PHAR 1123PHARMACEUTICAL CHEMISTRY II

STRUCTURE, BONDING, AND ORGANIC REACTIONS

Faculty of PharmacyCyberjaya University College of Medical

Sciences

STRUCTURE AND BONDING

Learning Objectives

1. To discuss electronic structure of atoms.

2. To differentiate shells and orbitals.

3. To write the ground state electron configuration of a given element.

4. To name the two theory of chemical bonds.

5. To discuss valence bond theory and hybrid orbitals.

6. To draw skeletal structures.

7. To discuss polar covalent bonds.

ORGANIC CHEMISTRY

The study of the compounds of CARBON

The chemistry of carbon and only a few other elements(H, O, N, S, P and halogen)

• Video on organic chemistry• http://www.youtube.com/watch?v=JgNg5IQnGhM

Electronic Structure of Atoms

• Atoms are composed of 3 principal kinds of subatomic particles:

- protons

- neutrons

- electrons

• At the center of an atom, is the nucleus, which is a very tiny, extremely dense core. Electron

s

Nucleus (protons + neutrons)

Shells

• Electrons do not move freely.

• Confined to regions of space

called principle energy levels, or

shells.

• Numbered 1, 2, 3, and so forth

from the inside out.

• Each shell contain 2n2 electrons,

n = number of the shells.

Nucleus

12

3

Orbitals

• Shells are divided into

subshells: s, p, d, and f.

• Within this subshells,

electrons are grouped in

orbitals.

• Orbital: a region of space that

can hold 2 electrons.

Different p Orbitals

• The 3 different p orbitals within a given shell are oriented in space

along mutually perpendicular directions, (px, py, pz).

• The two lobes of each p orbital are separated by a region of zero

electron density (node).

• The separated lobes have different algebraic signs, + and –.

• The different signs of the lobes have important consequences with

respect to chemical bonding and chemical reactivity.

Electron Configurations

• Ground-state electron configuration (lowest energy arrangement)

can be predicted by following 3 rules.

Aufbau principle The lowest-energy orbitals fill up first.

Pauli exclusion principle

Electrons act as if they are spinning, which have

2 orientations, up and down. The two electrons

that occupy an orbital must be of opposite spin.

Hund’s ruleIf two or more empty orbitals of equal energy are

available, one electron occupies each with spins

parallel until all orbitals are half full.

• E.g. Ground-state configuration for carbon (6 electrons):

1s2 2s2 2px1 2py

1

or

1s22s22p2

Chemical Bonding Theory

• Atoms bond together because the compound that results is lower in

energy, more stable than separate atoms.

• Energy (heat) always flows out of the chemical system when a

chemical bond forms.

• The bonds are not oriented randomly, they have specific spatial

directions.

• E.g. methane – the 4 hydrogen to which carbon is bonded sit at the

corners of a regular tetrahedron, with carbon in the center.

• Solid lines: bonds in the plane of the

page.

• Heavy wedge line: bond coming out

of the page towards the viewer.

• Dashed line: bond receding back

behind the page, away from the

viewer.

Solid

Heavy wedge

Dashed

Electron Octet

• Electron octet in the atom’s valence shell impart special stability to

the noble gas elements.

• The chemistry of main-group elements is governed by their tendency

to take on the electron configuration of the nearest noble gas.

• Group 1 and group 17 will each lose and gain 1 electrons, forming

ions. The ions are held together in compounds (NaCl) by an

electrostatic attraction that is called ionic bond.

• For elements closer to the middle of the periodic table, e.g. carbon,

it would take too much energy to gain or lose 4 electrons to achieve

a noble gas configuration.

• Thus, carbon bonds to other atoms by sharing the electrons.

• The shared-electron bond is called a covalent bond.

• The neutral collection of atoms held together by covalent bonds is

called a molecule.

Lewis and Kekule Structures

Lewis Kekule

Electron-dot structures Line-bond structures

Valence electrons are represented

as dots

Two-electron covalent bond is

indicated as a line drawn between

atoms

• The number of covalent bonds an atom forms depends on how many

additional valence electrons it needs to reach a noble gas

configuration.

Element Noble gas Bonds formed

Hydrogen (1s) Helium (1s2) 1 bond

Carbon (2s2 2p2) Neon (2s2 2p6) 4 bonds

Nitrogen (2s2 2p3) Neon 3 bonds

Oxygen (2s2 2p4) Neon 2 bonds

Halogens 1 bond

• Valence electrons that are not used for bonding are called lone-pair electrons.

• E.g. nitrogen atom in ammonia.

Chemical Bonds

• Chemical bonds theory are to explain the forming of bonds between

atoms by electron sharing.

• 2 models:

valence bond theory

molecular orbital theory

• Both has their own strengths and weaknesses, thus they are used

interchangeably depending on the circumstances.

• Valence bond theory is the more easily visualized of the two.

Valence Bond Theory

• A covalent bond forms when two atoms approach each other closely

and a singly occupied orbital on one atom overlaps a singly occupied

orbital on the other atom.

• The electrons are now paired in the overlapping orbitals and are

attached to the nuclei of both atoms, thus bonding the atoms

together.

• E.g. Bonding in a hydrogen molecule

• The overlapping orbitals have the elongated egg shape if two spheres

were pressed together.

• If a plane were to pass through the middle of the bond, the

intersection of the plane and the overlapping orbitals would be a

circle.

• The H – H bond is cylindrically symmetrical.

• Such bonds are called sigma (σ) bonds.

• During the bond forming reaction, 436 kJ/mol energy is released.

• The product has less energy than the starting atoms.

• The product is more stable than the reactant.

• Bond strength : 436 kJ/mol.

• There is also an optimum distance between nuclei that leads to

maximum stability, called the bond length.

Stability of Covalent Bonds

• An electron pair occupies the region between two nuclei.

• This arrangement will shield the repulsive forces from one positively

charged nucleus to the other nucleus.

• At the same time, electron pair attracts both nuclei.

• The internuclear distance is fixed to within very narrow limits.

• This distance is the bond length and every covalent bond has a

definite bond length.

sp3 Hybrid Orbitals

• How are sp3 hybrid orbitals formed?

• An s orbital and 3 p orbitals can combine (hybridized) to form 4

equivalent atomic orbitals with tetrahedral orientation.

• The tetrahedrally oriented orbitals are called sp3 hybrids.

4

tetrahedral geometryexcited state

hybridize

4 identical sp3 orbitalszyx2p

2s4 sigma bonds requires 4 hybrid orbitals

CH4C

H

HH

H 2s

2p x y z

ground state

sp3

• Why should sp3 hybrid orbitals formed?

• The sp3 hybrid orbitals are unsymmetrical about the nucleus.

• One of the two lobes is much larger than the other and can

therefore overlap more effectively with an orbital from another atom

when it forms a bond.

• Thus sp3 hybrid orbitals form stronger bonds than do unhybridized s

or p orbitals.

• The asymmetry of sp3 orbitals arises because of the difference in

algebraic signs of each lobes.

• When a p orbital hybridizes with an s orbital, the positive p lobe

adds to the s orbital but the negative p lobe subtracts from the s

orbital.

• E.g. the formation of methane.

Each C – H bond has a strength of 436 kJ/mol and a length of 109 pm. The bond angle is 109.5o.

Ethane

• The same kind of orbital hybridization for methane also accounts for

the bonding together of carbon atoms into chains and rings.

• This makes possible the many millions of organic compounds.

• Ethane is the simplest molecule containing a carbon-carbon bond.

Each C – H bond has a strength of 423 kJ/mol and a length of 109 pm. Each C – C bond has a strength of 376 kJ/mol and a length of 154 pm. The bond angles is ~109.5o.

sp2 Hybrid Orbitals

• Another possibility of hybridization.

• 2s orbital combines with only two of the three available 2p orbitals.

• Three sp2 hybrid orbitals result, one 2p orbital remains unchanged.

• The three sp2 orbitals lie in a plane at angles of 120o to one another,

the remaining p orbital perpendicular to the sp2 plane.

• E.g. ethylene.

pz

H

HC

H

H

CH2=CH2

C

3 sigma bonds requires 3 hybrid orbitals

2s

2p x y z

ground state

2s

2p x y z3 identical sp2 orbitals

hybridize

excited statetrigonal planar geometry

3sp2 pz

sp2 hybrid carbon atom

• When two sp2 hybridized carbons approach each other, they form a

σ bond by sp2-sp2 head-on overlap.

• The unhybridized p orbitals approach with the correct geometry for

sideways overlap, forming pi (Π) bond.

• The combination of an sp2-sp2 σ bond and a 2p-2p Π bond results in

the sharing of four electrons and the formation of a carbon-carbon

double bond.

• H atoms form s bonds with four sp2 orbitals.

• H–C–H and H–C–C bond angles of about 120°.

• C=C double bond in ethylene is shorter and stronger than the single

bond in ethane.

• Ethylene C=C bond length is 134 pm (C–C 154 pm).

• The carbon-carbon double bond is less than twice as strong as a

single bond.

• This is due to the overlap in the π part of the double bond is not as

effective as the overlap in the σ part.

sp Hybrid Orbitals

• A carbon 2s orbital hybridizes with only a single p orbital.

• Two sp hybrid orbitals result, two p orbitals remain unchanged.

• The two sp orbitals are oriented 180o apart (linear molecule) on the x-

axis, while the remaining two p orbitals are perpendicular on the y-axis

and the z-axis.

• E.g. acetylene.

• Two sp hybrid orbitals from each C form sp–sp s bond.

• pz orbitals from each C form a pz–pz bond by sideways overlap

and py orbitals overlap similarly.

• Hybridization video• http://www.youtube.com/watch?v=SJdllffWUqg&feature=related

Molecular Orbital Theory

• Covalent bond formation arises from a combination of atomic orbitals

on different atoms to form molecular orbitals.

• Molecular orbital describes a region of space in a molecule where

electrons are most likely to be found.

Drawing Chemical Structures

• 2-Methylbutane

• Kekule structure:

• Condensed structure: CH3CH2CH(CH3)2

• Skeletal structure:

Rules for drawing skeletal structures:

1. Carbon atoms are not usually shown. Instead, a carbon atom

is assumed to be at each intersection of two lines (bonds)

and at the end of each line. Occasionally, a carbon atom

might be indicated for emphasis or clarity.

2. Hydrogen atoms bonded to carbon are not shown. Since

carbon always has a valence of 4, we mentally supply the

correct number of hydrogen atoms for each carbon.

3. Atoms other than carbon and hydrogen are shown.

Electronegativity

• The two known chemical bonds: ionic and covalent.

• However, most bonds are neither fully ionic nor fully covalent but

are somewhere between the two extremes.

• Called polar covalent bonds, which means the bonding electrons

are attracted more strongly by one atom than the other so that the

electron distribution between atoms is not symmetrical.

The symbol δ means partial charge, either partial positive for the

electron-poor atom or partial negative for the electron rich atom.

• Bond polarity is due to differences in

electronegativity (EN).

• EN: the intrinsic ability of an atom to attract

the shared electrons in a covalent bond.

• Fluorine is the most electronegative element

(EN = 4.0) and cesium is the least

electronegative element (EN = 0.7).

• Inductive effect: the shifting of electrons in a

σ bond in response to the electronegativity of

nearby atoms.

Dipole Moments

• Molecules as a whole are also often polar.

• The polarity caused by:

- the net sum of individual bond polarities

- lone-pair

• The measure of net molecular polarity is called the dipole moment, , expressed in debyes (D).

• 1 D = 3.336 x 10-30 coulomb meters (C . m)

= Q x r, where

Q = magnitude of the charge at either end of the molecular dipoles r = distance between the charges

Formal Charges

Resonance

• The true structure is intermediate between the two, they are called

resonance forms.

• The only difference between resonance forms:

- the placement of the Π bond.

- nonbonding valence electrons.

ORGANIC REACTIONS

Learning Objectives

1. To classify organic reactions into different kinds and mechanism.

2. To identify nucleophile and electrophile.

3. To discuss electrophilic addition reaction.

Chemical Reactions

• All reactions (in lab or living organisms), follow the same “rules”.

• Reactions in living organisms look more complex, with the

involvement of enzymes.

• The principles governing all reactions are the same.

• E.g. biosynthesis of prostaglandin H2.

Organic Reactions

ORGANIC REACTIONS

What kind? How?

- Addition

- Elimination

- Substitution

- Rearrangements

Mechanism:

- Radical

- Polar

Addition Reaction

• Two reactants add together to form a single product with no atoms left-over.

Elimination Reaction

• A single reactant split into two products, often with formation of a small molecule, e.g. water.

• The opposite of addition reaction.

Substitution Reaction

• Two reactants exchange parts to give two new products.

Rearrangement Reaction

• A single reactant undergoes a reorganization of bonds and atoms to

yield an isomeric (compounds that have the same molecular formula

but different structure) product.

Mechanisms

• Reaction mechanism: an overall description of how a reaction occurs.

• A complete mechanism include:

- what takes place at each stage of a chemical transformation.

- the rate of each steps.

- all reactants used and products formed.

Bond Breaking

Arrowheads with a “half” head (“fish-hook”) indicate homolytic and

homogenic steps (called ‘radical processes’).

Arrowheads with a complete head indicate heterolytic and

heterogenic steps (called ‘polar processes’).

Symmetrical bond-breaking

Unsymmetrical bond-breaking

Bond Making

Symmetrical

Unsymmetrical

Radical Reactions

• Not as common as polar reactions.

• Important in some industrial processes and in numerous biological pathways.

• A radical is highly reactive – contains an atom with an odd number of electrons (usually seven) in its valence shell.

• A radical can achieve a valence-shell octet through several ways:

- radical substitution reaction

- radical addition reaction

• In industry, radical substitution reaction is used for the chlorination of methane.

• The substitution reaction is the first step in the preparation of dichloromethane and chloroform.

• 3 types of steps: initiation, propagation, and termination.

Initiation

• Homolytic formation of a few reactive chlorine radicals by irradiation

of a small number of chlorine molecules with ultraviolet light.

Propagation

• Reaction with molecule to generate radical.

• The overall process is called a chain reaction.

A radical will collide with a methane molecule and abstract a hydrogen atom.

Cycles back and repeats the first propagation step.

Termination

• Combination of two radicals to form a stable product.

• Occur infrequently.

Polar Reactions

• Polar reactions occur because of the electrical attraction between

positive and negative centers on functional groups in molecules.

• Bond polarity is the result of an unsymmetrical electron distribution in

a bond.

• Due to the difference in electronegativity of the bonded atoms.

*Carbon is always positively polarized except when bonded to a metal

• Polar bonds can also result from the interaction of functional groups with acids or bases.

• E.g. methanol.

• In neutral methanol, the C atom is electron-poor due to the electronegative O that attracts the electrons in the C – O bond.

• On protonation of the methanol oxygen by an acid, a full positive charge on oxygen attracts the electrons in the C – O bond much more strongly.

• This makes the C much more electron-poor.

• The fundamental characteristic of all polar organic reactions is that

electron-rich sites react with electron-poor sites.

• Bonds are made when an electron-rich atom shares a pair of

electrons with an electron-poor atom, and bonds are broken when one

atom leaves with both electrons from the former bond.

• Has a negatively polarized,

electron-rich atom.

• Can form a bond by donating a

pair of electrons to a positively

polarized, electron-poor atom.

• Neutral or negatively charged.

• E.g. ammonia, water,

hydroxide ion, chloride ion.

• Has a positively polarized,

electron poor atom.

• Can form a bond by accepting

a pair of electrons from a

nucleophile.

• Neutral or positively charged.

• E.g. acids, alkyl halides,

carbonyl compounds.

Nucleophile Electrophile

Some nucleophiles and electrophiles. Electrostatic potential maps identify the nucleophilic (red) and electrophilic (blue) atoms.

Polar Reaction: Addition of HBr to Ethylene

• A typical polar reaction – addition reaction of an alkene with hydrogen bromide.

• The reaction is an example of electrophilic addition reaction.

• The reaction begins when the alkene donates a pair of electrons

from its C=C bond to HBr to for a new C – H bond plus Br-.

• One curved arrow begins at the middle of the double bond (source

of electron pair) and points to the H atom in HBr (the atom to which

the bond will form).

• A second curved arrow begins in the middle of the H-Br bond and

points to the Br, indicating that the H-Br bond breaks.

• The electrons remain with the Br atom, giving Br-.

• When one of the alkene carbon atoms bonds to the incoming

hydrogen, the other carbon atom (having lost its share of the

double-bond electrons) is left with six valence electrons, thus it is

positively charged.

• The positively charged species is called a carbocation, is an

electrophile that can accept an electron pair from nucleophilic Br-

anion.

• The curved arrow shows the electron pair movement from Br- to the

positively charged carbon.

• This will form a C-Br bond and yield the addition product.

Using Curved Arrows

• Curved arrows are a way to keep track of changes in bonding in polar reaction.

• The arrows track electron movement.

• Electrons always move in pairs.

• Charges change during the reaction.

• One curved arrow corresponds to one step in a reaction mechanism.

Rules for Using Curved Arrows:

1. The arrow goes from the nucleophilic source (Nu: or Nu:-) to the

electrophilic sink (E or E+).

2. The nucleophilic site can be neutral or negatively charged.

3. The electrophilic site can be neutral or positively charged.

4. Octet rule must be followed.

QUESTIONS??

THANK YOU