Group 17 Halogens: Trends and Reactivity

• Group 17 elements are known as halogens

Colours and Trends in Volatility

• Halogen colours darken down the group at room temperature
• Boiling points increase down the group because intermolecular forces get stronger
• Fluorine is a pale yellow gas
• Chlorine is a green gas
• Bromine is a red-brown liquid
• Iodine is a grey solid

Volatility and Van Der Waals Forces

• Fluorine has the lowest melting and boiling points, making it the most volatile in Group 17
• This is because it has the weakest van der Waals forces
• The number of electrons and the size of molecules increase going down Group 7, resulting in stronger temporary dipoles and van der Waals forces
• More energy is needed to overcome these forces, therefore volatility decreases down Group 17

Reactivity as Oxidising Agents

• Halogens gain an electron to form negative ions
• Reactivity decreases down the group because it becomes harder to gain an electron
• Electron shielding and atomic radius increase down the group, meaning there is weaker attraction for incoming electrons
• Oxidising agents gain electrons, halogens act as oxidising agents but become less oxidising down the group
• Oxidising strengths are reflected in displacement reactions with halide ions
• Chlorine is the strongest oxidising agent and iodine is the weakest

Displacement Reactions

• Chlorine will displace bromide and iodide ions
• The ionic equations are:
  • Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂
  • Cl₂ + 2I⁻ → 2Cl⁻ + I₂
• Bromine will displace iodide ions
  • The ionic equation is: Br₂ + 2I⁻ → 2Br⁻ + I₂
• Iodine will not react with chloride or bromide ions
• A halogen will displace a halide from a solution if the halide ion is below it in the periodic table
• Displacement reactions can be identified by a change in the solution's colour
• Chlorine solution is colorless
• Bromine solution is orange
• Iodine solution is brown
• For example:
  • When chlorine is added to potassium bromide, the solution changes from colorless to orange: Cl₂ + 2KBr → Br₂ + 2KCl

Reaction with Hydrogen

• Halogens react with hydrogen to form hydrogen halides which shows that reactivity decreases down Group 17
• The standard reaction equation for the reaction with hydrogen (where X is the halogen) is: X₂ + H₂ → 2HX
• Fluorine reacts explosively with hydrogen to form hydrogen fluoride gas, even in cold atmospheres
• Chlorine reacts with hydrogen when lightly heated or exposed to sunlight
• Bromine reacts with hydrogen when heated with a flame
• Iodine only partially reacts with hydrogen when constantly heated because an equilibrium is set up: I₂ + H₂ ⇌ 2HI

Thermal Stability of Hydrides

• Thermal stability is how easily a hydrogen halide breaks into its constituent elements when heated
• Hydrogen fluoride and hydrogen chloride don't split into hydrogen and the halogen when heated under lab conditions and are very thermally stable
• Hydrogen bromide splits into hydrogen and bromine when heated
• Hydrogen iodide splits into hydrogen and iodine more easily than hydrogen bromide
• Thermal stability of the hydrides decreases down Group 17 because the covalent bonds are weaker since larger halogen atoms result in bonding pairs being further from the nucleus and weaker attraction

Bond Energies

• Thermal stability of halogens decreases down the group
• This relates to bond energies/enthalpies
• Bond enthalpies of hydrogen halides decrease down Group 17 because halogen size increases, meaning less energy is needed to break the covalent bond

Bond Enthalpies of Halogens

• Bond enthalpies of halogen molecules decrease from Cl₂ to I₂
• Larger molecules have bonding pairs that are less attracted to the nucleus, covalent bonds can therefore be easily broken

Reactions with Silver Nitrate Followed by Aqueous Ammonia

• Silver nitrate solution is used to test for halide ions:
  • Add nitric acid to remove interfering ions.
  • Add silver nitrate solution (AgNO₃).
  • Observe the precipitate formed
• The standard equation is: Ag⁺(aq) + X⁻(aq) → AgX(s)
• Specific observations include:
  • Fluoride ions - no precipitate
  • Chloride ions - white precipitate
  • Bromide ions - cream precipitate
  • Iodide ions - yellow precipitate
• Aqueous ammonia is added to verify the precipitates:
• Chloride precipitate will dissolve in dilute NH₃
• Bromide precipitate will dissolve in concentrated NH₃
• Iodide precipitate is insoluble in dilute and concentrated NH₃

Reactions with Concentrated Sulfuric Acid

• All halide ions produce a hydrogen halide when reacting with concentrated sulfuric acid, but a secondary reaction can occur depending on the halide

Reactions of NaCl and NaF with H₂SO₄

• NaF + H₂SO₄ → NaHSO₄ + HF
• NaCl + H₂SO₄ → NaHSO₄ + HCl
• HF and HCl are misty fumes but not strong enough reducing agents to react further

Reactions of NaBr with H₂SO₄

• NaBr + H₂SO₄ → NaHSO₄ + HBr
• HBr misty fumes are produced but HBr is a reducing agent and reacts with H₂SO₄ producing choking SO₂ gas and brown Br₂ fumes: 2HBr + H₂SO₄ → Br₂ + SO₂ + 2H₂O
• Redox reaction occurs

Reactions of NaI with H₂SO₄

• NaI + H₂SO₄ → NaHSO₄ + HI
• HI misty fumes are produced and HI is a strong reducing agent, reacting with H₂SO₄ and producing SO₂
• HI is a very strong reducing agent, so SO₂ is further reduced to H₂S (smells like rotten eggs):
  • 2HI + H₂SO₄ → I₂ + SO₂ + 2H₂O
  • 6HI + SO₂ → H₂S + 3I₂ + 2H₂O

Reactions of Chlorine with Aqueous Sodium Hydroxide

• Disproportionation reactions involve an element being both oxidised and reduced. Chlorine reacts with cold dilute sodium hydroxide: 2NaOH(aq) + Cl₂(g) → NaClO(aq) + NaCl(aq) + H₂O(l) Chlorine is reduced from 0 in Cl₂ to -1 in NaCl and oxidised from 0 in Cl₂ to +1 in NaClO (sodium chlorate(I) solution), which is bleach, used for water treatment, textile and paper bleaching, and cleaning Chlorine reacts with hot concentrated sodium hydroxide: 6NaOH(aq) + 3Cl₂(g) → 5NaCl(aq) + NaClO₃(aq) + 3H₂O(l) Chlorine has been reduced from 0 in Cl₂ to -1 in NaCl and oxidised from 0 in Cl₂ to +5 in NaClO₃

Chlorine in Water Purification

• Chlorine is used to kill bacteria in water purification
• It reacts with water in a disproportionation reaction and produces chloride and chlorate ions
• The reaction produces HCl, so an alkali is added to reduce the acidity: Cl₂ + H₂O ⇌ 2H⁺ + Cl⁻ + ClO⁻
• Chlorate ions kill bacteria, making water safe to drink or swim in
• Prevents reinfection, algae growth, and bad tastes, smells and discoloration in the long term
• Chlorine is toxic and can react with organic matter to cause cancer but the benefits are agreed to outweigh the risks

Industrial Importance and Environmental Significance of Halogens

• Chlorine and chlorate ions kill bacteria, preventing disease and infection
• The polymer poly(chloroethene), PVC, contains one chlorine atom in each polymer unit and is used for windows and drain pipes
• Plasticiser can be added to PVC to make it more flexible and extend its use to other things

Halogenated Hydrocarbons

• Halogens react with alkanes and create halogenoalkanes
• Chlorofluorocarbons (CFCs) are halogenoalkanes with all hydrogen atoms replaced by chlorine and fluorine
• CFCs were previously used as coolants, solvents, and propellants. However, they damage the ozone layer, therefore are banned
• HFCs (hydrofluorocarbons) which don't contain chlorine now replace them

Damage To The Ozone Layer

• Ozone (O₃) is the layer in the upper atmosphere which absorbs a lot of UV radiation from the sun
• Ozone forms when an oxygen molecule reacts with an oxygen free radical in the presence of UV light: O₂ → O + O, O + O₂ → O₃
• Chlorine free radicals are formed when CFCs are broken down by UV radiation: CCl₃F → Cl + CCl₂F
• Chlorine free radicals then react with and break down ozone into oxygen:
  • Cl + O₃ → O₂ + ClO
  • ClO + O₃ → 2O₂ + Cl
• Overall: 2O₃ → 3O₂

Explore Group 17 elements, the halogens, focusing on color changes and volatility trends. Learn how boiling points increase due to stronger intermolecular forces. Understand the reactivity of halogens as oxidizing agents and how it decreases down the group.