Atomic Theory

Explain how cathode-ray observations supported the existence of electrons.

Describe Thomson’s plum-pudding model of the atom.

Explain how Millikan’s oil-drop experiment established the electron’s charge.

Use the alpha-scattering results to infer that atoms contain small, dense nuclei.

Describe Rutherford’s nuclear model of the atom.

Explain how evidence for the neutron resolved a limitation of the early nuclear model.

Compare the relative charge, mass, and location of protons, neutrons, and electrons.

Explain why atomic number, rather than atomic mass, defines an element.

Interpret a nuclide symbol in terms of protons, neutrons, and electrons.

Distinguish isotopes of an element using particle diagrams and nuclide notation.

Explain conceptually how a mass spectrometer provides evidence for isotopes.

Interpret relative-abundance diagrams to explain why atomic masses are weighted averages.

Represent the formation of positive and negative ions as electron loss or gain.

Distinguish an atom, an isotope, and an ion.

Explain qualitatively how nuclear attraction and proton repulsion affect nuclear stability.

Distinguish alpha, beta, and gamma radiation by composition, charge, and penetrating ability.

Represent alpha and beta decay using balanced nuclear symbols.

Interpret a half-life graph as evidence of probabilistic radioactive decay.

Distinguish a nuclear transformation from a chemical reaction.

Explain nuclear fission as the splitting of a heavy nucleus and nuclear fusion as the joining of light nuclei.

Compare common benefits and risks of radioactive isotopes using atomic-scale reasoning.

Relate wavelength and frequency using visual representations of electromagnetic waves.

Explain the relationship between light frequency and photon energy without calculus.

Explain how the photoelectric effect supports the particle-like description of light.

Interpret a line spectrum as evidence that atoms have discrete energy levels.

Use an energy-level diagram to explain atomic emission and absorption.

Describe the Bohr model of hydrogen in terms of quantized electron energies.

Identify observations that the Bohr model explains and phenomena it cannot explain.

Explain wave-particle duality for electrons using the electron-diffraction experiment.

Explain qualitatively why the uncertainty principle rules out precise electron trajectories.

Distinguish a quantum orbital from a classical orbit.

Interpret an electron probability cloud as a map of likely electron locations.

Distinguish shells, subshells, and orbitals using nested visual models.

Identify the shapes and orientations represented by s and p orbitals.

Explain the meaning of electron spin and the two-electron capacity of an orbital.

Apply the Aufbau principle to order orbitals by filling energy.

Apply the Pauli exclusion principle when placing electrons in orbitals.

Apply Hund’s rule when filling orbitals of equal energy.

Construct ground-state electron configurations for representative elements.

Construct orbital-box diagrams for representative elements.

Distinguish core electrons from valence electrons in an electron configuration.

Use an electron configuration to locate an element’s period and broad region of the periodic table.

Infer recurring chemical behavior from shared valence-electron patterns.

Explain shielding and effective nuclear charge using a visual atomic model.

Explain the general trend in atomic radius across a period and down a group.

Explain the general trend in first ionization energy across a period and down a group.

Explain the general trend in electronegativity across a period and down a group.

Predict whether a representative atom tends to gain or lose electrons.

Compare the sizes of an atom and its common ion.

Use atomic structure to explain why periodic trends have exceptions.

Explain chemical bonding as a lower-energy arrangement of valence electrons and nuclei.

Predict ionic or covalent bonding from the valence behavior of the participating atoms.

Represent simple ionic compounds as repeating lattices of oppositely charged ions.

Represent electron sharing in simple covalent molecules with Lewis structures.

Use electronegativity differences to explain bond polarity.

Explain metallic bonding using a model of delocalized valence electrons.

Relate ionic, molecular covalent, and metallic structures to characteristic physical properties.

Compare the Dalton, Thomson, Rutherford, Bohr, and quantum-mechanical models by their structures and evidence.

Explain why scientific models of the atom changed without concluding that earlier evidence was useless.

Explain how atomic spectroscopy can identify an element.

Distinguish microscope images of atomic-scale surfaces from literal pictures of isolated atoms.

Correct common misconceptions about atomic emptiness, electron motion, and probability clouds.

Construct a visual evidence chain linking experimental observations to the modern atomic model.

Deliver a coherent explanation of atomic theory that connects history, evidence, structure, periodic behavior, and bonding.