Atomic Structure
Electron Configuration
Atomic Structure
Electron Configuration
z
z
z
3d
E
3p
3rd SHELL
3s
2p
2nd SHELL
x
x
y
y
2px
x
y
x
y
2py
2pz
z
z
2s
2p
2nd SHELL
x
y
z
2s
1st SHELL
x
y
1s
2s
Atomic Structure
Ground State Configuration
Lowest energy orbitals fill first
Atomic Structure
Ground State Configuration
Eg. Sodium. 11 electrons
(Aufbau principle)
3s
E
Each
orbital can contain up to two
electrons
2px,y,z
2s
• electrons have opposite spins
(Pauli exclusion principle)
Orbitals
of equal energy fill evenly
(Hunds rule)
1s
Na: 1s2 2s2 2p6 3s1
1
Covalent Bonding
Remember:
Covalent Bonding
Sharing
of electrons
• Eg. Carbon is 1s2 2s2 2p2 (half full)
H
+ 4xH
C
A pair of bonding electrons is represented as a
straight line
H C H
H
H
+ 2xH
O
H
B
Na Mg
Cr Mn
Cu Zn
Pd
Lewis
structure
Kekule
structure
C N O F
Al Si P S Cl
Ti
K
H
H C H
H O H
Li
H
H C H
H
Se Br
Sn
I
Hg
Os
Alkynes
Organic Chemists Shorthand
Hydrogens are ignored and carbons are corners and terminals
H
H
H
H
H
H
H
H
H
H
H
H
H
H
H
H
H
H
H
H
H
H
Quick and convenient
2
Polar Covalent Bonds
Polar Covalent Bonds
continuum
of possibilities between
ionic and covalent based on ‘degree of
sharing’
X
X
δ+
δ-
+
_
X
Y
X
Y
Electronegativity
H
Li
B
Na Mg
Cr Mn
Ti
K
Cu Zn
Pd
Polar Covalent Bonds
Examples
δ+
H
C
MgBr
δ-
H
H
H
H
H
K=0.8 , H=2.1 , C=2.5 , O=3.5 , F=4.0
Polar Covalent Bonds
Dipole Moment
overall polarity of a molecule
• sum of all individual polarities
Example
δ−
Hδ+
δ
Cl
H
C
Eg.
−
δδ+
C
I
ionic
polar
covalent
δ = ‘just a little bit’
“partial charge”, not a “full charge”
H
Se Br
Sn
Hg
Os
covalent
C N O F
Al Si P S Cl
δ
O
Hδ
+
N
δH +
Hδ
+
Cl
+
−
δ
δ+
H
−
δ
Cl
δ− Cl
C
Cl δ−
H
H
3
Intermolecular Bonding
Intermolecular Bonding
Dipole-dipole
Molecules with dipole moment
Attraction of permanent partial charges
MEDIUM strength
Hydrogen Bonding
Strong form of dipole-dipole interaction
Require H bonded to highly
electronegative atom (N, O or Halogen)
• Eg Propanone (acetone)
• Eg. Ethanol
δ−
δ−
+
δ
δ−
+
O
O
δ
O
δ−
+
+
Hδ
H+
O
δ
δ
Intermolecular Bonding
Structures in 3D!!!
Van der Waals forces
Non polar molecules
‘Instantaneous’ dipole arises due to
electrons random movements in atoms
δ H
WEAK
GAS
Alkynes as Pharmaceuticals
anti-cancer compounds
COOH
+
C
+
δ
C
H
OH
O
HN
O
δ−
H
−
δ
H
H
+
δ
H
δ−
C H
H
H
H
one split second
O
LIQUID
LIQUID
• Eg. Methane
−
δ
+
H
δ
OCH3
H
H
Dynemicin A
−
δ
H
C
H
H
OH
O
OH
next split second
4
Structure
Structure
V.S.E.P.R. – Valence Shell Electron Pair
Repulsion
– Valence Shell Electron
Pair Repulsion
V.S.E.P.R.
• Useful for predicting structure
• Electron pairs repel
• Shape reflects this. eg METHANE
• eg AMMONIA
H
H
C
H
3s
E
2px,y,z
2s
109.5º
H
tetrahedral
H
H
107o
VSEPR predicts 109.5º
Structure
– Valence Shell Electron
Pair Repulsion
lone
• eg WATER
– Valence Shell Electron
Pair Repulsion
V.S.E.P.R.
• Multiple bonds
pairs
121o
H
3s
C
O
H
105
H
VSEPR predicts 109.5º
118
H
C
C
H
C
H
H
C
H
o
H
H
o
180 o
H
C
H
1s
Oxygen
H
Nitrogen
V.S.E.P.R.
2px,y,z
2s
N
1s
Structure
E
lone
pair
C
H
C
H
VSEPR predicts 120o
5
Covalent Bonding
Combining Atomic orbitals
Sharing of electrons requires orbital
overlap
singly occupied atomic orbitals of
individual atoms overlap
2p
Hybridisation
Hybrid Orbitals
Carbon ‘re-organises’ its orbitals for
covalent bonding
New orbitals derived from a mathematic
combination of s and p atomic orbitals
Why does it bother?
1s
Hybridisation
Electron
2p
x
configuration of carbon:
z
y
1s
C
carbon
‘promotes’ a 2s electron to the
2pz orbital
2p
2s
H
Hybridisation
H
Remember:
Only SINGLY
occupied orbitals can
overlap
2s
1s
x
y
z
H H
C
H
C
HH H H
happy H
Happy ☺
Not quite this simple!
6
Hybridisation
one
s and three p orbitals hybridise to
give four sp3 orbitals
Hybridisation
one s and three p orbitals hybridise to give
four sp3 orbitals
+
Hydridisation
Why?
sp3 orbitals allow better overlap and form
stronger bonds
(stronger bonds = more stable compounds)
Structure
sp3 hybridised carbon based molecules are
TETRAHEDRAL.
New single bonds are called sigma bonds
Bond angles 109.5º
H
109.5
the concept of hybridisation also explains
the structure of carbon based molecules
109.5
H
C
H
H
Solid wedge
(towards you)
sp3 hybridised
Dashed wedge
(away from you)
7
Structure
Structure
Ethane
sp3 hybridised
Ethane
sp3 hybridised
• four sigma bonds for each carbon
Hybridisation
sp2 hybridisation
One s orbital and
two p orbitals
hybridise to give
three sp2 orbitals
H
H
C
H
H
C
H
H
Structure
sp2 hybridisation
Bonding
s
px,y
pz
• Each carbon has three sigma (σ) bonds
• And one pi (π) bond
3 sp2 hybridised
Trigonal planar
8
Structure
sp2 hybridisation
Ethene
Structure
Hybridisation
The double bond itself = one σ and one π bond
H
• 5 sigma (σ) bonds
• one pi (π) bond
H
C
C
H
H
Flat
H
H
C
C
H
H
Alkenes
Alkenes
Introduction
Hydrocarbons containing one or more
C-C double bond
General formula: CnH2n
Degree of Unsaturation
Alkenes are said to be ‘unsaturated’
• double bonds = unsaturation
• rings also = unsaturation
H
H
H
H
H
C
C
H
C
H H
C
C
H
H
C
C
H H
H
H
H
C
H H
H H
Propene
Cyclopentene
H
"Saturated"
CnH2n+2
C
C
H
H
H
H
H
C
C
H H
H
C
H
H
H
H
C
C
C
H
H
C
C
H
H
H
H
"Unsaturated"
CnH2n
9
Alkenes
Alkenes
Degree of Unsaturation
Degree of unsaturation gives the
number of double bonds and rings
• Refered to as the number of ‘double bond
equivalents’ (or DBEs)
Stereoisomers
Rotation about C-C double bond would
require breaking the π bond
Does not occur under normal
conditions
O
rotate
Myrcen
(oil of bay)
α -Pinene
(Turpentine)
Carvone
(spermint oil)
Structure
Stability
Cis-Trans Isomerism
Restricted rotation gives rise to
cis/trans isomerism
• same side = cis (Z)
• opposide sides = trans (E)
H 3C
CH 3
C
cis
H 3C
CH 3
H
H 3C
H 3C
C
H
C
H
C
H
CH 3
H
C
H
(E) form has steric strain minimized
∴ more stable
C
C
H
Trans
C
H
CH 3
trans
cis
trans
10
Hybridisation
sp hybridisation
One s orbital and
one p orbital
hybridise to give
two sp orbitals
Structure
px
s
py,z
sp hybridisation
Bonding
• Each C has two sigma (σ) bonds
• two pi (π) bonds
2 sp hybridised
Structure
sp hybridisation
Ethyne
• two pi (π) bonds
Hybridisation and structure
Hybridisation ‘reshapes’ atomic orbitals for
better bonding
sp3 (4σ bonds)
H C
C H
Linear
• Tetrahedral
• Eg methane, ethane
sp2 (3σ +1π bond)
• Flat, planar
• Eg. Ethene
sp (2σ +2π bond)
sp3
sp3
sp2
sp3
sp
sp
sp2
Cl
OH
• Linear
• Eg. ethyne
11
Reactions and Mechanisms
Compounds
Acid Base
•
•
•
•
Formal Charges
Formal charges
Lowry–Bronsted theory
Ka and pKa
Lewis theory
with an unusual number of e-
(a covalent bond is two shared electrons, each atom
“owns” one electron)
Example
Reaction Types
H
• Addition, substitution, elimination and rearrangement
Reaction Mechanisms
H
H
N
H
N "owns" 5eN has 5 valence eno charge
H
H
N
+
H
N "owns' 4eN has 5 valence e1+ charge
Acids and Bases
Acids and Bases
The Lowry-Bronsted Definition
an acid is a substance which donates a
proton (H+)
a base is a substance which accepts a proton
(H+)
Example
Acidity Constant (Ka)
acids differ in their H+ donating ability
measured based on their ability to
donate H+ to water
H-A
+
acid
H Cl
+
-:B
A:-
base
conjugate
base
O H
Cl
+
B-H
conjugate
acid
H-A + H 2O
A- + H 3O+
position of the eq. relates to acid strength
•equilibrium
favors the rhs ⇒ strong acid
•equilibrium favors the lhs ⇒ weak acid
+ H O H
12
Acids and Bases
quantified by measuring the eq. constant
+
Ka =
[H3O+][A-]
[HA]
Acid
pKa
Ka
CH3CH2OH
16.0
10-16
H2O
15.7
10-15.7
CH3COOH
4.8
10-4.8
HNO3
-1.3
HCl
-7.0
-
[H3O ] [A ]
[HA] [H2O]
+
Ka =
Relative Acid Strengths
A- + H 3O+
H-A + H2O
Keq =
Acids and Bases
-
[H3O ] [A ]
[HA]
often expressed as pKa (= -log Ka)
Acids and Bases
The Lewis Definition
a lewis acid is an electron pair acceptor.
acceptor
a lewis base is an electron pair donor.
donor
Example
101.3
increasing
acid
strength
107
Acids and Bases
Lewis acids require a unfilled low energy
orbital
No electrons,
H+ ∴ empty s orbital
F
F
Only six electrons,
∴ empty 2p orbital
B
F
Lewis bases require a lone pair of
electrons
H+ +
lewis acid
NH3
lewis base
H NH3
H
H
O
N
S
CH3
CH3
H
H
H
13