Why does Fluorine have a less negative electron gain enthalpy than Chlorine, despite being more electronegative?

This is a classic exception to the general trend. While Fluorine is indeed the most electronegative element, its exceptionally small atomic size leads to a very high electron density in its $2p$ subshell. When an incoming electron attempts to enter this already crowded shell, it experiences significant inter-electronic repulsion from the existing electrons. This repulsion partially counteracts the attraction from the nucleus, making the overall energy released (electron gain enthalpy) less negative compared to Chlorine. Chlorine, being larger, has a more diffuse electron cloud, reducing this repulsion and allowing for a more efficient accommodation of the incoming electron, thus releasing more energy.

Explain the trend in bond dissociation enthalpy for halogens, especially the anomaly of Fluorine.

Generally, as atomic size increases down a group, bond length increases, and bond dissociation enthalpy decreases. So, one might expect F-F to have the highest bond dissociation enthalpy. However, the observed trend is $Cl_2 > Br_2 > F_2 > I_2$. The F-F bond is unusually weak. This anomaly is attributed to the small size of fluorine atoms. The lone pairs of electrons on each fluorine atom are in close proximity, leading to strong inter-electronic repulsion. This repulsion weakens the F-F bond significantly, making it easier to break compared to the Cl-Cl and Br-Br bonds where such repulsion is less pronounced due to larger atomic size.

How does the physical state of halogens change down the group, and what is the reason?

At room temperature, Fluorine ($F_2$) and Chlorine ($Cl_2$) are gases, Bromine ($Br_2$) is a volatile liquid, and Iodine ($I_2$) is a solid. This systematic change is due to the increasing strength of van der Waals forces (specifically London dispersion forces) as we move down the group. As the atomic size and the number of electrons in the halogen molecule increase, the polarizability of the electron cloud increases, leading to stronger instantaneous dipole-induced dipole interactions. These stronger intermolecular forces require more energy to overcome, resulting in a progressive increase in melting and boiling points, and thus a change from gaseous to liquid to solid state.

Why do halogens show varying oxidation states, and which halogen is an exception?

Halogens typically exhibit a -1 oxidation state in most compounds because they readily gain one electron to complete their octet. However, except for Fluorine, other halogens (Cl, Br, I) can exhibit positive oxidation states such as +1, +3, +5, and +7. This is possible because these heavier halogens have vacant d-orbitals in their valence shell. Electrons from the p-orbitals, and even s-orbitals, can be promoted to these empty d-orbitals, allowing for the formation of more covalent bonds and thus higher positive oxidation states, especially when bonded to more electronegative elements like oxygen or fluorine. Fluorine, lacking d-orbitals, can only show a -1 oxidation state.

Describe the trend in oxidizing power of halogens.

Halogens are strong oxidizing agents, meaning they readily accept electrons from other species. Their oxidizing power decreases as we move down the group from Fluorine to Iodine. This trend is directly related to their decreasing tendency to gain an electron (less negative electron gain enthalpy) and decreasing electronegativity. A halogen higher in the group can oxidize halide ions of halogens lower in the group. For example, $F_2$ can oxidize $Cl^-$, $Br^-$, and $I^-$ ions, while $Cl_2$ can oxidize $Br^-$ and $I^-$ ions, but not $F^-$ ions. This demonstrates Fluorine as the strongest oxidizing agent among the halogens.

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Definition Deep Explanation Core Principles NEET Importance Problem Strategy Practice Problems MCQ Practice Quick Revision

The Group 17 elements, collectively known as halogens, exhibit distinct and predictable trends in their physical and chemical properties as one moves down the group from fluorine to astatine. These trends are primarily governed by the increasing atomic number, which leads to a systematic increase in atomic size and nuclear charge, accompanied by a corresponding increase in the number of electron s…

Quick Summary

Group 17 elements, the halogens (F, Cl, Br, I, At), are highly reactive non-metals with $ns^2np^5$ valence electron configurations. As we move down the group, atomic and ionic radii increase due to the addition of new electron shells.

This leads to a decrease in ionization enthalpy and electronegativity. Electron gain enthalpy generally becomes less negative, but Fluorine is an exception, having a less negative value than Chlorine due to inter-electronic repulsion in its small size.

Bond dissociation enthalpy also shows an anomaly, with $F_2$ being weaker than $Cl_2$ and $Br_2$ due to lone pair repulsion. Physical states change from gas (F, Cl) to liquid (Br) to solid (I) due to increasing van der Waals forces, which also cause melting and boiling points to rise.

Density and colour intensity also increase down the group. Chemically, reactivity and oxidizing power decrease from F to I. Fluorine only shows a -1 oxidation state, while others can exhibit positive states (+1, +3, +5, +7) due to vacant d-orbitals.

Understanding these trends and their exceptions is crucial for NEET.

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Key Concepts

Electron Gain Enthalpy Anomaly of Fluorine

While halogens generally have highly negative electron gain enthalpies (exothermic process), Fluorine's value…

Bond Dissociation Enthalpy Trend and

F_2

Exception

For diatomic halogen molecules ( $X_2$ ), the bond dissociation enthalpy generally decreases as atomic size…

Oxidizing Power and Displacement Reactions

Halogens are strong oxidizing agents, meaning they readily accept electrons. Their oxidizing power decreases…

Atomic/Ionic Radii: — Increase down group (

F < Cl < Br < I

Ionization Enthalpy: — Decrease down group (

F > Cl > Br > I

Electron Gain Enthalpy: — Generally less negative down group; Anomaly:

Cl > F > Br > I

(magnitude).

Electronegativity: — Decrease down group (

F > Cl > Br > I

). F is highest.

Bond Dissociation Enthalpy ($X_2$): — Generally decrease; Anomaly:

Cl_2 > Br_2 > F_2 > I_2

Physical State: — Gas (

F_2, Cl_2

)

ightarrow

Liquid (

Br_2

)

ightarrow

Solid (

I_2

Melting/Boiling Points: — Increase down group.

Colour: — Pale yellow (

F_2

)

ightarrow

Greenish-yellow (

Cl_2

)

ightarrow

Reddish-brown (

Br_2

)

ightarrow

Violet/Dark grey (

I_2

). Intensity increases.

Oxidation States: — 1 for all; Cl, Br, I also +1, +3, +5, +7 (due to d-orbitals). F only -1.

Reactivity: — Decrease down group (

F_2 > Cl_2 > Br_2 > I_2

Oxidizing Power: — Decrease down group (

F_2 > Cl_2 > Br_2 > I_2

Thermal Stability (HX): — Decrease down group (

HF > HCl > HBr > HI

Acidic Strength (HX): — Increase down group (

HF < HCl < HBr < HI

For the anomalous bond dissociation enthalpy of halogens, remember: 'Clever Brothers Fail Instead' for the order $Cl_2 > Br_2 > F_2 > I_2$ .

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