(6A2X0) Introduction to
Chemistry & Chemical
Technology
Full Notes on Chemistry Part!
1
,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 34
-C5. Alkenes, Thermodynamics and Kinetics 52
-C6. The Stereochemistry of Addition Reactions 62
-C7. The Reactions of Alkynes 72
-C8. Aromaticity and Electronic Effects 81
-C9. Substitution and Elimination Reactions of Alkyl Halides 101
2
, 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
Z
number:
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.
3
,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.
6
,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.
7
,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 p*
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.
8
,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 sp 2 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.
9
, 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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