Chapter 1: Organic Chemistry and ReactionsMarshahida Mat YashimFakulti Kejuruteraan KimiaUiTM Terengganu

Course Learning Outcomes

The student should be able to: •Define of homologous series• Give examples and nomenclature of

organic compounds such as alkanes, alkyl substituents, cycloalkanes, alkyl halides according to IUPAC systems.

•Classify of alkyl halides, alcohols, ethers and amides

• Identify physical properties of alkanes, alkyl halides, alcohols, ethers and amides

Subtopic•1.1 Structure and bonding •1.2 Hybridization •1.3 σ and π bonds •1.4 Bond polarity •1.5 Formal Charges •1.6 Chemical structures •1.7 Resonance •1.8 Classifications, Aliphatic, Aromatic and

Derivatives •1.9 Organic Reactions and Mechanisms

Introduction

•“Organic” – until mid 1800’s referred to compounds from living sources (mineral sources were “inorganic”)

•Today, organic compounds are those based on carbon structures and organic chemistry studies their structures and reactions▫Includes biological molecules, drugs, solvents,

dyes▫Does not include metal salts and materials

(inorganic)▫Does not include materials of large repeating

molecules without sequences (polymers)

•An atom consists of a nucleus of protons and neutrons surrounded by electrons. Each of the elements in the periodic table is classified according to its atomic number, which is the number of protons in that element's nucleus.

•A neutral atoms have the same number of electrons and protons, but they can have a varying number of neutrons.

•Charged ions :- Cation and Anion•Cation is positively charged and has a fewer

electrons than its neutral form.•Anion is negatively charged and has more

electron than its atoms.

•Electrons are held most tightly with nucleus. Electrons are first added to the shells closest to the nucleus. Each shell contains a number of subshells called orbitals.

•4 different kinds of orbital s,p,d,f. (have particular shape and occupies certain shape)

1.1 Structure & Bonding•Chemical bonds : attraction between the

atoms that hold compounds together.•Atoms without noble gas configuration

generally reacts to produce such configuration to increase stability.

•“Through bonding, atoms attain a complete outer shell of valence electron and atom attain a stable noble gas configuration of electron.”

•Two major types of bonding : ionic and covalent bond

•Ionic Bonds : formed by the transfer of one or more electrons from one atom to another to create ions;

•Covalent Bond : a bond that results when atoms share electrons.

•Ionic Bonds : Resulting ions are held together by extremely strong electrostatic interaction. Positively charged cation formed from the element on the left side attracts a negatively charged anion formed from the element on the right side.

•Covalent Bond : Occurs with elements like C which would otherwise have to gain or lose several electron to form an anion with complete valence shell.▫Covalent bond is a two electron bond and a

compound with covalent bonds called molecule.▫Covalent bond also form between 2 elements

from same side, such H2, Cl2, CH4

Structure•The different ways to draw organic

molecules include Kekulé (straight-line), Condensed Formulas, and Bond-Line Formulas (zig-zag).

•Kekulé structures are similar to Lewis Structures, but instead of covalent bonds being represented by electron dots, the two shared electrons are shown by a line.

•Lone pairs remain as two electron dots, or are sometimes left out even though they are still there.

•A condensed formula is made up of the elemental symbols. The order of the atoms suggests the connectivity.

•If an atom that doesn't have a complete octet by the time you reach the next atom, then it's possible that there are double or triple bonds.

•E.g.: COOH means C(=O)-O-H instead of CH3-C-O-O-H because carbon does not have a complete octet and oxygens

•Bond-Line (a.k.a. zig-zag) Formulas : This formula is full of bonds and lines, and because of the typical (more stable) bonds that atoms tend to make in molecules, they often end up looking like zig-zag lines.

•It is usually not drawing the H's that are bonded to carbon, we do draw them in if they are connected to other atoms besides carbon

•Dashed-Wedged Line Structure: A widely used way of showing the 3D structure of molecules

•The placement of different atom could yield different molecules even if the molecules formulas were exactly the same.

1.2 Hybridization •Hybrization is the combination of two or

more atomic orbitals to form the same number of hybrid orbitals, each having same shape and energy.

•The 's' orbital is spherical about the nucleus and the 'p' orbitals are like double headed balloons arranged along the axis of (imaginary) three dimensional coordinates

•The 'p' orbitals are oriented at 90º to one another and yet there are few molecules that show a bond angle of 90º

sp3 hybridization : methane (CH4)

• In order to fulfill the octet rule, carbon must use its 4 valence electrons when bonding to other atoms.

•However, only unpaired electrons can bond. That means that the two paired electrons occupying the 2s orbital must become unpaired before they can bond. Since the energy gap between the 2s and 2p orbitals is very small, one of the 2s electrons can be promoted to the empty 2p orbital

hybridized

•The resulting shape is then a tetrahedron, where the carbon nucleus is at the center and the orbitals point to the corners of the tetrahedron.

•The ideal angle between orbitals is then 109.5 degrees.

Other Hybridization Patterns-sp and sp2

•(I) One 2s orbital and three 2p orbitals form four sp3 hybrid orbitals.

•(II) One 2s orbital and two 2p orbitals form three sp2 hybrid orbitals

•(III) One 2s orbital and one 2p orbital form two 2 sp hybid orbitals.

•Example:

•Hybrid orbital NH3

•Hybrid Orbital H2O

1.3 σ and π bonds •atomic orbitals (pure or hybrid) of different

atoms overlap to form covalent bonds, they may approach each other in two major ways: head to head, or sideways

•When orbitals approach each other in a head to head fashion, the resulting covalent bonds are called sigma bonds (σ bonds)

•The bonds between the sp3 orbitals of hybridized carbon and the s orbitals of hydrogen in methane are also example of sigma bonds.

•Two sp3 carbons can also overlap to form a C–C sigma bond where two sp3 orbitals overlap head to head,such as in the formation of the ethane molecule:

•Two p orbitals approach each other sideways when two sp2 hybridized carbon atoms approach each other to bond. The bond formed by the sp2 orbitals is a sigma bond, and the bond formed by the p orbitals is called a pi bond (π bond)

1.4 Bond polarity • Electronegativity is a measure of atom’s attraction

for electron in a bond.• It is used as a guideline to indicate whether the

electron in a bond are equally shared or unequally shared.

• Bond polarity is a useful concept for describing the sharing of electrons between atoms.

• A non-polar covalent bond is one in which the electrons are shared equally between two atoms

• A polar covalent bond is one in which one atom has a greater attraction for the electrons than the other atom. If this relative attraction is great enough, then the bond is an ionic bond .

•A general rule of thumb for predicting the type of bond based upon electronegativity differences:

•If the electronegativities are equal (i.e. if the electronegativity difference is 0), the bond is non-polar covalent

•If the difference in electronegativities between the two atoms is greater than 0, but less than 2.0, the bond is polar covalent

•If the difference in electronegativities between the two atoms is 2.0, or greater, the bond is ionic

• To determine whether a molecule has a net dipole : (1) Use electronegativity differences to identify all of polar bond the direction of the polar dipoles.(2) Determine the geometry around individual atoms by counting groups and decide if individual dipoles cancel or reinforce each other in space.

1.4 Bond polarity•Electronegativity is a measure of an

atom’s attraction for electrons in a bond.

Electronegativity values for some common elements

1.4 Bond polarity

• A carbon—carbon bond is nonpolar. The same is true whenever two different atoms having similar electronegativities are bonded together.

• C—H bonds are considered to be nonpolar because the electronegativity difference between C and H is small.

1.4 Bond polarity•Example: In the C—O bond, the electrons

are pulled away from C (2.5) toward O (3.4), the element of higher electronegativity.

•The bond is polar, or polar covalent. The bond is said to have dipole; that is, separation of charge.

To determine whether a molecule has a net dipole : (1) Use electronegativity differences to identify all of polar bond the direction of the polar dipoles.(2) Determine the geometry around individual atoms by counting groups and decide if individual dipoles cancel or reinforce each other in space.

1.4 Bond polarity•A polar molecule has either one polar

bond, or two or more bond dipoles that reinforce each other. An example is water:

δ+means the indicated atom is electron deficient.δ-means the indicated atom is electron rich.

1.4 Bond polarity•A nonpolar molecule has either no polar

bonds, or two or more bond dipoles that cancel. An example is carbon dioxide:

Try this…

•Decide whether the molecules represented by the following formulas are polar or nonpolar.

•CH2O

•BF3

•NH3

•CH2Cl2

•CCl4

1.5 Formal Charges • It is possible to assign positive and negative

charges to atoms in Lewis structures. These formal charges often give insight into chemical reactivity. The sum of all the formal charges must equal the total charge (if any) on the Lewis structure.

• The formal charge on an atom is obtained by subtracting the number of valence electrons that "belong" to the atom in its bonded state from the number of valence electrons in the neutral free atom.

1.5 Formal Charges • Assignment of Valence Electrons in the Bonded

State(1) Electrons in covalent bonds are shared equally by the two atoms held together by the bond.(2) Nonbonding electrons are assigned completely to the atom where they are located.

1.5 Formal Charges : Example

1.5 Formal Charges : Example

Table of Common Formal Charges

1.7 Resonance• Some molecules cannot be adequately represented by

a single Lewis structure. For example, two valid Lewis structures can be drawn for the anion (HCONH)¯. One structure has a negatively charged N atom and a C-O double bond; the other has a negatively charged O atom and a C-N double bond.

• These structures are called resonance structures or resonance forms. A double headed arrow is used to separate the two resonance structures.

Resonance Theory• Regarding the two resonance forms of (HCONH)-shown

below, it should be noted that:

• Neither resonance structure is an accurate representation for (HCONH)-. The true structure is a composite of both resonance forms and is called a resonance hybrid.

• The hybrid shows characteristics of both structures.• Resonance allows certain electron pairs to be

delocalized over two or more atoms, and this delocalization adds stability.

• A molecule with two or more resonance forms is said to be resonance stabilized.

Resonance TheoryThe following basic principles of resonance theory

should be kept in mind:• 1. Resonance structures are not real. An

individual resonance structure does not accurately represent the structure of a molecule or ion.

Only the hybrid does.

• 2. Resonance structures are not in equilibrium with each other. There is no movement of electrons from one form to another.

• 3. Resonance structures are not isomers. Two isomers differ in the arrangement of both atoms and electrons, whereas resonance structures differ only in the arrangement of electrons.

Drawing Resonance Structure• RULE 1: Two resonance structures differ in the

position of multiple bonds and non bonded electrons. The placement of atoms and single bonds always stays the same.

• RULE 2: Two resonance structures must have the same number of unpaired electrons.

Drawing Resonance Structure• RULE 1: Resonance structures must be valid Lewis

structures. Hydrogen must have two electrons and no second-row element can have more than eight.

• Curved arrow notation is a convention that is used to show how electron position differs between the two resonance forms.

• Curved arrow notation shows the movement of an electron pair. The tail of the arrow always begins at the electron pair, either in a bond or lone pair. The head points to where the electron pair “moves.”

Drawing Resonance Structure

Drawing Resonance Structure

In the two examples above, a lone pair is located on an atom directly bonded to a double bond.

Drawing Resonance Structure

• In the above examples, an atom bearing a (+) charge is bonded either to a double bond or an atom with a lone pair

Resonance Hybrid•A “better” resonance structure is one that

has more bonds and fewer charges.

Drawing Resonance Hybrid

Resonance Theory

Whenever a molecule or ion can be represented by two or more Lewis structures that differ only in the positions of electrons:

(1) No single resonance structure is a correct representation of the molecule. None will be in complete accord with the properties of the molecule.

(2) The actual molecule or ion will be better represented by a hybrid of all the resonance structures.

Resonance Hybrid• We can draw a Lewis-like structure that provides

a better description of the actual character of the carbonate ion by blending the resonance structures into a single resonance hybrid:▫Draw the skeletal structure, using solid lines for

the bonds that are found in all of the resonance structures.

▫Where there is sometimes a bond and sometimes not, draw a dotted line.

▫Draw only those lone pairs that are found on every one of the resonance structures. (Leave off the lone pairs that are on one or more resonance structure but not on all of them.)

C

O

O O

: :

:::

:: :

I

C

O

O O

: :

:

:: :

II

::

C

O

O O

: :

:

::: :

:

III

Lewis structures I, II and III are the resonance structures of the carbonate ion.

The double-headed arrow is specifically used to indicate resonance structures that contribute to the hybrid.

The real electronic structure of the carbonate anion is a hybrid of the three resonance structures.

hybrid

C

O

O O

there are three equivalentbonds in the carbonate anion

Each carbon-oxygenbond is 1 and 1/3 bonds.

Each oxygen has a formal charge of 2/3 -.

Structure of the Carbonate Ion

: :

:::

:: :

Any single Lewis structure shows two types of carbon-oxygen bonds: two single and one double.

C

O

O O

This equivalency is explained by resonance theory.

X-ray structural studies show that C-O bonds are longer than C=O bonds: 1.43 vs 1.20 Å.

All three carbon-oxygen bonds are the same length in the carbonate ion: 1.28 Å.

C

O

O O

hybrid

C

O

O O

C

O

O OC

O

O OC

O

O O

: : : : : :

: :

: :

: :: :::

:: : : :

:

I II III:

:

Examples

The carbonate ion: CO32-

Three equivalent C-O bonds. Each oxygen has a formal charge of 2/3 -.

hybrid

C

O

O O

N:O:

:

O:: :

-

-O::

=+ N

:O:

O:: :-O:

=+

::-N:O:

:O: :

-

=O: ::-

+

Examples

The nitrate ion, NO3-

The Lewis structures that satisfy the octet rule are:

Note the formal charges.

hybrid

+N

O

O O

There are three equivalent N-O bonds and each oxygen has a charge of 2/3-.

1.8 Classifications, Aliphatic, Aromatic and Derivatives

• In organic chemistry, compound composed of carbon and hydrogen are divided into two classes: aliphatic and aromatic compounds.

• Aliphatic compounds: open-chain compounds and ring compounds that are chemically similar to open-chain compounds. Alkanes, alkenes, alkynes, dienes, alicyclics, etc.

• Aromatic compounds: unsaturated ring compounds that are far more stable than they should be and resist the addition reactions typical of unsaturated aliphatic compounds. Benzene and related compounds.

• Aliphatic compound can be cyclic and acyclic. • They also can be saturated and unsaturated.• Saturated-max num of H per C.

A reaction mechanism is a detailed description of the bonding changes as a reaction proceeds. The reaction mechanism also includes the many important principles of organic chemistry. A plausible reaction mechanism must be consistent with the principles of organic chemistry.

Organic Reactions and their Mechanisms

Four General Categories of Organic Reactions

.Organic reactions tend to fall into four categories: substitutions, additions, eliminations, rearrangements

1.9 Organic Reactions and Mechanisms

Substitutions

In a substitution reaction, one atom or group replaces another in a structure. This type of reaction is commonly observed in saturated hydrocarbons and aromatics.

an alkyl halide

+ Na+ -OH + Na+ X-R-X R-OHH2O

an aromatic+ Br2

+ HBrAr-HFe

Ar-Br

The mechanisms of the above two substitution reactions are completely different.

Additions

.

Addition reactions are found in organic compounds with multiple bonds: alkenes, alkynes, carbonyl-containing compounds. In this reaction, the-component of the multiple bond is lost as new bonds are formed to the carbon (or other atomic) centers

+ Br2

Bromine adds to the alkene (ethene).

C C

H

H

H

H

C C

H

H

Br Br

H

H

Eliminations

.

These reactions are the reverse of addition reactions. In an elimination reaction, a molecule loses atoms or groups from adjoining carbon (or other atomic) centers, forming a multiple bond

C C

H

H

H

H

+ KOH + K+ Br- + H2OC C

H

H

H Br

H

H

The above reaction is a dehydrohalogenation, loss of HBr, of an alkyl halide to form an alkene.

Rearrangements

.In rearrangement reactions, there is a reorganization of the atoms or groups in a structure

H+

.In the presence of acid, the alkene on the left rearranges to the alkene on the right

Reaction Mechanisms and Chemical Intermediates

.

Reaction mechanisms are detailed descriptions of changes at the molecular level as reactants become products. Often the reactions involve a sequence of steps with one or more chemical species called intermediates that are formed and consumed

.Chemical intermediates typically are not stable structures that can be put in a bottle. Many exist for very short times (10-6 - 10-9 seconds)

We will explore how structural and electronic influences affect the stability of chemical intermediates and, thereby, control the path that reactions follow.

Bond Making and Bond Breaking Processes: Heterolysis and Homolysis

.A covalent bond may break by either of two different processes: heterolysis or homolysis

A : B +ions

+ -A B .

Double-barbed arrowis used to show movementof an electron pair

Homolysis (Gr: homo-"the same" + lysis)

Heterolysis (Gr: hetero- "different" + lysis-"cleavage")

A : B . +

radicals

A B. ..

Single-barbed arrowis used to show movementof a single electron

Bond Heterolysis

.

Bond heterolysis typically is found with polar covalent bonds where the intrinsic electronegativities promote polarization of the bonding electrons

A : B ++ -A B

polar covalent bond

Assisted Bond Heterolysis

.Bond heterolysis may be assisted by a second molecule or chemical species that can either donate or accept an electron pair.

A : B C

+C A : B

+ B+ -C : A C donates an electron

pair in forming bond to Aas B departs as anion.

A : B D+

+A : B

D+ +A+ B : D D+ accepts an electron pairas it bonds to B releasing A+.