(6M1X0) Organic chemistry
Full Notes!
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,Contents INDEX Page number
-C1. General Chemistry. Electronic Structure and Bonding 3
-C2. Acids and Bases 15
-C3. An Introduction to Organic Compounds 24
-C4. Isomers and the Arrangement of Atoms in Space 33
-C5. Alkenes, Thermodynamics and Kinetics 44
-C6. The Stereochemistry of Addition Reactions 55
-C7. The Reactions of Alkynes 72
-C8. Aromaticity and Electronic Effects 81
-C9. Substitution and Elimination Reactions of Alkyl Halides 96
-C10. Reactions of Alcohols, Ethers, Amines and Epoxides 126
-C11. Radicals 149
-C12. UV-Vis, Mass Spectroscopy and Infrared Spectroscopy 162
-C13. NMR Spectroscopy 194
-C14. Reactions of Carboxylic Acids and its Derivatives 210
-C15. Reactions of Aldehydes and Ketones 233
-C16. Reactions of Benzene and Substituted Benzene 248
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, General Chemistry. Electronic Structure and Bonding
Distribution of Mass and Charge in an Atom:
Most of the mass is concentrated in the centre of an atom, in the nucleus. The
nucleus is made up of protons and neutrons. While the electrons move around
in regions of space called orbitals. The number of protons and electrons are
the same as the number of positive charges must be equal to the negative
charges.
Proton number and nucleon number:
To calculate the number of protons, neutrons and electrons we used the
proton number and nucleon number. Nucleon Number or
mass number:
The number of protons +
A electrons.
Symbol of the element.
X
Proton Number or atomic number:
Z
The number of protons which equals
the number of electrons.
Isotopes: Are atoms of the same element with the same number of protons
and electrons but different number of neutrons. The chemical properties of an
element depend on the number of electrons in the outer electron shell. As
isotopes of the same element have the same number of electrons, they have
the same chemical properties.
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,Electronic Configuration:
Electrons are arranged in energy levels called shells. Each shell is described by
a quantum number.
The main quantum numbers are 1,2,3 and 4.
As the quantum number increases the energy of the shell increases. This is
because the lowest energy level, quantum number 1, is the closest to the
nucleus. Therefore, the shells further away from the nucleus have more
energy.
Inside the shell there are subshells: s, p, d and f.
Orbital: Is a region in space where there is a maximum probability of finding an
electron.
Each orbital can hold 2 electrons which spin in opposite directions
When electrons are placed in a set of orbitals of equal energy, they occupy
them singly and then pairing takes place.
Electrons placed in opposite direction, both negatively charged, create a spin
to reduce repulsion.
Completely filled or half-filled are more stable (reduced repulsion).
Shapes of Orbitals:
s – orbital:
Has spherical shape, it increases in size as quantum number
increases.
p – orbital:
Dumbbell shape.
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,Allocating Electrons:
Electrons are allocated in order of increasing energy. This is the order:
Chromium and Copper are exceptions, one of the electrons is promoted.
Electronegativity: Is a measure of the ability of an atom to pull electrons
toward itself.
Dipole Moments of Bonds:
There are polar and non-polar covalent bonds which depend on the
electronegativity. Polar covalent bonds have a dipole, as they have a positive
and a negative end. The size of the dipole is determined by the dipole moment.
The dipole moment is equal to the product of the size of the charge and the
distance between charges. The unit for dipole moment is the Debye, D.
𝐷𝑖𝑝𝑜𝑙𝑒 𝑀𝑜𝑚𝑒𝑛𝑡 = 𝑆𝑖𝑧𝑒 𝑜𝑓 𝐶ℎ𝑎𝑟𝑔𝑒 × 𝐷𝑖𝑠𝑡𝑎𝑛𝑐𝑒 𝑏𝑒𝑡𝑤𝑒𝑒𝑛 𝑐ℎ𝑎𝑟𝑔𝑒𝑠
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,Lewis Structure:
First count the total amount of valence electrons you have. Then, draw single
bonds with neighbouring atoms to get a starting structure. Then, count the
remaining valence electrons you have. Place the remaining electrons by using
the octet rule and seeing if there are missing bonds and lone pairs. Finally,
check the formal charge. It is the difference between the number of valence
electrons an atom has when it is not bonded to other atoms and the number it
“owns” when it is bonded.
Sigma Bonds:
They are formed by the overlap of 2 s-orbitals. The electrons in the bond are
symmetrically distributed. As the two orbitals begin to overlap, energy is
released because the electron in each atom is attracted to its own nucleus and
to the nucleus of the other atom. The more the orbitals overlap, the more the
energy decreases, until the atoms are so close that their positively charged
nuclei begin to repel each other. This repulsion causes a large increase in
energy.
Minimum energy is achieved at a certain distance between nucleus which is
known as bond length.
When the bond forms it releases energy. The same amount of energy is
required to break the bond, which is known as bond dissociation energy.
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,Bonding and Antibonding Molecular Orbitals:
Orbitals are conserved. The number of molecular orbitals formed must equal
the number of atomic orbitals combined. When two atomic orbitals overlap in
order to form a covalent bond, two molecular orbitals are formed one lower in
energy and one higher in energy than the atomic orbitals. Orbitals can combine
in a constructive way (Bonding) or in a destructive way (antibonding). An
antibonding orbital is indicated by an asterisk.
The bonding molecular orbital has lower energy and is more stable due to the
higher electron density between the nuclei. This is because in bonding orbitals
the electrons are found in the centre of both overlapped orbitals, where they
can more easily attract both nuclei simultaneously. The electrons in this case
assist in bonding.
Whereas, in antibonding orbitals electrons are found anywhere except
between the nuclei. This is because a nodal plane lies between the nuclei.
Therefore, these electrons detract from bonding.
Hence, in covalent bonds the electrons occupy the bonding orbital which is the
lowest energy orbital. Here, they are attracted to the positively charged nuclei
and this electrostatic force of attraction is what gives the covalent bond its
strength. Therefore, the strength of the covalent bond increases as the orbital
overlap increases.
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,Forming Pi-bonds:
When two p atomic orbitals overlap, the side of one orbital overlaps the side of
the other. The side-to-side overlap of two parallel p orbitals forms a pi bond.
The side-to-side overlap of two in-phase p atomic orbitals is a constructive
overlap and forms a p bonding molecular orbital, whereas the side-to-side
overlap of two out-of-phase p orbitals is a destructive overlap and forms a p*
antibonding molecular orbital. The p bonding molecular orbital has a nodal
plane that passes through both nuclei. The pi* antibonding molecular orbital
has two nodal planes.
Hybridization:
sp3 Orbitals:
This form when one carbon hybridizes an s orbital with 3 p orbitals. This forms
4 sp3 orbitals which are degenerate. They all have the same energy. The sp3
orbitals are more stable than a p orbital and less stable than an s orbital.
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,Bonding in Methane:
The carbon is bonded to 4 other atoms. The four sp3 orbitals adopt a spatial
arrangement that keeps them as far away from each other as possible. They do
this because electrons repel each other. Therefore, they form a tetrahedral
shape and hence, all the angles are 109.5 degrees.
E.g. Methane
The C-C bonds are formed by the overlap of 2 sp3 orbitals. The C-H bonds are
formed by the overlap of a sp3 and an s orbital.
Also, methane is non-polar due to the similar electronegativity of the carbons
and hydrogens.
Bonding in Ethane:
The carbons are bonded to 4 other atoms. The carbons use 4 sp3 orbitals to
bond. The C-C bond is formed by the overlap of 2 sp3 orbitals. The C-H bonds
are formed by the overlap of a sp3 and an s orbital as in methane. Hence, each
carbon in ethane is tetrahedral and its bond angles are 109.5 degrees.
Also, ethane is a non-polar molecule.
sp2 hybridization:
This forms when one s orbital combined with 2 p orbitals. Hence, each carbon
hybridizes 1 s orbital and 2 p orbitals, to form 3 sp2 degenerate orbitals and
one unhybridized p orbital is also left. Double bonds are formed by a sigma
bond and a pi bond. The sigma bond is formed by the overlap of 2 sp2 orbitals
from each carbon and the pi bond is formed by the side-to-side overlap of 2 p
orbitals. Hence, sp2 hybridization is done to form double bonds.
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, Bonding in Ethene:
The carbons are bonded to 3 other atoms. The carbons form 3 sp2 degenerate
orbitals and one unhybridized p orbital is left. To minimize electron repulsion,
the 3 sp2 orbitals need to get as far from each other as possible and hence they
form a trigonal planar structure, with bond angles of 120 degrees. The
unhybridized p orbital is perpendicular to the plane defined by the axes of the
sp2 orbitals.
The double bond in ethene contains a sigma bond, which is formed by the
overlap of 2 sp2 orbitals. The other bond in the double bond is a pi bonds,
which is formed by the side-to-side overlap of 2 p orbitals.
Also, the C-H bonds in ethene are formed by the overlap of a sp2 orbital and an
s orbital.
Shape of one carbon.
Shape of full molecule of ethene.
Ethene is a nonpolar molecule, with a slight accumulation of negative charge
above the carbons. All 6 atoms of ethene including the ones bonded to the
carbon with the double bond lie on the same plane.
For example, here all atoms with an asterisk lie
on the same plane. This is because they are
bonded to the C=C double bond.
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