Organic Chemistry

Meaning of “Organic”?

Initially scientists believed there was a special “force” in living organisms -this was assumed the unique component of organic material

In 1828 Wöhler synthesized urea (a known component of organic material) from inorganic ammonium cyanate

NH4 OCN !

H2N

O

NH2urea

Later Justus von Liebig, a noted 19th century chemist, declared the synthesis of urea “the very first beginning of the actual scientific organic chemistry”

In a subsequent paper Wöhler and Liebig wrote: “sugar, salicin, and morphium will be produced artificially. Of course, we do not

know the way yet by which the end result may be reached since the prerequisite links are unknown to us from which these materials will be develop -

however, we will get to know them.”

HO O OH

HN

glucose salicin morphine

Organic Chemistry

OHOHO

OHOH

OH

OHOHO

OOH

OH

HO

Organic Chemistry

1828

1835

1862

“Organic chemistry just now is enough to drive one mad. It gives the impression of a primeval, tropical forest full of the most remarkable things, a monstrous and

boundless thicket, with no way of escape, into which one may well dread to enter.”

“This investigation has yielded an unanticipated result that reaction of cyanic acid with ammonia gives urea, a noteworthy result in as much as it provides an example of

the artificial production of an organic, indeed a so-called animal, substance from inorganic substances.

“I am afraid I shall have to give up my trade: I am far too inert to keep up with organic chemistry, it is becoming too much for me, although I may boast of having contributed

something to its development. The modern system of formulae is to me quite repulsive.”

Consider the effect organic chemistry had on even a pioneer as Wöhler:

Initial Excitement!

Realization of the complexity

Acceptance of Failure?

Hopefully the class will allow the initial excitement without a subsequent acceptance of failure (at least not for 34 years!)

Organic chemistry is now considered the chemistry of carbon

**Not only carbon! There are few forms of only pure carbon

(no other elements present)

Some examples of compounds with only carbon present:

graphite diamond

C60 spherical ball carbon nanotube

Organic Chemistry

Carbon can form strong bonds with not only carbon but with other elements (e.g. hydrogen, nitrogen, oxygen, halogens)

Depending upon the order of the bond connections (both constitution and configuration) millions of compounds, all with potentially different properties, can be prepared

Why is this so important?

Almost every component of living organisms involve organic chemistry (proteins, enzymes, lipids, fats, carbohydrates, nucleic acids, etc.)

And the metabolic and other interactions in the body involve organic reactions

Organic Chemistry

In Organic Chemistry we study how compounds react

During a reaction old bonds are broken and new bonds form

Bonds form – two atoms share electrons Bonds break – two atoms no longer share electrons

Therefore, if we know more about the electrons we can understand organic chemistry

What do we want to know?

Where are the electrons?

How tightly are the electrons held in a “bond”?

We already know many of these basic concepts from general chemistry

Organic Chemistry

Background of an atom

An atom consists of three types of particles: Proton (positively charged)

Neutron (neutral) Electron (negatively charged)

The number of protons determines the element (also is the atomic number) Carbon therefore has 6 protons located in the nucleus

Usually the nucleus also contains an identical number of neutrons as protons If the number is different it is called an isotope

These two particles have similar mass (~1830 times greater than an electron)

Ironically the word “atom” comes from the Greek language meaning “indivisible”, as it was thought to be the smallest particle that did not have smaller constituents

Consider a Carbon Atom

Nucleus – means “kernel of a nut”

Nucleus size is ~2 fm (1 fm = 10-15 m), atom size is ~1 Å (1 Å = 10-10 m)

For an uncharged 12C atom, there are 6 protons, 6 neutrons and 6 electrons

Therefore the nucleus, which is responsible for ~3600/1 parts of the mass, only encompasses ~1 x 10-15 part of the volume (remember V = 4/3 πr3)

Shell of atom (region where electrons reside)

Nucleus of atom (region where protons and neutrons reside)

Nucleus of atom 6 protons (red), 6 neutrons (blue) -held together by nuclear forces

1 Å

1 fm

Electrons

Unlike the protons and neutrons which are in the nucleus (a relatively fixed point) we cannot say with certainty where an electron is located at a certain time

(Heisenberg uncertainty principle)

What we can say is that ‘on time average’ the electrons are located in orbitals (regions of space)

*much bigger region than the nucleus

The electrons are not randomly placed, nor do they reside in simple circular orbits

Schrödinger developed a formula that could describe the properties of the electrons (called Ψ)

This wave function, Ψ, mathematically describes the shape of an orbital where the electrons resides and the square of the wave function, Ψ2, is proportional to the probability of finding

an electron in a given volume.

Electrons

Each electron is described by a set of four quantum numbers (Pauli principle: no two electrons may have the same values of all four quantum numbers)

n

principal quantum number Integral values, n = 1, 2, 3, etc.

Indicative of “shell” the electron resides

Atoms where n=1 (H and He) are first shell

Atoms where n=2 (Li, Be, B, C, N, O, F, Ne) are second shell, and so on

The higher the value of n, the greater the average distance of the electron from

nucleus and thus the greater the electron’s energy

l

Related to shape of orbital Integral values, l = 0, 1, 2, …(n-1)

Therefore if n=1, l must be 0 If n=2, l can be either 0 or 1

If n=3, l can be 0, 1 or 2

Each value of l represents a different orbital shape l=0 is s orbital l=1 is p orbital l=2 is d orbital l=3 is f orbital

(naming or orbitals derives from description from early spectroscopists of metal alkali lines as sharp, principal, diffuse and fundamental)

Electrons

Each electron is described by a set of four quantum numbers (Pauli principle: no two electrons may have the same values of all four quantum numbers)

m1 s

Related to orientation of the orbital in space (x, y, z coordinates in three dimensions)

Has integral values –l,…., 0, …., +l Thus if l=0 then m1 must be 0

If l=1, then m1 can be -1, 0, or +1 If l=2, then m1 can be -2, -1, 0, +1, +2

Orbitals of the same shell (n) and shape (l) have the same energy regardless of m1

(thus all 3 p orbitals in the same shell have the same energy, 2px=2py=2pz)

Spin quantum number

Can only have two values (-1/2 or +1/2)

By convention, the two possible spin states are typically indicated by either having the

spins pointing up or down

With two electrons, if the spins are opposite they are stated to have paired spins, if the

spins are pointing in the same direction then they are stated to have unpaired spins

paired unpaired

Electronic Configuration

If the four quantum numbers are known for each electron, then the electrons can be placed into their respective orbitals by knowing the relative energy difference

Since the first row atoms (H and He) have only one orbital (1s) it is fairly easy to write the electronic configuration for these atoms (either have one electron or two paired electrons)

Consider the second row atoms The second row has a total of 5 orbitals (1s, 2s, 2px, 2py, 2pz)

1s

2s

2p

Energy The first shell orbitals are lower in

energy than second shell and p orbitals in the 2nd shell are higher in energy than s orbitals in the second shell

With Carbon (1s22s22p2), however, the question arises as to where the second 2p electron goes Hund’s rule: for a given electronic configuration, the state with the greatest number of

unpaired spins has the lowest energy Therefore carbon is 1s22s22px

12py1 where the spins are parallel

Shape of Orbitals

Solutions of the Schrödinger equation (EΨ=ĤΨ) will describe the wave equation (Ψ), since Ψ2=probability of finding electron it also describes the shape of orbitals

!2

r (distance from nucleus)

4!r2"2

r (distance from nucleus)

radial density plot radial probability plot

never goes to zero

1s orbital

three-dimensional shape

Shape of Orbitals

As the number of shells increases, the presence of nodes increases (nodes are regions of space where there is zero probability of finding an electron)

The number of nodes is equal to one less than the principal quantum number (therefore for n=1 there are 0 nodes, for n=2 there is 1 node, for n=3 there are 2 nodes)

All second shell orbitals thus have 1 node

2s 2px 2py 2pz

P orbitals have node at the nucleus Actual shape is

different than typically stylized “dumbbell”

shaped p orbitals

Bonding

Octet rule – atom is most stable if its outer shell of electrons is filled

Therefore an atom will give up, accept, or share electrons in order to achieve a filled outer shell

Often the inner shell of electrons is ignored (not counted for the octet rule)

Outer shell of electrons is also referred to as ‘valence’ electrons

Called ‘octet rule’ since second row atoms (which comprise organic compounds) have 8 electrons in their outer shell

Type of Bonds

If atoms give up or accept electrons an ion is formed

Consider lithium (3 total electrons)

There is only one valence electron

Li

1s22s1

1s

2s

2p

Lithium will therefore easily lose an electron to create a lithium ion (which has a filled outer shell)

The energy required to remove an electron is referred to as the ionization potential

Since lithium readily loses an electron to leave an atom with a complete outer shell it is called electropositive

Type of Bonds

Li

1s22s1

1s

2s

2p

1s

2s

2p

Li

1s2

Lose an electron

Halogens, on the other hand, readily gain an electron to complete the outer shell (energy released by adding an electron to an atom is called the electron affinity)

Consider fluorine (9 electrons, 7 valence electrons)

The outer shell needs one electron to be filled

Type of Bonds

1s

2s

2p

Gain an electron

1s

2s

2p

F

1s22s22p5

F

1s22s22p6

In a covalent bond, however, electrons are shared between two atoms (not lose or accept as in the formation of ions)

The sharing of electrons can allow both atoms to fill the outer shell

Consider F2 (fluorine gas)

Each fluorine atom needs one electron to fill the outer shell

One atom cannot donate an electron and have both atoms with a filled outer shell

Type of Bonds

F F F F

This Fluorine atom has a filled octet

This Fluorine atom is lacking 2 electrons in

outer shell!

If donated an electron

An alternative is to share two electrons (one from each atom) between both atoms

Represent each electron with a dot (called Lewis dot structures)

(only show valence electrons for a Lewis dot structure)

Both atoms now have 8 electrons in the outer shell (therefore octet rule is obeyed)

F

Fluorine has 9 total electrons, but only 7 valence electrons

F F F

Shared electrons

Type of Bonds

Differences between Ionic and Covalent Bonding

Lithium Fluoride Forms ionic bond by lithium ‘donating’ an electron to fluorine

Each outer shell is filled, but no electron density between two atoms

Fluorine gas Forms covalent bond by sharing electrons

Li F FLi

F F FF

Different Ways to Draw Organic Compounds

Organic chemistry has a shorthand for drawing compounds (need a way to indicate what atoms are connected by these covalent bonds)

Lewis Structures

Other representations

All structures shown represent the same compound (propane)

C C C

H

H

H

H

H

H

H

H

H CH

HCH

HCH

HH

CH3CH2CH3

Polar Bonds

All covalent bonds shown so far have been between identical atoms

What happens if a covalent bond is formed between two different elements?

*Both atoms do not need to share the electrons equally

Even though the electrons are shared, the electrons can be closer, on time average, to one nucleus than the other

How to predict where the electrons are located?

C C F F

Electronegativity Tables

Linus Pauling first established values to associate with each element (there have been many different values computed, but the trend is the same)

Elements toward the upper right hand of the periodic table are more electronegative

Also can predict the relative electronegativity of two atoms by their relative placement in the periodic table

H (2.3) Li (0.9) Be (1.6) B (2.1) C (2.5) N (3.1) O (3.6) F (4.2)

Cl (2.9)

Br (2.7)

I (2.4)

The numbers are a relative indication of how much the electrons are ‘attracted’ to a certain atom

As the number becomes larger, the electrons are attracted more by that atom

H3C OH2.5 3.6

Formal Charges

A formal charge represents a full charge on the atom (assuming no polarity of the covalent bond)

To calculate:

Formal charge = (group number) – (nonbonded electrons) – ½(shared electrons)

Use group numbers, not atomic numbers!

Consider ammonium:

Formal Charge = 5 – 0 – 1/2 (8) = +1

HN HH

H

Group number = 5 Nonbonded electrons = 0

Shared electrons = 8

Building Molecular Orbitals from Atomic Orbitals

Similar to a wave function that can describe the regions of space where electrons reside on time average for an atom, when two (or more) atoms react to form new bonds, the region

where the electrons reside in the new molecule are described by a new wave function

This new wave function describes molecular orbitals instead of atomic orbitals

Mathematically, these new molecular orbitals are simply a combination of the atomic wave functions

Hydrogen 1s atomic orbital

H-H bonding molecular orbital

Building Molecular Orbitals from Atomic Orbitals

An important consideration, however, is that the number of wave functions (molecular orbitals) resulting from the mixing process must equal the number of

wave functions (atomic orbitals) used in the mixing

In the case of H2, in addition to the new bonding molecular orbital obtained by adding the two atomic 1s orbitals, an antibonding orbital is obtained by subtracting the two atomic orbitals

H-H antibonding molecular orbital

node

Electronic Configuration for H2

Each Hydrogen 1s atomic orbital has one electron When two atomic orbitals mix, they produce two molecular orbitals

As the number of nodes increases, the energy of the orbital increases

The molecule has a total of two electrons and follow Aufbau principle and Pauli principle to fill electrons in molecule

Bond Strength

The bond strength for H2 is considered the amount of energy required to break the bond and produce two hydrogen atoms

Called the bond dissociation energy (BDE)

Homolytic bond cleavage

X Y X Y

Heterolytic bond cleavage

X Y X Y

Building Molecular Orbitals from Atomic Orbitals

When forming molecular orbitals of H2 from atomic orbitals from each hydrogen atom, hydrogen only has one electron in a 1s orbital (or a 1,0,0 orbital using n,l,m designation)

When building molecular orbitals from a 2nd row atom (like carbon) can use either the 2s (2,0,0) or 2p (2,1,-1; 2,1,0; or 2,1,1) atomic orbitals to form the bonds

When different orbitals interact, the overlap of the orbitals changes depending upon the direction of bond formation (both in degree of overlap and symmetry of the bond)

When two s orbitals interact, due to symmetry of orbital the direction of approach

is irrelevant

Overlap between orbitals of same phase leads to bonding region

Bonding MO (called σ bond)

When two p orbitals interact, if lobe with same phase is pointing toward each other a

bonding region can occur

Building Molecular Orbitals from Atomic Orbitals

Due to the unsymmetrical orientation of a p orbital, however, there are other possible orientations of approach

One bonding approach direction has the p orbitals on both atoms directed opposite to the approach direction

Still a bonding MO, but electron density is not symmetric about

internuclear axis (called π bond)

If s orbital approaches p orbital from side, however, there is no net overlap

The positive overlap (blue with blue) is exactly canceled with the negative overlap (blue with red),

thus there is no net overlap

The orbitals are said to be “orthogonal” to each other and thus do not mix

Building New Molecular Orbitals from Molecular Orbitals

In addition to building new molecular orbitals from adding atomic orbitals, new molecular orbitals can result from combining orbitals from two different molecules

using their molecular orbitals (the result of a reaction between two molecules)

For a given molecule there might be a multitude of molecular orbitals (the total number are due to the number of atoms in the molecule)

Hypothetical molecule that contains 6 molecular orbitals and 6 electrons

Would fill the orbitals by following Pauli exlusion (only 2 electrons per

orbital) and filling the lowest energy orbitals first

The orbitals are classified by whether they are “filled” or

“unfilled”

Occupied molecular orbitals (OMOs)

Unoccupied molecular orbitals (UMOs)

In addition the molecular orbital that is occupied that is highest in

energy is called the HOMO

The molecular orbital the is unoccupied that is lowest in energy

is called the LUMO

Building New Molecular Orbitals from Molecular Orbitals

When two molecules react (each with their own set of molecular orbitals) it is important to recognize which molecular orbital interaction determines the reaction

(if a chemist knows this then they can predict reactions)

Compound A Compound B

The number of OMOs and UMOs and energy placement of orbitals is dependent upon the

compound

Each molecular orbital in compound A however will react with each molecular

orbital in compound B

Whenever any two nonorthogonal orbitals interact they will create two new MOs, one higher in energy and one lower in energy

When two UMOs react, there is no change in energy as there are no electrons in the orbitals

When two OMOs react, there is an increase in energy due to the two higher energy electrons

outweighing any energy gain

There is only an energy gain when a OMO of one molecule interacts with an UMO on the

other molecule

Building New Molecular Orbitals from Molecular Orbitals

The amount of energy gain is also dependent upon how close in energy the two orbitals are before mixing

Consider mixing of two orbitals, one filled (OMO) and one unfilled (UMO)

If the OMO is identical in energy to the UMO there will be the maximum energy gain due to the best possible mixing of the orbitals

ΔE

As the OMO has a greater difference in energy to the UMO, the mixing will be less

and the energy gain will thus be lower

ΔE

Thus the energy gain is greatest in a reaction when the HOMO of one compound is closest in energy to the LUMO of the second compound

Building New Molecular Orbitals from Molecular Orbitals

When mixing any two orbitals therefore the two important considerations are the overlap between the two orbitals and the match in energy of the two orbitals before mixing

Considerations between mixing of orbitals are therefore: -when two nonorthogonal orbitals overlap and mix, they generate two new orbitals

(one higher in energy and one lower in energy) -the amount of energy shift upon mixing is greater with more overlap of the orbitals and

lower the further apart in energy the orbitals are before mixing -average energy of two new orbitals is slightly higher than average of original orbitals

(partly an artifact of electron-electron repulsion in higher energy orbital)

Compound A Compound B

Consider the original hypothetical compound A reacting with compound B

The most important interaction to consider is the HOMO of A reacting with the LUMO of B

(largest energy gain)

The energy gain from this interaction must be large enough to overcome the energy loss of each OMO mixing

with another OMO (which causes an energy loss)

Acid/Base Reactions

Important considerations for Organic chemistry:

What makes a compound acidic?

How do we determine acid strengths?

Need to first consider what definition of acid/base reactions we are using

Lavoisier was the first chemist who tried to characterize acids and bases

Lavoisier believed that oxygen was responsible for acidity -An acid is thus the oxidized form while the base is the “base” atomic state

Sulfuric acid

Phosphoric acid

Lavoisier in fact coined the term “oxygen” from the Greek words oxys = sour and genes = born, thus oxygen literally means “acid maker”

oxygenS H2SO4

P H3PO4oxygen

The Arrhenius definition is used to determine the pH of a solution

Kw = [H3O+][HO-] = 1 x 10-14

In neutral solution the concentration of [H3O+] and [HO-] are equal Therefore [H3O+] = [HO-] = 1 x 10-7 M

Acidic solutions have an excess of [H3O+] Therefore [H3O+] > 1 x 10-7 and [HO-] < 1 x 10-7

Acids – something that dissociates in water to give hydronium ion (H3O+) Bases – something that dissociates in water to give hydroxide ion (HO-)

Arrhenius Acid/Base Definition

Due to the magnitude, the acid strength is expressed in a logarithmic scale pH = -log10 [H3O+]

The Arrhenius definition is poor for organic compounds Since very few will dissociate into hydroxide ions

H Cl H Cl

Na OH Na OH

Brønsted-Lowry definition is better for organic compounds

Acid – any species that can donate a proton

Base – any species that can accept a proton

As can be seen from the reaction above, every Arrhenius acid and base are still considered acids and bases with the Brønsted-Lowry definition

Other compounds, however, would not be considered bases under Arrhenius which do qualify with Brønsted-Lowry

Brønsted-Lowry Acid/Base Definition

H ClNa OH H2O NaClacid base

NH3 H Cl Cl NH4base acid

Conjugate Acid and Base

An important concept with the Brønsted-Lowry acid/base definition is the resultant conjugate acid and base

(every acid becomes a conjugate base after the reaction)

H Cl H2O H3O Cl

base

acid

conjugate acid

conjugate base

Lewis Definition of Acids and Bases

G.N. Lewis postulated that an acid/base reaction need not involve a transfer of a proton

An acid/base reaction can refer to any reaction that involves the formation of new bonds

Lewis acid: a species that accepts a lone pair of electrons to form a new bond

Lewis base: a species that donates a lone pair of electrons to form a new bond

H ClH2O H3O Cl

base acid

This definition is far more general, but any acid or base in Brønsted/Lowry definition remains the same in the Lewis definition

Introduces new terms that are used in many organic reactions A Lewis base is called a nucleophile - “lover of nuclei”

A Lewis acid is called an electrophile - “lover of electrons”

General HOMO-LUMO Definition of Acids and Bases

When Lewis introduced his concept of nucleophiles reacting with electrophiles to describe acid/base reaction, the concept of molecular orbital theory was just beginning

Today Lewis’ definition is even more generalized by considering a filled molecular orbital (HOMO is the highest occupied molecular orbital) reacting with an empty molecular orbital

(LUMO is the lowest unoccupied molecular orbital)

HOMO

LUMO

reaction

Energy gain

product

The closer in energy the HOMO is to the LUMO, the greater the interaction (and more energy gain)

The compound reacting through the HOMO is the nucleophile, the compound reacting through the LUMO is the electrophile

nucleophile

electrophile

Evolution of Acid-Base Theories

Lavoisier (1789)

Arrhenius (1887)

Brønsted/Lowry (1923)

Lewis (1923)

HOMO-LUMO (1960’s)

Theory Acid Base

oxidized substance

substance to be oxidized

H+ source HO- source

H+ donor H+ acceptor

e-pair acceptor “electrophile”

e-pair donor “nucleophile”

unusually low LUMO

unusually high HOMO

increasing generality

Acid-Base Reactions

HO- H+ HO-H

Curved arrows designate e-pair shifts (not atomic motion)

Start arrow at e-pair location in starting material End arrow at e-pair location in product

Acid-Base Reactions

HO- H+ HO-H

NH3: H+ NH3-H

HO- BH3 HO-BH3

NH3: BH3 NH3-BH3

Arrhenius

Brønsted/Lowry

Lewis

Lewis

All of these acid-base reactions follow a similar arrow pushing mechanism

+  

-­‐  

+              -­‐  

Acid Strength

Organic acids are defined by the acid dissociation constant (Ka)

Similar to pH measurements, this quantity is expressed in the logarithmic form

pKa = -log Ka

The stronger the acid, the smaller the pKa

Don’t confuse pKa with pH

pKa is a constant for a given acid referring to the pH where half of the acid is ionized pH refers to the concentration of hydronium ions in solution

HA H3O AH2O

Ka

Ka =H3O A

HA

Resonance in Organic Compounds

What is resonance? (also called ‘delocalization’)

Look at a nitro group

The negative charge on the oxygen could be placed on either oxygen using Lewis structures

Which structure is correct?

It turns out neither structure is correct, but the charge is delocalized onto both oxygens -This process of being able to delocalize the charge onto more than one atom is called resonance

RNO2 R NO

O

R NO

OR N

O

O

(Resonance is a result of not being able to draw an accurate structure using one Lewis structure)

“Rules” of Resonance

1.  All resonance structures must be valid Lewis structures (e.g. cannot have 10 electrons on one carbon in one structure)

2.  Only electrons move (cannot move nuclei, only electrons

– usually double bonds or lone pairs connected through an extended p orbital system)

3. Number of unpaired electrons must be constant

How does resonance explain acidity?

Consider pKa of organic molecules

Both structures place a negative charge on oxygen after loss of proton, but the pKa difference is greater than 11

Resonance in Organic Compounds

CH3CH2OH CH3CH2O H

pKa

16

H3CO

OHH3C

O

OH 4.8

Resonates similar to nitro group

A carbonyl group is a common resonance source

Neither structure is correct, but rather the negative charge can be delocalized over both oxygen atoms

H3C

O

O H

H3C

O

OH3C

O

O-H

Once the acid is deprotonated, the negative charge is located on one oxygen

The charge can be delocalized (resonated) onto the other oxygen

Resonance structures are simply a result of one Lewis structure being incomplete in describing the location of electrons

A double headed arrow always means two (or more) resonance forms

Two arrows mean two chemically distinct structures

Resonance in Organic Compounds

Resonance of Acetate Anion

H3C

O

OH3C

O

O

H3C

O

O

Actual structure has charge equally on both oxygen atoms

Rotation of Acetate Group

H3CO

OLook down C-C bond H3C

O

O

H3CO

O

Neither structure is correct

Actual structure has charge “delocalized” over both oxygen

atoms equally

Comparison of Electron Density for Ethoxide versus Acetate anion

The excess negative charge is more stable on the acetate anion that can resonate, thus the conjugate acid is more acidic

Resonance in Organic Compounds

CH3CH2OH CH3CH2O H3CO

OHH3C

O

OH3C

O

O

Important to Remember: Not all resonance structures need to contribute equally

If two resonance structures are not of equal energy, then they will not contribute equally to the actual structure

This leads to major and minor contributors

Resonance in Organic Compounds

NH

H

H

HN

H

H

H

H

Octet rule obeyed on all atoms

Positive charge on less electronegative

atom

Actual “hybrid” structure

Resonance Forms

NH

H

H

HN

H

H

H

H

Deeper blue color on carbon (less charge on carbon)

Resonance Forms

H3C CH3

O

H3C CH3

O

H3C CH3

O

H3C CH3

O

major minor inconsequential Only a resonance form if spins are

paired

Not all resonance forms need to contribute equally, but rather a weighting factor is given to each resonance form depending upon its importance to the actual form

Typically forms with more bonds are more important (often due to forms with less bonds typically have atoms without an octet)

Placing charge on the more electronegative atoms is important

The number of paired electrons must be constant

Factors affecting stability of resonance structures:

- Placement of charge

When the only difference is the location of formal charge, structure is more stable when anion is placed on more electronegative atom

Resonance Forms

H3C

O

CH2 H3C

O

CH2

major minor

- Amount of Charge

Also related to number of bonds in a structure

While structure with four formal charges shown is a “valid” resonance form, if structure is dramatically higher in energy then it is practically an irrelevant resonance form

Resonance Forms

inconsequential

- Octet rule is important

Having all atoms with a filled octet rule is more stable than a resonance form that only has 6 electrons in one outer shell

Even if this requires a positive formal charge to be placed on a more electronegative atom

Second row atoms are always more stable with a filled outer shell

Resonance Forms

O O

Octet rule not obeyed major

Curved arrows represent movement of electrons

As already observed in acid/base reactions, a curved arrow indicates movement of electrons

Arrows always show where electrons are moving

Formal charges on atoms are a result of electrons moving

Resonance Forms

O O O

Drawing resonance structures properly is an aid to predict location of electrons

Remember actual structure is a hybrid of all relevant resonance forms

These resonance forms allow a chemist to predict where excess electron density is located in a molecule

Excess negative charge is located on three carbon atoms, not on all five equally

Resonance Forms

Chemical properties of molecule are not like one resonance form

Have already observed this with acidity difference between ethanol and acetic acid

Empirical Evidence for Resonance

H3CO

OHH3C

O

OH

HCH3CH2OH CH3CH2O

pKa

16

4.8

Observe also with dipole values for compounds with resonance

N OH3C

H3CH3C

N N O N N O

4.47 D 1.30 D

Interatomic distances do not correspond to single, double, or triple bonds

H3C CH3 H2C CH2

1.54 Å 1.32 Å

1.39 Å

Empirical Evidence for Resonance

Each C-C bond is part single and part double bond