Chapter 2: Electrons in Atoms

Electrons in Atoms chemistry chapter delves into the detailed arrangement of electrons within an atom, moving beyond simple planetary models to explore principal quantum shells (designated by 'n'), sub-shells, and atomic orbitals. Electrons are restricted to existing within these discrete energy levels. Each principal shell is divided into sub-shells, categorized as s, p, d, and f, where the energy of electrons typically increases in the order s (lesser than) p (lesser than) d within a shell. An atomic orbital is defined as the three-dimensional region around the nucleus where there is a high probability of finding one or two electrons. The maximum electron capacity is determined by the number of orbitals in each sub-shell: one orbital for s (holding 2 electrons), three for p (holding 6 electrons), and five for d (holding 10 electrons). Structurally, s orbitals are spherical, while p orbitals are represented as dumbbell-shaped lobes arranged orthogonally along three axes. Establishing the ground state electronic configuration involves filling these sub-shells sequentially according to increasing energy, noting exceptions like potassium, where the 4s sub-shell fills before the 3d sub-shell. Furthermore, electrons fill orbitals within the same sub-shell individually before pairing up, minimizing repulsion, a process known as spin-pair repulsion. This detailed configuration helps define a free radical as a species containing one or more unpaired electrons. A core concept connecting electronic structure to elemental properties is ionization energy (IE), defined as the minimum energy required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous positive ions. Successive ionization energies (IE1, IE2, etc.) increase progressively, with significant jumps indicating that an electron is being removed from a principal quantum shell located much closer to the nucleus, enabling the deduction of an element's shell structure and group placement. The magnitude of ionization energy is influenced by four factors: the magnitude of the nuclear charge, the distance of the outer electron from the nucleus, the degree of shielding by inner electron shells, and spin-pair repulsion. These factors also explain periodic trends: IE generally increases across a period because increasing nuclear charge pulls electrons closer, while shielding remains constant. IE decreases down a group because the outer electron is further away and shielding increases, outweighing the nuclear charge increase. The chapter also addresses atomic and ionic radii trends, which are similarly explained by nuclear attraction and shielding. Finally, when d-block elements form ions, they exhibit a unique behavior where electrons are lost from the 4s sub-shell before the 3d sub-shell, despite the 4s filling first.