Study Notes on Atomic Structure and Models

Nature of Matter

1.2 Atomic Structure

Lesson Overview

• Date: 12th September 2025
• Learning Goals:   - Describe the nuclear model of the atom.   - State the assumptions & limitations of the nuclear model.   - Describe the Bohr model of the atom.   - State the assumptions & limitations of the Bohr model.   - Describe the orbital model of the atom.   - State the assumptions & limitations of the orbital model.   - State the properties of protons, neutrons & electrons.

Definition of Matter

• Matter: Anything that takes up space and has mass.
• Composed of tiny invisible particles (Particulate nature of matter):   - Atoms   - Molecules   - Ions

Understanding Models

• What is a Model?   - A simplified representation of reality used to understand, explain, or predict phenomena.   
• Purpose of Using Models:   - Facilitate comprehension of complex systems or concepts.   - Assist in the prediction of outcomes.

Atomic Models to Examine:

• Nuclear Model
• Bohr Model
• Orbital Model

Timeline of Atomic Models

• Dalton's Model of The Atom (1808):   - Matter consists of indivisible atoms.   - Atoms can arrange in different combinations to form various compounds.   
• Plum Pudding Model (1904) by J.J. Thomson:   - Indicated particles within the atom but lacked experimental proof.

Historical Context

Greek Philosophers

• Proposed that matter was made up of small indivisible particles (~400 BC).

Dalton's Atomic Theory (1808)

• Proposing insights by John Dalton:   1. All matter comprises very small particles called atoms.   2. Atoms are indivisible.
• Critique: Dalton's theories began to be challenged by the end of the nineteenth century.

Discovery of Cathode Rays

William Crookes (1875)

• Conducted experiments passing electric current through gases using a vacuum tube.
• Observed rays emerging from the cathode that cast a shadow of a Maltese cross.
• Named these rays cathode rays.

J.J. Thomson (1897)

• Expanded on Crookes’ findings:   - Passed cathode rays through parallel metal plates.   - Observed the behavior in uncharged vs. charged scenarios.     - Uncharged plates: Undeflected ray.     - Charged plates: Ray deflected towards the positive plate.
• Conclusion: Cathode rays consist of negatively charged particles. Discovered the electron.

Plum Pudding Model

• Proposed in 1898 by J.J. Thomson, depicting an atom as a positively charged sphere with randomly embedded electrons.
• Model explained atomic neutrality but lacked empirical evidence.

The Nuclear Model of the Atom

Rutherford’s Experiment (1909)

• Studied alpha particle scattering by gold foil.   - Observations:     - Most alpha particles passed through (empty space).     - Some deflected at large angles (indicating a small, dense nucleus).     - Few reflected back (suggesting density and positive charge in the nucleus).
• Conclusion: Key characteristics of the nucleus:   - Small, positively charged, and dense.   - Proton discovered as a part of the nucleus.

Assumptions of the Nuclear Model

1. Atoms contain a small dense nucleus.
2. Most of the atom is empty space.
3. Electrons scattered around in an electron cloud surrounding the nucleus.

Limitations of the Nuclear Model

• Challenge: Like charges repel each other; thus, why does the nucleus not disintegrate?
• Additional query: How do electrons remain in stable orbits without spiraling into the nucleus?

James Chadwick: Discovery of the Neutron

• Bombarded beryllium with alpha particles, observing neutrons could displace protons.
• Negative charge of neutrons made them difficult to detect; utilized paraffin wax as a detector.
• Neutron discovered in this process.

Limitations of the Nuclear Model Post-Discovery

• Despite electrons moving around the nucleus, the model fails to clarify how they avoid crashing into the nucleus due to attraction.

The Bohr Model of the Atom

Niels Bohr's Contributions

• Expanded upon Rutherford's nuclear model, offering a clearer arrangement of electrons:   - Electrons revolve in fixed paths (orbits) around the nucleus, termed energy levels (n).   - Lowest energy level is identified as n = 1, followed sequentially by n = 2, etc.   - Energy within each orbit is quantized (fixed amount of energy).

Assumptions of Bohr’s Theory

• While in a fixed energy level, electrons neither lose nor gain energy, explaining their stability.
• Atoms, ideally, exist in the ground state with electrons at the lowest energy. Example: hydrogen at n = 1.
• Upon energy absorption, an electron can jump to a higher energy level, resulting in an excited state.

Energy Levels and Transitions

1. Energy Transition: Absorption leads to a higher state; emission releases energy as a photon when it reverts.
2. Energy relationship: $E_2 - E_1 = hf$ where $h$ is Planck's constant and $f$ frequency of emitted light.
3. Emission Line Spectra from these transitions are unique to each element.

Electron Transitions Explained

• Balmer Series (visible lines): Electrons fall to n=2.
• Paschen Series (infrared): Electrons to n=3.
• Lyman Series (ultraviolet): Electrons to n=1.

Unique Emission Spectra

• Each element has distinct electron configurations leading to unique transitions.

Energy Sublevels

• Emission spectra revealed complexities:   - Lines originally thought singular actually consist of closely spaced lines (as found by refined spectrometers).   - Introduced subdivisions: s, p, d, f sublevels that refine understanding of atomic structure.

Advanced Theories and Limitations

Louis De Broglie

• Proposed wave-particle duality; challenged fixed electronic paths.

Heisenberg’s Uncertainty Principle

• Confirmed that position and velocity of electrons cannot be simultaneously measured.
• Suggested rethinking of electron behavior in atoms, defining probabilities rather than certainties.

Limitations of Bohr’s Model

• Effective for single-electron systems like hydrogen but inadequate for complex atoms:   - Inability to account for sublevel presence.   - Failed to integrate wave-particle duality, leading to incorrect assumptions of fixed paths and energies.