General Chemistry 2: (Learn Visually Through Animation)

In this course, we present a systematic introduction to the fundamental principles of quantum mechanics that form the basis of modern atomic theory. The course begins with the Heisenberg uncertainty principle, which establishes the intrinsic limits on simultaneously determining the position and momentum of an electron, emphasizing the departure from classical descriptions of particle motion. This is followed by the Schrödinger equation, the central equation of quantum mechanics, which provides a mathematical framework for describing the wave-like behavior of electrons in atoms. By solving the Schrödinger equation for atomic systems, students learn how quantized energy levels arise and how these solutions lead to four quantum numbers—principal, angular momentum, magnetic, and spin—that uniquely describe an electron’s energy, orbital shape, spatial orientation, and intrinsic spin.

The next chapter is devoted to atomic orbitals, focusing on both their geometric shapes and their relative energy levels. The characteristics and graphical representations of the s, p, d, and f orbitals are examined in detail, highlighting differences in symmetry and electron density distribution. Orbital energies are first analyzed for the hydrogen atom, where orbitals sharing the same principal quantum number are degenerate. The discussion is then extended to many-electron atoms, where electron–electron repulsion and shielding effects break this degeneracy, leading to the experimentally observed ordering of orbital energies and providing a foundation for understanding atomic structure and periodic trends.