Organic chemistry A

Chapter 1

Introduction

By Prof. Dr.

Adel M. Awadallah

Islamic University of Gaza

Chapter 1

Bonding and IsomerismOrganic chemistry is the chemistry of the compounds of carbon

Atoms consist mainly froma) Nucleus: (containing Protons and Neutrons)

Protons (positive particles, Atomic Number)

Neutrons (Neutral particles)

Protons + Neutrons (Atomic weight)

b) Electrons: negatively charged particles

Electronic Configuration (review)1) Shells (n = 1,2,3,4, ….)2) (Subshells , s, p, d, f)3) (orbitals: rgions of space around the nucleus containing

electrons)4) each orbital contains only 2 electrons with different spins

SubshellNumber of orbitals (electrons)Name of orbitals

S1 (2)s

P3 (6)Px, py, pz

d5 (10)Dxy, dyz, dxz, dz2,

dX2

-y2

Examples:Filled Shells

Play almost no role in chemical reaction

Valence Electrons

They are the outermost electrons and they are mainly involved in chemical bonding

Ionic CompoundsThey are formed by the transfer of one or more valence

electrons from one atom to another

Electropositive atoms: give up electrons and form cations.

Electronegative atoms: accept electrons and form anions

Ionic compounds: are composed of positively charged cations and negatively charged cations

Covalent compoundsA covalent bond: is formed when two atoms share one or more electron

pairs. A molecule consists of two or more atoms joined by covalent bonds

Bond energy: is the energy necessary to break a mole (6.022 x 1023)of covalent bonds.

Bond length: is the average distance between two covalently bonded atoms

• Valence and bonding in organic compounds

O

O

Cl

N

CH3

H

H

H

H

H

• Polarity of Bonds: depends on the electronegativity difference

• Examples:

• Polarity of Molecules: depends on the sum of the polarity of bonds (geometry)

• Isomers: different compounds having the same molecular formula

CH3CH2OH CH3OCH3

ClI

Br

H

ICl

Br

HCH3 CH3CH3CH3

CH3

Et

Cl

Cl HH

CH3

Et

Cl

ClH

H

ClH

Cl H

ClH

H Cl

Cl H

ClH

Cl Cl

HH

IsomersDifferent compounds that have the same molecular formula

Constitutional isomers Stereoisomers

Atoms are attached to one another in different ways

Atoms have the same constitution but different arrangement in space

Enantiomers Diasteriomers

Non superimposable mirror- image stereoisomers

Non mirror-image stereoisomers

Conformational Configurational

Interconvertable by rotation

Interconvertable by bond breaking

SR

cis- trans-

cis trans

• Abbreviated structural formulas

CH3CH2CH2OH OH OH

H

HH

H

H

H

H

CH2

CH2

CH2

CH2

CH2

= =

=

• Formal Charges: are the charges that each atom carries, and can be calculated as follows

• Formal charge = Valence electrons – bonds – electrons

• Example:

• Resonance: arises whenever we can write two or more structures for a molecule with different arrangements of electrons but identical arrangement of atoms

• Arrows:

CH3 CH3

Reaction

equilibrium reaction

Resonance

Movement of two electrons

Movement of one electron

straight arrow

double headed straight arrow

Pair of straight arrows with half heads

Curved arrow

Fishhook

2CH3.

• The orbital view of bonding

+

H. H. H - H

H.

+

F. H - F

+

F.F. F - F

• Orbital picture of Methane and ethane

•

The bond formed by end-to-end overlap is called a sigma bond.

Bond Angles in Methane

• Bonding in Ethylene (ethene)

1s2

sp2 sp2 sp2

p

A pi bond is one in which the electrons in the p orbitals are held above and below the plane of the molecule.The sigma bond is stronger than the pi bond.A double bond is formed from a sigma bond and a pi bond, and so it is stronger than a single bond.

• Bonding in acetylene (ethyne)

• Classification of organic compounds according to functional groups

OHO

R H

O

SHS

R OH

O

R OR

O

R R

O

R X

O

R NH2

O

R NH2 R CN

Alkane Alkene Alkyne Aromatic

Alcohol Thiol Ether Thioether

Aldehyde Ketone Carboxylic acid Ester Acid halide Amide

Amine Nitrile

Intramolecular Forces

Bond Dissociation Energy: The amount of energy consumed or liberated when a bond is broken orformed

Intermolecular ForcesIntermolecular forces: attractive forces between molecules.

•Intramolecular forces: hold atoms together in a molecule.(chemical bonding)

•Intramolecular forces stabilize individual molecules, whereas intermolecular forces are primarily responsible for the bulk properties of matter (i.e. melting point and boiling point).

•Intermolecular vs Intramolecular•41 kJ to vaporize 1 mole of water (inter)

•930 kJ to break all O-H bonds in 1 mole of water (intra)•Generally, intermolecular forces are much weaker than intramolecular

forces .•The boiling points of substances often reflect the strength of the

intermolecular forces operating among the molecules. At the boiling point, enough energy must be supplied to overcome the attractive intermolecular forces between liquid molecules. The same is true for melting a solid

• Types of Intermolecular forces:• 1- Dipole-Dipole Forces: are attractive forces that act between polar

molecules•

• The larger the dipole moments, the greater the force.

• The Hydrogen Bond: is a special type of dipole-dipole interaction between the hydrogen atom in a polar N-H, O-H, or F-H bond and an electronegative O, N, or F atom.

• Aــــ H---B or Aــــ H---A• A & B are N, O, or F• The average energy of a hydrogen bond is quite large for a dipole-dipole

interaction (up tp 40KJ/mol). Thus, hydrogen bonds are a powerful force in determining the structure and properties of many compounds.

• 2- Ion-Dipole Forces: Attractive forces between an ion (either a cation or an anion) and a polar molecule.

•

• 3- Dispersion (London) force: arise as a result of temporary dipoles induced in the molecules or atoms.

• - act between all molecules.• - only force between nonpolar molecules and noble gas

atoms.• - strength depends primarily on size of molecule, very weak

for small molecule but fairly strong for large ones. (Shape plays a role too.)

- +

Na+

+ -

I -

- +

-

-

(a)

+

Cation Induced dipole

- +

-

Induced dipole

(b)

- +

(c)

Dipole

•Melting points• Melting is the change from the highly ordered arrangement

of particles in the crystalline lattice to the more random arrangement that characterizes a liquid

•Mp NaCl = 801

•Mp CH4 = -183

Boiling point: The temperature at which the vapor pressure of a liquid equals the external pressure

Boiling involves the breaking away from the liquid of individual molecules or pairs of oppositely charged ions

Associated liquids

(Hydrogen bonded liquids)

Have relatively high bp.

HF boils 100 oC higher than HCl

H2O boils 160 oC higher than H2S

Solubility: Like dissolve likeIonic Solutes are dissolved in water or very polar

solvents. They form ion-dipole bonds

Non-ionic solutes: solubility depends on polarityMethane (London forces ) in CCl4 (London forces )

Methanol (hydrogen bonding) in water (hydrogen bonding)

Acids and bases:Arrhenius Definition Acid: H+ donor, Base: OH- donor

HCl + NaOH → NaCl + H2O

Lowry-Brønsted Acids and Bases:

Acid: H+ donor , Base: H+ acceptor

HCl + :OH- → Cl- + H2O

HCl + :NH3 → Cl- + NH4+

Lewis Definition: Acid: e acceptor, base: electron donor

H+ + :OH- → H2O

F3B + :NH3 → F3B:NH3