Group 7

Trend in atomic radius of the halogens

• Atomic radius increases down the group

• As you go down the group, there is more shells and more shielding, so a weaker attraction between the nucleus and outer electrons

Why is the ionic radius of halogens bigger than the atomic radius?

The halide ion (X^-) has the same number of protons but more electrons than the atom (X) so the attraction between nucleus and outer electron is weaker

Trend in boiling point of halogens

• Increases down the group

• The halogens become less volatile (evaporate less easily)

• This is because as you go down the group

  • Molecules (X₂) are larger with more electrons

  • Vdws intermolecular forces between molecules are stronger

  • More energy Is needed to overcome them

Trend in electronegativity of halogens

• Electronegativity decreases down the group

• As you go down the group

• There are more shells and more shielding

• So the attraction between the nucleus and the pair of electrons in the covalent bond is less

Trend in first ionisation energy of halogens

• First ionisation energy decreases down the group

• As you go down the group

• There are more shells and more shielding

• So the attraction between the nucleus and outer electron is less

• Less energy is needed to remove it

Appearance of Cl₂

• Green

• Gas at room temp

• Pale green in aqueous solution (polar solvent)

• Pale green layer formed in cyclohexane (non-polar solvent)

Appearance of Br₂

• Red/brown

• Liquid at room temp

• Orange in aqueous solution (polar solvent)

• Orange top layer forms in cyclohexane (non-polar solvent)

Appearance of I₂

• Black

• Solid at room temp

• Brown in aqueous solution (polar solvent)

• Pink/purple top layer forms in cyclohexane (non-polar solvent)

Why are halogens more soluble in non-polar solvents (cyclohexane)?

The halogens are non-polar, so dissolve better in non-polar solvents

What do redox reactions of halogens provide evidence for?

The trend in oxidising powers of the halogens (halogens X₂ are oxidising agents (they accept electrons)

Redox reaction of Cl₂ with KBr

• Forms potassium chloride and bromine

  • Chlorine is reduce, Cl₂ is the oxidising agent

    • Cl₂ + 2e^- —> 2Cl^-

  • Bromine is oxidised, Br^- is the reducing agent

    • 2Br^- —> Br₂ + 2e^-

• First write the full equation, then shorten to ionic and then half equations

• Orange solution forms (Br₂)

Redox reaction of Cl₂ with KI

• Forms potassium chloride and iodine

  • Chlorine is reduce, Cl₂ is the oxidising agent

    • Cl₂ + 2e^- —> 2Cl^-

  • Iodine is oxidised, I^- is the reducing agent

    • 2I^- —> I₂ + 2e^-

• First write the full equation, then shorten to ionic and then half equations

• Brown solution forms (I₂)

Redox reaction of Br₂ with KI

• Forms potassium bromide and iodine

  • Bromine is reduce, Br₂ is the oxidising agent

    • Br₂ + 2e^- —> 2Br^-

  • Iodine is oxidised, I^- is the reducing agent

    • 2I^- —> I₂ + 2e^-

• First write the full equation, then shorten to ionic and then half equations

• Brown solution forms (I₂)

Further reactions of halogens with no change

• When KCl is reacted with Br₂ it remains orange

• When KCl and KBr are reacted with I₂ they remain brown

Why are Br₂ and I₂ less good at acting as oxidising agents?

• An oxidising agent is species that accepts electrons, halogens act as oxidising agents by accepting electrons to form halide ions

• As you go down group 7,

  • There are more shells and more shielding

  • So the nuclear attraction to the outer shell is less

  • Electrons are less readily accepted

• X₂ is able to oxidise Y^- is X if higher up the group than Y

Disproportionation definition

A redox reaction in which the same element is both oxidised and reduced

Disproportionation (redox) reaction of chlorine and water

• equilibrium is establish between Cl₂ and H₂O, and on the other side HCl and HClO

• Chlorine is both oxidised and reduced

• Chlorine is pale green solution, HCl turns blue litmus red, HClO (chloric (I) acid) bleaches litmus paper

• Used of chlorine in water treatment

  • Benefit: kills bacteria

  • Risks: toxic in large quantities

• In presence of sunlight : hydrochloric acid and oxygen are produced, chloric (I) acid decomposes in sunlight to form the products

Disproportionation (redox) reaction of chlorine and cold dilute sodium hydroxide solution

• Produces sodium chlorine, sodium chlorate (I) and water

• Cl₂ +2NaOH —> NaCl + NaClO + H₂O

• Green colour of chlorine fades and smell is less pungent

• Chlorine is both oxidised and reduced

• Used to kill bacteria and as a bleach

Test for aqueous halide ions

• Determined from a precipitation reaction using acidified silver nitrate solution

• Cl^- ions form white precipitate (AgCl)

• Br^- ions form cream precipitate (AgBr)

• I^- ions form yellow precipitate (AgI)

• F^- ions do not form precipitate a precipitate as AgF is soluble in water- a colourless solution is seen

• Silver chloride is soluble in dilute ammonia, silver bromide in concentrated ammonia and silver iodide in neither (used to distinguish between precipitates)

What do redox reactions of halide ions provide evidence for?

• Halides are reducing agents (electron donors) as they donate electrons to form halogen molecules

• Halide ions act as reducing agents to different extents- this can be investigated by studying the reaction of halide ions with concentrated sulfuric acid

Reaction of solid halide salts e.g. KCl, NaBr, with concentrated sulfuric acid

• If sulfuric acid is reduced, there are a few possible reduction products

• If it is reduced, X^- will be oxidised to the halogen molecule X₂

Possible products of reduction of concentrated sulfuric acid

• Sulfuric acid has an oxidation number of +6

• SO₂ sulfur (IV) oxide, +4, appearance = choking gas (acidic, toxic)

• S sulfur, 0, yellow solid

• H₂S hydrogen sulfide, -2, gas with a smell of bad eggs

What are the two types of reaction when a solid halide salt reacts with concentrated sulfuric acid?

• Acid-base: sulfuric acid acts as an acid (H⁺) donor, X^- is the base so accepts electrons

• Redox: sulfuric acids acts as oxidising agent (e^- acceptor) (X^- acts as reducing agent)

Observations of NaCl and H₂SO₄

• Acid-base: steamy white fumes HCl

• Also forms NaHSO₄

• No redox reaction because no element has changed its oxidation state

• Cl^- (base) cannot reduce sulfuric acid (acid) so chloride ions are a weak reducing agent

Observations of NaBr and H₂SO₄

• Acid-base: steamy white fumes HBr and also forms KHSO₄

• Sulfuric acid: H⁺ donor, Br^-: H⁺ acceptor

• Redox: Br^- ion reduces sulfuric acid to SO₂, Br^- ion is reducing agent (donates electrons to sulfuric acid), sulfuric acid is oxidising agent (accepts electrons)

• Observations: orange fumes of Br₂ (product of oxidation), choking gas SO₂ (product of reduction)

Observations of NaI and H₂SO₄

• Acid-base: steamy white fumes HI and NaHSO₄ also forms

• Sulfuric acid: H⁺ donor, I^-: H⁺ acceptor

• Redox: I^- reduces sulfuric acid to SO₂, S, H₂S

• SO₂: choking gas (product of reduction), black solid I₂ (product of oxidation)

• S: yellow solid (product of reduction), black solid I₂ (product of oxidation)

• H₂S: gas with a smell of bad eggs (product of reduction), black solid I₂ (product of oxidation)

• H₂S is the lowest oxidation number of sulfur

• Iodide ions are powerful reducing agents, the sulfuric acid acts as an oxidising agent (electron acceptor) in each of these redox reactions

Why are some halides better at being reducing agents than others?

• A reducing agent donates electrons (causes gain)

• Halide ions donate electrons to form X₂

• The reducing strength of halides increases down the group

  • As you go down the group, the halide ion has more shells and more shielding so the nuclear attraction to the outer shell is less and electrons are more readily lost

Summary of the strength of oxidising agents and reducing agents

• The halogens are oxidising agents- strength increases up the group and electrons are gained more readily

• The halides are reducing agents- strength increases down the group as electrons are lost more readily