Electronic Configuration — Revision Notes

⚡ 30-Second Revision

Aufbau Principle: — Fill lowest energy orbitals first ($1s < 2s < 2p < 3s < 3p < 4s < 3d ...$).
Pauli's Exclusion Principle: — Max 2 electrons per orbital, opposite spins ($\uparrow\downarrow$).
Hund's Rule: — For degenerate orbitals, fill singly with parallel spins first, then pair up.
Quantum Numbers: — $n$ (shell), $l$ (subshell, $0$ to $n-1$), $m_l$ (orbital orientation, $-l$ to $+l$), $m_s$ (spin, $\pm 1/2$).
Group 2 Configuration: — $[\text{Noble Gas}] ns^2$.
Electron Removal (Ions): — Remove from highest 'n' shell first (e.g., $4s$ before $3d$).
Exceptions: — Cr ($[Ar] 3d^5 4s^1$), Cu ($[Ar] 3d^{10} 4s^1$).

2-Minute Revision

Electronic configuration is the electron arrangement in orbitals, governed by three rules. The Aufbau principle dictates filling orbitals from lowest to highest energy, often remembered by the $(n+l)$ rule where $4s$ fills before $3d$.

Pauli's exclusion principle states that each orbital holds a maximum of two electrons with opposite spins. Hund's rule ensures that degenerate orbitals are first singly occupied with parallel spins before pairing.

Group 2 elements (alkaline earth metals) are characterized by their $ns^2$ valence configuration, meaning two electrons in their outermost 's' orbital. This configuration explains their tendency to lose these two electrons to form stable $+2$ ions.

Remember that for ions, electrons are removed from the outermost principal shell (highest 'n') first. Also, be aware of exceptions like Chromium and Copper, which achieve extra stability through half-filled or fully-filled d-orbitals.

5-Minute Revision

To master electronic configuration, start with the fundamental rules. The Aufbau principle guides the filling order: $1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s$. Remember the $(n+l)$ rule to confirm this order.

For example, $4s$ ($n+l=4$) fills before $3d$ ($n+l=5$). Pauli's exclusion principle limits each orbital to two electrons with opposite spins ($\uparrow\downarrow$). Hund's rule applies to degenerate orbitals (like $p_x, p_y, p_z$), stating that each orbital is singly occupied with parallel spins before any pairing occurs.

For instance, Nitrogen ($Z=7$) is $1s^2 2s^2 2p_x^1 2p_y^1 2p_z^1$, not $1s^2 2s^2 2p_x^2 2p_y^1$.

Group 2 elements (Be, Mg, Ca, Sr, Ba, Ra) all share the characteristic valence configuration of $ns^2$. For example, Calcium ($Z=20$) is $[Ar] 4s^2$. This $ns^2$ configuration explains their metallic nature and strong tendency to lose both valence electrons to form $+2$ ions ($M^{2+}$), achieving a stable noble gas configuration. For example, $Mg \rightarrow Mg^{2+} + 2e^-$.

Crucially, when forming cations, electrons are removed from the orbital with the highest principal quantum number ($n$) first. For instance, for Iron ($Z=26$, $[Ar] 3d^6 4s^2$), to form $Fe^{2+}$, the two electrons are removed from $4s$, resulting in $[Ar] 3d^6$.

Finally, be mindful of exceptions to the Aufbau principle, notably Chromium ($Z=24$: $[Ar] 3d^5 4s^1$ instead of $3d^4 4s^2$) and Copper ($Z=29$: $[Ar] 3d^{10} 4s^1$ instead of $3d^9 4s^2$). These exceptions occur due to the enhanced stability of half-filled ($d^5$) or fully-filled ($d^{10}$) subshells. Practice writing configurations for various elements and their ions, and correlate them with periodic properties.

Prelims Revision Notes

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Definition: — Electronic configuration is the distribution of electrons in atomic orbitals.
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Governing Principles:

* Aufbau Principle: Electrons fill orbitals in increasing order of energy. The order is $1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, ...$. Use the $(n+l)$ rule: lower $(n+l)$ fills first; if $(n+l)$ is same, lower 'n' fills first.

(e.g., $4s$ ($n+l=4$) before $3d$ ($n+l=5$)). * Pauli's Exclusion Principle: An orbital can hold a maximum of two electrons, and they must have opposite spins ($\uparrow\downarrow$). No two electrons in an atom can have identical sets of all four quantum numbers.

* Hund's Rule of Maximum Multiplicity: For degenerate orbitals (same energy, e.g., $p_x, p_y, p_z$), electrons occupy each orbital singly with parallel spins before any orbital is doubly occupied.

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Quantum Numbers:

* n (Principal): Energy level/shell ($1, 2, 3, ...$). * l (Azimuthal/Angular): Subshell shape ($0$ to $n-1$). $l=0 \rightarrow s, l=1 \rightarrow p, l=2 \rightarrow d, l=3 \rightarrow f$. * **$m_l$ (Magnetic):** Orbital orientation ($-l$ to $+l$). Number of orbitals in a subshell is $2l+1$. * **$m_s$ (Spin):** Electron spin ($\pm 1/2$).

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Group 2 Elements (Alkaline Earth Metals):

* Valence electronic configuration: $ns^2$. * Examples: Be ($[He] 2s^2$), Mg ($[Ne] 3s^2$), Ca ($[Ar] 4s^2$), Sr ($[Kr] 5s^2$), Ba ($[Xe] 6s^2$), Ra ($[Rn] 7s^2$). * Tendency: Readily lose 2 valence electrons to form stable $+2$ ions ($M^{2+}$), achieving noble gas configuration.

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Electronic Configuration of Ions:

* For cations, remove electrons from the orbital with the highest principal quantum number ($n$) first. If 'n' is the same, remove from the highest 'l'. (e.g., for transition metals, $4s$ electrons are removed before $3d$ electrons).

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Exceptions to Aufbau Principle:

* Chromium (Cr, Z=24): $[Ar] 3d^5 4s^1$ (not $3d^4 4s^2$) - due to stability of half-filled d-orbital. * Copper (Cu, Z=29): $[Ar] 3d^{10} 4s^1$ (not $3d^9 4s^2$) - due to stability of fully-filled d-orbital.

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Unpaired Electrons: — Determine by drawing orbital diagrams and applying Hund's rule. Important for magnetic properties (paramagnetic if unpaired, diamagnetic if all paired).

Vyyuha Quick Recall

To remember the Aufbau filling order (up to $5p$): Some People Say People Say Don't Play So Damn Poorly.

S — (1s)
S — (2s) P (2p)
S — (3s) P (3p)
S — (4s) D (3d) P (4p)
S — (5s) D (4d) P (5p)

For Group 2 elements: Better Make Careful Studies Baby Radiant. (Beryllium, Magnesium, Calcium, Strontium, Barium, Radium)