Orbitals

• s-orbitals: These are spherical in shape and exist in every principal shell starting from $n=1$. As the principal quantum number increases, the s-orbital becomes larger and the electron spends more time further from the nucleus.

• p-orbitals: These have a dumbbell shape and appear from the second shell ($n=2$) onwards. There are always three p-orbitals in a p-subshell, oriented at right angles to each other along the x, y, and z axes ($p_x, p_y, p_z$).

• d and f orbitals: These are more complex in shape and appear in higher energy shells. A d-subshell contains five orbitals (total 10 electrons), while an f-subshell contains seven orbitals (total 14 electrons).

Comparison of Structural Levels

• Orbital vs. Subshell: A common point of confusion is the difference between a 'p-orbital' and a 'p-subshell'. A p-subshell is the collection of all three p-orbitals found in a specific energy level.

• The 4s vs. 3d Rule: Always remember that the 4s subshell is lower in energy than the 3d subshell and fills first. However, once the 3d subshell contains electrons, the 4s electrons are actually lost first during ionisation.

• Box Notation Accuracy: When drawing electrons in boxes, always fill every box in a subshell with one 'up' arrow before adding any 'down' arrows. This demonstrates your understanding of Hund's Rule.

• Spin Representation: Ensure your arrows in the same box point in opposite directions. Failing to show opposite spins is a frequent cause of lost marks regarding the Pauli Exclusion Principle.

• Sanity Check: If asked for the maximum electrons in a shell, use the $2n^2$ formula. For $n=3$, the answer is $2(3)^2 = 18$ electrons ($2$ in 3s, $6$ in 3p, $10$ in 3d).