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:
- Cl2 + 2KBr \rightarrow Br2 + 2KCl
- Ionic Equations of Displacement Reactions:
- Chlorine displaces bromide and iodide ions:
- Cl2 + 2Br^- \rightarrow 2Cl^- + Br2
- Cl2 + 2I^- \rightarrow 2Cl^- + I2
- Bromine displaces iodide ions:
- Br2 + 2I^- \rightarrow 2Br^- + I2
- Iodine does not react with chloride or bromide ions.
- Chlorine displaces bromide and iodide ions:
Reactions with Hydrogen
Reaction Trend:
- Halogens react with hydrogen to form hydrogen halides; reactivity decreases down Group 17.
Standard Reaction Equation:
- X2 + H2 \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:
- I2 + H2 \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 Cl2 to I2.
- 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.
- Add nitric acid to the halide ion solution to remove interfering ions (e.g., carbonate ions).
- Add a few drops of silver nitrate solution (AgNO_3).
- Observe the precipitate formed.
- Silver nitrate solution tests for halide ions in a solution.
- Standard Equation:
- Ag^+{(aq)} + X^-{(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 (H2SO4) to produce a hydrogen halide (HX), but secondary reactions vary the products depending on the halide.
- Reaction of NaCl and NaF with H2SO4:
- NaF + H2SO4 \rightarrow NaHSO_4 + HF
- NaCl + H2SO4 \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 + H2SO4 \rightarrow NaHSO_4 + HBr
- Misty fumes of HBr are produced.
- HBr is a strong reducing agent and reacts with H2SO4:
- 2HBr + H2SO4 \rightarrow Br2 + SO2 + 2H_2O
- Products include choking gas (SO2) and brown fumes of Br2.
- Reaction of NaI with H2SO4:
- NaI + H2SO4 \rightarrow NaHSO_4 + HI
- Misty fumes of HI are produced.
- HI is a very strong reducing agent and reacts with H2SO4:
- 2HI + H2SO4 \rightarrow I2 + SO2 + 2H_2O
- Because HI is a strong reducing agent, SO2 is further reduced to H2S (rotten egg smell):
- 6HI + SO2 \rightarrow H2S + 3I2 + 2H2O
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) + Cl2(g) \rightarrow NaClO(aq) + NaCl(aq) + H2O(l)
- Chlorine is reduced from 0 in Cl2 to -1 in NaCl and oxidized from 0 in Cl2 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) + Cl2(g) \rightarrow 5NaCl(aq) + NaClO3(aq) + 3H_2O(l)
- Chlorine is reduced from 0 in Cl2 to -1 in NaCl and oxidized from 0 in Cl2 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.
- Cl2 + H2O \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
- O2 + O \rightarrow O3
- Oxygen molecules react with oxygen free radicals (produced by UV light):
- CFCs and Chlorine Free Radicals:
- Chlorine free radicals are formed when CFCs are broken down by UV radiation:
- CCl2F2 \rightarrow Cl "+ CCl_2F^*
- Chlorine free radicals are formed when CFCs are broken down by UV radiation:
- Ozone Depletion:
- Chlorine free radicals react with ozone, destroying the ozone layer by the following mechanism.
- Cl + O3 \rightarrow O2 + ClO
- ClO + O3 \rightarrow 2O2 + Cl
- Chlorine free radicals react with ozone, destroying the ozone layer by the following mechanism.
- Overall Reaction:
- 2O3 \rightarrow 3O2