Physics XII - Chapter 12: Atoms
Practice Atoms Class 12 MCQs. Learn atomic models, Bohr's theory, spectra, energy levels, and Rutherford's experiment for board & competitive exams.
Quick Revision Box
- Rutherford Model: Nuclear atom with electrons orbiting nucleus; planetary model.
- Alpha Particle Scattering: Gold foil experiment revealing atomic nucleus.
- Impact Parameter: Perpendicular distance from nucleus to initial velocity direction.
- Electron Orbits: Only certain stable orbits allowed in Bohr model.
- Bohr's Postulates: Quantized angular momentum and non-radiating orbits.
- Angular Momentum Quantization: mvr = nh/(2π); n = 1, 2, 3...
- Bohr Radius: rₙ = (ε₀h²n²)/(πme²) = 0.529 × 10⁻¹⁰ n² m; radius of nth orbit.
- Orbital Velocity: vₙ = e²/(2ε₀hn) = (2.18 × 10⁶)/n m/s; decreases with n.
- Energy Levels: Eₙ = -13.6/n² eV; negative indicating bound state.
- Ionization Energy: Minimum energy to remove electron from ground state; 13.6 eV for hydrogen.
- Excitation Energy: Energy needed to move electron from ground state to excited state.
- Line Spectra: Discrete wavelengths emitted by atoms; atomic fingerprint.
- Rydberg Formula: 1/λ = R(1/n₁² - 1/n₂²); predicts hydrogen spectral lines.
- Rydberg Constant: R = 1.097 × 10⁷ m⁻¹; fundamental constant in atomic spectra.
- Lyman Series: UV region; n₁ = 1, n₂ = 2,3,4...
- Balmer Series: Visible region; n₁ = 2, n₂ = 3,4,5...
- Paschen Series: IR region; n₁ = 3, n₂ = 4,5,6...
- Brackett Series: Far IR; n₁ = 4, n₂ = 5,6,7...
- Pfund Series: Far IR; n₁ = 5, n₂ = 6,7,8...
- De Broglie Explanation: Stable orbits have integer number of electron wavelengths.
- Atomic Spectra: Unique to each element; used in chemical analysis.
Basic Level Questions
Chapter Summary
Atoms takes us on an extraordinary journey into the heart of matter, exploring the intricate architecture of the building blocks that make up everything around us. This chapter reveals how scientists peeled back layer after layer of mystery to understand the astonishing structure of the atom - a tiny solar system with a massive nucleus at the center and electrons whizzing around in precisely defined orbits.
The adventure begins with Rutherford's gold foil experiment, where alpha particles bouncing back from thin gold foil revealed the stunning truth that atoms are mostly empty space with a tiny, dense nucleus at their center. This discovery overturned the prevailing "plum pudding" model and set the stage for one of physics' most beautiful theoretical constructions - Niels Bohr's quantum model of the atom.
Bohr's genius lay in blending classical physics with revolutionary quantum ideas, proposing that electrons can only occupy certain special orbits where their angular momentum is quantized. The mathematics that emerges is breathtakingly elegant: specific energy levels, precise orbital radii, and the magical number 13.6 electron volts that defines hydrogen's ionization energy. Most wonderfully, Bohr's model perfectly explains the mysterious line spectra that had puzzled scientists for decades - the unique fingerprint of light that each element emits when excited.
The spectral series - Lyman, Balmer, Paschen, and beyond - become our window into the quantum dance of electrons jumping between energy levels. Each transition tells a story of energy absorbed or emitted, creating the beautiful colors that identify elements in distant stars and laboratory samples. This chapter reminds us that the same quantum rules governing hydrogen atoms in laboratories also orchestrate the nuclear furnaces of stars, connecting the microscopic world with the cosmic scale in a symphony of universal physical laws.