Group 7 halogens Flashcards
(27 cards)
Colour of halogen gases-F2 Cl2 Br2 and I2-
Fluorine- very pale yellow gas. It is highly reactive
Chlorine: greenish, reactive gas, poisonous in high concentrations
Bromine: brown orange liquid, that gives off dense brown/orange poisonous fumes
Iodine: shiny grey solid sublimes to purple gas. can be brown when is displaced witheither bromine or chlorine
Trend in melting point and boiling point of halogens (4)
- Increase down the group
- As the molecules become larger they have more electrons and so have larger van der waals forces between the molecules.
- As the intermolecular forces get larger more energy has to be put into break the forces.
- This increases the melting and boiling points
Trend in electronegativity for halogens
- Electronegativity is the ability for an atom to attract electrons towards itself in a covalent bond.
- This decreases as we go down the group. The
atoms get larger and the distance between the positive nucleus and bonding electrons increases. There is also more shielding. - The atomic radii increases down the group due to the increasing number of shells.
- The nucleus is therefore less able to attract the bonding pair of electrons.
- The displacement reactions of halide ions by halogens.
A halogen that is a strong oxidising agent will
displace a halogen that has a lower oxidising
power from one of its compounds.
What are oxidisng agents
The oxidising strength decreases down the group.
Oxidising agents are electron acceptors.
colour of halogens in a solution
The colour of the solution in the test tube shows which free halogen is present in solution.
Chlorine =very pale green solution (often colourless),
Bromine = yellow solution
Iodine = brown solution (sometimes black solid
present)
Half equtions for bromine and chlorine
2Br⁻ (aq)→ Br₂ (aq)+ 2e⁻
Cl₂ (aq)+ 2e⁻ → 2Cl⁻ (aq)
- The reactions of halide ions with silver nitrate
This reaction is used as a test to identify which halide ion is present.
The test solution is made acidic with nitric acid, and then silver nitrate solution is added dropwise.
Fluorides produce no precipitate
Chlorides produce a white precipitate
Ag+(aq) + Cl⁻(aq) → AgCl(s)
Bromides produce a cream precipitate
Ag+(aq) + Br⁻(aq) → AgBr(s)
Iodides produce a pale yellow precipitate
Ag+(aq) + I⁻(aq) → AgI(s)
Sodium carbonate reaction with nitric acid
The role of nitric acid is to react with any carbonates present to prevent formation of the precipitate
Ag2CO3. This would mask the desired observations
2 HNO3 + Na2CO3 –> 2 NaNO3 + H2O + CO2
Silver hlide prewciptate reaction iwth ammonia
The silver halide precipitates can be treated with ammonia solution to help differentiate between them if the colours look similar:
Silver chloride dissolves in dilute ammonia to form a complex ion
AgCl(s) + 2NH₃(aq) → [Ag(NH₃)₂]⁺ (aq) + Cl⁻ (aq)
Colourless solution
Silver bromide dissolves in concentrated ammonia to form a complex ion
AgBr(s) + 2NH₃(aq) → [Ag(NH₃)₂]⁺ (aq) + Br⁻(aq)
Colourless solution
Silver iodide does not react with ammonia - it is too insoluble.
- The reaction of halide salts with concentrated sulfuric acid
The halides show increasing power as reducing agents as one goes down the group.
- This can be clearly demonstrated in the various reactions of the solid halides with concentrated sulfuric acid.
Fluoride and Chloride
F- and Cl ions are not strong enough reducing agents to reduce the S in H2SO4. No redox reactions occur. Only acid-base reactions occur.
Sodium Fluoride and chloride reactions with sulfuric acid
NaF(s) + H₂SO₄(l) → NaHSO₄(s) + HF(g)
Observations: White steamy fumes of HF are evolved.
NaCl(s) + H₂SO₄(l) → NaHSO₄(s) + HCl(g)
Observations: White steamy fumes of HCl are evolved.
H2SO4 plays the role of an acid (proton donor).
Explanation of differing reducing power of halides
- A reducing agent donates electrons.
- The reducing power of the halides increases down group 7
- They have a greater tendency to donate electrons.
- This is because as the ions get bigger it is easier for the outer electrons to be given away as the pull from the nucleus on them becomes smaller.
Bromide
Br- ions are stronger reducing agents than Cl- and F- and after the initial acid-base reaction, the bromide ions reduce the sulfur in H2SO4
from +6 to + 4 in SO2
Acid base and redox step for bromide and give observations
Acid- base step: NaBr(s) + H₂SO₄(l) → NaHSO₄(s) + HBr(g)
Redox step: 2 H⁺ + 2 Br⁻ + H₂SO₄ → Br₂(g) + SO₂(g) + 2 H₂O(l)
Observations: White steamy fumes of HBr are evolved.
orange fumes of bromine are also evolved and a colourless, acidic gas SO2
Iodide
- I- ions are the strongest halide reducing agents. They can reduce the sulfur from +6 in H2SO4 to + 4 in SO2 , to 0 in S and -2 in H2S.
acid base (1) redox (3) Equations for iodide
Nal(s) + H₂SO₄(l) → NaHSO₄(s) + HI(g)
2 H⁺ + 2I⁻ + H₂SO₄ → I₂(s) + SO₂(g) + 2 H₂O(l)
6 H⁺ + 6I⁻ + H₂SO₄ → 3 I₂ + S(s) + 4 H₂O(l)
8 H⁺ + 8I⁻ + H₂SO₄ → 4 I₂(s) + H₂S(g) + 4 H₂O(l)
Observation for the redox and acid base reactions with iodine (4)
Observations:
White steamy fumes of HI are evolved. Black solid and purple fumes of Iodine are also evolved
A colourless, acidic gas SO2
A yellow solid of sulfur
H2S (Hydrogen sulfide), a gas with a bad egg smell,
Oxidation and reduction half equations for iodine
Ox 1⁄2 equation 2I⁻ → I₂ + 2e⁻
Re 1⁄2 equation H₂SO₄ + 2H⁺ + 2e⁻ → SO₂ + 2H₂O
Re 1⁄2 equation H₂SO₄ + 6H⁺ + 6e⁻ → S + 4H₂O
Re 1⁄2 equation H₂SO₄ + 8H⁺ + 8e⁻ → H₂S + 4H₂O
What is disproportionation
the name for a reaction where an element simultaneously oxidises and reduces.
reversible reaction of Chlorine with water:
Chlorine with water:
Cl2(g) + H2O (l) ⇌ HClO (aq) + HCl (aq)
If some universal indicator is added to the solution it will first turn red due to the acidity of both reaction products.
- It will then turn colourless as the HClO bleaches the colour.
Reaction with water in sunlight chlorine
- If the chlorine is bubbled through water in the
presence of bright sunlight a different reaction
occurs.
2Cl₂ + 2H₂O → 4H⁺ + 4Cl⁻ + O₂ - The same reaction occurs to an equilibrium mixture of chlorine water when standing in sunlight.
- The greenish colour of chlorine water fades as the Cl2 reacts and a colourless gas (O₂) is produced.
Use of chlorine
- Chlorine is used in water treatment to kill bacteria.
- It has been used to treat drinking water and the water in swimming pools.
- The benefits to health of water treatment by chlorine outweigh its toxic effects.
