Discovery of Sub-atomic Particles and Atomic Models

Discovery of Sub-atomic Particles

• Michael Faraday's Contribution (1830): Electricity passing through an electrolyte leads to chemical reactions at the electrodes, indicating a particulate nature of electricity.
• Cathode Ray Experiment (1850s): Conducted by Faraday and others in cathode ray tubes at low pressure and high voltage.
  • Cathode Rays: Stream of particles (electrons) emitted from the cathode moving towards the anode.
  • Observations:
  1. Cathode rays are not visible but can be detected by phosphorescent materials (like zinc sulfide).
  2. In the absence of fields, they travel in straight lines.
  3. Negatively charged particles; do not depend on the cathode material.
• J.J. Thomson (1897): Measured the charge-to-mass ratio of electrons using magnetic and electric fields, resulting in the value rac{e}{m_e} = 1.758820 imes 10^{11} ext{ C kg}^{-1}.
• Robert Millikan's Oil Drop Experiment (1906-14): Determined the charge of the electron as -1.6 imes 10^{-19} C and the mass as 9.1094 imes 10^{-31} kg.

Discovery of Protons and Neutrons

• Canal Rays: Discovered while experimenting with cathode ray tubes, showed the existence of positively charged particles (protons).
• Chadwick (1932): Proposed neutrons found by bombarding beryllium.
• Protons Characteristics: Different charges/masses depending on the gas in the cathode ray tube, whereas neutrons are neutral.

Atomic Models

• John Dalton (1808): Proposed atomic theory considering atoms as indivisible.
• Thomson's Model (1898): Proposed "Plum Pudding" model where electrons were embedded in a positive sphere. Failed due to lack of stability evidence.
• Rutherford's Model (1911): Discovered through alpha particle scattering that most of the atom is empty and mass is concentrated in a small nucleus with electrons orbiting.
  • Contributions of Rutherford's Experiment:
  1. Most α-particles pass through foil, hence atoms are largely empty space.
  2. Deflections indicate concentrated positive charge in the nucleus.

Bohr's Model of the Atom

• Niels Bohr (1913): Proposed electrons move in defined orbits around nucleus with quantized energy levels. Can jump between levels by absorbing/emitting energy.
  • Observations:
  1. Only certain orbits are allowed, corresponding to certain energy levels.
  2. Can calculate the radius and energy of electron in orbits.

Quantum Mechanical Model

• Wave-Particle Duality: Developed by Schrödinger who proposed an equation that includes wave properties of particles.
• **Quantum Numbers: **
  • Principal Quantum Number (n): Indicates energy level and size.
  • Azimuthal Quantum Number (l): Indicates shape of orbital.
  • Magnetic Quantum Number ($m_l$): Orientation.
  • Spin Quantum Number ($m_s$): Direction of electron spin.
• Electron Configuration: Electrons are filled into orbitals following the Aufbau principle, Pauli exclusion principle, and Hund’s rule.

Important Principles

• Pauli Exclusion Principle: No two electrons can have the same set of quantum numbers.
• Hund’s Rule: Every orbital in a subshell gets one electron before any gets two.
• Aufbau Principle: Electrons fill orbitals starting from the lowest energy level upwards.

Key Features of Quantum Mechanical Model

• Energy levels are quantized.
• Uncertainty principle states position and momentum cannot be precisely measured simultaneously.
• Wave Function ($
  extstyle{ar{r}}$): Probability of finding an electron in a certain region.
• Atomic Orbitals: Defined by a set of three quantum numbers and have specific shapes (s, p, d, f).

Conclusion

• The successive refinement of atomic models from Dalton's indivisible atom, Thomson's Plum Pudding Model, Rutherford's Nuclear Model to Bohr's planetary model and finally to Quantum Mechanical Model provide a comprehensive understanding of atomic structure and behavior.