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The Basic – Bonding and Molecular Structure
Organic Chemistry and Life
1.1 The Development of Organic
Chemistry as Science
 Organic compounds: compounds that could be obtained
from living organisms
 The 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
sources
 Occur as a salts
Atomic Orbitals
Atomic Orbitals
S - orbital
p- orbital
Electrons Configuration
 Show
Orbital Diagram and Electron
Configuration
 Electrons configuration: H,
He, C, Mg
 Aufbau Principle: fill lowest
orbital first to full capacity,
then next
1.2 The structural Theory of Organic
Chemistry
 Atoms in organic compounds can form a fixed number of
bonds using their valence electrons
1.2 The structural Theory of Organic
Chemistry
 A carbon atom can use one or more of its valence electrons
to form bonds to other carbon atoms
1.3 Isomers: The Importance of
Structural Formulas
 Constitutional isomers – non identical compounds with
same molecular formula
 Do not necessary share similar properties
1.4 Ionic Bonds
Occurs 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
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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
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Metallic elements – low
electronegativities
Halogens and other
elements in upper righthand corner of periodic
table – high
electronegativity
Polarity
 Polar covalent bonds –
the bonding electrons are
attracted somewhat more
strongly by one atom in a
bond
 Electrons are not completely
transferred
 More electronegative atom:
δ- . (δ represents the partial
negative charge formed)
 Less electronegative atom: δ+
Lewis Structures
 represents 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
Rules
 Duet Rule: sharing of 2 electrons
 E.g H2
 H:H
 Octet Rule: sharing of 8 electrons
 Carbon, oxygen, nitrogen and fluorine always obey this rule in
a stable molecule
 E.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
Bonds
 Use 6N + 2 Rule
 N = number of atoms other than Hydrogen
 If
 Total valence – (6N + 2) = 2

 1 double bond
 Total valance e- - (6N + 2) = 4

 two double bonds or 1 triple bond
Examples
 Write the Lewis structure of CH3F, ClO3-, F2
1.7 Formal Charges
 Difference between the number of outer-shell electrons
“owned” by a neutral free atom and the same atom in a
compound
Examples
 Determine the formal charge for each atom in the following
molecules
 NH4+
 NO2 CO32-
Resonance
 Whenever a molecule or ion can be represented by two or
more Lewis structures that differ only in the position of the
electrons
 None of these resonance structures will be a correct
representation for the molecule or ion
 The actual molecule or ion will be better represented by a
hybrid or hypothetical structures
 Represented by a double headed arrows (
)
Examples
Resonance - stabilization
 The more covalent bonds a structure has, the more stable it
is
Resonance-stabilization
 Structure in which all the atoms have a complete valence
shell of electrons are especially
Resonance stabilization
 Charge separation decrease stabilization
 Resonance contributors with negative charge on highly
electronegative atoms are stable ones with negative charge
on less or nonelectronegative atoms
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.
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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 3D 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)
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Intensity of radiant energy is
proportional to amplitude
Wave function
1.10 Atomic orbital
 Heisenberg 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 Orbitals
 Two types of atomic of atomic orbitals are combined as
they come close to each other
 Hybridization: blending
combination of atomic orbitals to form new
orbital
 Carbon has three possible molecular orbitals
 sp3
sp2
sp
Orbitals repsonsible for creating the
covalent bonds
 2 special names for covalent bonds of organic molecules
Sigma (σ) bond
Pi (π) bond
Created when “head on”
overlap occurs of orbitals
Created when “side on”
overlap occurs of orbitals
3
sp
molecular orbitals
 sp3 orbitals responsible for creating all “single bonds” of
all organic molecules
  alkanes
Examples
sp2 molecular orbitals
 All sp2 molecular orbitals responsible for creating all double
bonds in organic molecules
  alkenes
Examples
1.13B – Cis –Trans Isomerism
 Which of the following alkene can exist as cis-trans isomers?
Write their structure
sp molecular orbitals
 All sp orbitals responsible for creating all triple bonds of
organic molecules
  alkynes
Examples
Examples
 Draw a bonding picture for the following molecule, showing
all π, σ – bonds using σ-framework and π-framework
Molecular Orbitals
 Two types:
 Bonding molecular
orbitals
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Contains both electrons
in the lowest energry
state or ground state
Formed by intereaction
of orbitals with same
phase signs
Increases the propability
Molecular orbitals
 Antimolecular orbitals
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Contains no electrons in
the ground state
Formed by intereaction
of orbitals with opposite
phase signs
Result with nodes
Molecular orbitals
Shape of Molecules
 VSEPR Theory
 Valence shell electron pair repulsion
 Bond 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
Shape of Molecules
Molecular shapes
 VSEPR method can be used to predict the shapes of
molecules containing multiple bonds
 Assume that all electrons of a multiple bond act as one unit
Examples
 Use VSEPR theory to predict the geometry of each of the
following molecules and ions
 SiF4
 BeF2
1.17 Representation of Structural
formulas
 Structual formula for propyl alcohol
Dash Structure
 Atoms are joined by single bonds can rotate relatively freely
with respect to one another
Dash Formula - Isomerism
Condensed Structural Formulas
 All hydrogen atoms are written immediately after the carbon
that they’re attached
Bond-Line Formulas
 Hydrogen and carbon atoms will not appear in the formula
 Each end of the line represents carbon atom
Examples
 For each of the following, write a bond line formula