond strength of the
COCO
bond in ethanol, 358 kJ. By analogy to the comparison of bonds
strengths in ethane to ethene, we can imagine that this difference
in bond strengths results from double bonding in
CO2CO2.
Indeed, a Lewis structure of
CO2CO2
in which only single electron pairs are shared (Figure 14) does not obey the octet rule, but one in which
we pair and share the extra electrons reveals that double bonding
permits the octet rule to be obeyed (Figure 15).
Figure 14
Figure 15
A comparison of bond lengths is consistent
with our reasoning: the single
COCO
bond in ethanol is 148 pm, whereas the double bond in
CO2CO2
is 116.
Knowing that oxygen atoms can double-bond, we
can easily account for the structure of formaldehyde. The strength
of the
COCO
bond in
H2
COH2
CO
is comparable to that in
CO2CO2,
consistent with the Lewis structure here.
Figure 16
What about nitrogen atoms? We can compare the
strength of the
CN
CN
bond in
HCNHCN,
880 kJ, to that in methyl amine, 290 kJ. This dramatic
disparity again suggests the possibility of multiple bonding, and
an appropriate Lewis structure for
HCNHCN
is shown here.
Figure 17
We can conclude that oxygen and nitrogen
atoms, like carbon atoms, are capable of multiple bonding.
Furthermore, our observations of oxygen and nitrogen reinforce our
earlier deduction that multiple bonds are stronger than single
bonds, and their bond lengths are shorter.
As our final examples in this section, we
consider molecules in which oxygen atoms are bonded to oxygen
atoms. Oxygen-oxygen bonds appear primarily in two types of
molecules. The first is simply the oxygen diatomic molecule,
O2O2,
and the second are the peroxides, typified by hydrogen peroxide,
H2O2H2O2.
In a comparison of bond energies, we find that the strength of the
OO bond in
O2O2
is 499 kJ whereas the strength of the OO bond in
H2O2H2O2
is 142 kJ. This is easily understood in a comparison of the Lewis
structures of these molecules, showing that the peroxide bond is a
single bond, whereas the
O2O2
bond is a double bond, shown here.
Figure 18
We conclude that an oxygen atom can satisfy
its valence of 2 by forming two single bonds or by forming one
double bond. In both cases, we can understand the stability of the
resulting molecules by in terms of an octet of valence
electrons.
Interpretation of Lewis Structures
Before further developing our model of
chemical bonding based on Lewis structures, we pause to consider
the interpretation and importance of these structures. It is worth
recalling that we have developed our model based on observations of
the numbers of bonds formed by individual atoms and the number of
valence electrons in each atom. In general, these structures are
useful for predicting whether a molecule is expected to be stable
under normal conditions. If we cannot draw a Lewis structure in
which each carbon, oxygen, nitrogen, or halogen has an octet of
valence electrons, then the corresponding molecule probably is not
stable. Consideration of bond strengths and bond lengths enhances
the model by revealing the presence of double and triple bonds in
the Lewis structures of some molecules.
At this point, however, we have observed no
information regarding the geometries of molecules. For example, we
have not considered the angles measured between bonds in molecules.
Consequently, the Lewis structure model of chemical bonding does
not at this level predict or interpret these bond angles. (This
will be considered here.) Therefore, although the Lewis
structure of methane is drawn as shown here.
Figure 19
This does
not imply that methane is a flat molecule, or
that the angles between
CHCH
bonds in methane is 90°. Rather, the structure simply reveals
that the carbon atom has a complete octet of valence electrons in a
methane molecule, that all bonds are single bonds, and that there
are no non-bonding electrons. Similarly, one can write the Lewis
structure for a water molecule in two apparently different ways,
shown here.
Figure 20
However, it is very important to realize that
these two structures are
identical in the Lewis model, because both show
that the oxygen atom has a complete octet of valence electrons,
forms two single bonds with hydrogen atoms, and has two pairs of
unshared electrons in its valence shell. In the same way, the two
structures for Freon 114 shown here are also
identical.
Figure 21
These two drawings do not represent different
structures or arrangements of the atoms in the bonds.
Finally, we must keep in mind that we have
drawn Lewis structures strictly as a convenient tool for our
understanding of chemical bonding and molecular stability. It is
based on commonly observed trends in valence, bonding, and bond
strengths. These structures must not be mistaken as observations
themselves, however. As we encounter additional experimental
observations, we must be prepared to adapt our Lewis structure
model to fit these observations, but we must never adapt our
observations to fit the Lewis model.
Extensions of the Lewis Structure Model
With these thoughts in mind, we turn to a set
of molecules which challenge the limits of the Lewis model in
describing molecular structures. First, we note that there are a
variety of molecules for which atoms clearly must bond in such a
way as to have more than eight valence electrons. A conspicuous
example is
SF6SF6,
where the sulfur atom is bonded to six F atoms. As such, the S atom
must have 12 valence shell electrons to form 6 covalent bonds.
Similarly, the phosphorous atom in
PCl5PCl5
has 10 valence electrons in 5 covalent bonds, the Cl atom in
ClF3ClF3
has 10 valence electrons in 3 covalent bonds and two lone pairs. We
also observe the interesting compounds of the noble gas atoms,
e.g. XeO3XeO3,
where noble gas atom begins with eight valence electrons even
before forming any bonds. In each of these cases, we note that the
valence of the atoms S, P, Cl, and Xe are normally 2, 3, 1, and 0,
yet more bonds than this are formed. In such cases, it is not
possible to draw Lewis structures in which S, P, Cl, and Xe obey
the octet rule. We refer to these molecules as "expanded
valence" molecules, meaning that the valence of the central
atom has expanded beyond the expected octet.
There are also a variety of molecules for
which there are too few electrons to provide an octet for every
atom. Most notably, Boron and Aluminum, from Group III, display
bonding behavior somewhat different than we have seen and thus less
predictable from the model we have developed so far. These atoms
have three valence shell electrons, so we might predict a valence
of 5 on the basis of the octet rule. However, compounds in which
boron or aluminum atoms form five bonds are never observed, so we
must conclude that simple predictions based on the octet rule are
not reliable for Group III.
Consider first boron trifluoride,
BF3BF3.
The bonding here is relatively
simple to model with a Lewis structure if we allow each valence
shell electron in the boron atom to be shared in a covalent bond
with each fluorine atom.
Figure 22
Note that, in this structure, the boron atom
has only six valence shell electrons, but the octet rule is obeyed
by the fluorine atoms.
We might conclude from this one example that
boron atoms obey a sextet rule. However, boron will form a stable
ion with hydrogen,
BH4-BH4-,
in which the boron atom does have a complete octet. In addition,
BF3BF3
will react with ammonia
NH3
NH3
for form a stable compound,
NH3
BF3NH3
BF3,
for which a Lewis structure can be drawn in which boron has a
complete octet, shown here.
Figure 23
Compounds of aluminum follow similar trends.
Aluminum trichloride,
AlCl3AlCl3,
aluminum hydride,
AlH3AlH3,
and aluminum hydroxide,
Al(OH)3Al(OH)3,
all indicate a valence of 3 for aluminum, with six valence
electrons in the bonded molecule. However, the stability of
aluminum hydride ions,
AlH4-AlH4-,
indicates that Al can also support an octet of valence shell
electrons as well.
We conclude that, although the octet rule can
still be of some utility in understanding the chemistry of Boron
and Aluminum, the compounds of these elements are less predictable
from the octet rule. This should not be disconcerting, however. The
octet rule was developed in Section 3 on the basis of the observation
that, for elements in Groups IV through VIII, the number of valence
electrons plus the most common valence is equal to eight. Elements
in Groups I, II, and III do not follow this observation most
commonly.
Resonance Structures
Another interesting challenge for the Lewis
model we have developed is the set of molecules for which it is
possible to draw more than one structure in agreement with the
octet rule. A notable example is the nitric acid molecule,
HNO3HNO3,
where all three oxygens are bonded to the nitrogen. Two structures
can be drawn for nitric acid with nitrogen and all three oxygens
obeying the octet rule.
In each structure, of the oxygens not bonded to
hydrogen, one shares a single bond with nitrogen while the other
shares a double bond with nitrogen. These two structures are not
identical, unlike the two freon structures in Figure 12, because the atoms are bonded
differently in the two structures.
Review and Discussion Questions
Problem 1
Compounds with formulae of the form
CnH2n+2CnH2n+2
are often referred to as "saturated" hydrocarbons.
Using Lewis structures, explain how and in what sense these
molecules are "saturated."
Problem 2
Molecules with formulae of the form
CnH2n+1CnH2n+1
(e.g.
CH3CH3,
C2H5C2H5)
are called "radicals" and are extremely reactive. Using
Lewis structures, explain the reactivity of these molecules.
Problem 3
State and explain the experimental evidence
and reasoning which shows that multiple bonds are stronger and
shorter than single bonds.
Problem 4
Compare
N2N2
to
H4N2H4N2.
Predict which bond is stronger and explain why.
Problem 5
Explain why the two Lewis structures for Freon
114, shown in Figure 21Figure 21, are identical.
Draw a Lewis structures for an isomer of Freon 114, that is,
another molecule with the same molecular formula as Freon 114 but a
different structural formula.
Comments, questions, feedback, criticisms?
Discussion forum
*
Join the discussion »
Send feedback
*
E-mail the author
*
E-mail Connexions
Creative Commons License
More about this content: Metadata | Version History | Cite This Content
This work is licensed by John S. Hutchinson. See the Creative Commons License about permission to reuse this material.
Last edited by Raymond Wagner on Jul 25, 2007 12:14 pm GMT-5.