Group 7 Flashcards
Halogens observed colours
F2= pale yellow gas
CL2= Aq- pale green Non polar- pale green
Br2= Aq- yellow/brown. Non polar- orange/brown
I2= Aq- brown Non polar- purple/solid black
Describe the structure and bonding in the halogens, explain why the melting points increase down the group
Simple covalent molecules.
Intermolecular forces existing between halogen molecules are Van der Waals forces as the diatomic molecules are non-polar.
I2 is a bigger molecule with more electrons so the van der waal forces are larger and harder to overcome.
Why is fluorine the most electronegative halogen
A shortest atomic radius and the least shielding will help withdraw electrons more strongly.
Chemical reactions of halogens
Halogens get reduced, so are oxidising agents
Displacement Reactions
(Reacting Chlorine with Sodium Bromide)
(Reacting Chlorine with Potassium Fluorine)
the element HIGHEST UP Group 7 should form a HALIDE ION, whereas the halogen LOWEST DOWN Group 7 should form a HALOGEN MOLECULE
Reacting Chlorine with Sodium Bromide
Cl2(aq) + 2NaBr(aq) → Br2(aq) + 2NaCl(aq)
= Cl2 + 2Br- → Br2 + 2Cl-
Reacting Chlorine with Potassium Fluorine
Cl2(aq) + 2KF(aq) → No Reaction
If Br2 forms as a product the mixture will turn a brown/orange colour.
If I2 is produced and black solid will form or purple solution will be seen.
Halide Ions
Halide ions are oxidised, so are reducing agents
Fluoride
F- is not a powerful enough reducing agent, and so will undergo an acid-base reaction with H2SO4.
NaF(s) + H2SO4(aq) → NaHSO4(s) + HF(g)
Observations: steamy fumes of Hydrogen Fluoride gas is release. HF gas is extremely dangerous and is used to etch glass.
Chloride
Cl- is also not a powerful enough reducing agent, and so will also undergo an acid-base reaction with H2SO4.
NaCl(s) + H2SO4(aq) → NaHSO4(s) + HCl(g)
Observations: Steamy fumes of Hydrogen Chloride gas is release
Bromide
Stronger after initial acid-base reaction, bromide ions reduce S in H2SO4 from +6 to +4 in SO2 (H2SO4 =oxidising agent)
Acid-base step: NaBr(s) + H2SO4(aq) → NaHSO4 (s) + HBr(g)
Redox: 2H+ + 2Br- + H2SO4 → Br2(g) + SO2 (g) + 2H2O(l)
Overall: 2NaBr + 3H2SO4 → 2NaHSO4 + SO2 + Br2 + 2H2O(l)
Observations: Acidic gas formed and brown fume of Br2 gas appears
Iodide
Strongest reducing agent.
Reduces Sulfur from +6 in H2SO4 to +4 in SO2 to 0 in S and -2 in H2S
Acid-base reaction: NaI(s) + H2SO4(aq) → NaHSO4(s) + HI(g)
a) Reducing S from +6 to +4 in SO2
2H+ + H2SO4 + 2I- → SO2(g) + 2H2O(l) + I2(s)
An acidic gas forms and a black solid of Iodine forms
b) Reducing to 0 in element S
6H+ + H2SO4 + 6I- → S(s) + 4H2O(l) + 3I2(s)
Yellow solid sulfur produced and black solid of Iodine forms
c) Reducing to -2 in H2S
8H+ + H2SO4 + 8I- → H2S (g) + 4H2O(l) + 4I2(s)
Gas smelling of bad eggs and black solid of Iodine forms
Testing for Halide Ions
- Add Nitric acid (HNO3) to remove carbonate ions
- CO32- + 2H+ → H2O + CO2
Ammonium ions
NH4+ + OH- → NH3 + H2O
NH3 turns litmus blue
Uses of chlorine
Low conc. safe for humans, kills bacteria
Reaction with water:
Cl2(g) + H2O(l) → HClO(aq) + HCl(aq)
- CL2 goes through disproportionation
- (HClO) is an oxidising agent and kills bacteria by oxidising them. HClO is also a bleach.
- chlorine is used to treat drinking water, as well as used to keep swimming pools clean.
In presence of bright sunlight:
2Cl2(g) + 2H2O(l) → 4HCl(aq) + O2(q) = Pale green to colourless
- chlorine is rapidly lost from outdoor swimming pools and so needs to be constantly added.
Alternative:
NaClO(s) + H2O(l) ⇌ Na+(aq) + -OH(aq) + HClO(aq)
Reaction with Alkali:
Cl2(g) + 2NaOH(aq) → NaClO(aq) + NaCl(aq) + H2O(l)
- NaClO = ox ag used in household bleach
