These MP Board Class 11th Chemistry Notes for CChapter 2 Structure of Atom help students to get a brief overview of all the concepts.

## MP Board Class 11th Chemistry Notes Chapter 2 Structure of Atom

→ Atom: Smallest particle which has all properties of the element is called atom.

→ Atomic structure : Distribution of fundamental constituent particles of atom is called atomic structure.

→ Cathode rays : Negatively charged rays which move from cathode to anode in discharge tube. Anode rays: Rays of positively charged particles which move opposite the cathode in discharge tube are called anode rays.

→ Electron (^{0}_{1}e): Particle with unit negative charge 1.60 × 10^{-19} coulomb and mass 91 × 10^{-31} kg.

→ Proton (^{1}_{1}e) : Fundamental particle of atom with unit positive charge 1.60 × 10^{-19} coulomb and mass 1.67 × 10^{-27} kg. Its mass is nearly equal to mass of hydrogen atom.

→ Neutron (^{1}_{1}n) : Fundamental particle of atom which has no charge. Mass of neutron is 1.6747 × 10^{-27} kg. It is heavier than proton.

→ Nucleus: Central part of atom is called nucleus. Radius of nucleus is 10“23 cm. Total mass of atom and positive charge is in the nucleus.

→ Atomic number: Number of proton in the riucleus of atom or number of electrons in atom is . called atomic number.

→ Mass number : Sum of neutron and proton present in nucleus is called mass number.

→ Isotopes : Different atoms of elements which have same atomic number but different atomic mass are called isotopes.

→ Shell or Orbit : Electrons rotate in stable and definite circular orbits around the nucleus. These are called energy levels or shell or orbits.

→ Relation between frequency (υ) and wavelength (λ) υ = \(\frac{c}{\lambda}\) (Where c is velocity of light = 3 x 10^{8}ms^{-1})

→ Einstein equation E = mc^{2}

Planck equation E = hυ = \(\frac{h c}{\lambda}\) (Where h = Planck’s constant = 6.626 x 10^{-34} Js)

→ Photoelectric Effect hυ = hυ_{0} + \(\frac { 1 }{ 2 }\)mv^{2}

Where hυ = Energy of Striking photon

hυ_{0} = w_{0} = Work function

\(\frac { 1 }{ 2 }\)mv^{2} = Kinetic energy of Ejected electron

(Where R = Rydberg constant = 1.09678 × 107 m^{-1})

→ Rydberg formula \(\bar{v}=\frac{1}{\lambda}=\mathrm{RZ}^{2}\left[\frac{1}{n_{1}^{2}}-\frac{1}{n_{2}^{2}}\right]\), n_{2} > n_{1}

(Where R = Ryberg constant = 1.09678 × 10^{7} m^{-1})

n_{1} =1, n_{2} = 2,3,4………………………….. UV region

n_{1} = 2, n_{2} = 3,4,5………………….. Visible region

→ Frequency of the absorbed or emitted radiation at two unstable states of transmission v = \(=\frac{\Delta E}{h}=\frac{E_{2}-E_{1}}{h}\) (Where E_{1} and E_{2} are the energies of lower and higher states)

→ Energy of stable state E_{n} = -R_{H} \(\left(\frac{1}{n^{2}}\right)\) where n = 1,2,3…………………..

→ Stable energy of H and species similar to H (Like : He^{+},Li^{2+},Be^{3+}) i.e., one electron species ) E_{n} = -2.18 x 10^{-18}\(\left(\frac{z^{2}}{n^{2}}\right)\) J

→ Radius of n^{th} orbit r_{n} = \(\frac{n^{2} a_{0}}{Z}\) [where a_{0} (Bohr’s radius of H) = \(\frac{h^{2}}{4 \mathrm{~A}^{2} m e^{2} k}\) = 0.529Å

→ de-Broglie equation \(\lambda=\frac{h}{m v}=\frac{h}{\sqrt{2 m(\mathrm{KE})}}\)

→ Angular momentum mvr = \(\frac{n h}{2 \pi}\)

→ Heisenberg’s uncertainty principle Δx . Δp ≥ \(\frac{h}{4 \pi}\) or Δx . Δp ≥ \(\frac{h}{4 \pi}m\)

[Where Δx and Δv are uncertainty in position and principle]

→ Velocity in n shell v_{n} = 2.182 × 10^{6} × \(\frac{z}{n}\).

→ Ionisation Energy (IE)_{H} = ∆E ∝\(z^{2}\left[\frac{1}{n_{1}^{2}}-\frac{1}{n_{2}^{2}}\right]\)

→ Position of an electron in an atom is determined by four quantum numbers (n, l, in. s

(i) n (Principal Quantum number) = 1, 2, 3, 4 ………………n

(ii) l (Azimuthal Quantum number) = 0, 1, 2 ………….. (n – 1)

(iii) m (Magnetic Quantum number) = – l to + l

(iv) s (Spin Quantum number) = \(+\frac{1}{2}\) or \(-\frac{1}{2}\)

→ Determination of subshell = nl

→ l also expresses the shape of orbital surrounded by electron : s – Spherical, p – dumbell, d – Double dumbell, f- Complex.

→ Total value of m – (2l + 1) = Number of orbitals in subshell = Number of spectrum lines in Magnetic or Electric field.

→ In an atom, l angular node, (n – l – 1) radial node i.e., Total nodes are (n – 1).

→ Electrons are filled in the various orbitals on the basis of the following rules :

(i) Aufbau’s Principle, (ii) Hund’s Rule, (iii) Pauli’s Exclusion Principle.

→ Electronic configuration of _{24}Cr = [Ar]3r^{5}4s^{1}

→ Electronic configuration of _{29}CU = [Ar]3d^{10}4s^{1}

→ Completely filled and half filled orbitals are more stable due to same symmetry and maxi¬mum energy exchange.

→ Energy of electron due to which it is bound to the nucleus = Work function (w_{0}) of metal.

→ Subshell: In a shell all electrons do not have same energy so shells are divided in different subshells. These subshells are s, p, d and f

→ Orbital: Space around the nucleus where probability to find a electron is maximum is called orbital. ‘

→ Spectrum: On jumping of electrons from higher energy level to lower energy level, obtained lines on photographic plate from the produced light is called spectrum.

→ Visible spectrum : It is a part of electromagnetic radiation which can be seen by our eyes. Wavelength range is 4000 Å to 7500 Å.

→ Invisible spectrum : The wavelength range of 7500 Å to 3 x 1.0^{6}Å and before violet up to 4000A, which we cannot see is called invisible spectrum.

→ Hund’s rule : Pairing of electrons in orbitals of equivalent energy occurs when there is no vacant orbital.

→ Aufbau’s principle: Electrons are filled in subshells in increasing order of energy. Electrons are filled first in the shell whose (n + l) value is less. If (n + l) value is same, then electron will go to that shell whose in n value is lower.

→ Pauli’s exclusion principle: No two electrons in an atom have similar four quantum numbers