• Understand the historical development of atomic theory.
• Explain the significance of J. J. Thomson's experiments in the context of atomic structure.
• Describe the differences between Thomson's and Rutherford's models of the atom.
• Analyze the limitations of Rutherford's model in explaining atomic stability and spectra.
• Summarize Bohr's model of the hydrogen atom and its postulates.
• Apply Bohr's quantization condition to calculate energy levels and transitions in hydrogenic atoms.
• Discuss the implications of de Broglie's hypothesis on the wave nature of electrons.
• Evaluate the applicability of Bohr's model to multi-electron atoms and its limitations.

Chapter Twelve: Atoms

12.1 Introduction

• By the nineteenth century, evidence supported the atomic hypothesis of matter.
• In 1897, J. J. Thomson discovered that atoms contain negatively charged electrons, making them electrically neutral overall.
• Thomson's model (1898): Positive charge is uniformly distributed, with electrons embedded like seeds in a watermelon (plum pudding model).
• Subsequent studies revealed a different arrangement of charges.

12.2 Key Models of the Atom

Thomson's Model

• Atom as a spherical cloud of positive charges with electrons embedded.

Rutherford's Model

• Most mass and positive charge concentrated in a tiny nucleus.
• Electrons revolve around the nucleus.
• Issues with Rutherford's model:
  • Predicts instability due to spiraling electrons.
  • Cannot explain characteristic line spectra of different elements.

Bohr's Model

• Proposed to explain line spectra and stability of hydrogenic atoms.
• Introduced three postulates:
  1. Electrons revolve in stable orbits without emitting energy.
  2. Angular momentum is quantized: L = nh/2π (n = principal quantum number).
  3. Electrons can transition between orbits, emitting or absorbing photons.
• Total energy quantized: Eₙ = -13.6 eV/n².
• Ground state energy of hydrogen atom is -13.6 eV.

12.3 De Broglie's Explanation

• De Broglie proposed that electrons have wave properties, explaining Bohr's quantization of angular momentum.
• Standing waves form in electron orbits, leading to quantized energy levels.

12.4 Limitations of Bohr's Model

• Applicable only to hydrogenic atoms; fails for multi-electron atoms.
• Cannot explain relative intensities of spectral lines.

12.5 Summary of Key Points

1. Atoms are electrically neutral, containing equal positive and negative charges.
2. Thomson's model depicts a cloud of positive charge with embedded electrons.
3. Rutherford's model has a nucleus with electrons revolving around it.
4. Bohr's model introduces quantized orbits and energy levels for hydrogenic atoms.
5. The wave nature of electrons explains quantization in Bohr's model.
6. Bohr's model has limitations and is replaced by quantum mechanics for complex atoms.