INVISIBLE MECHANICS
"don't be a normie, i guess"
INVISIBLE MECHANICS
1 .Introduction
Chemical bonding is the attractive force that holds atoms together to
form molecules, crystals, and other stable structures. It is caused by the
interaction of the electrons of the atoms involved. There are three main
types of chemical bonds: ionic, covalent, and metallic.
Atoms can achieve the octet rule by forming chemical bonds, they are
more stable when their valence electrons are arranged in certain ways,
gains stability and decreases reactivity.
Octet rule
Octet rule: Each (Lewis, Kosset) atom tores to obtain the octet state.
Contraction of Octet Slate: Central atom is electron deficient or
does not have an octet state. eg. Bex2, BX3, ALX34
Expansion of Octet State: Central atom has more than 8 electrons
due to empty d orbitals. Eg PCl3, SF6, ICl3
Exceptions
other exceptions:
Transition (Cr3+, Fe 2+)
Pseudo Inert gas configuration (Zn2+)
Odd electron species (NO, NO2)
Irrler halogen compound (IF7, Brf3)
Compounds of Xenon (Xef2, Xef4)
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1 .Introduction
Types of Bonds
Types of Bond : {bond strength decreasing order}
Ionic Bond
Covalent Bond
Coordinate Bond
Hydrogen Bond
Van der Waals Bond
Ionic Bond
Max number of electrons transferred by a metal to a non metal is
three, as in the case of AlF3
A Non directional bond
Condition for formation of ionic bond
A) process. H = +ve
B) Metal must have low ionisation energy
C) Non metal must have high electron affinity
D) Ions must have high lattice energy
E) Cation should be large with low electronegativity.
F) Anion must be small with high electronegativity.
1. Ionic compounds are hard
2. Hardness is directly proportional to electrostatic force of attraction
3. Hardness is directly proportional to charge on ion
4. Hardness is inversely proportional to ionic radius
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Ionic Bonding
Ionic compound has high value of BP, MP & boiling point
Ionic compound shows isomorphism
Conductors in fused, melted, aqueous state
Show fast ionic reaction
Don't show space isomerism
Electrical conductivity in solution or when melted. Ionic compounds
conduct electricity in solution or when melted because the ions are
free to move around and carry an electric current. However, solid
ionic compounds do not conduct electricity because the ions are fixed
in place.
Water solubility. Many ionic compounds are soluble in water because
the water molecules can solvate the ions, meaning that they surround
the ions and hold them in solution.
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Covalent Bonding
Such chemical bonds are formed by sharing of electrons between the
elements of almost same electronegativity or between the elements
having less difference in electronegativity.
e.g. Formation of O2 molecule
Covalent bonds are directional, meaning that the atoms involved in
the bond must be oriented in a specific way. This is because the
shared electrons are attracted to the nuclei of both atoms. The shape
of a molecule is determined by the arrangement of the covalent
bonds between its atoms.
Covalent compounds have a number of characteristic properties
Softness and ductility. Covalent compounds are soft because the
molecules are not tightly packed together. They are also ductile,
meaning that they can be easily drawn into wires.
Low melting and boiling points
Insolubility in water. Covalent compounds are generally insoluble
in water because they are not attracted to the polar water
molecules.
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Modern Concept of Covelent Bond
Valence Bond Theory
It explains bond formation in terms of overlapping of orbitals, e.g.
the formation of H2 molecule from two hydrogen atoms involves the
overlap of 1s-orbital of two H-atoms which are singly occupied.
Because of orbital overlap, electron density between the nuclei
increases which helps in bringing them closer.
Sigma (σ) Bond
It is the result of end to end overlapping or axial
overlapping between s-s, p -p, s-p orbitals, e.g. single bond.
The electron density accumulates between the centres of the atoms
being bonded.
Pi(π)Bond
It is formed by incomplete, sidewise or parallel overlapping of
orbitals. Double bond has one σ-bond and one π-bond. Triple bond
has two π-bonds and one σ-bond.
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Pauli and Slater’s theory
The Paul-Slater theory is a theory of valence bond theory (VBT)
that explains the hybridization of atomic orbitals. It was
developed by Linus Pauling and John C. Slater in the 1930s.
The Paul-Slater theory is based on the following principles:
Atomic orbitals hybridize to form new orbitals that are more
directional and have lower energy than the original orbitals.
The hybridization scheme is determined by the number of sigma
bonds and lone pairs of electrons around the atom.
The hybridized orbitals are arranged in a tetrahedral,
octahedral, or trigonal bipyramidal geometry, depending on the
hybridization scheme.
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Bond Length
Bond length is the distance between the nuclei of two atoms that
are bonded together. It is one of the most important properties of a
chemical bond. Bond length depends on a number of factors,
including:
The type of bond: Ionic bonds are generally longer than
covalent bonds. This is because ionic bonds are formed by the
transfer of electrons from one atom to another, while covalent
bonds are formed by the sharing of electrons.
The electronegativity of the atoms: The greater the difference
in electronegativity between two atoms, the longer the bond
length. This is because the more electronegative atom will pull
the electrons closer to its nucleus, making the bond longer.
The hybridization of the atoms: The hybridization of the atoms
can also affect bond length. Hybridized orbitals are more
directional than unhybridized orbitals, which can lead to shorter
bond lengths.
The bond order: The bond order is a measure of the number of
electrons that are shared between two atoms. The higher the
bond order, the shorter the bond length.
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Bond Energy
Bond energy is the amount of energy required to break a chemical
bond. It is a measure of the strength of a chemical bond. Bond
energy is typically expressed in kilojoules per mole (kJ/mol).
The bond energy of a chemical bond depends on a number of
factors, including:
The type of bond: Ionic bonds are generally stronger than
covalent bonds. This is because ionic bonds are formed by the
transfer of electrons from one atom to another, while covalent
bonds are formed by the sharing of electrons.
The electronegativity of the atoms: The greater the difference
in electronegativity between two atoms, the stronger the bond
energy. This is because the more electronegative atom will pull
the electrons closer to its nucleus, making the bond stronger.
The hybridization of the atoms: The hybridization of the atoms
can also affect bond energy. Hybridized orbitals are more
directional than unhybridized orbitals, which can lead to
stronger bond energies.
The bond order: The bond order is a measure of the number of
electrons that are shared between two atoms. The higher the
bond order, the stronger the bond energy.
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Bond Angle
Bond angle is the angle between two covalent bonds that originate
from the same atom. It is a measure of the three-dimensional shape
of a molecule. Bond angle is typically expressed in degrees (°).
The bond angle of a molecule depends on factors, including:
The type of bond: Ionic bonds do not have bond angles.
Covalent bonds have bond angles that are determined by the
hybridization of the atoms involved in the bond.
The hybridization of the atoms: The hybridization of the atoms
determines the arrangement of their valence orbitals.
The size of the atoms: The larger the atoms, the greater the
bond angle. This is because larger atoms have larger valence
orbitals, which require more space.
The presence of lone pairs: Lone pairs are electrons that are not
involved in bonding. Lone pairs repel bonding electrons, which
can increase the bond angle.
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Types Of Overlapping
Sigma overlap occurs when two orbitals overlap head-on, along
the internuclear axis. The overlap is symmetrical about the
internuclear axis, and the resulting bond is also symmetrical.
Sigma bonds are the strongest type of covalent bond.
Pi overlap occurs when two orbitals overlap sideways, above and
below the internuclear axis. The overlap is not symmetrical about
the internuclear axis, and the resulting bond is also not
symmetrical. Pi bonds are weaker than sigma bonds.
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Polarizisation And Fajan’s Rule
The partial covalent character of an ionic bond has been explained
by Fajan as follows:
When a cation approaches an anion, the electrons cloud of the anion
is deformed which is called polarisation. Greater the polarisation,
more is the covalent character.
Covalent character
polarising power of cation
Covalent character is inversely proportional to size of cation
Polarisability of anion
size of anion
∝
∝
Coordinate Covalent (Dative) Bond
Coordinate Covalent (Dative) Bond
A bond in which complete pair of shared electrons is
contributed by the same element. Atom contributing its extra
electron pair is called donor while other atom is called acceptor
and the bond is represented by an arrow ( ) from donor to
acceptor.
→
→ H⁺
NH4⁺ : NH3 
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Dipole Moment
Dipole moment is a measure of the polarity of a molecule. It is defined
as the product of the magnitude of the charge and the distance
between the centers of the positive and negative charges. The dipole
moment of a molecule is a vector quantity, meaning that it has both a
magnitude and a direction. The direction of the dipole moment points
from the negative charge to the positive charge.
Dipole moments are caused by the uneven distribution of electrons in
a molecule. When a molecule has a dipole moment, it is said to be
polar. Polar molecules are attracted to each other by dipole-dipole
interactions. Dipole-dipole interactions are one of the main types of
intermolecular forces.
The magnitude of the dipole moment of a molecule depends on the
following factors:
The electronegativity of the atoms in the molecule: The greater
the difference in electronegativity between two atoms, the larger
the dipole moment of the bond between them.
The geometry of the molecule: The geometry of the molecule
determines how the polarity of the individual bonds is distributed.
For example, water has a large dipole moment because the
oxygen atom is more electronegative than the hydrogen atoms
and the molecule has a bent geometry.
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Dipole Moment
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Hydrogen Bonding
Hydrogen bonding is a special type of intermolecular force that
occurs between molecules that have a hydrogen atom bonded to a
highly electronegative atom, such as oxygen, nitrogen, or fluorine.
Intermolecular hydrogen bonding is the hydrogen bonding that
occurs between different molecules.
Hydrogen bonding is caused by the uneven distribution of electrons
in the molecule. The hydrogen atom has a partial positive charge,
while the electronegative atom has a partial negative charge. This
creates a dipole in the molecule, which is attracted to the opposite
dipole in another molecule. It is a weak interaction between H and a
highly electronegative and small sized atom.
Intermolecular H Bonding
Intermolecular H Bond : Formed between two or more different
molecules of the same or different types.
intermolecular hydrogen bonding also plays a role in the following:
The structure of proteins
The double helix structure of DNA
The transport of water in plants
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Intermolecular H Bonding
Effects of intermolecular hydrogen bonding:
High boiling and melting points: Substances with
intermolecular hydrogen bonding have high boiling and
melting points. This is because the hydrogen bonds must be
broken in order for the molecules to separate, and this
requires a lot of energy.
Viscosity: Substances with intermolecular hydrogen
bonding tend to have high viscosity. This is because the
hydrogen bonds make it difficult for the molecules to flow
past each other.
Surface tension: Substances with intermolecular hydrogen
bonding have high surface tension. This is because the
hydrogen bonds at the surface of the liquid are attracted to
each other, which creates a strong cohesive force.
Solubility: Substances with intermolecular hydrogen
bonding are generally soluble in other substances with
intermolecular hydrogen bonding. This is because the
hydrogen bonds between the different molecules can break
and reform, allowing the molecules to mix together.
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Intramolecular H Bonding
Formed within the molecule
Effects of intramolecular hydrogen bonding:
Shape: Intramolecular hydrogen bonding can restrict the rotation of
functional groups within a molecule, which can lead to specific
shapes. For example, the intramolecular hydrogen bond in salicylic
acid (aspirin) is responsible for its flat, planar structure.
Physical properties: Intramolecular hydrogen bonding can affect the
physical properties of molecules, such as boiling point, melting point,
and solubility. For example, molecules with intramolecular hydrogen
bonding tend to have higher boiling points and melting points than
similar molecules without intramolecular hydrogen bonding. This is
because the hydrogen bonds must be broken in order for the
molecules to separate, and this requires a lot of energy.
Chemical properties: Intramolecular hydrogen bonding can also
affect the chemical properties of molecules, such as acidity and
reactivity. For example, the intramolecular hydrogen bond in amides
makes them less acidic than other types of carbonyl compounds. This
is because the hydrogen bond ties up the lone pair of electrons on
the nitrogen atom, making it less available to donate to a proton.
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Resonace And Resonance Energy
Resonance: Sometimes, molecules are represented by many
structural formulas that are canonical structures or resonating
structures. It’s observed due to delocalization of pi electrons.
Resonance Energy(RE): Energy of most stable canonical
structure- Resonance Hybrid Energy.
RE
number of canonical structures
RE stability
RE
1/reactivity
RE = Expected heat of hydrogenation - Actual heat of
hydrogenation
∝
∝
∝
Due to high RE, benzene is quite stable and undergoes
electrophilic substitutions. It does not undergo addition
reactions, although it has double bonds(due to delocalization of
pi electrons or resonance).
Stability of canonical structures:
A non polar structure is always more stable than a polar
structure.
Greater the number of covalent bonds greater will be the
stability.
The canonical structure in which positive on electropositive
atom and negative charge on electronegative atom is more
stable.
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Resonace And Resonance Energy
The canonical structure in which each atom has an octet is more
stable.
If like charges are closer, the structure is unstable.
Types of resonance:
1. Isovalent: canonical structures have same number of bonds and
same type of charges.
2. Heterovalent: canonical structures have different number of bonds
and charges
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Hybridisation
(Pauling and Slater): Intermixing or redistribution of energy among
two or more half filled, fully-filled, incompletely filled or empty
orbitals of comparable energy to form same number of hybrid
orbitals of identical energy.
A hybrid bond is always a sigma bond.
A hybrid bond is always stronger than a non-hybrid bond.
Hybridization is directly proportional to overlapping.
Hybridization increases stability.
Hybridization occurs in the central atom in a molecule.
Hybridization does not occur in isolated atoms but in bonded
atoms.
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Hybridisation
Finding the type of hybridization:
1. For covalent compunds and ions, count the number of valence
electrons and (+/-) charge, to find a particular value (N).
2. Divide N to get quotient X(number of bp electrons)
If N is between 2 and 8, divide by 2
If N is between 10 and 56, divide by 8
If N is 58 or more, divide by 18.
If any remainder is left, divide to get quotient Y.
3. If X or X+Y=2=sp
If = 3 = sp²
If = 4 = sp³
If = 5 = sp³d
If = 6 = sp³d²
If = 7 = sp³d³
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Hybridisation
Finding geometry of covalent compounds:
Total number of electron pair around central atom which gives
hybridization
Chemical Bonding
Lone pair = P-N
P: electron pair around central atom
N: number of atoms surrounding central atom or number of bond pair
of electrons.
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Hybridisation in Complexes
VSEPR Theory
Valence Shell Electron Pair Replusion Theory:(Nyholm and
Gillispie)
According to the theory, besides hybridization, the nature of
electrons around the central atom also decide the shape of
molecule.
There may be two types of electrons around. The central
atom, i.e., bond pair and lone pair.
These electrons undergo repulsion and decreasing order of
repulsion is lp-lp>lp-bp>bp-bp.
Due to this repulsion, shape of molecules gets distorted and
bond angle changes.
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Rule for determination of total no. of hybrid orbitals
Rule for determination of total number of hybrid orbitals:
1. Count the number of central atom and peripheral atom’s valence
electrons
2. Divide the number by 8.
3. If N is the number, then let N/8 = 8a + b.
4. Here, a is the number of σ bonds, b is the number of nonbonding electrons
Born-Lande Equation:
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Molecular Orbital Theory
by Hund and Mulliken based on LCAO(Linear combination of
atomic orbitals) model.
Atomic orbitals undergo linear combination to form same
number of molecular orbitals, if they fulfill the following
conditions:
1. Atomic orbitals must have comparable energies
2. Atomic orbitals must overlap linearly for enough and effective
overlapping.
3. Atomic orbitals must have same symmetry along with the
major molecular axis.
Molecular orbitals are formed due to constructive and
destructive interference of atomic orbitals.
Constructive interaction of orbitals between orbital lobes
having same wave function produce bonding MOs like 𝝈,π
and Δ. These are HoMOs (highest occupied MOs.
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Molecular Orbital Theory
Destructive interaction between orbitals having different sign of 𝚿
produces antibonding MOs(LuMOs), Lowest Unoccupied MOs. eg
𝝈*,π* and Δ*.
Facts related to HoMOs and LuMOs:
1. Energy: LuMOs > HoMOs
2. Wavelength: LuMOs < HoMOs
3. LuMOs have nodal planes, HoMOs may or may not have nodal
planes.
4. Electrons contribute force of attraction in HoMOs while they
contribute repulsion in LuMOs.
5. Like atomic orbitals, MOs also follow Pauli Exclusion Principle, Hund’s
Rule, Aufbau Principle.
The MO obtained by the addition of atomic orbitals is of lower
energy and is called bonding orbital. The MO obtained by
subtraction of atomic orbitals is of higher energy and is called an
anti-bonding orbital.
Bond order =
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Molecular Orbital THeory
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Molecular Orbital THeory
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Molecular Orbital THeory
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Some Typical Bonds
ODD Electron Bond
Odd electron bond: These include one electron bond and three
electron bond. Eg: H₂⁺. The bond is half as strong as a sared electron
pair bond.
NO and NO₂ are examples of add molecules having three e bonds.
Three electron bond is formed when two atoms have nearly same
electronegativity. 3e bonds are also half as strong as a normal bond.
Back Bonding
Back bonding: If among the bonded atoms, one atom has a vacant
orbital and another has excess of e, then a sort of π bonding takes
place between the two. If this is between p orbitals of the two, this
is pπ-pπ back bonding. These bonds are most efficient when the
atoms are very small and the orbitals involved of the two are of
same energy level.
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Some Typical Bonds
Banana Bond
3 centre 2 electron bonding or bridge bonding
Conditions for banana bonding:
1.There must be presence of covalent bonding.
2.Central atom must have electron deficiency and vacant orbitals.
3.Central atom is surrounded by less than 8 electrons
4.There must be presence of dimeric or polymeric form of the hybrid
molecule.
5.If the molecules have deficiency of 2 electrons then dimeric form
will be used.
6.There must be combination of 3-atom and 2 electrons.
7.Central atom must have sp3 hybridization
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Determining Lattice Energy Of Ionic Compound
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