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Loading home…Halogens are Group VIIA (17) elements: F, Cl, Br, I, At.
They exist as diatomic molecules X₂.
Fluorine and chlorine are gases; bromine is a liquid; iodine is a solid.
Cl₂ is greenish-yellow; Br₂ is red-brown; I₂ is dark grey/violet vapour.
Boiling points increase down the group as London forces strengthen.
They are non-metals with seven valence electrons (ns² np⁵).
Fluorine has the smallest atomic radius among the halogens.
It has the highest electronegativity and a very strong tendency to gain electrons.
The F–F bond is relatively weak compared with the strong bonds fluorine forms with other elements.
These factors make F₂ the strongest oxidizing agent in the group.
Reactivity of halogens decreases down the group: F₂ > Cl₂ > Br₂ > I₂.
Fluorine attacks many substances vigorously, including some that resist chlorine.
Oxidizing power decreases down the group.
A more reactive halogen displaces a less reactive halide from solution.
Cl₂ displaces Br⁻ and I⁻; Br₂ displaces I⁻; I₂ does not displace Cl⁻ or Br⁻.
Example: Cl₂ + 2KBr → 2KCl + Br₂.
This trend matches decreasing electronegativity and oxidizing strength.
Displacement reactions are classic halogen test-tube experiments.
H₂ + Cl₂ → 2HCl (explosive in sunlight; controllable with heat/light).
Chlorine water: Cl₂ + H₂O ⇌ HCl + HOCl (hypochlorous acid).
HOCl is responsible for bleaching by oxidation.
In sunlight, HOCl can decompose releasing oxygen.
Moist chlorine therefore bleaches litmus; dry Cl₂ does not bleach as effectively.
HCl(aq) is a strong acid; HOCl is a weak acid and oxidizing agent.
Chlorine: water disinfection, PVC manufacture, bleaching, solvents, and many organochlorine chemicals.
Iodine: thyroid hormone production (essential nutrient), antiseptics, and analytical iodometry.
Iodized salt prevents iodine-deficiency disorders.
Chlorine must be handled carefully because it is toxic.
Iodine sublimes on heating to a violet vapour.
Both elements illustrate useful but hazardous halogen chemistry.
Sublimation is the change of a solid directly to vapour without melting (under the conditions used).
Iodine crystals on gentle heating give a violet vapour.
On a cold surface the vapour deposits as shiny dark crystals.
This is used to purify iodine in the laboratory.
The process is reversible: vapour ⇄ solid.
It is a characteristic physical property of iodine among the common halogens.
Halogens occupy Group 17 with valence configuration ns² np⁵.
They need one electron to complete an octet, so they form X⁻ ions and covalent compounds.
Atomic and ionic radii increase down the group.
Electronegativity and oxidizing power decrease from F to I.
Melting and boiling points increase down the group; F₂/Cl₂ gases, Br₂ liquid, I₂ solid.
Colours deepen: pale yellow F₂, greenish-yellow Cl₂, red-brown Br₂, violet vapour of I₂.
Bond energy of X₂ varies; F–F is unexpectedly weak, contributing to fluorine’s reactivity.
They form hydrogen halides HX; acid strength in water increases HF < HCl < HBr < HI.
Critical point: HI is the strongest hydrohalic acid because H–I is the weakest bond.
HF is a weak acid in water due to the very strong H–F bond (and other aqueous factors).
Halogens show displacement reactions reflecting relative oxidizing strength.
They react with metals to form ionic halides and with non-metals to form covalent halides.
Interhalogen compounds also exist (e.g. ICl).
Exam answers should list physical states, trends, and at least two reaction types.
Emphasize the science lock on HX acid strength vs H–X bond strength.
Mention toxicity and careful handling, especially of Cl₂ and F₂.
Relate London forces to boiling-point trend of X₂ molecules.
Give fluorine’s extreme electronegativity as a key highlight.
Conclude with everyday links: PVC (Cl), Teflon (F), iodized salt (I).
Remember the board science lock: among HX acids, HI is strongest and HF is weak because of the strong H–F bond.
Displacement shows relative oxidizing power of F₂, Cl₂, Br₂ and I₂.
A halogen higher in the reactivity series oxidizes halide ions lower in the series.
Chlorine water added to KBr(aq) liberates bromine: Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂.
Chlorine also displaces iodide: Cl₂ + 2I⁻ → 2Cl⁻ + I₂ (brown/violet colour).
Bromine displaces iodide: Br₂ + 2I⁻ → 2Br⁻ + I₂.
Iodine cannot displace bromide or chloride under normal aqueous conditions.
Fluorine is so reactive it is not usually demonstrated in simple aqueous displacement labs.
Colour changes help identify the halogen set free.
These reactions are redox: halogen molecules are reduced to X⁻; halide ions are oxidized to X₂.
Standard electrode potentials rationalize the order F₂ > Cl₂ > Br₂ > I₂.
In qualitative analysis, chlorine water and starch are used in iodide tests.
Organic solvents can extract Br₂/I₂ giving characteristic layer colours.
Exam answers need word equations, ionic equations, and observations.
State clearly which halogen is more reactive in each pair.
Link to periodic trend of oxidizing ability down the group.
Warn that concentrated/reactive conditions can give other products.
Displacement is evidence-based ranking, not memorized colour alone.
Include one full worked example with states.
Summarise the rule: more reactive halogen displaces less reactive halide.
Remember the board science lock: among HX acids, HI is strongest and HF is weak because of the strong H–F bond.
Alkali metals react vigorously with halogens to form ionic salts MX.
Example: 2Na + Cl₂ → 2NaCl; similar reactions occur with F₂, Br₂, I₂.
The products are colourless crystalline ionic halides (unless impure).
With hydrogen, halogens form hydrogen halides: H₂ + X₂ → 2HX.
The reaction with fluorine is explosive; with chlorine it needs light/heat; iodine reacts less readily.
HX acidity in water: HI > HBr > HCl > HF (HF weak).
Do not confuse this order with H–X bond strength or with halogen reactivity.
Chlorine reacts with water to give HCl and HOCl; bleaching is due to oxidation by HOCl/O.
Bromine dissolves giving similar but weaker bleaching solutions.
Iodine is sparingly soluble in water but dissolves in KI(aq) as I₃⁻.
Fluorine reacts violently with water, producing oxygen and HF (not a simple lab demo).
Metal + halogen reactions are classic redox salt preparations.
Hydrogen halide gases dissolve to form hydrohalic acids.
Exam answers should give balanced equations for each class of reaction.
Stress safety: Cl₂ and F₂ are hazardous.
Relate aqueous bleaching to oxidation, not to “chlorine colour”.
Mention that dry Cl₂ does not bleach litmus effectively.
Connect HX acid strength science lock once more for board marks.
End with uses that follow from these reactions (salts, HCl industry, disinfection).
Remember the board science lock: among HX acids, HI is strongest and HF is weak because of the strong H–F bond.
Chlorine is prepared industrially by electrolysis of brine (chlor-alkali process).
In the laboratory, MnO₂ + concentrated HCl (heat) liberates Cl₂.
Cl₂ is a greenish-yellow, pungent, poisonous gas, denser than air.
It is moderately soluble in water, forming chlorine water (HCl + HOCl).
It supports neither combustion of ordinary fuels nor breathing; it is an oxidant.
With cold dilute NaOH: Cl₂ + 2NaOH → NaCl + NaOCl + H₂O (chloride + hypochlorite).
With hot concentrated NaOH: 3Cl₂ + 6NaOH → 5NaCl + NaClO₃ + 3H₂O (chloride + chlorate).
These are disproportionation reactions of chlorine.
NaOCl solutions are used in household bleach.
Chlorine disinfects drinking water and swimming pools.
It is used to make PVC, solvents, pesticides, and many intermediates.
Bleaching of moist coloured materials occurs by oxidation.
Cl₂ turns moist blue litmus red then bleaches it.
Exam answers should include prep, physical properties, NaOH reactions, and uses.
Write both cold and hot NaOH equations correctly.
State observations for chlorine water tests.
Mention toxicity and the need for fume control.
Link hypochlorite bleaching to oxidation chemistry.
Conclude with industrial importance of the chlor-alkali industry.
Remember the board science lock: among HX acids, HI is strongest and HF is weak because of the strong H–F bond.
Hydrogen halides are HX compounds: HF, HCl, HBr, HI.
Bond strength of H–X decreases down the group: H–F is strongest; H–I is weakest.
In aqueous solution, acid strength increases: HF ≪ HCl < HBr < HI.
HI is the strongest acid among the common hydrogen halides.
HF is a weak acid because the H–F bond does not ionize extensively in water.
High H–F bond energy and strong hydrogen bonding in HF(aq) contribute to this behaviour.
Students often wrongly pick HF as strongest because F is most electronegative — that is incorrect for aqueous acid strength.
Electronegativity predicts bond polarity, not the observed Ka order of HX(aq).
HCl, HBr and HI are strong acids and ionize essentially completely in dilute water.
HF etches glass (reaction with SiO₂) despite being a weak acid.
Boiling points of HX are anomalous for HF due to hydrogen bonding.
Preparation: NaCl + H₂SO₄ → HCl; similar routes exist for other HX with care.
Uses: HCl in industry and labs; HF in etching and fluoride chemistry; HI in organic reductions/iodinations.
Exam answers must state the acid-strength order and the bond-strength reason.
Give the science lock explicitly: strongest hydrohalic acid is HI, not HF.
Write dissociation: HX + H₂O → H₃O⁺ + X⁻ for strong acids.
Contrast with partial ionization of HF.
Mention that gas-phase acidity trends can be discussed separately if asked.
Finish with one clear warning against the fluorine–strongest-acid misconception.
Remember the board science lock: among HX acids, HI is strongest and HF is weak because of the strong H–F bond.