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Topic 2: Atomic Structure

2.1 The Nuclear Atom
Essential Idea:
  • The mass of an atom is concentrated in its minute, positively charged nucleus.
  • Atoms contain a positively charged dense nucleus composed of protons and neutrons (nucleons).
  • Negatively charged electrons occupy the space outside the nucleus.
  • The mass spectrometer is used to determine the relative atomic mass of an element from its isotopic composition.
Applications & Skills:
  • Use of the nuclear symbol notation to deduce the number of protons, neutrons and electrons in atoms and ions.
  • Calculations involving non-integer relative atomic masses and abundance of isotopes from given data, including mass spectra.
Nature of Science:
  • Evidence and improvements in instrumentation—alpha particles were used in the development of the nuclear model of the atom that was first proposed by Rutherford.
  • Paradigm shifts—the subatomic particle theory of matter represents a paradigm shift in science that occurred in the late 1800s.
PPT: 2.1 Atomic Models
Atomic Structure Timeline
Video: Cathode Ray Tube Experiment (Thompson Model)
Video: Gold Foil Experiemtn (Rutherford)
2.1 Relative masses,charges and positions of the subatomic particles
2.1.3 Define mass number (A), atomic number (Z) and isotopes
The Bohr Model of the Atom
2.2 Electron Configuration
Essential Idea:
  • The electron configuration of an atom can be deduced from its atomic number.
  • Emission spectra are produced when photons are emitted from atoms as excited electrons return to a lower energy level.
  • The line emission spectrum of hydrogen provides evidence for the existence of electrons in discrete energy levels, which converge at higher energies.
  • The main energy level or shell is given an integer number, n, and can hold a maximum number of electrons, 2n2.
  • A more detailed model of the atom describes the division of the main energy level into s, p, d and f sub-levels of successively higher energies.
  • Sub-levels contain a fixed number of orbitals, regions of space where there is a high probability of finding an electron.
  • Each orbital has a defined energy state for a given electronic configuration and chemical environment and can hold two electrons of opposite spin.
Applications & Skills:
  • Description of the relationship between colour, wavelength, frequency and energy across the electromagnetic spectrum..
  • Distinction between a continuous spectrum and a line spectrum.
  • Description of the emission spectrum of the hydrogen atom, including the relationships between the lines and energy transitions to the first, second and third energy levels.
  • Recognition of the shape of an s atomic orbital and the px, py and pz atomic orbitals.
  • Application of the Aufbau principle, Hund’s rule and the Pauli exclusion principle to write electron configurations for atoms and ions up to Z = 36.
Nature of Science:
  • Developments in scientific research follow improvements in apparatus—the use of electricity and magnetism in Thomson’s cathode rays.
  • Theories being superseded—quantum mechanics is among the most current models of the atom.
  • Use theories to explain natural phenomena—line spectra explained by the Bohr model of the atom.
PPT: 2.2 Electron Configuration
Emission Spectra Short
Video: Emission Spectra of Atoms
Interactive Periodic Table of Emission Spectra
Virtual Flame Test Lab
Aufbau principle, Hund’s rule and the Pauli exclusion principle to write electron configurations
Electron Configuration and Orbital Notation