Electronic Configuration — Revision Notes

⚡ 30-Second Revision

Definition: — Distribution of electrons in orbitals.
Rules:

* Aufbau: Fill lowest energy orbitals first ($1s < 2s < 2p < 3s < 3p < 4s < 3d dots$). * Pauli: Max 2 electrons/orbital, opposite spins ($m_s = pm 1/2$). * Hund's: Maximize unpaired electrons in degenerate orbitals.

Alkali Metals (Group 1):

* General EC: $[Noble,Gas] ns^1$ * Valence electrons: 1 (in 's' orbital) * Key Property: Readily lose 1 electron to form $M^+$ (achieve noble gas configuration). * Consequences: Low ionization enthalpy, high electropositivity, +1 oxidation state, high reactivity.

2-Minute Revision

Electronic configuration is the systematic arrangement of electrons in an atom's orbitals, governed by three key principles: the Aufbau principle (electrons fill lowest energy orbitals first), Pauli's exclusion principle (each orbital holds max two electrons with opposite spins), and Hund's rule (degenerate orbitals are singly occupied before pairing).

For Group 1 elements, the alkali metals (Li, Na, K, Rb, Cs, Fr), their defining characteristic is the $ns^1$ valence shell configuration. This means they possess a single, loosely held electron in their outermost 's' orbital.

This configuration is responsible for their extremely low first ionization enthalpies, making them highly electropositive and strong reducing agents. They readily lose this single electron to form stable unipositive ions ($M^+$) with a noble gas configuration, exhibiting a characteristic +1 oxidation state.

Reactivity increases down the group as the valence electron becomes easier to remove due to increasing atomic size and shielding.

5-Minute Revision

Electronic configuration is the blueprint of an atom's electron arrangement, crucial for understanding its chemical behavior. It's built upon quantum numbers ($n, l, m_l, m_s$) defining orbital properties and three rules:

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Aufbau Principle: — Electrons occupy orbitals in increasing order of energy. Remember the $(n+l)$ rule: lower $(n+l)$ fills first; if equal, lower 'n' fills first. Example: $4s$ ($n+l=4$) fills before $3d$ ($n+l=5$).
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Pauli Exclusion Principle: — Each orbital can hold a maximum of two electrons, which must have opposite spins ($+1/2$ and $-1/2$). No two electrons in an atom can have the same four quantum numbers.
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Hund's Rule: — For degenerate orbitals (same energy, e.g., $p_x, p_y, p_z$), electrons fill singly with parallel spins first, before pairing up. This maximizes stability.

Alkali Metals (Group 1): These elements (Li, Na, K, Rb, Cs, Fr) are defined by their unique valence electronic configuration: $ns^1$. This means they have one electron in their outermost 's' orbital.

Lithium (Li, Z=3): — $1s^2 2s^1$ or $[He] 2s^1$
Sodium (Na, Z=11): — $1s^2 2s^2 2p^6 3s^1$ or $[Ne] 3s^1$
Potassium (K, Z=19): — $1s^2 2s^2 2p^6 3s^2 3p^6 4s^1$ or $[Ar] 4s^1$

Implications of $ns^1$:

Low Ionization Enthalpy: — The single valence electron is loosely held and easily removed, leading to very low first ionization energies. This value decreases down the group due to increasing atomic size and shielding.
High Electropositivity: — They readily lose this electron to form stable $M^+$ ions with a noble gas configuration, making them highly electropositive and strong reducing agents.
Oxidation State: — They exclusively exhibit a +1 oxidation state, as removing a second electron would require breaking into a stable noble gas core, which is energetically prohibitive.
Reactivity: — Their high tendency to lose the valence electron makes them extremely reactive, with reactivity increasing down the group.

Example: Why is the second ionization enthalpy of Sodium much higher than its first ionization enthalpy?

First IE: — $Na ([Ne] 3s^1) \rightarrow Na^+ ([Ne]) + e^-$. This involves removing the single $3s^1$ valence electron, which is relatively easy due to shielding and distance from the nucleus.
Second IE: — $Na^+ ([Ne]) \rightarrow Na^{2+} ([He] 2s^2 2p^5) + e^-$. This involves removing an electron from a very stable, completely filled noble gas configuration ($[Ne]$). This requires a tremendous amount of energy, hence the second ionization enthalpy is exceptionally high.

Prelims Revision Notes

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Electronic Configuration (EC): — The arrangement of electrons in atomic orbitals. It dictates an element's chemical properties.
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Quantum Numbers:

* $n$ (Principal): Energy level/shell ($1, 2, 3, dots$). * $l$ (Azimuthal): Subshell/shape ($0 dots n-1$; $0=s, 1=p, 2=d, 3=f$). * $m_l$ (Magnetic): Orbital orientation ($-l dots +l$). * $m_s$ (Spin): Electron spin ($pm 1/2$).

1
Rules for Filling Orbitals:

* Aufbau Principle: Fill lowest energy orbitals first. Order: $1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s$. * Pauli Exclusion Principle: Max 2 electrons per orbital, with opposite spins. * Hund's Rule: For degenerate orbitals, fill singly with parallel spins before pairing.

1
Alkali Metals (Group 1): — Li, Na, K, Rb, Cs, Fr.

* General EC: $[Noble,Gas] ns^1$. (e.g., Li: $[He] 2s^1$, Na: $[Ne] 3s^1$, K: $[Ar] 4s^1$). * Valence Electrons: One electron in the outermost 's' orbital. * **Key Properties due to $ns^1$:** * Low First Ionization Enthalpy: Easy to remove the single valence electron.

Decreases down the group. * High Electropositivity: Strong tendency to lose electron and form positive ions. * Oxidation State: Exclusively +1 (losing one electron gives stable noble gas configuration).

* Reactivity: Very high, increases down the group. * Metallic Character: High, increases down the group.

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Important Note: — The second ionization enthalpy of alkali metals is extremely high because it involves removing an electron from a stable, noble gas core configuration.
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Practice: — Be able to write EC for elements up to Z=30, identify elements from EC, and correlate EC with periodic trends (size, IE, reactivity).

Vyyuha Quick Recall

All People Have Single Valence Electrons (for Alkali Metals):

Aufbau
Pauli
Hund's
S — orbital (valence electron)
Valence electron (one)
Electrons (easily lost)