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

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.”

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

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

Understanding Organic Chemistry is Learning about Bonding and the location of Electrons

O

H

H H

H

H H

Consider a molecule called acetone

Shorthand drawing All atoms indicated 3-dimensional drawing

OO

H

H H

H

H H

Electron density plot

wedge and dash lines have meaning

By understanding, and predicting, the properties and orientation of bonds Organic reactions can be predicted without memorization

Compare these compounds:

Why does carbon form such a variety of bonds?

Realize position in the periodic table

Organic Chemistry focuses on the second row of the periodic table

Carbon is in the exact middle of the second row

This position allows carbon to form various bonds **more specifically it allows carbon to share electrons easily to form bonds

Organic chemistry is fundamentally about what electrons do, and how they ‘behave’

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 chemsitry

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)

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)

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

As the number of electrons increase they reside in concentric shells where each shell contains subshells known as atomic orbitals of different properties and shapes

shell atomic orbital

1 s 2 s,p 3 s,p,d 4 s,p,d,f

The different types of atomic orbitals have a characteristic number of orbitals

s 1 p 3 d 5 f 7

Organic chemistry is primarily concerned with only the first two shells (therefore only the s and p atomic orbitals)

Also due to only two spin states for an electron Only two electrons can be placed in each orbital

(Pauli exclusion principle)

Orbitals are filled by electrons starting in the lowest energy orbital

Carbon can therefore only have a maximum of 8 electrons in its outer (2nd) shell 4 atomic orbitals (1s and 3p) with 2 electrons per orbital

Electronic Configuration for Carbon through Fluorine

When filling degenerate orbitals electrons will go into different orbitals before pairing in the same orbital (Hund’s rule)

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

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 energy

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

Halogens, on the other hand, readily gain an electron to complete the outer shell

Consider fluorine (9 electrons, 7 valence electrons)

The outer shell needs one electron to be filled

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

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)

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

Representation of Organic Structures

Lewis Structures As already shown with F2 each valence electron

is represented by a dot

Two dots between two atoms represents a single bond (two shared electrons)

Consider methane (a molecule with CH4 molecular formula)

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)

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?

Electronegativity Tables

Linus Pauling established values to associate with each element

Elements toward the upper righthand 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 (2.2) Li (1.0) Be (1.6) B (1.8) C (2.5) N (3.0) O (3.4) F (4.0)

Cl (3.2)

Br (3.0)

I (2.7)

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

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

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

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

The Arrhenius definition is used to determine the strength of an acid The strength is determined by how easily the molecule dissociates

in water to give a hydronium ion

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

Due to the magnitude, the acid strength is expressed in a logarithmic scale

pH = -log10 [H3O+]

Therefore in neutral solutions pH = 7

Acidic solutions pH < 7

Basic solutions pH > 7

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

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

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”

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

How to Predict the Relative Strength of Acids

- Common point is the ability to stabilize a negative charge (molecules that can handle more excess electron density after deprotonation are stronger acids)

1.  Amongst atoms of similar size, the atom with a greater electronegativity will be a stronger acid

2. Bigger atoms will be stronger acids

Consider size of atom where charge is located

This trend usually is relevant when comparing atoms in the same column (as the atom becomes larger going down a column,

the excess negative charge is more stabilized)

3.  Polar bonds near anion source can stabilize negative charge (inductive effect)

Consider acetic acid derivatives:

Electron withdrawing groups can pull electron density away from another region of molecule (this “through bond” effect is called inductive)

4. Resonance can lower electron density on a given atom

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

“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

A carbonyl group is a common resonance source

The negative charge can therefore be delocalized over both oxygen atoms

Resonance of Acetate Anion

Rotation of Acetate Group

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

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

Actual “hybrid” structure

Remember also that number of paired electrons must be constant

Factors to affect 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

- 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

- 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

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

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

Empirical Evidence for Resonance

Chemical properties of molecule are not like one resonance form

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

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