Week 1 - Building Blocks

8/13/2019 Week 1 - Building Blocks

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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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Organic Chemistry/Evans 2

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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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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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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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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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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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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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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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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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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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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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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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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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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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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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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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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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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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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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
http://www.youtube.com/watch?v=nX9gjQYBh5Y


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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/

Week 1 - Building Blocks

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