Chemistry 2.3- Group 7, the halogens

What are the colours and states of the halogens?

What are the colours and states of the halogens?

Fluorine- pale yellow gas

Chlorine- pale green gas

Bromine- brown liquid/orange vapour

Iodine- grey solid/purple vapour

How and why does the atomic radius vary down the group?

• As you go down the group, the elements have more electron shells

• So the atom radii increase

How do melting and poiling points change down the group?

Increase

• Atoms have more electrons, meaning they can form greater temporary and induced dipoles (greater van der Waals forces)

How and why does the electronegativity vary down the group? How does this affect the trend in reactivity?

• As you go down the group, the elements have more shells, so the atomic radii increase

• This means that the outer shell is further away from the nucleus, and experiences more shielding (like charges repelled by the inner electron shells)

• So the bonding pair of electrons is less attracted, and electronegativity decreases going down the group

Because of this, reactivity of halogens decreases down the group

How can halogens and halides act as oxidising agents and reducing agents?

Halogens- oxidising agents:

• Halogens can oxidise atoms by removing an electron

• They gain the electron, so are themselves reduced

• Halogens are electron acceptors, so they are oxidising agents

Halide ions- reducing agents:

• Halides can reduce atoms by donating an electron

• They lose the electron, so are themselves oxidised

• Halide ions are electron donors, so they are reducing agents

What are the trends in oxidising and reducing power down the group?

Halogens- oxidising power:

• Down the group, the atomic radii increase

• The outer shell electrons are further from the nucleus and experience more shielding

• This means that it is harder for the atoms to accept (gain) an electron

• Oxidising power decreases down the group

Halide ions- reducing power:

• Down the group, the ionic radii increase

• The outer shell electrons are further from the nucleus and experience more shielding

• This means that it is easier for the ions to donate (lose) aan electron

• Reducing power increases down the group

Explain the displacement reactions of halogens in terms of oxidising power

• Oxidising power decreases down the group as the halogens are less easily reduced (harder for bigger atoms to gain electrons)

• This means that the halogens at the top of the group are stronger oxidising agents, so they can oxidise a lower halide, displacing it from its metal halide compound

Eg. chlorine is higher up than bromine, so it is a stronger oxidising agent and can oxidise bromide ions, meaning it will displace bromine from sodium bromide

• Cl₂ + 2NaBr → 2NaCl + Br₂

• Ionic equation= Cl₂ + 2Br^- → 2Cl^- + Br₂

Describe the reaction of sodium chloride with sulfuric acid

An acid-base reaction occurs: H₂SO₄(l) + NaCl(s) → HCl(g) + NaHSO₄(s)      

• So white misty fumes produced (hydrogen chloride), and damp blue indiciator paper turns red (acidic)

The chloride ions are too weak of a reducing agent to reduce the sulfuric acid

Describe the reaction of sodium bromide with sulfuric acid

An acid-base reaction occurs: H₂SO₄(l) + NaBr(s) → HBr(g) + NaHSO₄(s)  

• So, white misty fumes produced (hydrogen bromide), and damp blue indicator paper turns red (acidic)

The bromide ions produced are strong enough reducing agents to reduce the sulfuric acid to sulfur dioxide: 2HBr(g) + H₂SO₄(l) → Br₂(g) + SO₂(g) + 2H₂O(l)

• So orange fumes produced (bromine gas), and orange dichromate paper turns green (sulfur dioxide present)

Describe the reaction of sodium iodide with sulfuric acid

An acid-base reaction occurs: H₂SO₄ (l) + NaI (s) → HI (g) + NaHSO₄ (s)   

• So white misty fumes produced (hydrogen iodide), and damp blue indicator turns red (acidic)

The iodide ions are strong enough reducing agents to reduce sulfuric acid to sulfur dioxide: 2HI (g) + H₂SO₄ (l) → I₂ (g) + SO₂ (g) + 2H₂O (l)

• So purple fumes produced (iodine gas), and orange dichromate paper turns green (sulfur dioxide present)

They are also strong enough to reduce sulfuric acid to solid sulfur: 6HI (g) + H₂SO₄ (l) → 3I₂ (g) + S (s) + 4H₂O (l)

• So yellow solid produced (solid sulfur)

They are also strong enough to reduce sulfuric acid to hydrogen sulfide: 8HI (g) + H₂SO₄ (l) → 4I₂ (g) + H₂S (s) + 4H₂O (l)

• So white lead ethanoate paper turns black (hydrogen sulfide present)

Explain the reactions of sodium halides with sulfuric acid in terms of reducing power

• Reducing power increases down the group as the halides are more easily oxidised (easier for bigger atoms to lose electrons)

• This means that the halide ions are better at reducing sulfuric acid as you go down the group, and can reduce it to a lower oxidation state

Chloride can’t reduce H₂SO₄ (sulfur stays +6)

Bromide can reduce H₂SO₄ to SO₂ (+4)

Iodide can reduce H₂SO₄ to SO₂ (+4), to S (0), and to H₂S (-2)

How can we test for the halides?

• Add nitric acid and silver nitrate solution

• If a precipitate is formed, a halide ion is present

  • Chloride = white precipitate

  • Bromide = cream precipitate

  • Iodidie = yellow precipiate

Often it is hard to distinguish between these colours, so adding ammonia solution is used as a follow-up test, by observing whether the precipitate dissolves

• Chloride = dissolves in dilute ammonia

• Bromide = dissolves only in concentrated ammonia

• Iodidde = doesn’t dissolve in concentrated ammonia

Why is silver nitrate solution used to test for halides?

Silver chloride, silver bromide and silver iodide are insoluble, with different coloured precipitates

• Ag⁺ (aq) + X^- (aq) → AgX (s)

Why is silver nitrate acidified in the halide ion test?

Nitric acid will react with any carbonate or hydroxide ions, as silver carbonate and silver hydroxide are insoluble, so they would also give a positive test

This works because:

• OH^-_(aq) + H⁺_(aq) → H₂O_(l)

• CO₃²⁻_(aq) + H⁺_(aq) → CO₂_(g) + H₂O_(l)

How does chlorine react with water? Why is this useful?

• In a disproportionation reaction, chlorine (oxidation state 0) is both reduced to -1 in hydrochloric acid, and oxidised to +1 in chloric (I) acid

• Cl₂ (g) + H₂O (l) → HCl (aq) + HClO (aq)

We use this reaction to sterilise drinking water and pools because HClO dissociates to ClO^- (chlorate (I) ion), which is a dissinfectant

Why does chlorine have to be regularly topped up in pools?

• Remaining chlorine is depleted in sunlight to form HCl, and no longer acts as a dissinfectant

• 2Cl₂ (g) + 2H₂O (l) → 4HCl (aq) + O₂ (g)

How is it safe to add chlorine to drinking water?

Though chlorine is toxic, the benefits of adding it in small quanitities to water to sterilise it outweigh the risks

How do we use chlorine to produce bleach?

• Chlorine reacts with cold, dilute sodium hydroxide in a disproportionation reaction to form sodium chlorate (I), NaClO, the active ingredient in bleach

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