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Introduction to the Building Blocks
Lewis structures describe the electronic structure of organic
molecules held together by covalent bonds.1Being as systematic as
possible, we can identify a relatively small number of generalized building
blocks that may be combined with one another to compose the structure ofany organic compound. In addition to forming a foundation for the
construction of Lewis structures, the building blocks also allow us to
transfer the properties of a structural pattern found in one molecule to
another--recognizing the same generalized building block (or chain of
connected building blocks) in two different molecules often allows us to
make comparisons between structures that seem, on the surface,
unrelated. In essence, beginning with the building blocks allows us to
generalize the organization of organic compounds and reactions by
functional group, using an additional level of abstraction. In this section,
we'll learn the generalized and particular building blocks and work through
some simple examples of interpreting and constructing Lewis structures.
Stable atomic fragments of second-row elements bear eight or (less
commonly) fewer total electrons. This dictum is an important rule in organic
chemistry known as the octet rule, and all the structures we will see and
draw will follow this rule. To see the rule in action, let's investigate the Lewis
structures of several important organic compounds. Although the octet rule
does have a theoretical basis, we'll develop it inductively, since that's howthe rule came about in practice. Even today, newly synthesized organic
compounds put the octet rule to the test and continue to support it
inductively.
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Figure 1. Lewis structures of well-known organic compounds. Can
you identify structural patterns within and between structures? What do all
of the heavy atoms (non-hydrogens) have in common?
To understand the octet rule systematically, we need to develop
definitions for some key concepts. Firstly, we need to understand what an
atomic fragmentis--note that we'll use this term interchangeably with
building blockin the future. An atomic fragment includes a single nucleus
and any electrons associated with that nucleus in the form of bonds (solid,
wedged, or dashed lines) and non-bonding electrons (dots). In the figure
above, I have boxed the atomic fragment associated with the nitrogen in
acetonitrile. The building block includes the nitrogen nucleus itself, its
associated lone pair of electrons pointing to the right, and all six electronsinvolved in its triple bond to carbon.
Secondly, we need to define the total electron count (TEC)of a
building block. Total electron count simply refers to the total number of
electrons contained within a building block. In our example of the nitrogen
in acetonitrile, the total electron count is
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TEC= 2 non-bonding electrons + 6 bonding electrons = 8 total
electrons
Try calculating the total electron count for the other building blockswithin acetontrile (there are six BBs in all). What do you notice about the
total electron counts of the carbons' building blocks? The total electron
counts of both carbon atoms are also 8. In fact, the total electron counts of
all building blocks belonging to the second-row elements--at least in the
examples here--are 8. What about the hydrogens' building blocks?
Hydrogen is a first-row element and has an analogous rule of its own, the
duetrule. Hydrogen in organic compounds is characterized by one
building block (with a TEC of 2) involving a single bond to the hydrogen
nucleus.
Figure 2.Deconstruction of acetonitrile into individual atomic
fragments or building blocks. Each building block supports either the octet
or duet rule.
Armed with the insight that building blocks must conform to the octetrule (an idea, you should recall, that we developed inductively) and the
notion that building blocks are centered on nuclei, we can enumerate all the
possible building blocks within stable organic molecules. But to truly
generalize the building blocks, we must also see that from the perspective
of electron counting, there is no difference between a bond and a lone
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Organic Chemistry/Evans 4
pair--both "count" as two electrons. Since the octet and duet rules are our
only mandates, it follows that it should not matter whether we place a bond
or lone pair at a particular position within a building block; that is, bothare
valid building blocks. Consider the two building blocks in the figure below.
Ignore the formal charge on the bromine atom, which is irrelevant to ourcurrent discussion. These two particular building blocks are incarnations of
a single, more general form we will call the generalized building block
(shown in blue below the specific forms). Either a single bond (to a different
atom A) or a lone pair of electrons may be placed at the end of each single
line in the generalized form. Double and triple lines refer to double and
triple bonds and must point to atoms.
Figure 3.Both of the bromine building blocks shown here have the
same total electron count (8). We can generalize the building blocks by
recognizing the electron-counting equivalence of single bonds and lone
pairs. Note that the same ideas do not apply to double and triple bonds!
It will be useful to introduce the concept of an electron-pair domain
(EPD)before enumerating the generalized building blocks. Electron-pair
domains are regions within building blocks that enclose either non-bonding
electrons or multiple bonds between the same two atoms. In acetonitrile,
nitrogen's building block bears two electron-pair domains, one associated
with the lone pair and one with the triple bond. We can organize the
generalized building blocks according to the number of electron-pair
domains they possess, and this is done in the table below. We will revisit
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Organic Chemistry/Evans 5
electron-pair domains when we discuss molecular geometry in the next
section; for now, recognize that EPDs represent a convenient way to
organize the generalized building blocks.
Figure 4.The generalized building blocks of organic Lewis
structures. Nuclei are represented with X's, and formal charges have beendeliberately omitted because the identity of the central atom is not specified
in these generalized forms.
Using a systematic process of construction, we can generate all of
the particular building blocks from the generalized forms. Now, let's clarify
the process of construction and examine the most prominent particular
building blocks, which incorporate specific atoms and formal charge.
Broadly speaking, the process of generating particular building blocks
involves adding single bonds and lone pairs according to the patternspecified by the generalized building block, then replacing the central atom
with an organic atom (hydrogen, carbon, nitrogen, oxygen, or a halogen)
and adjusting its formal charge based on its valence electron count (see
below). Let's carry out the process with the triple-bond-containing
generalized building block to produce all of its particulars.
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In the first step, we replace the single-line "placeholders" in the
generalized building block with either single bonds or lone pairs. There is
only one placeholder in our building block, so there are two ways to carry
out this operation: we can replace it with a single bond or a lone pair. We
will omit the hypothetical atom "A" from this point on. Often the lines drawnto electron pairs are omitted, to avoid confusing the lines themselves with a
bond.
Figure 5.The first step of generating particular building blocks in an
exhaustive way. Replace any single-line placeholders with bonds tohypothetical atoms (A) or non-bonding electron pairs.
Step two involves the replacement of the central atom "X" with
reasonable possibilities. To determine what reasonable choices for the
central atom are, we must make use of the concept of formal charge.
Formal charge is defined as the difference between the valence electron
count (VEC)of the atom in the building block and the valence electron
count of the free, neutral atom. Valence electron count refers to the number
of electrons that the central atom itself brings to the building block;
assuming that each atom of a two-electron bond contributes one electron to
the bond, we may define valence electron count mathematically as
VEC= (# of bonds) + (# of lone pairs)*2
The figure below shows that the valence electron count of X depends
on whether we place a lone pair or a single bond within the building block.
In the former case, the valence electron count of X is 5; in the latter case,its VEC is 4.
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Figure 6.Valence electron count reflects the number of electrons X
brings to the building block, assuming equal electronic contributions from
the two atoms of a bond. Non-bonding electrons belong to the atom on
which they reside.
What does the building block's valence electron count have to do with
reasonable choices for X? An important principle that we can again support
inductively is Pauling's rule of charges: formal charges in organic
molecules never possess magnitude greater than 1. In combination with
this idea, VEC tells us the central atoms we may use within a building
block. Consider our example of the 5-electron triply bound building block at
the left of Figure 5. Since the formal charge of X must be either 1, 0, or +1
(according to Pauling's rule), atoms allowed to replace X must bear 4, 5, or6 valence electrons in their neutral forms. Appealing to the periodic table,
we can see that only carbon, nitrogen, and oxygen are reasonable choices.
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Figure 7.Particular building blocks containing a triple bond and lone
pair. Particulars containing formal charges of magnitude greater than 1 areunreasonable.
An analogous treatment of the building block containing a triple bond
and a single bond (at right of the figures above) reveals that boron, carbon,
and nitrogen are reasonable possibilities for the central atom in this case. A
VEC of 4 for this building block suggests that reasonable possibilities for X
must bear 3, 4, or 5 valence electrons in their neutral forms.
Figure 8.Particular building blocks containing a triple bond and
single bond.
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Watch The Building-block Formalism: Essential Concepts
Watch The Generalized Building Blocks
Before continuing on, revisit the table of generalized building blocks
above and try producing particulars for a different generalized buildingblock. By systematically describing the process of moving from the general
to the particular, our aim is to develop a conceptual framework for organic
structure grounded in fundamental principles, like the octet rule and
Pauling's rule. In subsequent discussions, we will use the foundation to
discuss the behavior and reactivity of different structural types (i.e., different
building blocks) in a general way. Although a full enumeration of the 31
most common particular building blocks is provided below, take a few
minutes to review the previous discussion before examining the figure
below. Make sure you understand how each particular building block is
related to its generalized form.
Figure 9.The 31 most common particular building blocks of organic
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Organic Chemistry/Evans 10
chemistry, with their associated general forms. A section of the periodic
table is provided to aid in determining the valence electron counts of
neutral atoms. Please note that lines pointing to lone pairs designate only
positions in space, not bonds!
At this point, we can address a few points related to the particular
building blocks as a whole. Since the valence electron counts of elements
in the same group are the same, building blocks incorporating the period 2
elements (B, C, N, O, F) possess analogues involving the non-metallic
period 3 (Al, Si, P, S, Cl), 4 (Br), and 5 (I) elements. A few seemingly
reasonable possibilities are missing from the figure above, including the
triply bound boron we identified previously. I have elected to leave out
building blocks involving multiply bound boron since multiple bonds to
boron are only observed under certain structural conditions, which are
beyond the scope of our discussion. We should recognize at this point that
building blocks are invariant to translation and rotation, which may change
their appearance "on paper," but do not alter their chemical identity. Avoid
"tunnel vision" as you look for building blocks within Lewis structures.
The ultimate endgame of our systematic approach is to see the forest
for the trees when drawing and interpreting Lewis structures. Armed with
the generalized building blocks, we can recognize important similarities
between structures that may otherwise appear completely unrelated.Consider the examples of the imine and ketone functional groups provided
in the figure below. At first glance, the two compounds below look very
different; however, focusing on the nitrogen and oxygen atoms (the key
points of difference), we see that both of their building blocks are
associated with the same general form.
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Organic Chemistry/Evans 11
Figure 10.Key points of difference in acetone (left) and its imine of
ethylamine (right) involve the same generalized building block. These
Lewis structures are analogous in an important way!
In the next section, we'll establish connections between the
generalized building blocks and molecular geometry. It is an important
maxim that particular building blocks involving the same general form
possess the same geometry, when we take into account the spatial
positions of all bonds and lone pairs. In general chemistry you may have
dealt with a variety of geometries, including "t-shaped," "bent," "linear,"
"pyramidal," and others. Organic structures involve only three fundamental
geometric arrangements: linear, trigonal planar, and tetrahedral.
Watch The Particular Building Blocks
***
Molecular Geometry
Organic molecules, like the macroscopic objects we see and use
every day, possess defined shapes and certain rotational degrees of
freedom. Molecular shape and function are inexorably related, and in this
section we'll develop a systematic theory of molecular shape that we can
use to reason from the basic connectivity of a molecule (that is, itsconstituent building blocks) to its three-dimensional shape. Future sections
will develop relationships between molecular shape, spatial properties, and
reactivity. Keep these relationships in mind as the endgame of our
discussion here--although the systematic approach described here may
seem dry or theoretical, we will rely on it as a foundation for future studies.
Most generally, molecular shape is a property of a molecule as a
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Organic Chemistry/Evans 12
whole, not its individual atoms or building blocks. However, in a wide
variety of cases, we can think about the shape of a molecule as a kind of
"sum" of the geometries of its individual building blocks. Ignoring situations
in which resonance and steric hindrance are important, which we'll explore
later, the geometry of a molecule is a direct function of the geometries of itscomponent building blocks. Thus, we see that an understanding of the
geometries of individual building blocks leads us directly to the shapes of a
variety of organic molecules. The way we drew the building blocks in the
last section presupposed a particular geometry; we'll now put those
geometries in the spotlight. Let's first turn our attention back to the
generalized building blocks and examine their shapes. It is an important
fact that the geometries of all particular building blocks derived from the
same generalized form are the same--geometry is a conserved property of
the generalized form!
Countless observations of building-block geometry using the
experimental method of x-ray crystallographyhave revealed that three
geometries absolutely dominate organic molecules: linear, trigonal planar,
and tetrahedral. Theory has validated the existence of these three
geometries; however, we'll develop the fundamental theory associated with
molecular geometry in a separate section. For now, we can justify the three
key geometries using the electron-pair domain concept and valence shell
electron pair repulsion theory (VSEPR theory). According to VSEPRtheory, building-block geometry depends only on the number of electron-
pair domains associated with the building block. Valid geometries place
distinct electron-pair domains as far away from one another as possible.
Thus, two EPDs associated with the same central atom will naturally point
in opposite directions and form a linear geometry. Three EPDs will point to
the corners of a planar triangle, as any deviation from planarity brings the
domains closer to one another. Four EPDs will point to the four corners of a
tetrahedron to get as far away from one another as possible.2The figure
below outlines the shape and bond angles of the three key moleculargeometries of organic chemistry. Observed bond angles often vary slightly
from these theory-based ideals.
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Figure 11.The key geometries of organic building blocks. It is
worthwhile to commit these shapes to memory.
Importantly, the content of the EPD itself does not affect geometry;
geometry depends only on the total number of EPDs in the building block.
We can see this most prominently in the two-EPD case, for which there are
two generalized building blocks. One involves a triple bond and a single
bond (or lone pair); the other involves two double bonds. Although the
content of the EPDs is different in these two cases, their geometries are the
same (linear) because the total number of EPDs is the same for both
building blocks. We can now supplement our generalized building block
table from the last section with geometries. Of course, the ways the
building blocks were drawn in the last section also reflect these ideas, but
we now have a justification for how they are drawn and some terminologyto describe them.
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Figure 12.The generalized building blocks of organic chemistry, with
geometries. Wedges and dashes are used to show bonds pointing toward
you and away from you, respectively.
The next section describes the dynamic building-block formalism. To
this point, we've looked at the building blocks from a static perspective,ignoring how they might interact with one another or change. For a number
of reasons, our current perspective is incomplete. In order to deepen the
correspondence between our model and reality, we need to consider the
dynamicbehavior of building blocks. Doing so will allow us to enumerate
the possibilities for organic chemical change in an exhaustive manner and
gain a deeper understanding of the origins of molecular geometry. Before
moving on, please make sure to read and understand the shorthand
conventions used by organic chemists in Lewis structures.
***
Dynamic Building Blocks
Our current building-block system, which describes molecules as
arrangements of static building blocks containing electrons localized in
electron-pair domains (EPDs), is incomplete. We can demonstrate its
incompleteness by exploring the relationship between apparent number of
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Organic Chemistry/Evans 15
EPDs and geometry in greater detail. The table we developed in the last
section implies that there is a direct relationship between number of EPDs
and geometry: 4 EPDs means tetrahedral, 3 means trigonal planar, and 2
means linear geometry. But crystal structures of a variety of organic
compounds have shown that this trend does not hold up under all structuralconditions. Certain atoms, like nitrogen within the amide functional group,
possess geometries that are inconsistent with their numbers of EPDs in
completely neutral, systematically built Lewis structures. Although we would
expect nitrogen in the figure below to be tetrahedral, in practice it is trigonal
planar.
Figure 13.The empirically observed geometry of amide nitrogens
differs from our expectation based on the building-block formalism. How
can we enrich our current system to account for inconsistencies like theone seen here?
A second difficulty of the static building-block formalism concerns its
inability to account for chemical change. Chemical reactions are the result
of reorganizations of electrons--using building-block terminology, we might
say that chemical reactions occur when building blocks change. Yet, with
just the building blocks themselves in hand, we can't make predictions
about how they might change into one another. Recognizing that our
current system incompletely describes reality, we need to advance ourunderstanding by considering the dynamicbehavior of building blocks. In
this section we'll explore and systematize how electrons within the building
blocks can move, both internally (to and from building blocks to which they
are directly connected) and externally (to and from building blocks in
entirely separate molecules).
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Like water flowing down a hill, electrons flow from regions of high
potential energy to low potential energy, or from sources to sinks. What is
"electronic potential energy"? Intuitive ideas that you may already have
concerning the energy of charged particles apply to electrons in this
context: Negatively charged electrons repel one another, and thus have high
potential energy when confined to a small space.
Negatively charged electrons closer to the positively charged nucleus
have lower potential energy
Electrons associated with atoms with high effective nuclear charge
have lower potential energy.
Remarkably, we can use the Lewis structure of an organic compound
as a reasonably reliable "map" of its electronic potential energy. Regions
within a molecule where electronic potential energy is high are called
electron sources(or electron donors), while regions where electronic
potential energy is low are called electron sinks(or electron acceptors).
The dynamic behavior of our building blocks can be completely described
by the idea that electrons move from electron sources to electron sinks. To
illustrate this concept graphically, let's identify the elements of Lewis
structures that represent electron sources and sinks, then illustrate how
sources and sinks interact within and between molecules. Now is a good
time to read and understand the curved-arrow formalism.Electron sources are concentrated regions of electron density in
molecules. Being exhaustive, we can say that any pair of electrons (bonds
or non-bonding lone pairs) may serve as an electron source. However,
more "localized" electrons tend to be better electron donors than less
localized electrons with more room to "spread their legs," or less electron-
electron repulsion. For this reason, non-bonding lone pairs, which are
localized on a single atom (according to our formalism thus far), tend to be
the best electron sources. Lone pairs are followed in reactivity by !bonds
(the second and third bonds of double/triple bonds) and "bonds (singlebonds), respectively. Thinking of a "bond as an electron source is
relatively rare, but still very important. In a later section, we will clarify these
trends and the labels used in the figures below using molecular orbital
theory. For now, it's important just to recognize electron sources within the
building blocks. Build your pattern recognition skills now to establish a solid
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foundation for learning later!
Figure 14.The three classes of electron sources within building
blocks. The labels n, !, and "correspond to non-bonding lone pairs,
multiple bonds, and single bonds.
Electron sinks are a bit more difficult to spot, as they don't correspond
to bonds or lone pairs within Lewis structures. An electron sink is an atom
or functional group with the ability to gain additional electrons. Using
building block terminology, we can identify two ways in which an electron
sink might gain electrons. The first involves an increase in total electron
count, and only applies to building blocks with fewer than eight total
electrons. Carbocations are famous for this type of electronic inheritance.
In the figure below, the six-electron carbocation building block becomes aneight-electron, tetrahedral building block after donation from a lone pair
associated with bromide anion. Donation from an internal electron source,
establishing a new double or triple bond, is also possible. In cases when
the electron-accepting building block has six or fewer total electrons, we
call the electron sink a("a" for atomic!).
Figure 15.Building blocks bearing seven or six total electrons can
inherit one or two electrons from an electron source to bring their total
electron counts up to eight. These examples illustrate donation from an
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Organic Chemistry/Evans 18
external (left) and internal (right) electron source. Generalized building
blocks for the atom gaining electrons are shown below each structure.
The examples above suggest that building blocks bearing fewer than
eight electrons tend to be good electron sinks, which is true. However, wewould be mistaken to conclude that onlyelectron-deficient building blocks
can serve as electron sinks. The vast majority of stable organic compounds
are composed of eight-electron building blocks, so in order to explain their
reactivity, we need to understand how these atomic fragments may serve
as electron acceptors. The key idea here is that electronegativeatoms
can take up pairs of electrons residing in bonds as localized lone pairs.
This is a type of electron acceptance available for eight-electron building
blocks. In Figure 16, we can see that moving a bond between two atoms
onto a single acceptor atom does not change the total electron count of the
acceptor. Now, examine the atom that gave up the electrons in the bond--it
bears six total electrons and a positive charge, indicating that it's able to
inherit two more electrons. In essence, the electronegative atom pulls
electrons toward itself, allowing the other atom in the bond to gain electrons
from somewhere else. Because the acceptor atom is electronegative, it's
able to bear a negative charge.
Figure 16.Oxygen as an electron sink in carbonyl compounds. Note
that the carbon that gives up electrons ends up with a total electron count
of 6, so it's able to accept electrons from a separate donor.
Electronegative atoms can inherit electrons from single or multiple
bonds. In all of these cases, although we call the electronegative atom the
sinkper se, the atom to which it's bound is the one that is actually able to
gain electrons from a source. The electronegative atom just gains electrons
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Organic Chemistry/Evans 19
internally. The figure below illustrates the three modes of electron
acceptance in which eight-electron building blocks can engage. Please
note that the figure below is not meant to feature full building blocks; atoms
X and Y may possess additional bonds and lone pairs.
Figure 17.Electronegative atoms Y as electron sinks. The labels "*
and !* indicate the nature of the bond whose electrons are given to Y.
We're now ready to identify electron sources and sinks in molecules,
to predict how electrons may flow within and between them. Within
molecules, sinks and sources adjacent to one another can interact. We use
curved arrows to represent interactions between sources and sinks within
molecules; curved arrows also depict the interconversion of equivalent
Lewis structures. In the next section, we'll explore the equivalence of Lewis
structures (called resonance) in more detail. Between molecules, electron
flow from sources to sinks describes the mechanisms of organic chemical
reactions, or chemical change. The distinction between internal electronic
interactions (resonance) and external electron flow (reactivity) is important,
because although both forms of electron movement look similar, thephenomena they represent are fundamentally different. The distinction can
be a source of confusion for students and experts alike!
Let's now explore at a few examples of resonance and reactivity.
Firstly, let's return to the amide functional group that we saw at the
beginning of this section. At this point, it should be clear that the amide
contains a good electron source (nitrogen's lone pair) next door to a good
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Organic Chemistry/Evans 20
electron sink (the C=O double bond). The left half of Figure 18 shows
curved arrows and resulting Lewis structures for donation from nitrogen
and acceptance by the carbonyl oxygen. It's apparent that that these
separate, isolated arrows present some problems for the amide. Just
drawing electron donation results in a disturbing ten-electron building blockin stark violation of the octet rule (no way!). On the other hand, just drawing
electron acceptance results in a questionable six-electron carbocation
building block (not best). Combining both of these movements into a single,
coupled movement of electrons from source to sink produces the best
alternative Lewis structure, which includes an octet of electrons on every
atom (best).
Figure 18.Internal interactions between an electron source
(nitrogen's lone pair) and sink (the C=O bond) in the amide functional
group. Notice that in the best alternative Lewis structure, we would expectnitrogen to be trigonal planar, not tetrahedral!
Pay close attention to the nitrogen atom's building block in the
original and best alternative Lewis structures. There seems to be a
geometric issue here--the alternative form suggests that the nitrogen
should be trigonal planar, but the original Lewis structure suggests that it
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Organic Chemistry/Evans 21
should be tetrahedral. We saw this issue at the beginning of this section,
but we can now see why the observed trigonal planar geometry makes
sense: donation of nitrogen's lone pair into the C=O electron sink
influences nitrogen's geometry. On your own, try confirming that a trigonal
planar geometry at nitrogen brings the lone pair closer to the !bondrelative to a tetrahedral geometry (note that the lone pair sits perpendicular
to the trigonal plane).
Acylium ions are interesting intermediates in several reactions, most
notably Friedel-Crafts acylation of aromatic compounds. The figure below
depicts an acylium ion and one of its alternative resonance forms. The
adjacent source and sink in this case are two !bonds: a CC double bond
and the CO triple bond. Using the terminology already developed to
describe sources and sinks, we can describe the electronic interaction
captured here as a !#!* interaction. The source is listed first, before the
arrow, and the sink after the arrow.
Figure 19.An alternative Lewis structure for an acylium ion, reflecting
a !#!* electronic interaction.
Halogen atoms are ubiquitous in organic chemistry, and are famous
as electron sinks. The reactivity of alkyl halides in the presence of electron
sources provide evidence that halogen atoms tend to be excellent sinks.
The SN2 substitution reaction involves the simultaneous donation of a pair
of electrons from a source and acceptance of a pair of "bonding electrons
by a halogen atom. The curved arrows in the figure below portray a
reaction mechanism, and are different from the internal arrows in thefigures above. While the alternative Lewis structures in the figures above
are simplifications of a single, more complex reality, the structures on either
side of the single-headed arrow in the figure below are truly unique
chemical species. Nonetheless, we can use similar notation to denote the
electronic interactions in all three figures. In the figure below, we can
describe the curved arrows as representing an n#"* interaction.
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Figure 20.Representing chemical change as electron flow in the SN2
substitution reaction. Notice the changes that occur in the building blocks
associated with sulfur and bromine as the reaction takes place.
Eliminations are a second important class of reactions often observed
for alkyl halides. As in substitution reactions, the halogen atom serves as
an electron sink in elimination reactions. E2 elimination involves thesimultaneous donation of electrons from a "bond and acceptance of
electrons by a halogen atom. This electronic interaction is internal, which
suggests that chemical change might not be taking place. However, a base
is required for the reaction to occur, so an external n#"* interaction also
plays a role in the mechanism.
Figure 21.Representing chemical change as electron flow in the E2
elimination reaction. These curved arrows involve a combination of external
and internal electronic movements; because of the external component, E2
elimination is certainly considered chemical change.
As you study these examples, keep in mind that our goal is the
systematization of organic chemical structure and change. This section has
introduced the three electron sources (n, ", !) and three electron sinks (a,
"*, !*) of organic chemistry. These sources and sinks form a complete
system for describing the structures of even-electron molecules and polar
chemical change, and eventually, we will connect the cryptic labels defined
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in this section to molecular orbital theory. For the time being, recognize our
system of dynamic building blocks as a way to classify, categorize, and
otherwise organize your knowledge of organic chemistry. You'll be exposed
to a vast collection of functional groups and reactions throughout this book,
but their similarities (and differences) can be understood in the light of thesystem developed in this section. I hope to demonstrate this point
throughout the remainder of the text.
In the next section, we will explore internal electronic interactions in
more detail and develop the theory of resonance forms. Like two paintings
of the same model, resonance forms are alternative representations of the
same physical molecule. Although we already know how to depict the
interconversion of resonance forms using internal curved arrows, in the
next section we'll develop heuristics for understanding what makes a
resonance structure "good" or "bad."
Watch The Dynamics of Building Blocks
Watch Structural Analogies
***
Resonance Theory
Organic molecules are completely defined by the connectivity and
spatial positions of their nuclei and the number of electrons they contain--
we'll call this the molecular identity principle. Molecules with the sameatoms, atomic positions, and total number of electrons are identical. This
does not imply, however, that a single Lewis structure can always fully
specify the identity of a molecule! In fact, many molecules (most, some
would argue) are best drawn using multipleLewis structurescontaining
different arrays of building blocks, but identical connectivity and total
number of electrons. Consider the two partially drawn molecules in the
figure below. A connectivity map and total number of valence electrons are
provided for each molecule. Based on the molecular identity principle, the
information provided in the figure below is sufficient to conclude that thetwo molecules are the same, no matter how we decide to "scatter" missing
valence electrons about each structure. Of course, the incomplete
molecules look identical now; the important point is that they will remain
identical, even if we end up distributing their remaining valence electrons in
different ways. According to the molecular identity principle, we already
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Organic Chemistry/Evans 24
have enough information to conclude that these molecules are the same,
without providing missing multiple bonds and lone pairs!
Figure 22.Connectivity maps, numbers of valence electrons, and
formal charges for two identical molecules. The molecular identity principletells us that no matter how we complete the two Lewis structures, both will
represent the same compound.
Let's now proceed to complete the two Lewis structures
independently. When given a connectivity map or !skeleton, the most
straightforward way to complete a Lewis structure is to...
1) decorate the skeleton with multiple bonds and lone electron pairs until
each building block conforms to the octet rule and all valence
electrons are accounted for;2) adjust the formal charges of any atoms bearing more or fewer
valence electrons than their neutral, elemental forms.
In our example, 16 valence electrons are accounted for by the "bonds
implied by the given skeleton. We are left with 8 electrons for multiple
bonds and lone pairs. At this point, we should recognize that the "skeleton
also implies a geometry for each atom, so we're limited in our choices of
generalized building blocks. Carbon 1 has tetrahedral geometry and since it
already bears 8 total electrons, we should leave carbon 1 alone. Carbon 2
appears to have trigonal planar geometry and three EPDs; to ensure that
C2 satisfies the octet rule, we must add a double bond to it. Carbon 3
appears to be in a similar situation--it likewise needs a double bond to
satisfy the octet rule--so we can draw a double bond between carbons 2
and 3 to address both of these issues at once. This leaves us with 6
valence electrons to place on the structure. To decide where to place these,
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Organic Chemistry/Evans 25
we should note that all of the building blocks we've established so far bear
8 total electrons, so it's unreasonable to add more electrons to the carbon
atoms. Only the oxygen atom doesn't satisfy the octet rule. Placing all 6
electrons on oxygen as three lone pairs, we see that its building block now
satisfies the octet rule. Oxygen bears 7 valence electronsin this buildingblock but only 6 when neutral; thus, the oxygen atom has a formal charge
of 1.
Figure 23.Building a Lewis structure using a "skeleton. As we add
lone pairs and multiple bonds to the skeleton, we check the total electron
count of each building block to ensure that it conforms to either the octet or
duet rule. We know we're done when all the building blocks conform to the
octet rule and all valence electrons are accounted for.
To construct an alternative possibility, we'll start with the "skeleton
again. Instead of establishing a double bond between carbons 2 and 3, let's
place a double bond between carbon 2 and oxygen to satisfy the octet rule
on carbon 2. We could then add two lone pairs to the oxygen atom to
establish an octet of electrons there. This leaves us with a pair of electrons
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Organic Chemistry/Evans 26
unaccounted for, which we can place on carbon 3 as a lone pair
perpendicular to the trigonal plane. Doing this seems inconsistent with the
trigonal planar geometry of carbon 3--are we "allowed" to place a lone pair
on the six-electron building block, while keeping its geometry trigonal
planar? The answer is yes, and we will soon see why this is acceptableunder certain conditions.
Figure 24.An alternative, equally correct Lewis structure may be built
by first establishing a double bond between carbon and oxygen, then
invoking an exotic building block to place a lone pair (and octet of
electrons) on carbon 3. Is this building block acceptable, or have we made
a mistake?
The two Lewis structures we've drawn are two distinct
representations of the same molecule. To demonstrate this, we can carry
out a straightforward set of experiments. Let's begin with two compounds,
each of which contains a trimethylsilyl (TMS) group. Upon treatment with
fluoride anion (in the form of tetrabutylammonium fluoride), the TMS cation
is displaced from each compound, leaving two anionic molecules behind. If
our two anions are different, we should expect two different products to
form upon treatment with methyl bromide--the methyl group may end up
attached to either oxygen or carbon 3, depending on the "location" of the
negative charge. But in practice, only the C3methyl isomer is observed
under both sets of reaction conditions. How on earth could the starting
material containing the oxygen-boundTMS group produce only a product
methylated on carbon 3?3
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Organic Chemistry/Evans 27
Figure 25.Experiments demonstrating that the two bracketed anions
are, in fact, representations of the same compound. Regardless of the
reagent used in step 2, the distribution of products is the same for bothstarting materials!
The simplest explanation posits that the two bracketed structures
actually represent the same compound. No matter what reagent we use in
the second step, the product distribution we observe is the same for both
starting materials. It is impossible to "independently generate" one structure
or the other. Yet the two structures do notinterconvert through an extremely
rapid equilibrium process--theory shows that they represent the same
underlying arrangement of nuclei and electrons! Each Lewis structure is an
incomplete description of the single moleculedefined by the "skeleton in
the first figure bearing 24 valence electrons (recall the molecular identity
principle). In reality, the single molecule is intermediate between the two
resonance forms, with partial properties of both (Figure 26).
Figure 26.The actual structure of our anion is a hybrid of its two
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Organic Chemistry/Evans 28
most important resonance forms. The $signs indicate partial charges
(between 0 and 1) on carbon and oxygen. The dotted lines indicate partial
double bonds (bond order between 1 and 2).
The two Lewis structures we just generated bear a specialrelationship: they're called resonance formsor resonance contributors
(since they "contribute" to the real structure of the molecule). In the
remainder of this section we'll develop the theory of resonance forms. Let's
begin with a key question: in light of the vast expanse of possible
resonance forms for many organic molecules, how can we distinguish
"good" resonance structures from poor ones? How do we know what
resonance forms to focus on?
The key to identifying important resonance forms is to recognize
active electron sources and sinks adjacent to one another in Lewis
structures. When active sources and sinks are next to one another,
resonance is important. We saw this idea in the Dynamic Building Blocks
section, when we discussed "alternative Lewis structures"--by now, you've
probably realized that those are just resonance forms! Review that section
if you need to jog your memory concerning the use of the curved-arrow
formalism to show the interconversion of resonance forms. At this point, we
need to establish a satisfactory definition of activity. What makes a source
or sink "active"? First of all, the type of source or sinkmatters greatly, asenergies vary as a function of source/sink type. Figure 27 outlines the
relative activity of the different types of sources and sinks--we'll explain this
ordering in a future section on frontier molecular orbital theory.
Figure 27.Activity as a function of source/sink type. This ordering is
explained in more detail in the chapter on Frontier Molecular Orbital Theory.
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Organic Chemistry/Evans 29
What can we pull from this figure? For one thing, resonance is almost
always irrelevant for compounds containing only "bonds, since these
compounds lack lone pairs, empty atomic orbitals, !bonds and !
acceptors. Conversely, resonance is almost always important when a non-bonding lone pair and an empty atomic orbital find themselves next to one
another, since these are the most active sources and sinks. The figure
below shows a classic example of this kind of resonance. The resulting
resonance structure is quite important, since every atom in it is neutral!
Figure 28.A lone pair on oxygen and an empty 2porbital on carbon
are adjacent, so resonance is critical here.
When source/sink type is the same but atom types differ,
electronegativityand (more generally) charge stabilizationare the key
factors that allow us toreason about the relative importance of resonance
structures.4The resonance form featuring more stable charges is the more
important contributor to its true molecule. Imagine pushing the C=O bond inFigure 28 toward carbon to generate a carbanionic resonance form.
Electronegativity helps us explain why this form is completely irrelevant to
the true nature of the carbonyl group, while pushing toward oxygen
produces a much more important resonance form. Now, consider the two
sets of resonance forms in the figure below. Which resonance form
containing + and charges is the greater contributor to the corresponding
real molecules?
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Organic Chemistry/Evans 30
Figure 29.Resonance structures of the carbonyl and imine functional
groups. Which dipolar resonance form is more important to its
corresponding real molecule?
Since oxygen is more electronegative than nitrogen, the dipolarresonance form of the carbonyl is more important than that of the imine.
We might illustrate this difference by saying that the true carbonyl group is
a 10:90 mixture of its resonance forms (say), while the true imine group is a
5:95 mixture of its contributors. Note the smaller contribution of the dipolar
form of the imine relative to the dipolar form of the carbonyl.
When judging multiple resonance forms of the same true molecule to
determine the most important resonance contributor, the ideas above
(source/sink type, charge stabilization) still apply.5In addition to these
concepts, we need to consider charge separationwithin the different
resonance forms. Separation of opposite charges is, generally, a bad thing
when it comes to resonance forms.6Thus, rather unsurprisingly, the neutral
resonance form of the carbonyl (lacking charge separation) is more
important than its dipolar form (Figure 28). More important resonance forms
have opposite charges closer to one another, other things being equal. On
the other hand, separation of like charges is good!
Let's summarize what we've seen so far. Better resonance
contributors involve... Less separation of charge (none, if possible!)
Negative charge on more electronegative atoms, and positive charge
on less electronegative atoms
These principles help you evaluate given resonance forms, but they
don't help you spot when resonance is important for neutral molecules.
"Active" sources and sinks are the key to doing that, and these simply
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Organic Chemistry/Evans 31
reflect the principles just outlined for important resonance contributors.
Fundamentally, active sources and sinks interact to yield important ("good")
resonance forms. Figure 30 summarizes activity trends for electron sources
and sinks.
Figure 30.Activity trends for electron sources and sinks potentially
involved in resonance.
Before closing this section, an important note about geometry is in
order. We've seen that sources and sinks must be adjacent to one another
in order to interact. Adjacency is not the only geometric requirement for
resonance, although situations where geometry is important are rare. We'llrevisit this issue in a future section on frontier molecular orbital theory, after
we've discussed orbital shapes. For now, we need only note that
resonance has important effects on a building block's geometry--
specifically, the geometry of lone-pair-bearing atoms involved in resonance
as electron sources. Finally, we can resolve the apparent geometric
problems described at the beginning of the previous section. For reasons
that will become clear later, trigonal planar atoms are more effective
electron sources than tetrahedral atoms. For this reason, atoms like amide
nitrogens (Figure 13) and carbon 3 in Figure 24 exhibit trigonal planar (nottetrahedral) geometry.
Watch Resonance and the Building Blocks
1. Many organic compounds also feature ionic bonds; however, we can
think of ionic bonds simply as electrostatic forces between oppositely
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Organic Chemistry/Evans 32
charged particles, which are themselves composed of covalent building
blocks. The building blocks are the loci of complexity in organic
compounds; thus, we focus on these in our systemization of organic
structure.
2. You might verify on your own that four EPDs in a square plane are closerthan those in a tetrahedral geometry!
3. Examples of this phenomenon abound in the chemical literature. Even
more compelling examples involve reactions that yield a mixtureof
products. Product ratios from independently generated resonance forms
are universally equal--if they aren't there is almost always an alternative
explanation for the difference.
4. To read more, visit this pageby James Ashenhurst.
5. We're often interested in a related problem worded slightly differently: the
"second-best" resonance contributor following the completely neutral form.
The reason is that the second-best resonance form of a molecule usually
reveals much of its reactivity.
6. James Ashenhurst calls this the Rule of Least Charges.
http://masterorganicchemistry.com/2011/12/08/evaluating-resonance-forms-1-the-rule-of-least-charges/http://masterorganicchemistry.com/2011/12/12/evaluating-resonance-structures-2-applying-electronegativity/Recommended