Group 17 Halogens Flashcards

Group 17: The Halogens

Physical Properties

• Group 17 elements are known as halogens.
• Colors and Volatility Trends:
  • Color darkens down the group at room temperature.
  • Boiling points increase down the group due to increasing intermolecular forces.
    • Fluorine: Pale yellow gas
    • Chlorine: Green gas
    • Bromine: Red-brown liquid
    • Iodine: Grey solid
• Volatility and van der Waals Forces:
  • Fluorine is the most volatile, having the lowest melting and boiling points due to weak van der Waals forces.
  • Down the group, the number and size of molecules increase, leading to stronger temporary dipoles and more van der Waals forces.
  • More energy is needed to overcome these forces, so volatility decreases down Group 17.

Chemical Properties

• Reactivity as Oxidizing Agents:
  • Halogens gain an electron to form negative ions during reactions.
  • Reactivity decreases down the group.
  • Electron shielding and atomic radius increase down the group, weakening the attraction between incoming electrons and the nucleus.
• Oxidizing Agents:
  • Halogens act as oxidizing agents, becoming less oxidizing down the group due to decreasing reactivity.
  • Oxidizing strengths are demonstrated by displacement reactions with halide ions.
• Displacement Reactions:
  • Chlorine is the strongest oxidizing agent, and iodine is the weakest (among Cl, Br, I).
  • General trend: A halogen displaces a halide from a solution if the halide ion is below it in the periodic table.
  • Displacement reactions are indicated by color changes in the solution:
    • Chlorine solution: Colorless
    • Bromine solution: Orange
    • Iodine solution: Brown
  • Example: Chlorine added to potassium bromide solution changes the solution from colorless to orange:
    • $Cl<em>2 + 2KBr \rightarrow Br</em>2 + 2KCl$
• Ionic Equations of Displacement Reactions:
  • Chlorine displaces bromide and iodide ions:
    • $Cl<em>2 + 2Br^- \rightarrow 2Cl^- + Br</em>2$
    • $Cl<em>2 + 2I^- \rightarrow 2Cl^- + I</em>2$
  • Bromine displaces iodide ions:
    • $Br<em>2 + 2I^- \rightarrow 2Br^- + I</em>2$
  • Iodine does not react with chloride or bromide ions.

Reactions with Hydrogen

• Reaction Trend:

  • Halogens react with hydrogen to form hydrogen halides; reactivity decreases down Group 17.

• Standard Reaction Equation:

  • $X<em>2 + H</em>2 \rightarrow 2HX$ (where X is the halogen)

• Specific Reactions:

  • Fluorine reacts explosively with hydrogen, even in a cold atmosphere.
  • Chlorine reacts with hydrogen when lightly heated or exposed to sunlight.
  • Bromine reacts with hydrogen when heated with a flame.
  • Iodine partially reacts with hydrogen when constantly heated; an equilibrium is established:
    • $I<em>2 + H</em>2 \rightleftharpoons 2HI$

Thermal Stability of Hydrides

• Definition:
  • Thermal stability is the ease with which a hydrogen halide breaks up into its constituent elements when heated.
• Stability Trend:
  • Thermal stability decreases down Group 17.
• Specific Stability:
  • Hydrogen fluoride (HF) and hydrogen chloride (HCl) are very thermally stable.
  • Hydrogen bromide (HBr) splits into hydrogen and bromine when heated.
  • Hydrogen iodide (HI) splits into hydrogen and iodine more easily than hydrogen bromide.
• Explanation:
  • Covalent bonds become weaker down the group, making them easier to break upon heating.
  • Larger halogen atoms mean the bonding pair is further from the nucleus, reducing attraction and weakening the bond.

Bond Energies

• Bond Enthalpies of Hydrides:
  • The thermal stability of halogens decreases down the group due to decreasing bond energies.
  • Bond enthalpies of hydrogen halides decrease down Group 17 because the size of the halogen increases.
  • Less energy is required to break the covalent bond between hydrogen and halogen.
• Bond Enthalpies of Halogens:
  • The bond enthalpies of the halogen molecules decrease from $Cl<em>2$ to $I</em>2$.
  • Larger molecules have bonding pairs further from the nucleus, leading to weaker attraction and easier bond breakage.

Reactions of Halide Ions

• Reactions with Silver Nitrate and Aqueous Ammonia:
  • Silver nitrate solution tests for halide ions in a solution.
    1. Add nitric acid to the halide ion solution to remove interfering ions (e.g., carbonate ions).
    2. Add a few drops of silver nitrate solution ($AgNO_3$).
    3. Observe the precipitate formed.
• Standard Equation:
  • $Ag^+<em>{(aq)} + X^-</em>{(aq)} \rightarrow AgX_{(s)}$
• Observations:
  • Fluoride ions: No precipitate.
  • Chloride ions: White precipitate.
  • Bromide ions: Cream precipitate.
  • Iodide ions: Yellow precipitate.
• Ammonia Test:
  • Chloride precipitate: Dissolves in dilute $NH_3$.
  • Bromide precipitate: Dissolves in concentrated $NH_3$.
  • Iodide precipitate: Insoluble in dilute and concentrated $NH_3$.

Reactions with Concentrated Sulfuric Acid

• All halide ions react with concentrated sulfuric acid ($H<em>2SO</em>4$) to produce a hydrogen halide (HX), but secondary reactions vary the products depending on the halide.
• Reaction of NaCl and NaF with H2SO4:
  • $NaF + H<em>2SO</em>4 \rightarrow NaHSO_4 + HF$
  • $NaCl + H<em>2SO</em>4 \rightarrow NaHSO_4 + HCl$
  • HF and HCl appear as misty fumes.
  • HF and HCl are not strong reducing agents, so no further reactions occur.
• Reaction of NaBr with H2SO4:
  • $NaBr + H<em>2SO</em>4 \rightarrow NaHSO_4 + HBr$
  • Misty fumes of HBr are produced.
  • HBr is a strong reducing agent and reacts with $H<em>2SO</em>4$:
    • $2HBr + H<em>2SO</em>4 \rightarrow Br<em>2 + SO</em>2 + 2H_2O$
  • Products include choking gas ($SO<em>2$) and brown fumes of $Br</em>2$.
• Reaction of NaI with H2SO4:
  • $NaI + H<em>2SO</em>4 \rightarrow NaHSO_4 + HI$
  • Misty fumes of HI are produced.
  • HI is a very strong reducing agent and reacts with $H<em>2SO</em>4$:
    • $2HI + H<em>2SO</em>4 \rightarrow I<em>2 + SO</em>2 + 2H_2O$
  • Because HI is a strong reducing agent, $SO<em>2$ is further reduced to $H</em>2S$ (rotten egg smell):
    • $6HI + SO<em>2 \rightarrow H</em>2S + 3I<em>2 + 2H</em>2O$

Reactions of Chlorine with Aqueous Sodium Hydroxide

• Disproportionation Reaction:
  • A reaction where an element is both oxidized and reduced.
• Reaction with Cold Dilute NaOH:
  • $2NaOH(aq) + Cl<em>2(g) \rightarrow NaClO(aq) + NaCl(aq) + H</em>2O(l)$
  • Chlorine is reduced from 0 in $Cl<em>2$ to -1 in NaCl and oxidized from 0 in $Cl</em>2$ to +1 in NaClO.
  • NaClO (sodium chlorate(I) solution) is bleach, used in water treatment, textile and paper bleaching, and cleaning.
• Reaction with Hot Concentrated NaOH:
  • $6NaOH(aq) + Cl<em>2(g) \rightarrow 5NaCl(aq) + NaClO</em>3(aq) + 3H_2O(l)$
  • Chlorine is reduced from 0 in $Cl<em>2$ to -1 in NaCl and oxidized from 0 in $Cl</em>2$ to +5 in $NaClO_3$.

Chlorine in Water Purification

• Mechanism:
  • Chlorine kills bacteria in water purification.
  • Chlorine reacts with water in a disproportionation reaction, producing chloride and chlorate ions.
    • $Cl<em>2 + H</em>2O \rightleftharpoons 2H^+ + Cl^- + ClO^-$
  • The reaction produces HCl, so an alkali is added to reduce acidity.
  • Chlorate ions kill bacteria, making water safe to drink or swim in.
• Benefits:
  • Kills dangerous microorganisms which could cause diseases.
  • Persists in the water, preventing long-term reinfection.
  • Prevents the growth of algae.
  • Removes bad tastes and smells.
  • Removes discoloration.
• Risks and Benefits:
  • Chlorine is toxic, leading to discussions about its use in water.
  • Chlorine can react with organic matter, forming potentially carcinogenic compounds.
  • The consensus is that the benefits outweigh the risks due to the small amount of chlorine added.

Industrial Importance and Environmental Significance

• Killing Bacteria:
  • Chlorine and chlorate ions are industrially important for preventing disease and infection.
• PVC:
  • Poly(chloroethene), PVC, contains one chlorine atom in each polymer unit and is used in windows and drain pipes due to its hardness.
  • Plasticizers can be added to PVC to increase flexibility for electrical cable insulation and clothing.
• Halogenated Hydrocarbons:
  • Halogens react with alkanes to form halogenoalkanes.
  • Chlorofluorocarbons (CFCs) are halogenoalkane molecules in which all hydrogen atoms have been replaced by chlorine and fluorine.
  • CFCs were used as coolants, solvents, and propellants but damage the ozone layer.
  • CFCs have been banned and replaced with hydrofluorocarbons (HFCs), which do not contain chlorine.

Damage to the Ozone Layer

• Ozone Layer:
  • Ozone ($O_3$) in the upper atmosphere absorbs UV radiation from the sun.
• Ozone Formation:
  • Oxygen molecules react with oxygen free radicals (produced by UV light):
    • $O_2 \rightarrow O + O$
    • $O<em>2 + O \rightarrow O</em>3$
• CFCs and Chlorine Free Radicals:
  • Chlorine free radicals are formed when CFCs are broken down by UV radiation:
    • $CCl<em>2F</em>2 \rightarrow Cl "+ CCl_2F^*$
• Ozone Depletion:
  • Chlorine free radicals react with ozone, destroying the ozone layer by the following mechanism.
    • $Cl + O<em>3 \rightarrow O</em>2 + ClO$
    • $ClO + O<em>3 \rightarrow 2O</em>2 + Cl$
• Overall Reaction:
  • $2O<em>3 \rightarrow 3O</em>2$