Electronic Configuration — Revision Notes

⚡ 30-Second Revision

Aufbau Principle: — Fill lowest energy orbitals first. Order: $1s, 2s, 2p, 3s, 3p, 4s, 3d, \dots$ (use $(n+l)$ rule). \n- Pauli's Exclusion Principle: Max 2 electrons per orbital, with opposite spins ($+\frac{1}{2}, -\frac{1}{2}$). \n- Hund's Rule: For degenerate orbitals, fill singly with parallel spins before pairing. \n- Exceptions: Cr ($[Ar] 3d^5 4s^1$), Cu ($[Ar] 3d^{10} 4s^1$) due to stability of half-filled/fully-filled d-orbitals. \n- Ions: For cations, remove electrons from highest 'n' shell first (e.g., $4s$ before $3d$ for transition metals). For anions, add to lowest available orbital. \n- Magnetic Properties: Paramagnetic (unpaired electrons), Diamagnetic (all paired electrons).

2-Minute Revision

Electronic configuration describes how electrons are arranged in an atom's orbitals. This arrangement is governed by three fundamental rules. The Aufbau principle dictates that electrons fill orbitals in increasing order of energy, often following the $(n+l)$ rule (e.

g., 4s fills before 3d). Pauli's Exclusion Principle states that no two electrons in an atom can have the same set of four quantum numbers, meaning each orbital can hold a maximum of two electrons with opposite spins.

Hund's Rule of Maximum Multiplicity specifies that for degenerate orbitals (like the three p-orbitals), electrons will first occupy each orbital singly with parallel spins before any pairing occurs, maximizing stability.

\n\nCrucial exceptions exist for elements like Chromium ($[Ar] 3d^5 4s^1$) and Copper ($[Ar] 3d^{10} 4s^1$), where an electron from the 4s orbital is promoted to the 3d orbital to achieve the enhanced stability of a half-filled ($d^5$) or completely filled ($d^{10}$) d-subshell.

For ions, remember that electrons are removed from the outermost shell (highest 'n' value) first when forming cations, especially for transition metals where 4s electrons are removed before 3d electrons.

The presence of unpaired electrons determines if a species is paramagnetic (attracted to a magnetic field) or diamagnetic (repelled).

5-Minute Revision

Electronic configuration is the systematic distribution of electrons in an atom's orbitals, which is key to understanding chemical behavior. The process is guided by three main principles: \n\n1. Aufbau Principle: Electrons fill orbitals from lowest to highest energy.

The $(n+l)$ rule helps determine this order: lower $(n+l)$ means lower energy. If $(n+l)$ values are equal, the orbital with lower 'n' is filled first. For example, 4s ($n+l=4$) fills before 3d ($n+l=5$).

\n2. Pauli's Exclusion Principle: Each atomic orbital can hold a maximum of two electrons, and these two electrons must have opposite spins ($+\frac{1}{2}$ and $-\frac{1}{2}$). This ensures every electron in an atom has a unique quantum state.

\n3. Hund's Rule of Maximum Multiplicity: For degenerate orbitals (orbitals of the same energy within a subshell, e.g., $p_x, p_y, p_z$), electrons will first occupy each orbital singly with parallel spins before any pairing occurs.

This minimizes electron-electron repulsion and maximizes exchange energy, leading to greater stability. For instance, for $2p^3$, it's $(\uparrow \_ ) (\uparrow \_ ) (\uparrow \_ )$, not $(\uparrow \downarrow \_ ) (\uparrow \_ ) (\_ \_ )$.

\n\nExceptions: The most important exceptions are Chromium (Cr, Z=24) and Copper (Cu, Z=29). Instead of the expected configurations ($[Ar] 3d^4 4s^2$ and $[Ar] 3d^9 4s^2$ respectively), they adopt configurations that provide extra stability due to half-filled ($d^5$) or completely filled ($d^{10}$) d-subshells: \n* Cr: $[Ar] 3d^5 4s^1$ \n* Cu: $[Ar] 3d^{10} 4s^1$ \n\nIons: \n* Cations: Electrons are removed.

For main group elements, from the outermost shell. For transition metals, electrons are removed from the $ns$ orbital *before* the $(n-1)d$ orbital, even though $ns$ was filled first (e.g., $Fe^{2+}$ is $[Ar] 3d^6$, not $[Ar] 3d^4 4s^2$).

\n* Anions: Electrons are added to the lowest available energy orbital. \n\nMagnetic Properties: \n* Paramagnetic: Contains one or more unpaired electrons (attracted to a magnetic field). \n* Diamagnetic: All electrons are paired (repelled by a magnetic field).

To determine this, draw orbital diagrams for the valence shell and apply Hund's rule.

Prelims Revision Notes

1
Electronic Configuration Notation: — Shorthand like $1s^2 2s^2 2p^6$. The superscript denotes the number of electrons in that subshell. Noble gas core notation (e.g., $[Ne]$) simplifies writing for larger atoms.\n2. Quantum Numbers: \n * $n$: Principal (shell, energy, size). $n=1, 2, 3, \dots$\n * $l$: Azimuthal (subshell, shape). $l=0 (s), 1 (p), 2 (d), 3 (f)$. Range: $0$ to $n-1$.\n * $m_l$: Magnetic (orbital orientation). Range: $-l$ to $+l$. Number of orbitals in a subshell is $2l+1$.\n * $m_s$: Spin (electron spin). $+\frac{1}{2}$ or $-\frac{1}{2}$.\n3. Rules for Filling Orbitals:\n * Aufbau Principle: Fill orbitals in increasing order of energy. Use the $(n+l)$ rule: lower $(n+l)$ is lower energy. If $(n+l)$ is same, lower 'n' is lower energy. \n * Order: $1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f < 5d < 6p < 7s < 5f < 6d < 7p$.\n * Pauli's Exclusion Principle: Maximum two electrons per orbital, with opposite spins.\n * Hund's Rule of Maximum Multiplicity: For degenerate orbitals (same subshell), fill each orbital singly with parallel spins before pairing electrons.\n4. Stability of Half-filled and Completely Filled Orbitals: $p^3, d^5, f^7$ (half-filled) and $p^6, d^{10}, f^{14}$ (completely filled) subshells are exceptionally stable due to: \n * Symmetry: Symmetrical distribution of electrons. \n * Exchange Energy: Maximized for parallel spins in degenerate orbitals.\n5. Exceptions: \n * Chromium (Cr, Z=24): $[Ar] 3d^5 4s^1$ (not $3d^4 4s^2$). \n * Copper (Cu, Z=29): $[Ar] 3d^{10} 4s^1$ (not $3d^9 4s^2$).\n6. Electronic Configuration of Ions: \n * Cations (positive ions): Remove electrons from the outermost shell (highest 'n' value) first. For transition metals, remove $ns$ electrons before $(n-1)d$ electrons. \n * Example: $Fe (Z=26): [Ar] 3d^6 4s^2 \implies Fe^{2+}: [Ar] 3d^6$. \n * Anions (negative ions): Add electrons to the lowest available energy orbital.\n7. Magnetic Properties: \n * Paramagnetic: Presence of one or more unpaired electrons. Attracted by a magnetic field. \n * Diamagnetic: All electrons are paired. Repelled by a magnetic field. \n * To determine, draw orbital diagrams for the valence shell and count unpaired electrons using Hund's rule.

Vyyuha Quick Recall

To remember the Aufbau filling order: 'Some People Don't Follow' (for s, p, d, f blocks). \nFor the $(n+l)$ rule: 'Nice Little Elephants' (N for n, L for l, E for Energy - lower sum, lower energy). \nFor Hund's Rule: 'Happy Hunters Have Houses' (Each orbital gets one electron 'house' before 'pairing' up).