Electronic Configuration, Oxidation States — Revision Notes

⚡ 30-Second Revision

General Electronic Configuration — $ns^2np^5$
Fluorine (F) — Always -1 oxidation state. No d-orbitals, highest electronegativity.
Chlorine (Cl), Bromine (Br), Iodine (I) — 1, +1, +3, +5, +7 oxidation states.
Reason for positive states — Presence of vacant d-orbitals for electron promotion.
Oxidation State Calculation — Sum of oxidation states = 0 (neutral compound) or charge (ion).
Example — Cl in $HClO_4$ is +7; Br in $BrF_5$ is +5; I in $IF_7$ is +7.

2-Minute Revision

Halogens (Group 17) are characterized by their valence electronic configuration of $ns^2np^5$, meaning they have seven valence electrons. This makes them highly reactive, with a strong tendency to gain one electron to achieve a stable octet, resulting in a common -1 oxidation state.

Fluorine is unique among halogens; due to its exceptionally high electronegativity and the absence of vacant d-orbitals in its second shell, it *always* exhibits a -1 oxidation state. It cannot expand its octet.

In contrast, heavier halogens like Chlorine, Bromine, and Iodine possess accessible vacant d-orbitals (3d, 4d, 5d respectively). This allows them to promote electrons to these d-orbitals, increasing the number of unpaired electrons and enabling them to form multiple covalent bonds.

Consequently, they can exhibit positive oxidation states of +1, +3, +5, and +7, particularly when bonded to more electronegative elements like oxygen or fluorine (e.g., in oxyacids and interhalogen compounds).

The stability of these higher oxidation states varies, generally increasing down the group for oxyanions but with specific exceptions for interhalogens.

5-Minute Revision

Let's consolidate the key aspects of electronic configuration and oxidation states for Group 17 elements, the halogens. All halogens (F, Cl, Br, I, At, Ts) share a general valence electronic configuration of $ns^2np^5$. This signifies seven valence electrons, making them highly electron-deficient and thus very reactive. Their primary goal is to gain one electron to achieve a stable noble gas configuration ($ns^2np^6$), which leads to their most common oxidation state of -1.

Fluorine's Anomaly: Fluorine is the absolute outlier. Its configuration is $1s^22s^22p^5$. Crucially, it lacks any d-orbitals in its valence shell (n=2). Coupled with its position as the most electronegative element, fluorine *exclusively* exhibits a -1 oxidation state. It cannot promote electrons or expand its octet to show positive oxidation states. This is a frequently tested concept.

Heavier Halogens (Cl, Br, I) and Variable Oxidation States: Chlorine ($[Ne]3s^23p^5$), Bromine ($[Ar]3d^{10}4s^24p^5$), and Iodine ($[Kr]4d^{10}5s^25p^5$) all possess vacant d-orbitals in their respective valence shells.

This allows for 'octet expansion'. Electrons from the s and p orbitals can be unpaired and promoted to these vacant d-orbitals, increasing the number of available unpaired electrons for bonding. This mechanism enables them to form 1, 3, 5, or 7 covalent bonds, leading to positive oxidation states of +1, +3, +5, and +7.

+1 — In $HOCl$ (hypochlorous acid) or $ClF$.
+3 — In $HClO_2$ (chlorous acid) or $ICl_3$.
+5 — In $HClO_3$ (chloric acid) or $BrF_5$.
+7 — In $HClO_4$ (perchloric acid) or $IF_7$.

Stability Trends: The stability of higher oxidation states generally increases down the group for oxyanions (e.g., $IO_4^-$ is more stable than $ClO_4^-$). However, for interhalogen compounds, the maximum oxidation state achieved can be limited by steric factors and bond energies (e.

g., $IF_7$ exists, but $BrF_7$ does not). Always remember to calculate oxidation states accurately using the rule that the sum of oxidation states in a neutral compound is zero, and for ions, it equals the charge.

This topic is foundational for understanding halogen chemistry.

Prelims Revision Notes

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General Electronic Configuration — All Group 17 elements (halogens) have a valence shell electronic configuration of $ns^2np^5$. This means 7 valence electrons.
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Reactivity — Due to $ns^2np^5$, halogens are one electron short of a stable octet ($ns^2np^6$). They have a strong tendency to gain one electron, making them highly reactive non-metals and strong oxidizing agents.
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Fluorine's Unique Behavior

* Electronic Configuration: $1s^22s^22p^5$. No d-orbitals in the n=2 shell. * Electronegativity: Highest (4.0 on Pauling scale). * Oxidation State: Always -1 in all its compounds. Cannot exhibit positive oxidation states or expand its octet.

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Heavier Halogens (Cl, Br, I) - Variable Oxidation States

* Presence of vacant d-orbitals (3d for Cl, 4d for Br, 5d for I) allows for electron promotion. * Ground State: $ns^2np^5$ (1 unpaired electron) $\rightarrow$ can form 1 bond (e.g., $HCl$), oxidation state -1 or +1 (e.

g., $HOCl$). * Excited States: Electrons from p and s orbitals can be promoted to vacant d-orbitals, increasing unpaired electrons: * 1st excited state: 3 unpaired electrons $\rightarrow$ +3 oxidation state (e.

g., $ClF_3$, $HClO_2$). * 2nd excited state: 5 unpaired electrons $\rightarrow$ +5 oxidation state (e.g., $BrF_5$, $HClO_3$). * 3rd excited state: 7 unpaired electrons $\rightarrow$ +7 oxidation state (e.

g., $IF_7$, $HClO_4$).

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Oxidation State Calculation Rules

* Sum of oxidation states in a neutral compound is 0. * Sum of oxidation states in an ion equals the charge of the ion. * Common values: H = +1 (except in metal hydrides), O = -2 (except in peroxides, superoxides, $OF_2$).

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Stability Trends

* For oxyacids/oxyanions, the stability of the highest oxidation state (+7) generally increases down the group (e.g., $ClO_4^- < BrO_4^- < IO_4^-$). * For interhalogen compounds, the maximum oxidation state achieved by the central halogen can be influenced by steric factors (e.g., $IF_7$ exists, but $BrF_7$ does not).

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Key Compounds — Remember common examples for each oxidation state: $HOCl$ (+1), $HClO_2$ (+3), $HClO_3$ (+5), $HClO_4$ (+7), $ClF_3$ (+3), $BrF_5$ (+5), $IF_7$ (+7).

Vyyuha Quick Recall

F-1, Cl-Br-I: 'D' for Diversity! (D-orbitals allow diverse oxidation states)