Bond Strength, Oxidizing Power, and Hydrogen Reactions of Halogens

13.3 Bond Strength of Halogen Molecules

• General Trend in Bond Strength: The bond strength of halogen molecules generally decreases moving down Group 17 from chlorine to iodine. This trend is attributed to the increase in atomic size as one progresses down the group.     * Effect of Atomic Size: As atomic size increases, bond lengths become longer, which inherently results in weaker bonds.     * Trend in Bond Energies: The bond energies of halogens show a gradual decrease from chlorine ($Cl_2$) to iodine ($I_2$).

• Exception: Fluorine ($F_2$): Fluorine is a notable exception to the group trend. Although chlorine has a stronger bond than bromine or iodine, fluorine's bond strength is relatively weak.     * Small Atomic Size: Fluorine atoms are exceptionally small.     * Electron-Electron Repulsion: Because of the small size of the fluorine atoms, the lone pairs of electrons ($lp - lp$) on adjacent fluorine atoms are very close to each other. This results in significant electron-electron repulsion ($lp - lp$ repulsion), which weakens the $F-F$ bond despite fluorine's high electronegativity.

• Table 13.3: Bond Energies of Halogen Molecules:     * $F_2$: $156\,kJ\,mol^{-1}$     * $Cl_2$: $243\,kJ\,mol^{-1}$     * $Br_2$: $193\,kJ\,mol^{-1}$     * $I_2$: $151\,kJ\,mol^{-1}$

• Visual Representation: Figure 13.2 depicts the $F_2$ molecule illustrating the lone pairs of electrons contributing to repulsion.

13.4 Relative Reactivities of the Halogens as Oxidizing Agents

• Nature of Halogens as Oxidizing Agents: All free halogens act as oxidizing agents during reactions with metals and most non-metals. When halogens form ionic compounds with metals, they gain electrons (undergo reduction) and are converted into negative halide ions ($X^-$).     * Reaction involving Sodium: $2Na(s) + Cl_2(g) \rightarrow 2NaCl(s)$

• Trend in Oxidizing Power: The oxidizing power of halogens decreases moving down the group. Fluorine possesses the highest oxidizing power, while iodine possesses the least.     * Decreasing Order: $F_2 > Cl_2 > Br_2 > I_2$

• Relationship to Electron Acquisition: The reactivity of halogens is directly linked to their ability to acquire an electron and form halide ions ($F^-$, $Cl^-$, $Br^-$, and $I^-$). Fluorine has the highest tendency to acquire an electron to form the fluoride ion.

• Displacement Reactions: A more reactive halogen can oxidize and displace a less reactive halide ion from its solution.     * Fluorine: Can oxidize and displace all other halide ions ($Cl^-$, $Br^-$, and $I^-$) to produce free halogens.         * $F_2(g) + 2Cl^-(aq) \rightarrow 2F^-(aq) + Cl_2(g)$         * $F_2(g) + 2Br^-(aq) \rightarrow 2F^-(aq) + Br_2(g)$         * $F_2(g) + 2I^-(aq) \rightarrow 2F^-(aq) + I_2(g)$     * Chlorine: Can oxidize and displace bromide ($Br^-$) and iodide ($I^-$) ions.     * Bromine: Can oxidize and displace iodide ($I^-$) ions.     * Iodine: Cannot oxidize any other halide ion.

• Standard Electrode Potential ($E^\circ$): The oxidizing power is related to standard electrode potential values ($E^\circ(X_2/X^-)$). Higher (more positive) values indicate a stronger oxidizing agent.     * Table 13.4: Standard Electrode Potential $E^\circ(X_2/X^-)$:         * $F_2$: $+2.87\,V$         * $Cl_2$: $+1.36\,V$         * $Br_2$: $+1.07\,V$         * $I_2$: $+0.54\,V$     * Observations: The electrode potential becomes less positive from fluorine to iodine, reflecting the decrease in oxidizing power.

• Factors Governing Oxidizing Power: The oxidizing power of halogens depends on several energetic factors:     * Energy of Dissociation: The energy required to break the diatomic molecule.     * Electron Affinity: The energy change when an atom gains an electron.     * Hydration Energies: The energy released when ions are hydrated in water.     * Heats of Vaporization: Specifically relevant for $Br_2$ and $I_2$.     * Criteria for High Oxidizing Power: A halogen will have high oxidizing power if it has low energy of dissociation, high electron affinity, and higher hydration energy of its ions.

Questions & Discussion

• Quick Check 13.2:     * Question (a): Why is the $F-F$ bond weaker than the $Cl-Cl$ bond although fluorine is the most electronegative element?     * Answer (a): This is due to the small size of fluorine atoms. The high electron-electron repulsion between the lone pairs situated on the small fluorine atoms weakens the bond, whereas chlorine atoms are larger, reducing this repulsion.     * Question (b): Is the reaction between $NaCl(aq)$ and $F_2$ gas possible?         * i) Give reason whether yes or no.         * ii) If yes, write the equation for this reaction.     * Answer (b):         * i) Yes, the reaction is possible because fluorine is a stronger oxidizing agent than chlorine and can displace chloride ions ($Cl^-$) from the solution.         * ii) $F_2(g) + 2NaCl(aq) \rightarrow 2NaF(aq) + Cl_2(g)$     * Question (c): What is the relationship between the oxidizing power of halogens and their standard reduction potential values?     * Answer (c): Oxidizing power is directly proportional to standard reduction potential values. A higher, more positive $E^\circ$ value correlates with a greater oxidizing power.

13.5 Reactions of the Halogens with Hydrogen

• General Reaction: Halogen elements react with hydrogen to produce hydrogen halides ($HX$).     * General Equation: $H_2 + X_2 \rightarrow 2HX$ (where $X_2 = F_2, Cl_2, Br_2, I_2$).

• Product Properties: Hydrogen halides are colourless gases. When dissolved in water, they form hydrohalic acids.

• Trend in Reactivity with Hydrogen: Reactivity decreases down the group from fluorine to iodine ($F_2 > Cl_2 > Br_2 > I_2$).

• Reactions by Halogen Type:     * Fluorine ($F_2$): Reacts explosively with hydrogen even at low temperatures and in the dark.         * Equation: $H_2(g) + F_2(g) \rightarrow 2HF(g)$         * Acid Formation: Dissolving this gas in water forms hydrofluoric acid.     * Chlorine ($Cl_2$): Reacts readily with hydrogen in the presence of UV light ($hv$) or a spark.         * Equation: $H_2(g) + Cl_2(g) \xrightarrow{hv} 2HCl(g)$         * Acid Formation: Dissolving colourless $HCl$ gas in water forms hydrochloric acid.     * Bromine ($Br_2$): Reacts with hydrogen only upon heating.         * Properties: $HBr$ gas is less reactive than $HCl$ and $HF$. It forms a strong hydrobromic acid in water.         * Thermodynamics: This is an exothermic reaction.