STRUCTURE AND REACTIONS OF ORGANIC MOLECULES – MODULE 1
Lecture 1
Learning Objectives
To know the shapes of the modern quantum mechanical orbitals
To be able to write down an electron configuration for an atom
To be able to draw a Lewis structure of a molecule
To understand resonance
Extra…
Atoms of the same element are alike in mass and size
Atoms of different elements have different masses and sizes
Chemical compounds are formed by the union of two or more atoms of difference elements.
Atoms combine to form compounds in simple ratios (1:1, 1:2), two elements maybe combined in
different ratios to form multiple compounds. BF 3 has an incomplete octet around the B atom
which changes its properties.
Atomic Structure and Bonding
Electrons can only circulate the nucleus at a certain distance, with certain allowed energies.
The lowest energy arrangement is closest to the nucleus, called the ground state. Electrons can
only exist at exact energy levels above this, not in-between in an excited state.
Electrons exist in orbitals- regions in space surrounding the nucleus where there is a high
probability of the electron existing.
Quantum Mechanical Atom
Electrons can be described as waves, the wave-function provides information about how the
electron wave is distributed in the atom- shape of electron distribution. Each wave-function has a
particular energy value and is characterised by three quantum numbers- these describe the different
types of orbitals in which an electron can exist.
Principle Quantum number, n: n=1,2,3 Describes the energy of the electron in an orbital
Shape (Azimuthal) Quantum number, l: s (sharp), p (principle), d (diffuse), f (fundamental.) l
designates the shape of an orbital.
Magnetic Quantum number, m: Indicates the number and orientation of a set of orbitals
s orbital come from groups of 1
p orbital come from groups of 3
d orbitals come from groups of 5
p and d orbitals are directional
The spin quantum number: Determines the at the electron in the orbital can have one of two
orientations
The bigger the orbital the higher the energy
Electron Configurations
Lowest energy (smallest) to highest energy (biggest):
1s 2s 2p 3s 3p 4s 3d 4p 5s etc…
Into these orbitals we can put this number of electrons:
1s 2s 2p 3s 3p 4s 3d 4p 5s etc…
2e- 2e- 6e- 2e- 6e- 2e- 10e- 6e- 2e-
, Lewis Structures (Chapter 8)
When individual atoms join together a molecule is formed. When atoms interact, they do so via their
outermost orbitals- valence electrons
When atoms join together the valance electrons of each atom are reorganized.
Ultimately the attractive forces win over the repulsion forces, allowing the strong and stable bond to
be formed.
Stable octet = 8 valance electrons for first row elements
Where possible maintain this octet.
Heavier elements can have an expanded octet or valence shell.
The “negative” sign on the corner of the molecule means gain an electron
The “positive” sign on the corner of the molecule means lose an electron
RULES:
1. Count valence electrons for each atom
2. Assemble bonding framework using single covalent bonds
3. Place three non-bonding pairs of electrons on each outer atom except H
4. Assign remaining valence electrons to inner atoms
5. Minimise formal changes on all atoms
Formal Charge = Valance Electrons on Free Atom - Electrons Assigned in Lewis Structure
We can minimise formal charge by converting lone electron pairs into bonding pairs.
Often the most electro-negative element is on the outside, or the first letter in the central atom.
Double bonds are shorter than single bonds
Resonance
A way of describing delocalized electrons within certain molecules where the bonding cannot be
expressed by one single Lewis structure.
Resonance is an important way to stabilise molecules, minimising the formal charge on an atom can
make it less reactive.
STRUCTURES:
Carboxylic acids can easily by deprotonated to form carboxylate anions.
Amides
HOMEWORK
PROBLEMS 6.52(A-D), 6.54, 6.56
PROBLEMS 8.6, 8.7, 8.38, 8.51, 8.53
,Lecture 2
Learning Objectives
To be able to use VSEPR theory to predict molecular shape
To have an understanding of valence bond theory
To have an understanding of why orbitals are hybridised
To be able to draw hybridised orbitals
TEXTBOOK : Chapter 9, Sections 9.1-9.6, Chapter 22, section 22.4
Molecular Shape
1. What is the three-dimensional shape of the molecule
2. What sorts of orbitals are the electrons in when the molecule is formed
VSEPR Theory (Valence Shell Electron Pair Repulsion)
Electrons have strong repulsion force between each other,
so the shape a molecule takes will be one where the
repulsion is minimised. The shape of the molecule is
determined by the arrangement of bonds and lone
electrons.
VSEPR makes no distinction between electron pairs in
single, double, or triple bonds.
Lone pairs repel more than bonded pairs because no
bonding pairs experience less nuclear attraction so their
electron domain can spread out more and thus repel more.
Valence Bond Theory
Valence bond theory states that in making a bond between two atoms, the atomic orbitals must be
half-filled so that the repulsion force may be minimised by the opposite spin of electron in each
orbital. These orbitals overlap to make a new orbital in which the pair of electrons exist.
o Sigma (σ) orbital
• This orbital is spherically symmetrical
about the internuclear axis with the ends
looking like a 'S' orbital.
When the valence electrons of an atom are in two different types of atomic orbitals (e.g. 2s 2p 2) a
hybrid orbital needs to be formed.
Only the 2p orbitals will be able to 'take in' another electron, only 2
bonds can be formed. In order to make bonds for this molecule you need
for half-filled orbitals.
, Here four bonds can be made as there is four half-filled orbitals to 'take in'
an electron. However, one of these bonds is at a different energy level
(and experimental evidence shows all four bonds are the same).
Four atomic orbital →Four hybrid orbitals
*CANNOT CREATE OR DESTROY ORBITALS*
Here, we have for half-filled orbitals all at the same
energy level (so the bonds will be equal and lie at a 109 o
angle to each other).
This can be achieved by mixing the wave-functions of the
2s and 2p orbital together to make a new type of orbital
or hybrid orbital. These hybrid orbitals can then be
overlapped with the orbital of another atom.
The energy level of the hybrid orbital will be an average of
the atomic orbitals involved (s goes up in energy, p goes
down in energy).
*HYBRID ORBITALS ALWAYS LEAD TO SIGMA BONDS*
p orbitals have one node
HOMEWORK
Problems 9.15, 9.16
Lecture 1
Learning Objectives
To know the shapes of the modern quantum mechanical orbitals
To be able to write down an electron configuration for an atom
To be able to draw a Lewis structure of a molecule
To understand resonance
Extra…
Atoms of the same element are alike in mass and size
Atoms of different elements have different masses and sizes
Chemical compounds are formed by the union of two or more atoms of difference elements.
Atoms combine to form compounds in simple ratios (1:1, 1:2), two elements maybe combined in
different ratios to form multiple compounds. BF 3 has an incomplete octet around the B atom
which changes its properties.
Atomic Structure and Bonding
Electrons can only circulate the nucleus at a certain distance, with certain allowed energies.
The lowest energy arrangement is closest to the nucleus, called the ground state. Electrons can
only exist at exact energy levels above this, not in-between in an excited state.
Electrons exist in orbitals- regions in space surrounding the nucleus where there is a high
probability of the electron existing.
Quantum Mechanical Atom
Electrons can be described as waves, the wave-function provides information about how the
electron wave is distributed in the atom- shape of electron distribution. Each wave-function has a
particular energy value and is characterised by three quantum numbers- these describe the different
types of orbitals in which an electron can exist.
Principle Quantum number, n: n=1,2,3 Describes the energy of the electron in an orbital
Shape (Azimuthal) Quantum number, l: s (sharp), p (principle), d (diffuse), f (fundamental.) l
designates the shape of an orbital.
Magnetic Quantum number, m: Indicates the number and orientation of a set of orbitals
s orbital come from groups of 1
p orbital come from groups of 3
d orbitals come from groups of 5
p and d orbitals are directional
The spin quantum number: Determines the at the electron in the orbital can have one of two
orientations
The bigger the orbital the higher the energy
Electron Configurations
Lowest energy (smallest) to highest energy (biggest):
1s 2s 2p 3s 3p 4s 3d 4p 5s etc…
Into these orbitals we can put this number of electrons:
1s 2s 2p 3s 3p 4s 3d 4p 5s etc…
2e- 2e- 6e- 2e- 6e- 2e- 10e- 6e- 2e-
, Lewis Structures (Chapter 8)
When individual atoms join together a molecule is formed. When atoms interact, they do so via their
outermost orbitals- valence electrons
When atoms join together the valance electrons of each atom are reorganized.
Ultimately the attractive forces win over the repulsion forces, allowing the strong and stable bond to
be formed.
Stable octet = 8 valance electrons for first row elements
Where possible maintain this octet.
Heavier elements can have an expanded octet or valence shell.
The “negative” sign on the corner of the molecule means gain an electron
The “positive” sign on the corner of the molecule means lose an electron
RULES:
1. Count valence electrons for each atom
2. Assemble bonding framework using single covalent bonds
3. Place three non-bonding pairs of electrons on each outer atom except H
4. Assign remaining valence electrons to inner atoms
5. Minimise formal changes on all atoms
Formal Charge = Valance Electrons on Free Atom - Electrons Assigned in Lewis Structure
We can minimise formal charge by converting lone electron pairs into bonding pairs.
Often the most electro-negative element is on the outside, or the first letter in the central atom.
Double bonds are shorter than single bonds
Resonance
A way of describing delocalized electrons within certain molecules where the bonding cannot be
expressed by one single Lewis structure.
Resonance is an important way to stabilise molecules, minimising the formal charge on an atom can
make it less reactive.
STRUCTURES:
Carboxylic acids can easily by deprotonated to form carboxylate anions.
Amides
HOMEWORK
PROBLEMS 6.52(A-D), 6.54, 6.56
PROBLEMS 8.6, 8.7, 8.38, 8.51, 8.53
,Lecture 2
Learning Objectives
To be able to use VSEPR theory to predict molecular shape
To have an understanding of valence bond theory
To have an understanding of why orbitals are hybridised
To be able to draw hybridised orbitals
TEXTBOOK : Chapter 9, Sections 9.1-9.6, Chapter 22, section 22.4
Molecular Shape
1. What is the three-dimensional shape of the molecule
2. What sorts of orbitals are the electrons in when the molecule is formed
VSEPR Theory (Valence Shell Electron Pair Repulsion)
Electrons have strong repulsion force between each other,
so the shape a molecule takes will be one where the
repulsion is minimised. The shape of the molecule is
determined by the arrangement of bonds and lone
electrons.
VSEPR makes no distinction between electron pairs in
single, double, or triple bonds.
Lone pairs repel more than bonded pairs because no
bonding pairs experience less nuclear attraction so their
electron domain can spread out more and thus repel more.
Valence Bond Theory
Valence bond theory states that in making a bond between two atoms, the atomic orbitals must be
half-filled so that the repulsion force may be minimised by the opposite spin of electron in each
orbital. These orbitals overlap to make a new orbital in which the pair of electrons exist.
o Sigma (σ) orbital
• This orbital is spherically symmetrical
about the internuclear axis with the ends
looking like a 'S' orbital.
When the valence electrons of an atom are in two different types of atomic orbitals (e.g. 2s 2p 2) a
hybrid orbital needs to be formed.
Only the 2p orbitals will be able to 'take in' another electron, only 2
bonds can be formed. In order to make bonds for this molecule you need
for half-filled orbitals.
, Here four bonds can be made as there is four half-filled orbitals to 'take in'
an electron. However, one of these bonds is at a different energy level
(and experimental evidence shows all four bonds are the same).
Four atomic orbital →Four hybrid orbitals
*CANNOT CREATE OR DESTROY ORBITALS*
Here, we have for half-filled orbitals all at the same
energy level (so the bonds will be equal and lie at a 109 o
angle to each other).
This can be achieved by mixing the wave-functions of the
2s and 2p orbital together to make a new type of orbital
or hybrid orbital. These hybrid orbitals can then be
overlapped with the orbital of another atom.
The energy level of the hybrid orbital will be an average of
the atomic orbitals involved (s goes up in energy, p goes
down in energy).
*HYBRID ORBITALS ALWAYS LEAD TO SIGMA BONDS*
p orbitals have one node
HOMEWORK
Problems 9.15, 9.16
