The Basic – Bonding and Molecular Structure

Organic Chemistry and Life

1.1 The Development of Organic Chemistry as ScienceOrganic compounds: compounds that

could be obtained from living organismsThe scientific study of the structure,

properties, composition, reactions, and preparation (by synthesis or by other means) of chemical compounds that contain carbon

Inorganic compounds: those came from non-living sourcesOccur as a salts

Atomic Orbitals

Atomic OrbitalsS - orbital p- orbital

Electrons ConfigurationShow

Orbital Diagram and Electron ConfigurationElectrons

configuration: H, He, C, Mg

Aufbau Principle: fill lowest orbital first to full capacity, then next

1.2 The structural Theory of Organic ChemistryAtoms in organic compounds can form a fixed

number of bonds using their valence electrons

C

Carbon atoms are tetravalent

O

Oxygen atomsare divalent

H Cl

Hydrogen and halogenare monovalent

1.2 The structural Theory of Organic ChemistryA carbon atom can use one or more of its

valence electrons to form bonds to other carbon atoms

C C C C C C

single bond double bond triple bond

1.3 Isomers: The Importance of Structural FormulasConstitutional isomers – non identical

compounds with same molecular formulaDo not necessary share similar properties

H C C O H

H

H

H

H

ethyl alcohol

H C O C

H

H

H

H

H

Dimethyl ether

1.4 Ionic BondsOccurs in ionic compound

Results from transferring electron

Created a strong attraction among the closely pack compound

Covalent Bonding Formation of a covalent Bond

Two atoms come close together, and electrostatic interactions begin to develop

Two nuclei repel each other; electrons repel each other Each nucleus attracts to electrons; electrons attract

both nuclei Attractive forces > repulsive forces; then

covalent bond is formed

Electronegativity Electronegativity

(EN): the ability of an atom in a molecule to attract the shared electron in a bond

Metallic elements – low electronegativities

Halogens and other elements in upper right-hand corner of periodic table – high electronegativity

PolarityPolar covalent bonds

– the bonding electrons are attracted somewhat more strongly by one atom in a bondElectrons are not

completely transferredMore electronegative

atom: δ- . (δ represents the partial negative charge formed)

Less electronegative atom: δ+

Lewis Structuresrepresents how an atom’s valence electrons

are distributed in a molecule Show the bonding involves (the maximum

bonds can be made)Try to achieve the noble gas configuration

RulesDuet Rule: sharing of 2 electrons

E.g H2 H : H

Octet Rule: sharing of 8 electronsCarbon, oxygen, nitrogen and fluorine always

obey this rule in a stable moleculeE.g F2, O2

Bonding pair: two of which are shared with other atoms

Lone pair or nonbonding pair: those that are not used for bonding

1 Lewis Structures of Molecules with Multiple BondsUse 6N + 2 Rule

N = number of atoms other than HydrogenIf

Total valence – (6N + 2) = 2 1 double bond

Total valance e- - (6N + 2) = 4 two double bonds or 1 triple bond

ExamplesWrite the Lewis structure of CH3F, ClO3

-, F2

1.7 Formal ChargesDifference between the number of outer-shell

electrons “owned” by a neutral free atom and the same atom in a compound

Formal charge = group # - unshared e- - share e-12

ExamplesDetermine the formal charge for each atom

in the following moleculesNH4

+

NO2-

CO32-

ResonanceWhenever a molecule or ion can be

represented by two or more Lewis structures that differ only in the position of the electronsNone of these resonance structures will be a

correct representation for the molecule or ionThe actual molecule or ion will be better

represented by a hybrid or hypothetical structures

Represented by a double headed arrows ( )

O

C

O O

23

-

23

-23

-

hybrid

O

C

O O

O

C

O O

O

C

O O

contributing resonance structures

Resonance - stabilizationThe more covalent bonds a structure has, the

more stable it is

H2C CH

CH

CH2 H2C CH

CH

CH2

H2C CH

CH

CH2

Resonance-stabilizationStructure in which all the atoms have a

complete valence shell of electrons are especially H2C O CH3 H2C O CH3

Here this carbonatom has only sixelectrons

Here the carbon atomhas eight electrons

Resonance stabilizationCharge separation decrease stabilizationResonance contributors with negative charge

on highly electronegative atoms are stable ones with negative charge on less or nonelectronegative atoms

H2C CH

Cl H2C CH

Cl

1.9 Quantum Mechanisms and Atomic Structure Schröndinger’s quantum mechanical model of

atomic structure is frame in the form of a wave equation; describe the motion of ordinary waves in fluids.

i. Wave functions or orbitals (Greek, psi , the mathematical tool that quantum mechanic uses to describe any physical system

ii. 2 gives the probability of finding an electron within a given region in space

iii. Contains information about an electron’s position in 3-D space

defines a volume of space around the nucleus where there is a high probability of finding an electron

say nothing about the electron’s path or movement

11.2 Electromagnetic Radiation Radiation energy – has

wavelike properties Frequency (υ, Greek nu) –

the number of peaks (maxima) that pass by a fixed point per unit time (s-1 or Hz)

Wavelength (λ, Greek lambda) – the length from one wave maximum to the next

Amplitude – the height measured from the middle point between peak and trough (maximum and minimum)

Intensity of radiant energy is proportional to amplitude

Constructive interfence of wave Destructive interference of wave

1.10 Atomic orbitalHeisenberg Uncertainty

Principle – both the position (Δx) and the momentum (Δmv) of an electron cannot be known beyond a certain level of precision

1. (Δx) (Δmv) > h 4π

2. Cannot know both the position and the momentum of an electron with a high degree of certainty

Molecular OrbitalsTwo types of atomic of atomic orbitals are

combined as they come close to each otherHybridization: blending combination of atomic orbitals

to form new orbitalCarbon has three possible molecular orbitals

sp3 sp2 sp

Orbitals repsonsible for creating the covalent bonds2 special names for covalent bonds of organic

moleculesSigma (σ) bond Pi (π) bond

Created when “head on” overlap occurs of orbitals

Created when “side on” overlap occurs of orbitals

sp3 orbitals responsible for creating all “single bonds” of all organic molecules alkanes

s sp3 sp3sp3

take pure atomic orbitalsand mix them

gives 4 molecular orbitals called sp3 orbitalsall equal in size, energy and shape

p sp3

Examples

H C

H

H

H

C

H

H

H

H

sp3C - sH

C C

H

H

H

H

H

H

sp2 molecular orbitals

All sp2 molecular orbitals responsible for creating all double bonds in organic molecules

alkenes

s sp2sp2 unhybridized orbital

p

hybridization

take pure s + 2p"mix them"

gives 3 molecular orbitals called sp2and lef t over p atomic orbital (unhybridized)sp2 orbitals equal in size, energy and shape

sp2p

C C

H

H

H

H

C C

H

H H

H

sp2C - sp2

C

sp2C - sH

C C

H

H

H

H

sp2C - sp2

C

framework

framework

1.13B – Cis –Trans IsomerismWhich of the following alkene can exist as cis-

trans isomers? Write their structure

H2C CHCH2CH3a.

b.H2C C(CH3)2

c. CH3CH CHCH3d. CH3CH2CH CHCl

sp molecular orbitals

All sp orbitals responsible for creating all triple bonds of organic molecules

alkynes

s sp sp unhybridized orbitalp

hybridization

p

take 1s + 1pmix them

give 2 molecular orbitals called sp and 2 unhybridized p atomic orbitals

H C C H

C CH H

spC - spC

spC - sH

H C C H

spC - spC

ExamplesDraw a bonding picture for the following

molecule, showing all π, σ – bonds using σ-framework and π-framework

C CH

H

H

C C H

Molecular OrbitalsTwo types:

Bonding molecular orbitals Contains both

electrons in the lowest energry state or ground state

Formed by intereaction of orbitals with same phase signs

Increases the propability

Molecular orbitalsAntimolecular

orbitals Contains no

electrons in the ground state

Formed by intereaction of orbitals with opposite phase signs

Result with nodes

Molecular orbitals

Shape of MoleculesVSEPR Theory

Valence shell electron pair repulsionBond angles and geometry

Steric number = # bond to - # lone pairs central atom to central atom

Rules: 1- Carbon will always be the central atom 2 – Double bond; triple bonds will count

as 1 bond

H C

H

H

H

109.5o

tetrahedral

C CC

HH

OH

H

H

H

120o 120o

trigonal planar

C CH H

180o

linear

O

H H

V-shaped or bent

109.5o

Molecular shapesVSEPR method can be used to predict the

shapes of molecules containing multiple bondsAssume that all electrons of a multiple bond

act as one unit

ExamplesUse VSEPR theory to predict the geometry of

each of the following molecules and ionsSiF4

BeF2

1.17 Representation of Structural formulasStructual formula for propyl alcohol

O H

H

H

H

H

H

H

H

Dash formulaor structural formula

CH3CH2CH2OH

condensed formula

OH

Bond-line formulaor stick f igure

Dash StructureAtoms are joined by single bonds can rotate

relatively freely with respect to one another

C

C

C

H

O

H

HH H

H H

H

C

C

C

O

H

H

H

H H

H H

H

C

C

C

OH

HH

H H

H H

H

12

3

Equivalent dash formula for propyl alcohol

C C C O

H

H

H H

H

H

H

C C C H

H

H

H H H

H

H

O

H

constitutional isomers

Condensed Structural FormulasAll hydrogen atoms are written immediately

after the carbon that they’re attached

H C C C H

H

H

H

Cl

H

H

H C C C C

H

H

H

H

H

H

H

H

O H

CH3CH2CH3

Bond-Line FormulasHydrogen and carbon atoms will not appear

in the formulaEach end of the line represents carbon atom

C C C O

H

H

H H

H

H

H

H

OH

H C C C H

H

H

H

Cl

H

H

Cl

bond-line

structural

ExamplesFor each of the following, write a bond line

formula(CH3)2NCH2CH3

H2C

H2C CH2

a. b.

c.

H3C

HC

CH

H2C

Cl

d.

CH3CHC CH

F