Atomic Structure Summary

Introduction

This chapter explores the fundamental structure of atoms, from early atomic models to modern quantum mechanical concepts. Understanding atomic structure is crucial for explaining chemical bonding, periodicity, and various physical properties of elements.

Atom is the smallest unit of matter that retains the properties of an element. Modern atomic theory is based on quantum mechanics and the concept of subatomic particles.

Fundamental Particles

Electron (e⁻)

Discovered by J.J. Thomson (1897)

Charge: -1.602 × 10⁻¹⁹ C

Mass: 9.109 × 10⁻³¹ kg

Specific Charge: 1.76 × 10⁸ C/g

Proton (p⁺)

Discovered by Goldstein (1886)

Charge: +1.602 × 10⁻¹⁹ C

Mass: 1.673 × 10⁻²⁷ kg

Specific Charge: 9.58 × 10⁴ C/g

Neutron (n)

Discovered by James Chadwick (1932)

Charge: 0

Mass: 1.675 × 10⁻²⁷ kg

Specific Charge: 0

Comparison of Fundamental Particles

Property	Electron	Proton	Neutron
Mass (kg)	9.109 × 10⁻³¹	1.673 × 10⁻²⁷	1.675 × 10⁻²⁷
Charge (C)	-1.602 × 10⁻¹⁹	+1.602 × 10⁻¹⁹	0
Specific Charge (C/g)	1.76 × 10⁸	9.58 × 10⁴	0
Relative Mass	1/1837	1	1

Atomic Models

Thomson's Plum Pudding Model

J.J. Thomson proposed that atoms are uniform spheres of positive charge with electrons embedded in them, similar to plums in a pudding.

Limitations: Could not explain Rutherford's scattering experiment results or atomic spectra.

Rutherford's Nuclear Model

Based on the gold foil experiment, Rutherford proposed that atoms have a small, dense, positively charged nucleus surrounded by electrons.

Radius of nucleus: r_n = r₀ × A^1/3 (where r₀ = 1.4 × 10^-13 cm)

Limitations: Could not explain the stability of atoms (electrons should spiral into nucleus) or atomic spectra.

Bohr's Atomic Model

Bohr combined classical physics with quantum concepts to explain hydrogen spectrum:

Angular momentum: mvr = nħ (where ħ = h/2π)

Radius of n^th orbit: r_n = 0.529 × n²/Z Å

Energy of n^th orbit: E_n = -13.6 × Z²/n² eV

Successes: Successfully explained hydrogen spectrum and calculated Rydberg constant accurately.

Quantum Mechanical Model

Dual Nature of Matter

De Broglie proposed that matter exhibits both particle and wave properties:

de Broglie wavelength: λ = h/p = h/mv

Heisenberg's Uncertainty Principle

It's impossible to simultaneously determine exact position and momentum of a particle:

Δx × Δp ≥ h/4π

Schrödinger Wave Equation

The fundamental equation of quantum mechanics:

ĤΨ = EΨ (where Ĥ is Hamiltonian operator)

Wave function (Ψ) describes the quantum state of a particle, and |Ψ|² gives the probability density of finding the particle.

Quantum Numbers

Quantum Number	Symbol	Values	Significance
Principal	n	1, 2, 3, ...	Energy level, size of orbital
Azimuthal	l	0 to (n-1)	Shape of orbital (s, p, d, f)
Magnetic	m	-l to +l	Orientation of orbital
Spin	s	+½, -½	Spin direction of electron

Orbital Shapes

s-orbitals: Spherical (l=0)
p-orbitals: Dumbbell-shaped (l=1)
d-orbitals: Double dumbbell or cloverleaf (l=2)
f-orbitals: Complex shapes (l=3)

Important Points to Remember

Key Points for JEE Main

Bohr's model works only for hydrogen-like atoms (one electron systems)
de Broglie wavelength is significant only for microscopic particles
Number of radial nodes = n - l - 1
Number of angular nodes = l
Maximum electrons in a shell = 2n²
Electron configuration exceptions: Cr, Cu, Mo, Ag, Au
Magnetic moment = √[n(n+2)] Bohr Magnetons (for n unpaired electrons)

Do's

Memorize quantum numbers and their significance

Practice numerical problems on Bohr's model

Understand the shapes of different orbitals

Learn electron configuration rules

Don'ts

Don't confuse Thomson and Rutherford models

Don't apply Bohr's model to multi-electron atoms

Don't ignore the limitations of each atomic model

Don't forget electron configuration exceptions

JEE Main Weightage

This chapter typically carries 2-3 questions in JEE Main, making it a high-weightage chapter. Questions often focus on quantum numbers, electronic configuration, and atomic models.

Chapter 2 Weightage in JEE Main

Weightage High (2-3 questions)