In simple terms
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Halogens: A Tale of Reactivity
Discover why halogens become less reactive as you descend Group 17, impacting the stability and acidity of the hydrogen halides. We'll also explore disproportionation, a unique redox reaction of chlorine.
Imagine a series of magnets trying to attract a loose paperclip (an electron). The first magnet (Fluorine) is small, its magnetic force is highly concentrated, and it grabs the paperclip with ease. As you move to bigger magnets down the line (Chlorine, Bromine, Iodine), their magnetic force is spread over a larger area and shielded by their own bulk, making them progressively weaker at attracting the paperclip. This is why fluorine is the most powerful oxidising agent.
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X₂ + 2e⁻ → 2X⁻ — oxidising power ↓ down group (F₂ most reactive).
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H–X bond enthalpy peaks at HF then ↓ — HF anomalously strong H-bonds.
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HX(aq) acid strength: HF < HCl < HBr < HI.
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Disproportionation: Cl₂ + OH⁻ → Cl⁻ + ClO⁻ (+ water) cold dilute.
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Full topic notes
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Halogen Reactivity and Oxidising Power
Reactivity for the halogens decreases as you go down the group. This is because their primary mode of reaction is by gaining an electron, acting as oxidising agents. Fluorine is the most powerful oxidising agent, and iodine is the weakest. This trend is explained by considering the factors affecting the attraction of an incoming electron: atomic radius and electron shielding. As you descend the group, the atomic radius increases, and there are more inner shells of electrons. This increased shielding and distance weakens the attraction between the positive nucleus and the electron being gained, making the element less reactive and a weaker oxidising agent.
General half-equation for halogen reduction:
This trend in oxidising power leads to a series of displacement reactions. A more reactive halogen will displace a less reactive halide ion from its salt solution.
(A colourless solution turns orange/brown)
(An orange/brown solution turns a darker brown)
Reactions such as do not occur as iodine is a weaker oxidising agent than chlorine.
Hydrogen Halides: Thermal Stability and Acidity
Hydrogen halides (HX) are formed when hydrogen reacts with halogens. Their thermal stability—their resistance to decomposition by heat—decreases down the group. This is directly related to the strength of the H-X bond. As the halogen atom size increases from F to I, the H-X bond length increases, and the bond enthalpy decreases. Consequently, less energy is required to break the H-I bond compared to the H-F bond, making hydrogen iodide the least thermally stable.
Example of thermal decomposition: (Purple iodine vapour is observed on heating).
In contrast, the acidity of the hydrogen halides in aqueous solution increases down the group. HF is a weak acid, while HCl, HBr, and HI are all strong acids, fully dissociating in water. While the H-F bond is the most polar, which might suggest it would be the most acidic, the deciding factor is the H-X bond enthalpy. The very high bond enthalpy of H-F means it is difficult to break the bond to release a proton () in solution. For HCl, HBr, and HI, the bonds are weaker, and they readily dissociate in water, making them strong acids.
Thermal stability trend:
Reason: H-X bond enthalpy decreases down the group.
Acidity trend (in water):
Reason: H-X bond enthalpy is the dominant factor; weaker bonds dissociate more easily.
A very common mistake is to confuse the trends for thermal stability and acidity. Remember: High bond energy means HIGH thermal stability but LOW acidity in solution. They are inversely related concepts for the hydrogen halides.
Disproportionation of Chlorine
Disproportionation is a specific type of redox reaction where an element from a single substance is simultaneously oxidised and reduced. Chlorine undergoes disproportionation when it reacts with cold, dilute aqueous sodium hydroxide. This reaction is important for the production of bleach.
Full Equation:
In this reaction, the oxidation state of chlorine changes from 0 in to -1 in sodium chloride (NaCl) and +1 in sodium chlorate(I) (NaClO). Since the oxidation state of chlorine has both decreased (reduction) and increased (oxidation), it is a disproportionation reaction. The ionic equation provides a clearer view of the species involved.
Ionic Equation:
Be precise with the conditions. The reaction of chlorine with hot, concentrated aqueous NaOH produces different products (chloride and chlorate(V), ). For AS level, the focus is on the cold, dilute conditions. Always state the conditions in your answer if they are relevant.
Worked examples
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A student mixes aqueous bromine, , with a solution containing both sodium chloride, , and potassium iodide, . Predict the observations and write the ionic equation for any reaction that occurs.
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Step 1: Identify the relative oxidising strengths. The order of oxidising power is . This means bromine is a stronger oxidising agent than iodine, but a weaker oxidising agent than chlorine. Step 2: Determine which displacement reactions are possible. Bromine can displace iodide ions () because it is more reactive. Bromine cannot displace chloride ions () because it is less reactive than chlorine. Step 3: State the observation. The initial solution containing the salts is colourless. When orange-brown aqueous bromine is added, the solution will turn a dark brown or black as iodine () is formed. The orange colour of the excess bromine may also be present. Step 4: Write the ionic equation. .
Explain why hydrogen chloride is a gas at room temperature but hydrogen fluoride is a liquid, despite HCl having a larger molar mass. Then, explain why HCl(aq) is a stronger acid than HF(aq).
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Part 1: Physical State. Fluorine is the most electronegative element, so the H-F bond is highly polar. This leads to the formation of strong hydrogen bonds between HF molecules. Hydrogen bonds are the strongest type of intermolecular force and require significant energy to overcome, resulting in HF having an unusually high boiling point (293 K) and being a liquid at room temperature. HCl has permanent dipole-dipole forces, which are weaker than hydrogen bonds. These weaker forces are easily overcome, so HCl is a gas at room temperature (boiling point 188 K). Part 2: Acidity. The strength of an acid in aqueous solution depends on the extent of its dissociation. The key factor for hydrogen halides is the H-X bond enthalpy. The H-F bond is very strong (568 kJ mol⁻¹) due to effective orbital overlap. This strong bond is difficult to break in water, so HF only partially dissociates and is a weak acid. The H-Cl bond is weaker (432 kJ mol⁻¹). This bond breaks much more readily in water, so HCl fully dissociates and is a strong acid. The strength of the bond enthalpy outweighs the effect of bond polarity in determining acidity.
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Define an oxidising agent in terms of electron transfer.
A species that accepts electrons from another species, causing the other species to be oxidised. The oxidising agent itself is reduced.
Key takeaways
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This trend in oxidising power leads to a series of displacement reactions. A more reactive halogen will displace a less reactive halide ion from its salt solution.
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(A colourless solution turns orange/brown)
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(An orange/brown solution turns a darker brown)
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Reactions such as do not occur as iodine is a weaker oxidising agent than chlorine.
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