Chapter 2: Atomic Structure and Periodicity Study Notes

Cengage Chemistry: An Atoms First Approach

Chapter 2: Atomic Structure and Periodicity

Historical Context and Key Figures
  • John Dalton (1803-1810): Early atomic theory. Proposed that:
    • Atoms are featureless spheres.
    • Differing elements differ in weight (e.g., Hydrogen = 1, Carbon = 5, Oxygen = 7… Lead = 207).
Modern Discoveries in Atomic Weights
  • The Cannizzaro Principle (1860):
    • The atomic weight of an element is the minimum weight in a volatile molecule.
    • Hydrogen's atomic weight is set as 1.
    • Example reactions:
    • $2H{2} + O{2} \rightarrow 2H_{2}O$.
Atomic Weights Established by J. Berzelius
  • List of atomic weights:
    • H = 1, Li = 7, Be = 9.4, B = 11, C = 12, N = 14, O = 16, F = 19…
    • Up to U = 240.
Döbereiner's Triads (1829)
  • Grouped elements demonstrating atomic weights could be averages:
    • E.g., Atomic weight of Br found to be an average of Cl and I:
    • $\frac{(35.470 + 126.470)}{2} \approx 80.470$.
Newlands' Law of Octaves (1865)
  • Grouped 56 known elements into 8 groups based on atomic weights and properties.
    • Validity limited to elements up to calcium.
The Periodic Table Development
  • Independently discovered by Dimitri Mendeleev and Lothar Meyer (1869):
    • Mendeleev arranged elements by atomic weight and chemical properties.
    • Meyer plotted atomic volumes against atomic weights, revealing periodicity with peaks corresponding to alkali metals.
Key Observations from Mendeleev’s Table
  1. When arranged by atomic weight, periodic properties emerge.
  2. Similar elements show closely related atomic weights.
  3. Common elements have lower atomic weights.
  4. Predictive power: Mendeleev’s table could suggest properties of undiscovered elements. Specifically:
    • Technecium (43), Promethium (61), etc.
    • Predictions reinforced the idea of periodicity.
Mendeleev's Notable Features:
  • Elements' order and gaps indicated undiscovered elements.
The Bohr Model
  • Introduced by Neils Bohr (1913) addressing earlier atomic models' limitations:
    • Model describes quantized electron orbits.
    • Energy levels defined as $E = - \frac{Z^{2}R_{H}}{n^{2}}$.
Quantum Theory Development
  • Max Planck quantized energy in 1900 with:
    • $E = h\nu$, where h = Planck’s constant.
  • Photoelectric Effect demonstrated light possesses particulate aspects.
    • Energy levels are quantized, explaining atomic stability and spectra.
  • Bohr's modifications led to a more complex understanding of the hydrogen atom's behavior in quantized terms, represented mathematically as:
    • $\Delta E = h
      u$ relates photon energy emitted/absorbed during electrons’ transitions.
Spectroscopy and Hydrogen's Emission Spectrum
  • Only specific wavelengths are emitted due to quantized energy levels:
    • Hydrogen emission lines fit for series of quantized states (Lyman, Balmer, etc.).
    • Rydberg formula utilized:
      1λ=R<em>H(1n</em>12−1n<em>22)\frac{1}{\lambda} = R<em>H \left(\frac{1}{n</em>{1}^2} - \frac{1}{n<em>{2}^2}\right) where $RH$ is Rydberg constant for hydrogen.
Quantum Mechanics Incorporation
  • Heisenberg Uncertainty Principle instructs that position and momentum cannot both be accurately measured, aiding the understanding of domains in atomic structure.
  • Solutions from Schrödinger’s equations yield orbitals defined by three quantum numbers (n, l, m_l).
Electron Configuration Principles
  • Aufbau Principle: States orbitals fill from lowest to highest energy.
  • Pauli Exclusion Principle: States that each orbital can hold a maximum of two electrons with opposite spins.
  • Hund’s Rule: Electrons fill degenerate orbitals singly with parallel spins before pairing.
Periodicity in the Periodic Table
  • Trends observed in atomic and ionic sizes.
    • Atomic radius decreases across a period due to increased Zeff.
    • Atomic radius increases down groups due to increasing n, which influences distance from nucleus.
  • Ionization energy trends track closely with atomic properties influenced by electron configuration and Zeff:
    • Ionization energy typically increases across periods but decreases down groups.
Ionic Sizes and Behavior
  • Cations are smaller than their neutral atoms; anions are larger due to electron-electron repulsions and shielding effects.
  • Values for ionization, electronegativity, and other reactive measures highlight how structures influence behavior in chemical equations.