Group 7 Tests, NH4 +, OH– , CO3 2– , SO42–

iodine colour and state room temp

grey-black solid, purple when sublimed

iodine + acidified AgNO3

pale yellow ppt AgI

chlorine colour & state at room temp.

pale green gas

bromine colour and state at room temp.

red liquid gives off orangey-brown fumes

fluorine colour and state at room temp.

pale yellow gas

does melting point increase or decrease going down the group and why?

increases - molecules become larger, more electrons in the van der waals = stronger IMFs

does electronegativity increase or decrease going down the group and why?

electronegativity decreases - going down the group the no. of shells increases so more repulsion → positive nucleus is further from electron

are halogens good or bad oxidising agents going down the group and why?

• oxidising power decreases going down the group

• halide ions are larger, weaker attraction of positive nucleus to electron it tries to gain

• a halogen that is a stronger oxidising agent (more readily reduced) will replace one that is weaker in a compound

chloride ions + acidified AgNO3

white ppt of AgCl

bromide ions + acidified AgNO3

cream ppt of AgBr

why is the AgNO3 acidified?

to remove any unreacted anions

silver bromide + dilute NH3

doesn’t dissolve still cream ppt

silver chloride + dilute NH3

dissolves colourless solution

silver iodide + dilute NH3

doesn’t dissolve still yellow ppt

silver iodide + conc. NH3

doesn’t dissolve still yellow ppt

silver bromide + conc. NH3

dissolves colourless solution

why isn’t HCl used to acidify AgNO3?

would form a white ppt

NaCl (chloride ions) + conc. H2SO4 observations & equation

white steamy HCl fumes

NaCl(s) + H₂SO₄(l) → NaHSO₄(s) + HCl(g)

→ acid base reaction, Cl- is not strong enough reducing agents to reduce S in H₂SO₄

NaBr (bromide ions) + conc. H2SO4 observations and equations

white steamy HBr fumes, oranges bromine fumes, SO2 colourless gas

acid-base step: NaBr(s) + H₂SO₄(l) → NaHSO₄(s) + HBr(g)

redox step: 2H⁺ + 2Br^- + H₂SO₄ → Br₂(g) + SO₂(g) + 2H₂O(l) reduction product

NaI (iodide ions) + conc. H2SO4

white steamy HI fumes, black solid & purple iodine fumes, H2S gas and bad egg smell, SO2 colourless gas, sulfur yellow solid

acid-base step: NaI(s) + H₂SO₄(l) → NaHSO₄(s) + HI (g) redox equations & reduction products:

2H⁺ + 2I^- + H2SO4 → I₂(s) + SO₂(g) + 2H2O(l)

6H⁺ + 6I^- + H2SO4 → 3I₂ + S(s) + 4H2O(l)

8 H⁺ + 8 I^- + H2SO4 → 4 I₂ (s) + H₂S(g) + 4 H2O(l)

chlorine + water to form an acidic solution

Cl2 (g) + H2O (l) ⇌ HClO (aq) + HCl (aq)

chlorine + water (in sunlight)

2Cl₂ + 2H2O → 4HCl + O2

why is it alright to use chlorine in water treatment?

kills bacteria, benefits outweigh toxic effects/risks

chlorine + cold dilute NaOH

Cl2 (aq) + 2 NaOH (aq) → NaCl (aq) + NaClO (aq) + H2O (l)

test for carbonate ions

add HCl (any dilute acid) → observe effervescence

bubble through limewater to test for CO2 → turns cloudy

test for presence of OH- ions

turns damp red litmus paper blue

test for presence of SO₄^2- ions

BaCl2+ HCl to acidify → white ppt of BaSO4

test for presence of NH₄⁺ ions

10 drops NH4Cl + 10 drops NaOH shake

warm mixture in water bath test fumes

fumes bleach damp red litmus paper → blue