NIOS Final Exam Focus

Lesson 2: Atomic Structure

Revise this chapter by remembering discoveries, models, spectrum formulas, Bohr theory, quantum numbers, orbital shapes, and electronic configuration rules.

1. Fundamental Particles

Atom is made of electron, proton, and neutron.
Electron has negative charge and very small mass.
Proton has positive charge and mass about 1836 times electron mass.
Neutron has no charge and mass nearly equal to proton.

2. Atomic Number and Mass Number

Atomic number Z = number of protons.
In a neutral atom, number of electrons = number of protons.
Mass number A = protons + neutrons.
Number of neutrons = A - Z.

A = Z + n

3. Isotopes and Isobars

Isotopes have same atomic number but different mass numbers.
Isobars have same mass number but different atomic numbers.
Isotopes have similar chemical properties because chemical behavior depends mainly on electrons.

4. Thomson and Rutherford Models

Thomson proposed plum pudding model: electrons embedded in a positively charged sphere.
Rutherford's alpha-particle scattering experiment proved that most of atom is empty space.
Nucleus is small, dense, positively charged, and contains most of the mass.
Rutherford model failed because it could not explain stability of atom.

5. Electromagnetic Radiation

Electromagnetic radiation travels as oscillating electric and magnetic fields.
It does not need a medium.
It travels with speed of light in vacuum.
Main terms: amplitude, wavelength, frequency, wave number, velocity.

c = νλ and wave number = 1/λ

6. Photon Energy Formula

Light also behaves like particles called photons.
Energy of photon is directly proportional to frequency.
Higher frequency means higher energy.

E = hν = hc/λ

7. Hydrogen Spectrum

Hydrogen gives a line spectrum, not a continuous spectrum.
Important series: Lyman, Balmer, Paschen, Brackett, Pfund.
Balmer series lies in visible region.
Lyman series lies in ultraviolet region.
Paschen, Brackett and Pfund lie in infrared region.

1/λ = R_H(1/n₁² - 1/n₂²)

8. Bohr's Model

Electrons move in fixed circular orbits called stationary orbits.
Electrons do not radiate energy while moving in these allowed orbits.
Energy is absorbed or emitted when electron jumps from one orbit to another.
Angular momentum is quantized.

mvr = nh/2π and E_f - E_i = hν

9. Key Bohr Result

Energy of orbit depends on principal quantum number n.
As n increases, energy increases and becomes less negative.
Bohr model explained hydrogen spectrum and atomic stability better than Rutherford model.
But it failed for multi-electron atoms.

E_n ∝ -1/n²

10. Wave-Particle Duality and Uncertainty Principle

Light shows both wave and particle nature.
de Broglie proposed that matter particles like electrons also have wave nature.
Heisenberg stated that exact position and exact momentum cannot be measured simultaneously.

λ = h/mv

11. Quantum Mechanical Model

Schrodinger wave equation gives wave function ψ.
ψ² gives probability of finding electron in a region.
Atomic orbital is the 3D region where probability of finding electron is maximum.
Orbit and orbital are different: orbit is fixed path, orbital is probability region.

12. Quantum Numbers

Principal quantum number n gives shell and energy.
Azimuthal quantum number l gives subshell and shape.
Magnetic quantum number m_l gives orientation.
Spin quantum number m_s gives spin direction.

13. Shapes of Orbitals

s orbital is spherical.
p orbitals are dumb-bell shaped.
d orbitals are cloverleaf shaped.
Nodal plane is region where probability of finding electron is zero.
Node is a zero-probability region.

14. Electronic Configuration Rules

Aufbau principle: electrons fill orbitals in increasing order of energy.
Pauli exclusion principle: no two electrons have same set of four quantum numbers.
Hund's rule: pairing starts only after all degenerate orbitals are singly occupied.

Order: 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p ...

15. Stability of Half-filled and Fully-filled Orbitals

Half-filled and completely filled subshells are extra stable.
Main reasons: symmetry and exchange energy.
Important exceptions: chromium and copper.
Cr = [Ar] 3d⁵ 4s¹, Cu = [Ar] 3d¹⁰ 4s¹.

16. NIOS Exam Traps

Do not mix up isotope and isobar.
Bohr model explains hydrogen spectrum, not all atoms.
Orbit is not same as orbital.
Remember region of spectrum for each hydrogen series.
Remember exceptions in electronic configuration like Cr and Cu.

17. High-Yield Questions to Practice

Compare Thomson and Rutherford models.
State and explain Bohr's postulates.
Derive or use photon energy and de Broglie equation numericals.
Differentiate line spectrum and continuous spectrum.
Explain quantum numbers and orbital shapes.
Write electronic configurations using Aufbau, Pauli and Hund rules.

Best scoring method: memorize definitions, one-line differences, and standard formulas with units.