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2.8Unit 2. Energetics, Group Chemistry, Halogenoalkanes and Alcohols

Group 7 — IAL Edexcel Modular (Chemistry) Revision Notes & Topical Questions

5 min read 12 parts Lesson 8 of 12 Unit quiz

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2.8 Group 7 — IAL EDEXCEL MODULAR (CHEMISTRY) notes: Explain trends in physical properties and reactivity down Group 7, and describe displacement, redox and disproportionation reactions of the halogens and the reactions of halide ions.Halogens — F, Cl, Br, I (and At), with outer …

  • Melting and boiling points
  • Volatility
  • Electronegativity
  • Oxidising power
  • Reactivity with hydrogen
  • Observations

Section 1 of 12The halogens and their physical trends

1Halogens
  • F, Cl, Br, I (and At), with outer configuration ns2np5
  • They exist as diatomic molecules X2, which are simple molecular structures.
2Colour
  • Colour darkens down the group
  • Chlorine is a pale green gas, bromine a red-brown liquid and iodine a grey-black solid giving purple vapour.
  • Fluorine is a pale yellow gas.
3Melting and boiling points
  • Melting and boiling points increase down the group
  • Molecules with more electrons have stronger London forces, so more energy is needed to separate them.
4Volatility
  • Volatility decreases down the group
  • Fluorine is the most volatile halogen and iodine the least volatile.

Section 2 of 12Electronegativity and oxidising power

Electronegativity
  • Electronegativity decreases down the group
  • Atomic radius and shielding increase, so an incoming electron is attracted less strongly.
Oxidising power
  • Oxidising power decreases down the group: F2 > Cl2 > Br2 > I2
  • A halogen reacts by gaining an electron to form X−, so its oxidation number falls from 0 to −1 and it is reduced.
Reactivity with hydrogen
  • H2 + X2 → 2HX becomes less vigorous down the group
  • It is explosive with fluorine and slow and incomplete with iodine.

Section 3 of 12Halogen displacement reactions

Displacement
  • Displacement a more reactive halogen oxidises the halide ion of a less reactive halogen
  • Chlorine oxidises bromide and iodide, bromine oxidises iodide only, and iodine oxidises neither.
Equations
  • Cl2(aq) + 2Br−(aq) → 2Cl−(aq) + Br2(aq).
  • Cl2(aq) + 2I−(aq) → 2Cl−(aq) + I2(aq).
  • Br2(aq) + 2I−(aq) → 2Br−(aq) + I2(aq).
Redox
  • Redox the halogen is reduced (0 to −1) and the halide ion is oxidised (−1 to 0)
  • The metal ion, such as Na+ or K+, is a spectator ion.
Observations
  • Observations bromine gives yellow-orange and iodine gives brown in aqueous solution
  • After adding an organic solvent such as hexane, bromine is orange or yellow-orange and iodine is purple in the organic layer.
  • The halogens are more soluble in the organic layer, so the colours are easier to see.

Section 4 of 12Halogens as oxidising agents

With metals
  • With metals halogens oxidise Group 1 and 2 metals to ionic halides
  • 2Na + Cl2 → 2NaCl, with Na going from 0 to +1.
  • Ca + Br2 → CaBr2, with Ca going from 0 to +2.
With iron(II)
  • With iron(II) chlorine and bromine oxidise Fe2+ to Fe3+
  • Br2 + 2Fe2+ → 2Br− + 2Fe3+.
  • Iodine is too weak an oxidising agent, and Fe3+ instead oxidises iodide: 2Fe3+ + 2I− → 2Fe2+ + I2.

Section 5 of 12Disproportionation of chlorine

1Disproportionation
Disproportionation the same species is both oxidised and reduced
2Cold dilute alkali (about 15 °C)
  • Cold dilute alkali (about 15 °C) gives chloride and chlorate(I), the active ion in bleach
  • Cl2 + 2NaOH → NaCl + NaClO + H2O.
  • Reduction
    • Cl2 + 2e- → 2Cl−;
      • oxidation: Cl2 + 4OH− → 2ClO− + 2H2O + 2e-.
  • Chlorine goes from 0 to −1 and 0 to +1.
3Hot alkali (about 70 °C)
  • Hot alkali (about 70 °C) gives chloride and chlorate(V)
  • 3Cl2 + 6NaOH → 5NaCl + NaClO3 + 3H2O.
  • Chlorine goes from 0 to −1 and 0 to +5.
4Chlorine in water
  • Cl2 + H2O ⇌ HCl + HClO
  • Chloric(I) acid sterilises water by killing bacteria.
  • HClO partly ionises to H+ and ClO−, and both HClO and ClO− kill bacteria.

Section 6 of 12Halide ions with concentrated sulfuric acid

  • General reaction
    • H2SO4 + X− → HX + HSO4−, producing toxic gases, so a fume cupboard is needed
    • Chloride gives misty fumes of HCl and no further reaction.
HalideObservationsSulfur change
NaClMisty fumes of HClNone: stays +6
NaBrMisty fumes of HBr, reddish-brown Br2, choking SO2+6 to +4
NaIMisty fumes of HI, purple I2, SO2, yellow sulfur, H2S with a bad-egg smell+6 to +4, 0 and −2
Halide ions with concentrated sulfuric acid
Equations for bromide and iodide
  • 2HBr + H2SO4 → Br2 + SO2 + 2H2O.
  • 2HI + H2SO4 → I2 + SO2 + 2H2O.
  • 6HI + H2SO4 → 3I2 + S + 4H2O.
  • 8HI + H2SO4 → 4I2 + H2S + 4H2O.
Reducing power of halide ions
  • Reducing power of halide ions increases down the group: I− > Br− > Cl−
  • Larger halide ions have outer electrons held less strongly because of greater distance and shielding.
  • They lose electrons more easily and reduce other substances.
With concentrated phosphoric acid
  • With concentrated phosphoric acid a halide gives the hydrogen halide without oxidation

Section 7 of 12Test for halide ions

Test for halide ions
Acidify the solution with dilute nitric acid to remove carbonate ions.
Add silver nitrate solution, forming a precipitate of the silver halide: Ag+(aq) + X−(aq) → AgX(s).
Add dilute and then concentrated ammonia solution to the precipitate.
Halide ionPrecipitate colourDilute NH3Concentrated NH3
Cl−WhiteDissolvesDissolves
Br−CreamInsolubleDissolves
I−Pale yellowInsolubleInsoluble
  • Fluoride
    • AgF is soluble, so acidified silver nitrate cannot be used to test for fluoride

Section 8 of 12Hydrogen halides

Hydrogen halides
Hydrogen halides
  • Hydrogen halides acidic gases that form misty fumes in moist air
  • HCl(g) → H+(aq) + Cl−(aq) in water.
With ammonia
  • With ammonia form white smoke of ammonium halide: NH3 + HCl → NH4Cl
Predicting fluorine and astatine
  • Predicting fluorine and astatine use the group trends but watch for exceptions
  • Astatine should be a dark solid and the weakest oxidising halogen.
  • HAt should dissolve in water as an acid like HI.

Section 9 of 12More from the specification

Element
  • A substance made of only one type of atom.
Sulfur dioxide
  • An acidic gas made when sulfur-containing fuels burn.

Section 10 of 12More on this topic

  • Halide tests with silver nitrate
    • chloride is white, bromide cream and iodide yellow
    • Dilute nitric acid is added first to remove carbonates.

Section 11 of 12Key terms

TermMeaning
HalogenAn element in Group 7 existing as diatomic molecules X₂.
Halide ionThe X⁻ ion formed when a halogen atom gains one electron.
Displacement reactionA reaction in which a more reactive halogen oxidises the halide ion of a less reactive halogen.
DisproportionationA reaction in which the same element is simultaneously oxidised and reduced.
Oxidising agentA species that accepts electrons and is itself reduced.
Reducing agentA species that donates electrons and is itself oxidised.
Chlorate(I) ionClO⁻, the active ion in bleach.
VolatilityHow easily a substance evaporates.

Section 12 of 12Exam technique

1
Saying reactivity increases down Group 7 — oxidising power decreases down the group
2
Saying halide reducing power decreases down the group — I⁻ is the strongest reducing agent
3
Forgetting to acidify with nitric acid before adding silver nitrate — carbonates would also precipitate
4
Describing hot alkali products as chlorate(I) — hot alkali gives chlorate(V)
5
Using the same answer for bromide and chloride with concentrated sulfuric acid — only bromide and iodide reduce the acid

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