week 1 - intro to orbitals

8/13/2019 Week 1 - Intro to Orbitals

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Beyond Lewis Structures

The building-block formalism of organic chemical structure, more

commonly called the "Lewis dot model," is a powerful schema for

organizing organic molecules and identifying similarities between them.

Coupled with the generalized stability trends, we can begin to reason fromthe seen to the unseen to make predictions about never-before-seen

chemical reactions. However, there are many observations that the Lewis

dot model cannot explain. Its explanatory power doesn't even come close

to its organizational strength. In fact, most stability trends find their origin

not in the Lewis dot model, but in a deeper theory of molecular structure:

molecular orbital theory. In this chapter, we'll develop a version of

molecular orbital theory useful for the student and everyday practitioner of

organic chemistry. Realize that we will notdisprove the Lewis dot model in

this chapter! Molecular orbital theory is meant to enhance, not replace the

Lewis dot model. Its organizational power still remains, and we will still rely

on the building-block formalism to make connections between analogous

structures in later discussions.

Before diving in to the details of molecular orbital theory, let's explore

some of the shortcomings of the Lewis dot model. What can't it explain,

and why is an additional theory necessary? Consider the process in Figure

1a, rotation about a carbon-carbon double bond. The Lewis dot model

stipulates that the four electrons in the double bond are shared betweenthe two carbon atoms, but it can't explain why rotation about the double

bond doesn't occur until the compound is heated to extremely high

temperatures. Molecular orbital theory reveals that carbon-carbon multiple

bonds cannot rotate without breaking.


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Organic Chemistry/Evans 2

Figure 1.Observations that the Lewis dot model does not explain

adequately. (a) Rotation about double bonds does not occur at room

temperature, even though rotation around unhindered single bonds is rapid

at room temperature. (b) One-step substitution reactions result in an

inversion of configuration (two atoms appear to switch places) at the

electrophilic carbon. (c) Carbocations substituted with more carbon atoms

are more stable than less substituted cations.

The single-step SN2 reaction in Figure 1b has properties that are not

addressed by the Lewis dot model, too. Displacement of bromide by

hydroxide leads exclusively to an inversion of configuration at the central

carbon atom. This observation suggests a particular trajectory for hydroxide

as the substitution takes place, but the Lewis dot model offers no dominant

path. The spatial aspects of the reaction are best explained by molecular

orbital theory--in fact, simple steric considerations may argue against the

observed configuration of the product!Finally, Figure 1c shows the relative stability of two carbocations.

More substituted cations are more stable than those that are less

substituted, other things being equal. Yet, aside from the somewhat hand-

wavy explanation of steric hindrance due to flanking CH bonds, the Lewis

dot model cannot account for the exceptional stability of the left-hand (tert-


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butyl) cation. Orbital interactions, which find their theoretical basis in

molecular orbital theory, elegantly explain this ubiquitous observation.

To account for all of the observations in Figure 1, we need a model

that suggests the spatial positions and energies of electrons within a

molecule. Molecular orbital theory provides this information--in fact, anorbitalis just an electronic container with an associated energy value. In

the next section, we'll introduce the fundamentals of the orbital concept.

Throughout the following sections, keep in mind our goal of describing in

detail the electronic structure and reactivity of organic molecules.

***

Introduction to Orbitals

In this section, we'll begin to extend the Lewis model by expanding

our conception of the electron. While the Lewis dot model stipulates that

electrons sit, isolated, on or between atoms, the quantum mechanical truth

is far more interesting: electrons may be delocalized over large regions of

space, and even across several atoms! Belonging to each electron in a

molecule is a function called an orbital, which describes all of its properties

according to the principles of quantum mechanics.

Strictly speaking, an orbital is a function over all space that specifies

the probability of finding an electron at a given point.1Orbitals for atoms

and molecules take into account the positions of nuclei and other electrons

nearby. In practice, probability values rapidly approach zero at certaindistances from the nucleus. For this reason, we can focus only on a "slice"

of space, a container that encloses all of the points for which the probability

of finding the electron is larger than some cutoff value.2Organic chemists

most often think of orbitals in this way, as containers in which electrons are

likely to be found. A single orbital can contain up to two electrons of

different electronic spin. We typically refer to the different spins of the

electron as "spin up" and "spin down" or +1/2 and 1/2. To depict the

occupancy, energy, and shape of orbitals, chemists use orbital diagrams.

A simple orbital diagram is shown in Figure 2. Notice that the number ofelectrons, energy, and appearance of the orbital are all shown on the orbital

diagram. A horizontal line is drawn to indicate the position of the orbital on a

vertical energy scale. Half arrows up and down are drawn to show how

many electrons are in the orbital, and what their spins are. Often, the shape

of the orbital will be omitted or drawn in a separate picture of the molecule


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under study.

Figure 2.The three essential features of every orbital, with an

example of an orbital diagram.

Orbitals can take on either positive or negative values over space.

This seems at odds with our earlier definition of the orbital as a distribution

of probabilities: how can a probability be negative? We now need to

expand our previous definition: the magnitudeof the orbital is a probability,

and its sign refers to a property called phase. Phase is important when

orbitals combine or overlap, because they do so like waves, enhancing oneanother in some regions and canceling one another out in others. In

pictures of orbital shapes, positive and negative phases are indicated either

with different colors or with shading. In Figure 2, for example, we might

choose the unshaded area to represent a region of positive phase and the

shaded area to represent a region of negative phase. Our exact choice is

arbitrary, but it's important to keep in mind that the shaded and unshaded

regions are places where the signs of the orbital are different. We'll revisit

phase shortly, when we discuss how orbitals combine with one another.

Orbitals are special functions that apply to the electron. But where do

the functions themselves come from? What's the origin of the shapes and

energies of orbitals? We won't follow this line of inquiry too far, except to

say that the orbital belonging to an electron depends on the electron's

environment. Designations like "atomic" and "molecular orbitals" reflect the

environment of the electron. Atomic orbitals belong to electrons in atoms,


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molecular orbitals belong to electrons in molecules, et cetera. The nature of

the environment and a quantum mechanical relation called the Schrdinger

equation are used to solve for the form of the orbital. We'll largely treat the

shapes and energies of atomic orbitals as axiomatic--that is, we'll present

them without proof. Such proofs are generally reserved for physicalchemists...and organic chemists are glad to give them up!

Watch The Orbital Concept

Let's begin with a look at the atomic orbitals (AOs)that characterize

the first- and second-row elements. There are five that are important for

organic chemistry: the 1s, 2s, 2px, 2py, and 2pzatomic orbitals. Their

shapes and relative energies are shown on an orbital diagram in Figure 3.

At the center of each 2porbital, where the nucleus sits, we can identify a

node, where the phase of the orbital changes. At the node, the value of the

orbital is zero. These five orbitals form a sort of scaffold, into which we can

add electrons to characterize the organic atoms.

Figure 3.The atomic orbitals on an orbital energy diagram, with
http://www.youtube.com/watch?v=PEkvPEt7oTk


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electronic configurations of the first- and second-row elements.

We use the term electron configurationto refer to how an atom's

electrons occupy the atomic orbitals. In the blue box in Figure 3.3, the

configurations of the first- and second-row elements are provided. The firstnumber in each orbital's name is its principal quantum number, a measure

of its energy (as principal quantum number goes up, energy does too). The

letters sandpindicate the orbital's subshell (an indicator of orbital shape

and energy). Finally, the superscripted 1's and 2's show the number of

electrons in each AO. Notice that electrons occupy the most stable orbitals

first, followed by more unstable orbitals. The three 2porbitals are all

degenerate--that is, they have the same energy. Degenerate orbitals are

filled one electron at a time, so that two electrons are not paired up until

they must be. Figure 4 illustrates how electrons fill the AOs of boron,

carbon, nitrogen, and oxygen. Notice that two electrons don't occupy the

same orbital until it's unavoidable. This method for filling orbitals with

electrons, which holds regardless of the type of orbital (atomic, molecular,

etc.) is called Hund's rule.

Figure 4. Orbital energy diagrams for B, C, N, and O atoms. Hund's

rule states that electrons should not be paired up until no other placement

is possible.


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Let's take a breather--what's the point of all this atomic orbital

mumbo-jumbo? With some notable exceptions, we won't deal directly with

the atomic orbitals once we've fully developed molecular orbital theory.

However, atomic orbitals are quite literally the building blocksof molecular

orbitals, which are our ultimate endgame. We'd like to understand theelectronic structures of molecules, but to do so, we need to recognize how

electrons are arranged in atoms. In particular, the molecular orbitals we

tend to care about are built from the valence atomic orbitals, an atom's

occupied orbitals with highest principal quantum number.3

Early in this section, we noted that the shape and energy of an

electron's orbital depends on the system of which it's a part. Atomic orbitals

pertain to atomic systems, molecular orbitals to molecular systems, etc. An

orbital is a solution to a problem that takes the system into account. Thus,

from first principles, we might imagine that molecular orbitals, the solutions

to molecular problems, are unrelated to atomic orbitals, the solutions to

atomic problems. Yet, remarkably, we find in practice that linear

combinations (weighted sums) of the atomicorbitals approximate solutions

to the problem of describing electrons in molecules. This approximation is

the cornerstone of the linear combinations of atomic orbitals-molecular

orbitals (LCAO-MO)method, which we'll explore in the context of

dihydrogen in the next section.

Watch The Atomic Orbitals

***

Molecular Orbitals of Dihydrogen

Our goal so far has been to describe the spatial positions and

energies of electrons in atoms and molecules, using the principles of

quantum mechanics as a foundation. Quantum mechanics sets up the

problem of determining an electron's orbital as a kind of physical

optimization problem: what spatial arrangement of electrons leads to the

most stable system? In the last section, we were introduced to the atomicorbitals, which are solutions for atomic systems of the first and second rows

of the periodic table. Thankfully, we do not need to start over at square one

to determine the orbitals of electrons in molecules! Molecular orbitals are

well approximated by linear combinations of the atomic orbitals associated

with the molecule's atoms.
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Keep our goal in mind throughout this section. We want to answer the

following questions:

Where are the electrons in a molecule located in space? What are

their energies?

We've already seen the containers that hold electrons in atoms--the

1s, 2s, and 2patomic orbitals. Molecular orbitals are constructed from

these building blocks, and have spatial and energetic properties that reflect

their construction from atomic orbitals. In fact, there are multiple ways to

carry out the process of building molecular orbitals, but we will adopt a

localized molecular orbital theoryapproach using the ideas of

hybridization and localized MOs. Localized MO theory breaks from the

"canonical" mold used by physical chemists, but will serve us well

throughout future discussions. Let's begin with a very simple case: the

molecular orbitals of the dihydrogen molecule, H2.

The hydrogen atom possesses a single electron in a 1satomic

orbital. When two hydrogen atoms come together, a molecule containing a

total of two electrons results (H2). Thus, we might imagine that only one

molecular orbital is needed to accommodate these two electrons. However,

this conclusion ignores the important fact that hydrogen's valence AO (1s)

is unfilled. Both the hydrogen atom and the dihydrogen molecule canaccept electrons. We could imagine giving one more electron to the

hydrogen atom to create the hydride anion, H. Then, we could combine

two Hatoms together to create the fantastical dianion, H22. Loading up

the valence atomic orbital of H with electrons shows us that we need two

molecular orbitals for the neutral H2molecule, even though the molecule

itself only possesses two electrons. H2can accept two additional electrons,

and we need a place to put those! In general, remembering that we need a

place to put any electrons that may enter the molecule, we can conclude

the following:

The total number of atomic orbitals possessed by all the atoms of a

molecule equals the molecule's total number of molecular orbitals.

Figure 5 below shows the molecular orbitals of dihydrogen at the


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center of an orbital energy diagram. The shapes of the molecular orbitals

are determined by quantum mechanical principles that we won't concern

ourselves with here. What we should notice is the relationship between

orbital shape and energy--notice that the MO with lobes of two different

phases (and a node in between) is higher in energy than the orbital lackingany sign changes. Mathematically, the higher-energy orbital is the result of

subtracting the two 1sAOs, and the lower-energy orbital is the result of

adding the two AOs. When we add two 1sorbitals of the same phase, they

reinforce one another in the region between the nuclei. When we add

orbitals of opposite phase, however (i.e., subtract two AOs of the same

phase), they tend to cancel one another out between the nuclei, and at

some point, a node results. Figure 5 is probably the simplest example of

the linear combinations of atomic orbitals-molecular orbitals (LCAO-MO)

method. The dotted lines from the atomic to the molecular orbitals indicate

contributions of the AOs to each MO.

Figure 5.A molecular orbital energy diagram for H2, with isolated H

atoms (and atomic orbitals) on the periphery and the molecule (and its

MOs) in the center.

Because the higher-energy orbital lacks electron density between the


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nuclei, we call that orbital antibonding. Antibonding orbitals have energies

that are higher than those of the isolated atoms. Bondingorbitals, on the

other hand, have enhanced electron density between the nuclei and lower

energies than the separated atoms. Bonding MOs involve constructive

overlapof AOs, overlap of two lobes of the same phase. Antibondingorbitals result from destructive overlap--when lobes of opposite phase

coincide in space.

The simple example of dihydrogen from this section has introduced

us to the idea that molecular orbitals are built as weighted sums and

differences of atomic orbitals. In future discussions, we'll create molecular

orbital diagrams like Figure 5 using other types of atomic orbital building

blocks that sit on nearby atoms. Let's end this section with a quandary.

We've seen that molecular orbitals may be built from the 1s, 2s, and 2p

atomic orbitals. We've also seen that the 2porbitals are at right angles to

one another. The angles of the directional 2porbitals seem inconsistent

with the bond angles we saw in the section on molecular geometry. How

can we reconcile these two ideas, which seem to be at odds with one

another? We'll see how the idea of hybridizationsolves this problem in the

next section.

Watch Linear Combinations of Atomic Orbitals

***Atomic & Molecular Orbitals in Organic Molecules

Applying the LCAO-MO method of the last section to large organic

molecules produces complex, delocalized molecular orbitals. Such MOs tell

us little about how we should expect molecules to behave. Furthermore,

since delocalized MOs bear little resemblance to the lines and dots of

Lewis structures, it can be difficult to make connections between the two.

To get around this problem, we'd like an orbital theory that yields localized

MOs that look and "feel" like the elements of Lewis structures we already

know and love. In this section, we'll learn the six localized molecularorbitals that correspond directly to bonds and lone electron pairs in Lewis

structures. Our goal is to describe the shapes and energies of electron

sources and sinks in more detail, to paint a more vivid picture of electronic

structure in organic molecules.

At the end of the last section we confronted a quandary: is it possible
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to reconcile the geometry of the 2patomic orbitals (90 bond angles) with

the observed geometries of organic compounds (tetrahedral, trigonal

planar, and linear)? The answer is a resounding "yes," thanks to Linus

Pauling's concept of hybridization. To account for the observed

geometries of organic compounds, Pauling proposed that beforecombiningwith orbitals on other atoms, atomic orbitals can "hybridize" to produce a

new set of atomic orbitalsfor bonding. The process of "hybridization" is

essentially a kind of on-atom linear combination. We can take bits and

pieces from the different atomic orbitals to construct the hybrid atomic

orbitals. Don't worry about how exactly this is done--it's important just to

recognize the final result of hybridization, the three sets of hybrid AOs in

Figure 6.

Figure 6.The three sets of hybrid orbitals and their correspondinggeneralized building blocks.

The relative energies of the hybrid atomic orbitals are straightforward

to understand.4In the last section, we saw that the 2patomic orbitals are

higher in energy than the 2sorbital. Thus, we might expect hybrid orbitals

made of a greater percentage of 2pAOs to be higher in energy than those


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with less 2pcharacter. Experiments support this idea--the stability of

electrons held in hybrid AOs follows this trend:

(highest energy) 2p> sp3> sp2> sp> 2s(lowest energy)

Stated another way, hybrid AOs of greater scharacterare lower in

energy than those with less scharacter. Does this lower energy

automatically imply greater stability? No!The lower-energy hybrids are

stable when filled, but unstable without electrons. For this reason, cations

and other unsaturated species that possess empty hybrid orbitals are

usually unstable.

There is a profound link between the nature of an atom's building

block within a molecule and its hybridization. From Figure 6, notice that

hybridization depends on its geometry: tetrahedral atoms are sp3-

hybridized, trigonal planar atoms are sp2-hybridized, and linear atoms are

sp-hybridized.5Furthermore, the number of !bonds on an atom

corresponds to the number of AOs used to form its hybrids: 4 !bonds = 1 s

+ 3p= sp3, et cetera. Multiple bonds, as we will soon see, can be

understood as arising from interactions between unhybridized, leftover 2p

orbitals. For instance, sp2hybridization leaves behind one 2porbital for

multiple (") bonding. spHybridization leaves behind two 2porbitals for "

bonding. Examining the building blocks corresponding to thesehybridization states, we see one and two multiple bonds, respectively.

That's not a coincidence!The sp2and sphybrid orbital sets are shown in

Figure 7 with their leftover 2porbitals.


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Figure 7.Hybridization may leave behind unused 2porbitals, if less

than four !bonds are needed. Leftover 2porbitals may be used for "

(multiple) bonding or to hold nonbonding electrons.

The hybrid atomic orbital sets will serve as our starting point for

thinking about the molecular orbitals of large organic compounds. Localized

molecular orbitals are built either from the hybrid AOs themselves, in

isolation (we call such molecular orbitals non-bonding), or from bonding

and antibonding combinations of the hybrids. At this stage, it's important to

realize that we can treat each !(single) bond in an organic compound like

we did the hydrogen atom, using hybrid orbitals in place of 1sorbitals. For

each !bond, there is a !bonding orbital, which holds the electrons of thebond, and a !* antibonding orbital, which reflects the ability of the bond to

break upon the addition of two more electrons. The bonding MO involves

constructive overlap of the hybrids, and the antibonding MO destructive

overlap. Figure 8 shows a simple orbital energy diagram for the carbon-

carbon !bond in ethane (C2H6).


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Figure 8.A simple orbital energy diagram for the CC bond of

ethane. The bond is associated with two molecular orbitals: a bonding !

orbital of low energy, and an antibonding !* orbital of high energy.

The !molecular orbital shows us where the electrons of the bond are

likely to be found in space. The energy of this orbital reflects the reactivityof the electrons in the bond as an electron source--a clean, intuitive idea.

But how should we interpret the empty antibonding orbital? What does an

antibonding orbital "mean" from the molecule's perspective? Put most

concretely: how does the nature of the !* orbital affect the bond's behavior

(structure and reactivity)? Spatially, the !* orbital shows us where incoming

electrons from a source are likely to go. Incoming electrons will tend to

approach the large lobes of the antibonding orbital. Like unfilled hybrid

AOs, antibonding MOs are most stable when high in energy. Summing up,

the nature of the !* orbital reflects the bond's potential as an electron sink!Since reactions require both an electron source and sink, we must keep the

importance of antibonding orbitals in mind when studying organic

reactivity--the sink is an essential piece of the puzzle.

!Bonding is advantageous for molecules because, as we can see

from Figure 8, the overall energy of electrons is lowered in the process. But


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!bonding also seems to present certain problems--for instance, how can

hybrids overlap to form multiple bonds? Using more than one hybrid orbital

to describe a multiple bond seems out of the question. Still, the 2porbitals

left behind on sp2- and sp-hybridized atoms don't look appropriately

positioned for orbital overlap. What gives? Evidently, the parallelarrangement of 2porbitals is good enough to establish a bond between

adjacent atoms. Figure 9 shows the idea for the simplest hydrocarbon that

contains a double bond, ethylene.

Figure 9.Multiple bonds are the result of "side on," parallel, or "-type

overlap of adjacent 2porbitals. Just as in the !bonding case, a bonding "

orbital and an antibonding "* orbital result from constructive and

destructive overlap.

Our spatial and energetic interpretations of the "and "* orbitals are

identical to those of the !and !* orbitals. To reiterate, the "orbital reflects

the potential of the double bond to serve as an electron source, and the "*

orbital reflects the potential of the double bond to serve as an electron sink.

Importantly however, the side-on overlap characteristic of "orbitals is

weaker than the head-on overlap involved in !bonding. As a result, "and


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"* MOs tend to be closer in energy to their atomic orbital building blocks

than !orbitals. Consequently, two important trends occur:

"Bonding MOs are higher in energy than !MOs.

"* Antibonding MOs are lower in energy than !* MOs.

Taken together, what do these trends suggest about the relative

reactivity of single and multiple bonds? In general, multiple bonds are more

reactive (both as sources andsinks) than single bonds. Electrons in "

bonds are higher in energy than electrons in !bonds; unfilled "* orbitals

are lower in energy than !* orbitals. In addition, we can see from Figures 8

and 9 that the spatial positions of "and !electrons differ. !Electron

density can be found along an axis connecting the atoms, while "electron

density is found above and below such an axis. In fact, "MOs possess a

node coinciding with the axis connecting their atoms.

Finally, we need to address nonbonding electrons and orbitals, which

are not involved in interactions with orbitals on adjacent atoms. In the

absence of resonance interactions, lone electron pairs can be found in

isolate hybrid AOs. We will refer to these as nmolecular orbitals (for

"nonbonding"). Figure 10 depicts three examples of n orbitals in common

compounds. Notice that each n orbital is characterized by a particular

hybridization, since the nMO is really just a hybrid atomic orbital. The norbital reflects the potential of the atom to serve as an electron source viaa

lone pair sitting on the atom.

Figure 10.Nonbonding n orbitals in common organic molecules.

Most commonly, n orbitals are just hybrid atomic orbitals.

Building blocks bearing fewer than 8 total electrons must possess an

empty nonbonding MO, which we call an aorbital (for "atomic"). What are


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the atomic orbital constituents of aMOs? Empty hybrid AOs present an

energetic problem: they're low in energy and thus tend to be unstable when

lacking electrons. To get around this problem, molecules without geometric

constraints adopt geometries that allow them to leave high-energy 2p

orbitals unfilled. Thus, in the vast majority of cases, aorbitals are justempty atomic 2porbitals. When you spot a building block with 6 or fewer

total electrons, take note that an empty aorbital is present on the atom!

This MO reflects the potential of the atom itself to serve as an electron sink.

Empty aorbitals are typically the lowest energy (and most unstable) unfilled

orbitals one finds in organic compounds. This fact is unsurprising when we

consider that building blocks bearing aorbitals lack an octet of electrons.

Figure 11 provides two examples of empty aorbitals: the typical a= 2p

case, and a case in which the empty orbital must be a hybrid AO (based on

the geometry of the building block, which demands sp2hybridization).

Figure 11.Empty atomic, aorbitals in organic molecules. aMOs are

most commonly just 2patomic orbitals; however, geometric constraints

may force a hybrid AO to be empty. Notice that both cationic building blocks

have 6 total electrons.

With the aorbital, we've reached the sixth and last of the localized

MO classes. Figure 12 provides a summary of the orbital shapes of the six

classes and the structural elements to which they correspond. These

shapes both confirm our intuition and suggest some intriguing new ideas.On the confirmatory side, notice that electron sources tend to be

concentrated between nuclei, where the Lewis model suggests we should

find electrons. More interestingly, the !* electron sink is primarily located

outside of the space between the nuclei. Incoming electrons will most likely

approach the outskirts of a bond, not between the atoms. Trigonal planar

atoms lacking an octet of electrons (thus bearing an aMO) will be


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approached by electrons perpendicular to the molecular plane. Consider

the "MOs--notice that they depend on a parallel alignment of 2porbitals.

Rotation away from the parallel alignment ruins "-type overlap, so double

bonds must remain planar. All kinds of interesting spatial ideas come to

light!

Figure 12.Shape and occupancy of the six classes of localized

molecular orbitals. The filled orbitals reflect the potential of the structure as

an electron source; the empty orbitals reflect the structure's potential as an

electron sink.

The relative energetics of these six classes are also extremely

important to keep in mind. These form a kind of stability trend that allows us

to quickly pinpoint the most reactive sources and sinks within a molecule.

We touched on this trend in Figure 13, but it's worth bringing up again here,

now that we've seen the shapes of the localized MOs. Figure 14 depicts

the most common relative energies of the localized MOs. Based on Figure

14, we should expect n orbitals to be the most reactive sources and a

orbitals to be the most reactive sinks. By identifying the localized MOs

present in a molecule and using Figure 14 to predict their relative energies,

we can predict how molecules will behave in a very powerful, general way.


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Figure 14.Typical relative energies of the localized molecular

orbitals. To drive the point home once more, low-energy unfilled orbitals

and high-energy filled orbitals are reactive.

Watch Localized Molecular Orbital Theory***

Effects of Electronegativity & Charge

Figures 12 and 14 lay out the shapes and relative energies of the six

classes of localized MOs. But a simple problem reveals that those figures

don't tell us the whole story. Consider Figure 15 below, which depicts a

good nucleophile (thiolate anion) in the presence of a compound containing

C=O and C=N bonds. Based on ideas from the last section, the "* orbitals

of the C=O and C=N bonds should be the electrophile's best electron sinks.

Can we predict which of the two sinks is better?
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Figure 15.Two possible courses of action: thiolate can add either to

the C=O or C=N bonds. Can we used localized MO theory to predict which

of the two "* orbitals is more reactive?

We can, if we first recognize that the identity of the atoms that make

up the two double bonds are different. Where we see a nitrogen in one

double bond, an oxygen sits in the other. We need to understand how atom

type affects the shapes and energies of molecular orbitals. Even more

specifically, we need to understand the relationship between

electronegativityand the properties of orbitals. How does electronegativity

affect the energies of the atomic orbitals? How are orbital shapes

influenced by electronegativity?

As we move from left to right across the periodic table, orbital

energies decrease. Rather intuitively, more electronegative atoms are

associated with lower energy orbitals. Another way of saying this is that

more electronegative atoms are more electrophilic, or that they tend to be

associated with more electrophilic orbitals. Put yet another way, more

electronegative atoms make worsenucleophiles (electron sources).

Whereas carbanions are extremely nucleophilic (electron-donating)

molecules, fluoride anion is hardly nucleophilic at all. Applying this idea to

the problem in Figure 15, we can predict that the "* orbital associated withthe C=O bond ("*CO) is lower in energy than the "* orbital associated with

the C=N bond ("*CO). Since this orbital is unfilled, the C=O bond ought to

be more reactive than the C=N bond.6

Molecular orbitals may involve identical atom types that differ only in

their charge. Thus, to finally complete our understanding of the effects of


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molecular structure on orbital energy and shape, we need to learn how

charge influences orbitals. Think of charge as a specialized case of

electronegativity, making use of the classic maxim that "opposite charges

attract and like charges repel." Positive charge on an atom boosts its

electronegativity over its neutral counterpart--positively charged atoms holdelectrons more tightly than neutral atoms (opposites attract). Negative

charge has the opposite effect, and lowers the electronegativity of the atom

relative to its neutral counterpart (like charges repel). Consider the example

in Figure 16.

Figure 16.Which of the two C=N double bonds ought to be more

reactive under these conditions?

Since positive charge increases electronegativity, we can think about

the positively charged nitrogen atom in the same way we thought about

oxygen in the first example from this section. It's more electronegative than

neutral nitrogen, so the energy of its "* orbital is lower. Consequently, the

C=N+double bond ought to be more reactive than the C=N double bond.

Organic chemists confirm predictions like these on a daily basis with

experiments!

Electronegativity's effects on orbital shape are intuitive to understandas well. Polarized MOs are built using atomic orbitals from atoms of very

different electronegativity, and consist of lobes of different sizes on each

atom. As we might expect, filled orbitals tend to have large lobes on

electronegative atoms, since electrons tend to spend their time around

atoms that are hungry for them. Unfilled orbitals, on the other hand, tend to

have larger lobes on less electronegative atoms. We will revisit this idea


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when we discuss frontier molecular orbital theory and the fundamentals of

reactivity.

For now, it's most important for us to understand how

electronegativity affects orbital energies. Our intuitive ideas about

electronegativity map nicely onto orbital energy ideas: electronegativeatoms possess low orbital energies and thus stabilize electrons. Treating

charge as a kind of electronegativity booster or restrictor, we can treat

charged atoms using this same idea. Armed with this principle, we can

easily compare the electronic viability of two possible reaction pathways,

even if the pathways in question differ only in the charge or element type of

the atoms involved.

1. Orbitals are interchangeably called "wavefunctions." We'll use the term

"orbital" whenever possible, but you should be aware that these two termsare often used interchangeably.

2. Because we think of orbital shapes as containers for electrons, we

commonly say that an electron is "in" an orbital.

3. For example, in Figure 3.4 the valence AOs are the 2slevel and all three

2plevels.

4. See the subsection addressing hybridization in the "Generalized Stability

Trends" section.

5. This relationship between geometry and hybridization always holds true.

However, it's important to remember that "geometry dictates hybridization,

not the other way around." (Scott Denmark) Hybridization follows from

geometry, and not necessarily from the apparent bonding network of the

atom.

Benzyne is a classic example of this idea: the triple bond of benzyne

suggests sphybridization for the two atoms involved; however, geometry

demands that they be at least approximately sp2-hybridized. The triple

bond of benzyne more closely resembles an sp2-hybridized diradical than a"classical" triple bond.

6. It's important to recognize this thought process as highly context

dependent. Say we treated the organic compound in Figure 15 with an acid

(HA) instead. Acids are electrophilic species, so we're interested in the

organic compound as a nucleophile. Comparing the norbitals on oxygen


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and nitrogen, we can conclude that the norbital on less electronegative

nitrogen is higherin energy than the norbital associated with oxygen.

Thus, nitrogen should be the better nucleophile (or base) in this context.

For every reaction you encounter, ask yourself which of the startingmaterials are nucleophiles (electron sources) and electrophiles (electron

sinks). As these examples show, mixing up these concepts can have

damaging effects on predictive ability.

week 1 - intro to orbitals

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