The P-Block Elements (Group 17 - Halogens)

Group 17 Elements

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Group 17 elements, known as the halogens, include Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I), Astatine (At), and Tennessine (Ts). They are a highly reactive group of non-metals.

Occurrence

Fluorine (F):

• Most electronegative element.
• Does not occur in free state due to its high reactivity.
• Found in nature mainly as insoluble fluoride minerals, such as fluorspar ($CaF_2$), cryolite ($Na_3AlF_6$), and fluorapatite ($Ca_5(PO_4)_3F$).

Chlorine (Cl):

• Does not occur in free state.
• Found abundantly in nature as common salt ($NaCl$) in seawater and rock salt deposits.
• Also found as other chlorides like potassium chloride ($KCl$), magnesium chloride ($MgCl_2$).

Bromine (Br) and Iodine (I):

• Occur in nature primarily as bromides and iodides in seawater and some mineral springs.
• Iodine is also found in tanah shi ($NaIO_3$) and Chile saltpeter ($NaNO_3$ containing $NaIO_3$).

Astatine (At):

• A rare radioactive element, formed by the radioactive decay of heavier elements. It occurs in trace amounts in nature.

Tennessine (Ts):

• A synthetic, highly radioactive element.

Electronic Configuration

General Configuration: The general valence shell electronic configuration of Group 17 elements is $ns^2np^5$.

• F: $[He] 2s^22p^5$
• Cl: $[Ne] 3s^23p^5$
• Br: $[Ar] 3d^{10} 4s^24p^5$
• I: $[Kr] 4d^{10} 5s^25p^5$
• At: $[Xe] 4f^{14} 5d^{10} 6s^26p^5$
• Ts: $[Og] 7s^27p^5$

Significance: The presence of seven valence electrons, one short of a stable octet, explains their strong tendency to gain one electron to form $-1$ ions ($X^-$) or share one electron to form covalent bonds.

Atomic Radii

Trend: Atomic radii increase down the group from F to I.

Reasons:

• Increase in the number of electron shells.
• Increased shielding effect of inner electrons.

Comparison to Group 16: Atomic radii are smaller than the corresponding elements in Group 16 of the same period due to the higher effective nuclear charge.

Covalent Radii vs. van der Waals Radii: For non-metals, covalent radii are typically discussed. For halogens, due to the weak interatomic forces, van der Waals radii are also considered and are larger than covalent radii.

Ionisation Enthalpy

Trend: First ionization enthalpies generally decrease down the group from F to I.

Reasons:

• Increase in atomic size.
• Increased shielding effect.

Comparison to Group 16: First ionization enthalpies are generally lower than those of Group 16 elements in the same period. This is because halogens have one less electron to achieve the stable octet compared to Group 16 elements.

Anomalous Behavior of Fluorine: Fluorine has a higher first ionization enthalpy than expected compared to other halogens. This is due to the small size and the strong attraction between the nucleus and the single valence electron.

Electron Gain Enthalpy

Trend: Electron gain enthalpy becomes less negative (less favorable) down the group from Cl to I.

• F: -328 kJ/mol
• Cl: -349 kJ/mol
• Br: -325 kJ/mol
• I: -295 kJ/mol

Anomalous Behavior of Fluorine: Fluorine has a less negative electron gain enthalpy than chlorine. This is attributed to the electron-electron repulsion in the small $2p$ subshell of fluorine, which makes it slightly difficult for the incoming electron to be accommodated.

Significance: The highly negative electron gain enthalpies indicate the strong tendency of halogens to accept an electron to form stable halide ions ($X^-$) with a noble gas configuration.

Electronegativity

Trend: Electronegativity decreases down the group from F to I.

• F: 3.98
• Cl: 3.16
• Br: 2.96
• I: 2.66

Fluorine's Electronegativity: Fluorine is the most electronegative element, strongly attracting electrons.

Physical Properties

States of Matter:

• Fluorine ($F_2$): Pale yellow gas.
• Chlorine ($Cl_2$): Greenish-yellow gas.
• Bromine ($Br_2$): Reddish-brown volatile liquid.
• Iodine ($I_2$): Violet volatile solid.
• Astatine ($At$): Radioactive solid.
• Tennessine ($Ts$): Synthetic radioactive element.

Diatomic Molecules: They exist as diatomic molecules ($X_2$) with single covalent bonds.

Bond Strength: The bond dissociation enthalpies of the $X-X$ bond decrease from Cl to I ($Cl-Cl > Br-Br > I-I$). Fluorine ($F-F$) has an anomalously low bond dissociation enthalpy due to repulsion between the lone pairs on the small fluorine atoms.

Melting and Boiling Points: Melting and boiling points increase down the group due to increasing van der Waals forces with increasing molecular size and mass.

Solubility: Fluorine and Chlorine are moderately soluble in water. Bromine is slightly soluble. Iodine is sparingly soluble in water but more soluble in solvents like KI solution (forming triiodide ion, $I_3^-$) due to complex formation.

Chemical Properties

1. Reactivity: Reactivity decreases down the group from F to I.

• Fluorine: Most reactive non-metal, reacts with almost all other elements, including noble gases. It is a very strong oxidizing agent.
• Chlorine: Also highly reactive and a strong oxidizing agent.
• Bromine: Less reactive than chlorine.
• Iodine: Least reactive among the stable halogens.

2. Oxidizing Agent: Halogens are strong oxidizing agents, readily accepting an electron to form halide ions ($X^-$). Their oxidizing strength decreases down the group ($F_2 > Cl_2 > Br_2 > I_2$).

• $F_2$ can oxidize $O_2$ and $N_2$.
• $Cl_2$ can displace $Br^-$ and $I^-$ from their salts.
• $Br_2$ can displace $I^-$ from its salts.
• $I_2$ cannot displace $F^-$ or $Cl^-$.

3. Reaction with Hydrogen: React directly with hydrogen to form hydrogen halides ($HX$).

• $F_2 + H_2 \rightarrow 2HF$ (highly exothermic, explosive reaction)
• $Cl_2 + H_2 \xrightarrow{UV \ light \ or \ heat} 2HCl$
• $Br_2 + H_2 \rightleftharpoons 2HBr$ (reversible)
• $I_2 + H_2 \rightleftharpoons 2HI$ (reversible, less favorable)

4. Reaction with Metals: React vigorously with most metals to form metal halides.

5. Reaction with Non-metals: Form interhalogen compounds.

6. Formation of Halides: Form halides with elements in various oxidation states.

7. Bleaching Properties: Chlorine and bromine act as bleaching agents in the presence of moisture due to their oxidizing action. Iodine has weak bleaching properties.

8. Acidic Nature of $HX$: Hydrogen halides ($HF$, $HCl$, $HBr$, $HI$) are acidic in aqueous solution. Their acidic strength increases down the group ($HF < HCl < HBr < HI$) due to the decreasing bond strength of the $H-X$ bond.

• $HF$ is unique because it forms hydrogen bonds and etches glass ($SiO_2$).

$SiO_2(s) + 4HF(aq) \rightarrow SiF_4(g) + 2H_2O(l)$