structure of atom

Structure of Atom

Introduction

  • Atoms are minute particles; fundamental building blocks of matter.

  • Atoms are not visible individually without sophisticated apparatus.

  • Question raised: What constitutes the atom and how are its particles arranged?

  • Basic structural units of matter; analogous to a building's structure (rooms = atoms).

Divisibility of Matter

  • Ancient Indian philosopher Maharishi Kanada proposed the division of matter leads to smaller particles, ultimately reaching paramanu, the smallest indivisible particles.

  • The atomic theory is fundamental in chemistry, as proposed by John Dalton in 1808, stating that atoms are indivisible.

  • Dalton's theory was later revised to incorporate the existence of subatomic particles: electrons, protons, and neutrons.

Advantages and Disadvantages of Dalton's Theory

Advantages

  • Successfully explained:

    • Law of Conservation of Mass

    • Law of Constant Composition

    • Law of Multiple Proportions

Disadvantages

  • Failure to account for the existence of subatomic particles.

Subatomic Particles

  • Subatomic particles are defined as particles smaller than an atom.

  • Electrons: Discovered via cathode rays experiments conducted by J. J. Thomson.

  • Cathode Ray Discharge Tubes: Experiment shows that gases conduct electricity under specific conditions (high voltage, low pressure).

Cathode Ray Experiment

  • Tubes constructed to study electrical discharge in a vacuum.

  • Current flows from cathode (-) to anode (+) and produces cathode rays.

  • Detection employed a phosphorescent screen (zinc sulfide).

_ Observations from Cathode Ray Experiment

  • Cathode rays contain negatively charged electrons.

  • Their behavior is independent of the types of gas present and the electrodes used.

Charge-to-Mass Ratio of Electrons

  • J.J. Thomson measured the charge-to-mass ratio of electrons (e/m) in 1897 using electric and magnetic fields.

  • Deflection depends on:

    1. Magnitude of negative charge.

    2. Mass of the particle.

    3. Strength of the electric and magnetic fields.

  • Thomson's findings led to an accepted charge-to-mass ratio of 1.75882 × 10^11 C/kg.

Discovery of Protons and Neutrons

Discovery of Protons

  • Michael Faraday's experiments indicated that atoms consist of both negatively charged electrons and positive particles.

  • Anode rays produced when a perforated cathode allows particles to flow from anode to cathode.

  • These positively charged particles were named protons.

Discovery of Neutrons :

  • Proposed by Rutherford, confirmed by James Chadwick in 1932 through experiments with beryllium and alpha particles.

  • Neutrons are neutral particles, contributing to atomic mass without charge.

Atomic Models

Thomson's Atomic Model (1898)

  • Proposed the atom as a uniform sphere of positive charge with atoms of negative charge (electrons) embedded within (like plum pudding).

  • Model failed due to discovery of the atomic nucleus.

Rutherford’s Nuclear Model (1911)

  • Based on alpha particle scattering experiments aiming to understand atomic structure.

  • Major observations:

    1. Most alpha particles pass through gold foil undeflected

    2. Few are scattered at small angles

    3. Very few bounce back nearly at 180°

  • Proposed that the nucleus is a small, dense, positively charged center of anatom, surrounded by electrons.

Limitations of Rutherford’s Model

  • Did not explain the stability of atoms nor electron arrangements around the nucleus.

  • Failed to address the electronicstructure and energy levels of electrons.

Bohr's Model (1913)

  • Niels Bohr modified Rutherford's model to introduce quantized electron orbits (stationary states) around the nucleus.

  • Key Postulates:

    1. Electrons occupy fixed orbits around the nucleus.

    2. Orbit energy levels are quantized.

    3. Electrons can only transition between these orbits by absorbing or emitting specific amounts of energy (quanta).

Support and Limitations of Bohr’s Theory

  • Supported spectral data for hydrogen atoms and isotopes.

  • Limitations include inability to explain complex spectra from multi-electron systems.

Quantum Mechanical Model

  • Combines concepts from earlier atomic models, the wave-particle duality, and uncertainty principles to describe atom behavior based on wave functions.

  • Schrödinger’s equation describes the quantum states of electrons in atoms.

Quantum Numbers

  • Four quantum numbers are needed to describe electron states:

    1. Principal quantum number (n) - energy level of electron.

    2. Azimuthal quantum number (l) - shape of orbital.

    3. Magnetic quantum number (m) - orientation of orbital.

    4. Spin quantum number (s) - orientation of electron spin.

Filling of Orbitals

  • Aufbau Principle: orbitals are filled in order of increasing energy.

  • Hund's Rule: electrons fill degenerate orbitals singly before pairing.

  • Pauli Exclusion Principle: no two electrons in an atom can have identical quantum numbers.

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