Atomic Structure and Quantum Numbers 4

Quantum Theory and the Electronic Structure of Atoms

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Chemistry Grade 11Teachers: Bader & BatoolDate: 9/23/2024

Page 2: Introduction to Atomic Structure

  • Atoms as Fundamental Units

    • All matter is composed of atoms.

    • Atoms are present in everything around us, including food, liquids, and even ourselves.

  • Historical Context

    • The modern concept of the atom began in the early 19th century with John Dalton's observations on element combinations.

Page 3: Discovery of Electrons

  • J.J. Thomson's Contribution

    • Discovered negatively charged particles (electrons) from metals under high voltage.

    • Proposed the "plum pudding" model, where electrons are embedded in a positively charged medium.

Page 4: Rutherford's Atomic Model

  • Rutherford's Findings (1911)

    • Introduced the concept of a dense, positively charged nucleus.

    • Proposed that electrons orbit the nucleus, similar to planets around the sun.

    • Suggested that most of the atom is empty space.

Page 5: Composition of the Atom

  • Nucleus and Electrons

    • Nucleus consists of protons and neutrons (nucleons).

    • Protons and neutrons have similar mass, while electrons are much lighter.

  • Particle Characteristics

    • Proton: Mass = 1, Charge = +1, Location = Nucleus

    • Electron: Mass = 1/1840, Charge = -1, Location = Electron cloud

    • Neutron: Mass = 1, Charge = 0, Location = Nucleus

Page 6: The Electromagnetic Spectrum

  • Forms of Electromagnetic Radiation

    • The spectrum includes various forms of energy, from low-energy radio waves to high-energy gamma rays.

    • Visible light is a small part of this spectrum.

Page 7: Key Terms in Electromagnetic Spectrum

  • Wavelength (λ)

    • Distance between consecutive crests or troughs.

  • Frequency (ν)

    • Number of waves passing a point per second.

Page 9: Line Emission Spectrum

  • Definition

    • A pattern of lines produced when light from glowing gas is separated through a prism.

Page 10: Absorption Spectrum

  • Definition

    • Spectrum of radiation that has passed through a medium, absorbing certain wavelengths.

Page 11: Relationship Between Wavelength and Frequency

  • Equation

    • ( c = ν \cdot λ ) (where ( c ) is the speed of light).

Page 12: Frequency Calculation Example

  • Example Calculation

    • For a wavelength of 650 nm, the frequency is calculated as:

      • Convert 650 nm to meters: ( 650 \times 10^{-9} m )

      • Calculate frequency: ( ν = \frac{c}{λ} = 4.62 \times 10^{14} Hz )

Page 13: Bohr’s Hydrogen Model

  • Quantization of Energy

    • Electromagnetic radiation exists in discrete packets called photons.

    • Energy of a photon is related to its wavelength and frequency.

Page 14: Bohr's Model Overview

  • Structure of Hydrogen Atom

    • Proton at the center with an electron in a circular orbit.

    • Each orbit has a fixed energy level.

Page 15: Energy Levels in Bohr's Model

  • Energy Expression

    • ( E = -R_H \left( \frac{1}{n^2} \right) ) where ( R_H ) is the Rydberg constant.

  • Electron Excitation

    • Electrons can move to higher energy levels when excited and emit photons when returning to lower levels.

Page 18: Emission Series

  • Hydrogen Emission Series

    • Lyman Series: n=1 (UV)

    • Balmer Series: n=2 (Visible and UV)

    • Paschen Series: n=3 (IR)

Page 21: Evolution of Atomic Theory

  • Changes in Atomic Structure Understanding

    • Bohr's model introduced the concept of energy levels and maximum electron capacity per shell:

      • 1st shell: 2 electrons

      • 2nd shell: 8 electrons

      • 3rd shell: 18 electrons

      • 4th shell: 32 electrons

      • 5th shell: 50 electrons

This note summarizes the key concepts and historical developments in atomic structure and quantum theory