# 12.2 Reactivity of Halogens This section explores the chemical reactivity of the halogens (Group 17 elements), focusing on their trends as oxidizing agents, their reactions with hydrogen, and the properties of the resulting hydrogen halides and halide ions. ## Reactivity and Oxidizing Power of Halogens The reactivity of halogens is defined by their ability to gain an electron to form a halide ion (e.g., $F^-$, $Cl^-$, $Br^-$, $I^-$). This tendency is highest for fluorine and decreases down the group. **Reactivity Trend:** Reactivity decreases down Group 17. This is because electronegativity decreases and atomic size increases, making it less favorable for larger atoms to attract and gain an electron. **Oxidizing Power:** Halogens are strong oxidizing agents because they readily accept electrons. Their oxidizing strength follows the same trend as their reactivity. The order of decreasing oxidizing power is: $$F_2 > Cl_2 > Br_2 > I_2$$ A more reactive halogen can displace a less reactive halide ion from its aqueous solution. **Displacement Reactions:** - Fluorine is the most reactive halogen and can displace chloride, bromide, and iodide ions. $$F_{2(g)} + 2KCl_{(aq)} \rightarrow 2KF_{(aq)} + Cl_{2(g)}$$ $$F_{2(g)} + 2KBr_{(aq)} \rightarrow 2KF_{(aq)} + Br_{2(l)}$$ $$F_{2(g)} + 2NaI_{(aq)} \rightarrow 2NaF_{(aq)} + I_{2(s)}$$ - Chlorine can displace bromide and iodide ions. - Bromine can only displace iodide ions. - Iodine is the least reactive and cannot displace any other halide ions. --- ## Reactions of Halogens with Hydrogen Halogens react with hydrogen gas to form gaseous hydrogen halides ($HX$). **General Reaction:** $$H_2 + X_2 \rightarrow 2HX \quad \text{(where } X = F, Cl, Br, I\text{)}$$ The vigor of this reaction decreases down the group as the reactivity of the halogen decreases. - **Fluorine ($F_2$):** Reacts explosively with hydrogen, even in cold and dark conditions. - **Chlorine ($Cl_2$):** Reacts explosively with hydrogen in the presence of sunlight or a flame. - **Bromine ($Br_2$):** Reacts with hydrogen only upon ignition, causing a mild explosion. - **Iodine ($I_2$):** Reacts with hydrogen partially and reversibly, even with continuous heating, forming an equilibrium. --- ## Properties of Hydrogen Halides ($HX$) ### Thermal Stability Thermal stability refers to the ability of a compound to resist decomposition upon heating. The stability of hydrogen halides depends on the strength of the H-X bond. **Trend:** The H-X bond strength decreases down the group as the halogen atom size increases and electronegativity decreases. **Order of Stability:** The energy required to break the H-X bond decreases from HF to HI. $$HF > HCl > HBr > HI$$ - HF is the most thermally stable, and HI is the least. ### Acidic Strength of Halogen Acids In aqueous solution, hydrogen halides form hydrohalic acids. Their strength is determined by how easily they can donate a proton ($H^+$). **Trend:** The acidic strength is inversely related to the H-X bond strength. A weaker bond allows for easier dissociation of the proton. **Order of Acidic Strength:** $$HI > HBr > HCl > HF$$ - HI is the strongest hydrohalic acid, and HF is the weakest. --- ## Reactivity of Halide Ions ($X^-$) Halide ions can act as reducing agents by losing an electron. Their reducing power increases down the group. **Reducing Power Trend:** As the ion size increases down the group, the attraction between the nucleus and the valence electrons weakens, making it easier for the ion to lose an electron. $$I^- > Br^- > Cl^- > F^-$$ - Iodide ($I^-$) is the strongest reducing agent, and fluoride ($F^-$) is the weakest. ### Reactions of Halide Ions with Concentrated Sulfuric Acid The reaction depends on the reducing strength of the halide ion. - **Fluoride ($F^-$) and Chloride ($Cl^-$):** These ions are weak reducing agents and do not reduce concentrated $H_2SO_4$. They undergo a simple acid-base reaction. $$\mathrm{NaF} + \mathrm{H_2SO_4} \rightarrow \mathrm{NaHSO_4} + \mathrm{HF}$$ $$\mathrm{NaCl} + \mathrm{H_2SO_4} \rightarrow \mathrm{NaHSO_4} + \mathrm{HCl}$$ - **Bromide ($Br^-$):** Strong enough to reduce concentrated $H_2SO_4$ to sulfur dioxide ($\mathrm{SO_2}$). $$\mathrm{NaBr} + \mathrm{H_2SO_4} \rightarrow \mathrm{NaHSO_4} + \mathrm{HBr}$$ $$2\mathrm{HBr} + \mathrm{H_2SO_4} \rightarrow \mathrm{Br_2} + \mathrm{SO_2} + 2\mathrm{H_2O}$$ - **Iodide ($I^-$):** The strongest reducing agent, capable of reducing concentrated $H_2SO_4$ to hydrogen sulfide ($\mathrm{H_2S}$). $$\mathrm{NaI} + \mathrm{H_2SO_4} \rightarrow \mathrm{NaHSO_4} + \mathrm{HI}$$ $$8\mathrm{HI} + \mathrm{H_2SO_4} \rightarrow 4\mathrm{I_2} + \mathrm{H_2S} + 4\mathrm{H_2O}$$ ### Reaction with Silver Nitrate (Silver Nitrate Test) This is a classic qualitative test to identify halide ions in an aqueous solution. Adding silver nitrate ($\mathrm{AgNO_3}$) solution produces a silver halide precipitate ($\mathrm{AgX}$) with a characteristic color and solubility in ammonia. **General Reaction:** $$\mathrm{Ag}^+_{(aq)} + \mathrm{X}^-_{(aq)} \rightarrow \mathrm{AgX}_{(s)}$$ | Halide Ion | Precipitate Color | Solubility in Aqueous Ammonia | | :--- | :--- | :--- | | **Chloride ($Cl^-$)** | White | Soluble | | **Bromide ($Br^-$)** | Cream (Off-white) | Partially soluble | | **Iodide ($I^-$)** | Yellow | Insoluble | ### Reaction of Chlorine with Sodium Hydroxide Chlorine undergoes a disproportionation reaction with sodium hydroxide, where it is simultaneously oxidized and reduced. The products depend on temperature and concentration. - **Cold, Dilute NaOH:** Forms sodium hypochlorite ($\mathrm{NaOCl}$). $$\mathrm{Cl_2} + 2\mathrm{NaOH} \rightarrow \mathrm{NaOCl} + \mathrm{NaCl} + \mathrm{H_2O}$$ - **Hot, Concentrated NaOH:** Forms sodium chlorate ($\mathrm{NaClO_3}$). $$3\mathrm{Cl_2} + 6\mathrm{NaOH} \rightarrow \mathrm{NaClO_3} + 5\mathrm{NaCl} + 3\mathrm{H_2O}$$ ### Reaction of Chlorine with Water Chlorine dissolves in water to form a mixture of hydrochloric acid ($\mathrm{HCl}$) and hypochlorous acid ($\mathrm{HOCl}$). $$\mathrm{Cl_2} + \mathrm{H_2O} \rightleftharpoons \mathrm{HCl} + \mathrm{HOCl}$$ Hypochlorous acid is a weak acid and a powerful oxidizing agent. It decomposes to produce nascent oxygen, which kills bacteria. This is the basis for using chlorine to disinfect water (chlorination). $$\mathrm{HOCl} \rightarrow \mathrm{HCl} + [O]$$ In sunlight, hypochlorous acid decomposes to form oxygen gas. $$2\mathrm{HOCl} \rightarrow 2\mathrm{HCl} + \mathrm{O_2}$$