(i) Similar to the other electron displacement effect, mesomeric effect is also classied into positive mesomeric effect (+M or +R) and negative mesomeric effect (-M of -R) based on the nature of the functional group present adjacent to the multiple bond.
(ii) Resonance is useful in explaining certain properties such as acidity of phenol. The phenoxide ion is more stabilised than phenol by resonance effect (+M effect) and hence resonance favours ionisation of phenol to form. H⁺ and shows acidity


The above structures shows that there is a charge separation in the resonance structure of phenol which needs energy, where as there is no such hybrid structures in the case of phenoxide ion. This increased stability accounts for the acidic character of phenol

Inductive effect (I):
(i) Inductive effect is defined as the change in the polarisation of a covalent bond due to the presence of adjacent bonds, atoms or groups in the molecule. This is a permanent phenomenon.
(ii) Let us explain the inductive effect by considering ethylchloride as example. The C-C bond in ethyl chloride is polar.
We know that chlorine is more electronegative than carbon, and hence it attracts the shared pair of electron between C-Cl in ethyl chloride towards itself. is develops a slight negative charge on Chlorine and a slight positive charge on carbon to which chlorine is attached.
To compensate it, the C₁ draws the shared pair of electron between itself and. C₂ . This polarisation effect is called inductive effect.
(iii) The magnitude of the charge separation decreases rapidly, as we move away from C₁ and is observed maximum for 2 carbons and almost insignicant after 4 bonds from the active group.
\(\overset { \delta }{ C } \overset { \delta + }{ { H }_{ 3 } } \longrightarrow \underset { 1 }{ C\overset { \delta + }{ { H }_{ 2 } } } \twoheadrightarrow \overset { \delta - }{ C } { l }_{ 2 }\)
It is important to note that the inductive effect does not transfer electrons from one atom to another but the displacement effect is permanent. The inductive effect represents the ability of a particular atom or a group to either withdraw or donate electron density to the attached carbon. Based on this ability the substituents are classified as +I groups and -I groups. Their ability to release or withdraw the electron through sigma covalent bond is called +I effect and -I effect respectively.
Highly electronegative atoms and atoms of groups which are carry a positive charge are electron withdrawing or -I group
Example : -F, -CI, -COOH, -NO₂, NH₂,
Higher the electronegativity of the substitutent, greater is the -I effect.
The order of the -I effect of some groups are given below :
NH³⁺ > NO₂ > CN > SO₃H > CHO > CO > COOH > COCI > CONH₂ > F > Cl > Br > I > OH > OR, NH₂ > C₆H₅ > H
Highly electropositive atoms and atoms are groups which carry a negative charge are electron donating or +I groups.
Example. Alkali metals, alkyl groups such as methyl, ethyl, negatively charged groups such as CH₃O^-, C₂H₅O^-, COO^- etc.
Lesser the electronegativity of the elements, greater is the +I effect. The relative order of +I effect of some alkyl groups is given below
\(-\mathrm{C}\left(\mathrm{CH}_3\right)_3>-\mathrm{CH}\left(\mathrm{CH}_3\right)_2>-\mathrm{CH}_2 \mathrm{CH}_3>-\mathrm{CH}_3\)
Let us understand the influence of inductive effect on some properties of organic compounds.
Reactivity :
When a highly electronegative atom such as halogen is attached to a carbon then it makes the C-X bond polar. In such cases the -I effect of halogen facilitates the attack of an incoming nucleophile at the polarised carbon, and hence increases the reactivity.

If a - I group is attached nearer to a carbonyl carbon, it decreases the availability of electron density-on the carbonyl carbon, and hence increases the rate of the nucelophilic addition reaction.
Aciditv of carborvlic acids :
When a halogen atom is attached to the carbon which is nearer to the carboxylic acid group, its -I effect withdraws the bonded electrons towards itself and makes the ionisation of H⁺ easy. The acidity of various chloro acetic acid is in the following order. The strength of the acid increases with increase in the -I effect of the group attached to the carboxyl group.
Tiichloro acetic acid > Dichloro acetic acid > Chloro acetic acid > acetic acid

Louis de Broglie proposed that all forms of matter showed dual character. To quantify this relation, he Jerived an equation for the wavelength of a matter wave. He combined the following two equations of energy of which one represents wave character (hu) and the other represents the particle nature (mc²).
Planck's quantum hypothesis:
E = hv      ....(1)
Einsteins mass-energy relationship:
E = mc²     .... (2)
From (1) and (2)
hv = mc²
hc/\(\lambda\) = mc²
\(\therefore \lambda ={h\over mc}\)        ...(3)
The equation (3) represents the wavelength of photons whose momentum is given by mc. (Photons have zero rest mass).
For a particle of matter with mass m and moving with a velocity v, the equation (3) can be written as
 \(\lambda ={h\over mv}\)       ....(4)
This is valid only when the particle travels at speed much less than the speed of Light.

(i) no. of electrons: 35 (given)
no. of protons : 35
(ii) Electronic configuration
1s² 2S² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁵
(iii) Last electron:

4P_x        4P_y     4p_z
last electron present in 4P_y orbital y
n = 4, l = 1 m₁ = either + 1 or -1 and s = -1/2

Statement : "No two electrons in an atom can have the same set of values of all four quantum numbers"
Explanation : It means that, each electron must have unique values for the four quantum numbers (n, l, m and s).
For the lone electron present in hydrogen atom, the four quantum numbers are: n = 1; l = 0; m = 0 and s = +1/2. For the two electrons present in helium, one electron has the quantum numbers same as the electron of hydrogen atom, n = 1.
l = 0, m = 0 and s = +1/2. For other electron, the fourth quantum number is different i.e., n = 1, l = 0, m = 0 and s = -1/2.
As we know that the spin quantum number can have only two values +1/2 and - 1/2, only two electrons can be accommodated in a given orbital in accordance with pauli exclusion principle.

Limitation of Bohr's atom model:
(a) The Bohr's atom model is applicable only to species having one electron such as hydrogen, Li²⁺ etc ... and not applicable to multi electron atoms.
(b) It was unable to explain the splitting of spectral lines in the presence of magnetic field (Zeeman effect) or an electric field (Stark effect).
(c) Bohr's theory was unable to explain why the electron is restricted to revolve around the nucleus in a fixed orbit in which the angular momentum of the electron is equal to nh/2π

It states that electron pairing in the degenerate orbitals does not take place until all the available orbitals contains one electron each. We know that there are three p orbitals, five d orbitals and seven f orbitals.
According to this rule, pairing of electrons in these orbitals starts only when the 4^th, 6^th and 8^th electron enters the p, d and f orbitals respectively. For example, consider the carbon atom which has six electrons. According to Aufbau principle, the electronic configuration is 1s², 2s², 2p². It can be represented as below,

In this case, in order to minimise the electron-electron repulsion, the sixth electron enters the unoccupied 2p_y orbital as per Hund's rule. i.e. it does not get paired with the fifth electron already present in the 2P_x orbital.

In this reaction two substituents are eliminated from the molecule, and a new C - C double bond is formed between the carbon atoms to which the eliminated atoms/groups are previbusly attached. Elimination reaction is always accompanied with chanle in hybridisation.
Example: n- Propyl bromide on reaction with alcoholic KOH gives Propene. In this reaction hydrogen and Br are eliminated.

a) Electrophiles:
Electrophiles are reagents that are attracted towards negative charge or electron rich center. They are either positively charged ions or electron deficient neutral molecules. All Lewis acids act as electrophiles.
Neutral molecules like SnCl₄ can also act as an electrophile, as it has vacant d-orbitals which can accommodate the electrons from others.

b) Nucleophiles:
Nucleophiles are reagents that has high affinity for electro positive centers. They possess an atom has an unshared pair of electrons, and hence it is in search for an electro positive centre where it can have an opportunity to share its elections to form a covalent bond, and gets stabilised.
They are usually negatively charged ions or electron rich neutral molecules (contains one or more lone pair of electrons). AII Lewis bases act as nucleophiles.

Electronic configuration of oxygen

ஃ 8^th electron present in 2p_x orbital and the quantum numbers are
n = 2,l = 1,m₁ = either + 1 or -1 and s = -1/2
Electronic configuration of chlorine 

15^th electron present in 3P_z orbital and the quantum numbers are n = 3, l = 1, m₁ = either +1 or -1 and ms = +1/2

Erwin Schrodinger expressed the wave nature of electron in terms of a differential equation. This equation determines the change of wave function in space depending on the field of force in which the electron moves. The time independent Schrodinger equation can be expressed as,
\(\overset { \wedge }{ H } \psi =E\psi \)          .........(1)
Where \(\overset { \wedge }{ H } \) is called Hamiltonian operator, \(\psi \) is the wave function and is a function of position coordinates of the particle and is denoted as \(\psi \) (x, y, z) E is the energy of the system
\(\overset { \wedge }{ H } =\left[ \frac { { -h }^{ 2 } }{ 8{ \pi }^{ 2 } } \left( \frac { { \partial }^{ 2 }\psi }{ { \partial x }^{ 2 } } +\frac { { \partial }^{ 2 }\psi }{ { \partial y }^{ 2 } } +\frac { { \partial }^{ 2 }\psi }{ { \partial z }^{ 2 } } \right) +V \right] \)
can be written as
\(\left[ \frac { { -h }^{ 2 } }{ 8{ \pi }^{ 2 }m } \left( \frac { { \partial }^{ 2 }\psi }{ { \partial x }^{ 2 } } +\frac { { \partial }^{ 2 }\psi }{ { \partial y }^{ 2 } } +\frac { { \partial }^{ 2 }\psi }{ { \partial z }^{ 2 } } \right) +V\Psi \right] =E\Psi \)
Multiply by \(\frac { 8{ \pi }^{ 2 }m}{ { -h }^{ 2 } } \)and rearranging
\(\frac { { \partial }^{ 2 }\psi }{ { \partial x }^{ 2 } } +\frac { { \partial }^{ 2 }\psi }{ { \partial y }^{ 2 } } +\frac { { \partial }^{ 2 }\psi }{ { \partial z }^{ 2 } } +\frac { 8{ \pi }^{ 2 }m }{ { -h }^{ 2 } } (E-V)\Psi =0\)          ........(2)
The above Schrodinger wave equation does not contain time as a variable and is referred to as time independent Schrodinger wave equation. This equation can be solved only for certain values of E, the total energy. i.e. the energy of the system is quantised. The permitted total energy values are called eigen values and corresponding wave functions represent the atomic orbitals.

The delocalisation of electrons of σ bond is called as hyper conjugation. It is a special stabilising effect that results due to the interaction of electrons of a σ-bond (usually C-H or C-C) with the adjacent, empty non-bonding p-orbital or an anti- bonding σ* or π^* orbitals resulting in an extended molecular orbital. Unlike electrometric eect, hyper conjugation is a permanent effect.
It requires an σ-CH group or a lone pair on atom like N, O adjacent to a π bond (sp² hybrid carbon). It occurs by the overlapping of the o-bonding orbital or the orbital containing a lone pair with the adjacent π-orbital or p-orbital.
Example 1 :

No bond resonance structures shown by propene are due to hyperconjugation
In propene, the σ-electrons of C-H bond of methyl group can be delocalised into the n-orbital of doubly bonded carbon as represented below.
In the above structure the sigma bond is involved in resonance and breaks in order to supply electrons for delocalisation giving rise to 3 new canonical forms. In the contributing canonical structures : (II), (III) & (IV) of propene, there is no bond between an carbon and one of the hydrogen atoms. Hence the hyperconjugation is also known as "no bond resonance" or "Baker-Nathan effect''. The structures (II), (III) & (IV) are polar in nature.
Example 2:
Hyper conjugation effect is also observed when atoms / groups having lone pair of electrons are attached by a single bond, and in conjugation with a π bond. The lone pair of electrons enters into resonance and displaces π electrons resulting in more than one structure.

Example 3 :
When electronegative atoms or group of atoms are in conjugation with a π - bond, they pull π - electrons from the multiple bond.

In case of carboc ations, greater the number of alkyl groups attached to the carbon bearing positive charge, greater is number of the hlper conjugate structure. thus the stability of various carbocations decreases in the order
3⁰ Carbocation > 2⁰ Carbocation > 1⁰ Carbocation.

All organic compounds having double or triple bond adopt addition reactions in which two substances unite to form a single compound. During the addition reaction the hybridization of the substrate changes as only one bond breaks and two new bonds are formed.
Example :
\(\underset { (Acetylene) }{ CH\equiv CH+{ CH }_{ 2 } } \longrightarrow \underset { (Bromethene) }{ { CH }_{ 2 }=CH\ominus Br } \) 

The resonance is a chemical phenomenon which is observed in certain organic compounds possessing double bonds at a suitable position. Certain organic compounds can be represented by more than one structure and they differ only in the position of
bonding and lone pair of electrons. Such structures are called resonance structures (canonical structures) and this phenomenon is called resonance. This phenomenon is also called mesomerism or mesomeric effect.
For example, the structure of aromatic compounds such as benzene and conjugated systems tike 1,3 - butadiene cannot be represented by a single structure, and their observed properties can be explained on the basis of a resonance hybrid.
In 1, 3 buta diene, it is expected that the bond between C₁ - C₂ and C₃ - C₄ should be shorter than that of C₂ - C₃, but the observed bond lengths are of same. This property cannot be explained by a simple structure in which two \(\pi\) bonds localised between C₁ - C₂ and C₃ - C₄ . Actually the \(\pi\) electrons are delocalised

These resonating structures are called canonical forms and the actual structure lies between these three resonating strucfures, and is called a resonance hybrid. The resonance hybrid is represented
Similar to the other electron displacement effect, mesomeric effect is also classified into positive mesomeric effect (+M or +R) and negative mesomeric effect (-M of -R) based on the nature of the functional group present adjacent to the multiple bond.
Positive Mesomeric Effect :
Positive resonance effect occurs, when the electrons move away from substituent attached to the conjugated system. It occurs, if the electron releasing substituents are attached to the conjugated system. In such cases, the attached group has a tendency to release electrons through resonance. These electron releasing groups are usually denoted as +R or +M groups.
Examples : -OH, -SH, -OR, -SR, -NH₂, -O- etc...
Negative Mesomeric Effect :
Negative resonance effect occurs, when the electrons move towards the substituent attached to the conjugated system. It occurs if the electron withdrawing substituents are attached to the conjugated system.

In such cases, the attached group has a tendency to withdraw electrons through resonance. These electron withdrawing groups are usually denoted as -R or -M groups.
Examples : NO₂, > C = O, -COOH, -C = N etc......
Resonance is useful in explaining certain properties such as acidity oi phenol. The phenoxide ion is more stabilised than phenol by resonance effect (+M effect) and hencp resonance tavours ionisation of phenol to form H⁺ and shows acidity.
The structures shows that there is a charge separation in the resonance structure of phenol which needs energy, where as there is no such hybrid structures in the case of phenoxide ion. This increased stability accounts for the acidic character of phenol.

Electronic configuration of Fe³⁺ : 1s² 2s² 2p⁶ 3s² 3p⁶ 4s⁰ 3d⁵
Electronic configuration of Ni²⁺ : 1s² 2s² 2p⁶ 3s² 3p⁶ 4s⁰ 3d⁸
Fe³⁺ has stable 3d⁵ half filled configuration.

n = 4 l = 0,1,2,3
4 sub shells s, p, d & f.
I = 0 m₁ = 0 + one 4s orbital.
I = 1 m₁ = -1, 0, + 1 \(\Rightarrow\) three 4p orbitals.
I = 2 m₁ = -2,:1, 0, +1, +2 \(\Rightarrow\) five 4d orbitals.
I = 3 m₁ = -3, -2, -1,0, +1, +2, +3 \(\Rightarrow\) seven 4f orbitals.
Over all 16 orbitals are possible.