Basic of Org Chem

8/8/2019 Basic of Org Chem

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1 Basics of Organic Chemistry

Organic chemistry is the study of molecules based on carbon

Several interesting organic molecules:

N

N N

N

O

O

O

OH

H

H

NH

N

N

N

O

O

SN

N

OO

N

OO

O

O

caffeinetetrahydrocannabinol (THC)

sildenafil citrate (Viagra)

cocaine

1.1 Electron configuration of atoms

Most of this is assumed knowledge from high school and first year chemistry; the following is

only be a brief review of this material

Ground-state electron configuration - the electron configuration of lowest energy for an atom,

molecule, or ion

When determining the ground-state electron configuration of an atom, three rules are used:

1. Aufbau principle orbitals are filled in order of increasing energy, from lowest to highest

2. Pauli exclusion principle no more than two electrons may be present in an orbital; if two

electrons are present then their spins must be paired

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3. Hunds rule when orbitals of equivalent energy are available but there are not enough electrons

to fill all of them completely, then one electron is added to each equivalent orbital before a second

electron is added to any one of them; this is especially true of the p-orbitals

Ground state electron configuration for carbon: 1s2

2s2

2px1

2py1

2pz0

1.2 Lewis structures

Valence shell the outermost electron shell of an atom; contains the valence electrons

Lewis structure of an atom the symbol of an element surrounded by a number of dots equal

to the number of electrons in the valence shell of the atom

N

OB

H

Li Be

Na Mg

He

IA IIA VA VIA VIIA VIIIAIIIA IVA

Cl

F

S

Ne

Ar

C

SiAl P

Table 1.1 Lewis structures for the first 18 elements of the Periodic Table.

1.3 Electronegativity and chemical bonds

Electronegativity a measure of the force of an atoms attraction for electrons that it shares

with another atom in a chemical bond

o Plays an extremely important role in the reactivity of various functional groups in

organic chemistry

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Figure 1.1 The electronegativity values for some atoms commonly found in organic chemistry. Values are expressed in

Pauling units.

In most cases carbon atoms are more electropositive than the other atoms to which they are

bonded, but you should always refer to the electronegativity values to be sure

two types of chemical bonds:

1. ionic bonds

2. covalent bonds

ionic bond a chemical bond resulting from the electrostatic interactions between a cation and

an anion

o generally occur between a metal and a non-metal

o the electronegativity of the two atoms involved differs by more than 1.9 Pauling units

o NaCl, MgSO4, KOH

Covalent bond a bond resulting from the sharing of two or more electrons between a set of

atoms

o Non-polar covalent bond a covalent bond between two atoms whose electronegativity

is different by less than 0.5 Pauling units

H2, CH4

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o Polar-covalent bond a covalent bond between two atoms whose electronegativities are

different by 0.5-1.9 Pauling units

Si-C, C-O, H-Cl

o

Two ways of depicting polar covalent bonds:

Draw a plus sign over the more electropositive atom and then draw an arrow to

the more electronegative atom through the plus sign

Use the lower-case Greek letter delta (means slight in chemistry speak) and

include the charge that the atom possesses

H Cl H Cl2.1 3.0

+ -

Figure 1.2 The two different ways of depicting the polarity of a bond.

Using this notation it is possible to determine whether a molecule is polar in nature (e.g. H2O)

or non-polar (e.g. CCl4)

o It is possible for a molecule to be non-polar yet still have polar bonds

1.4 Drawing Lewis structures

1. determine the number of valence electrons in the molecule or ion count the number of valence

electrons contributed by each atom

2. determine the connectivity of the atoms in the molecule usually needs to be determined

experimentally for all but the most simple molecules

3. connect the atoms with single bonds and arrange the remaining electrons such that each atom

has a complete outer shell

bonding electrons valence electrons involved in forming a covalent bond (i.e., shared electrons)

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non-bonding electrons valence electrons not involved in forming covalent bonds (a.k.a. lone

pairs)

Formal charge the charge on an atom in a polyatomic ion or molecule

H O+

H

H

O O

O

HO

N+ O

OHFC = 6 - (2+(0.5 x 6))

= 1 FC = 6 - (6+(0.5 x 1))= -1

FCN = 5 - (0+(0.5 x 8))

= 1FCO = 6 - (6+(0.5 x 2))

= -1

H ydrogen chloride

M ethane A mmonia

W ater

H O

H

H

H NH C

H

H

H C l

H

H

H 2O ( 8 )

N H 3 (8)C H 4 (8 )

H C l ( 8 )

Figure 1.3 The Lewis structures of some common molecules. The number of valence

electrons is shown in parentheses.

# of valence

electrons in

unbonded atom

all

unshared

electrons

one half

of all shared

electrons

+Formal

charge=

Figure 1.4 The formal charge on some common molecules used in organic chemistry.

Most atoms prefer to have an octet in their valence shell (why?)

There are however exceptions to the octet rule (hydrogen sulphide 8, dimethylsulfoxide

10, and sulphuric acid 12)

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1.5 Functional Groups

This section is simply an introduction to common functional groups found in organic chemistry

It is important to recognize these functional groups

OH

OH OH

NH2

N

H

N

R

Figure 1.5 Common functional groups in organic chemistry.

1.6 Resonance and the Movement of Electrons

Electrons are not static, not even when involved in covalent bonds in molecules

This movement of electrons is known asresonance

Generally involves the movement of electrons into, or out of, pi bonds

H

O

R R

O

R O

O

R R NH

O

R

Alcohols:

primary secondary tertiary

Amines:

primary secondary tertiary

Carbonyl Compounds:

aldehyde ketone ester amide

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Ethanoate ion

(Acetate ion)

C

O

O

C H 3C

O

O

C H 3

-

-

and

the oxygen atoms in the acetate ion are not different from each other

the negative charge is shared between the two oxygen atoms

the two molecules are resonance contributors

resonance only depicts the movement of electrons, not atoms

o sigma bonds remain intact

resonance imparts stability to a molecule

o the more resonance forms are possible, the more stable a molecule is (i.e., benzene)

Pauls Rule of Happy Molecules for the most part, organic molecules are happiest if they are

neutral, content if they possess partial charges, and usually ticked off (reactive) if they possess a

full formal charge

usecurly arrows to denote the movement of electrons

o understanding the use of curly arrows is extremely important in understanding organic

chemistry an excellent way to keep track of your electrons

N

O

O

N

O

O

-

-

Ethanoate ion

(equivalent

contributing

structures)

C

O

O

C H 3 C

O

O

C H 3

-

Nitrite ion

(equivalent

contributing

structures)

-

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reaction arrow- goesfrom reagents to products

curly arrow- depictsthe movement ofelectrons

resonance arrow- placedbetween two resonanceforms of a molecule

equilibrium arrow- placedbetween the reagents andproducts in equilibrium reactions

Figure 1.6 Arrows used in organic chemistry and what they mean.

1.7 Molecular orbital theory of covalent bonding

1.7.1 Atomic s and p orbitals

Atomic orbitals represent the probability of finding electrons in a given space around the

nucleus of an atom

o The part of an orbital in which there is zero probability of finding an electron is called a

node

Figure 1.7 The shape of the s atomic orbital.

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s orbitals are symmetric in all three planes

Figure 1.8 The shape of atomic p orbitals depicting the two ways of denoting the phase of the orbital

lobes and showing the nodal plane.

Figure 1.9 The three, mutually perpendicular p orbitals.

1.7.2 Formation of molecular orbitals

Molecular orbital (MO) theory the theory in which electrons in molecules occupy

molecular orbitals formed by the combination of the atomic orbitals of the atoms that make

up the molecule

Three rules for applying MO theory to the formation of covalent bonds:

1. the number of molecular orbitals formed is equal to the number of atomic orbitals combined

2. molecular orbitals are arranged in order of increasing energy

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3. molecular orbitals are filled with electrons using the same principles used for filling atomic

orbitals (Aufbau principle, Pauli exclusion principle, Hunds rule)

Figure 1.10 The combination of two 1s orbitals to form a bond.

Sigma () bond a bond in which the electron density is concentrated between the two nuclei

and along the axis joining them

Figure 1.11 The molecular orbital energy diagram for the hydrogen molecule, H2.

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Figure 1.12 The molecular orbital energy diagram of two atomic p orbitals coming together to form a sigma bond.

bonding molecular orbital molecular orbital formed between two atomic orbitals that have

the same phase

anti-bonding molecular orbital molecular orbital formed between two atomic orbitals that

have different phases

pi () bond a bond formed between the overlap of parallel p orbitals on adjacent atoms

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Figure 1.13 The molecular orbital energy diagram for the formation of a pi bond.

Sigma bonds are stronger than pi bonds

1.7.3 Hybridization of atomic orbitals

The second-period elements use 2s and 2p atomic orbitals to form molecular orbitals

Hybrid orbital orbital formed by the combination of two or more types of atomic orbitals

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3.7.1.1sp3 Hybrid orbitals

Figure 1.14 Four tetrahedral sp3 orbitals.

sp3 atomic orbitals are made up of one s orbital and 3 p orbitals

consist of two lobes, one larger than the other

the bond angle between the orbitals is 109.5

3.7.1.2sp2Hybrid orbitals

Figure 1.15 Orbitals contributing to the bonding structure of ethene. Unhybridized p orbital is not shown.

Consists of one s orbital and 2 p orbitals

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The third p orbital is not involved in hybridization and lies perpendicular to the plane of the sp 2

orbitals

A pair of sp2 orbitals involved in forming a bond and a pair of sp2 orbitals involved in

forming a bond

o sp2 orbitals involved in the formation of double bonds

The bond angle in sp2 systems is 120

3.7.1.3sp Hybrid orbitals

Figure 1.16 The orbitals contributing to the bonding structure of ethyne (a.k.a. acetylene).

sp orbitals result in one bond and two bonds to carbon

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o Produce carbon-carbon triple bonds

The bond angle between these orbitals is 180

Table 1.2 Bond lengths and bond strengths for ethane, ethene, and ethyne.

Molecule Bond Bond Length (pm) Bond Strength (kJ/mol)

Ethane C-C 153.2 368

Ethene (ethylene) C-C 133.9 611

Ethyne (acetylene) C-C 121.2 837

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