practice questions TOPIC 4

• Hydrogen fluoride

  Unusually high boiling temperature

  
• Electronegativity of fluorine

  Greater than that of bromine
• Fluorine atom has less shielding

  Greater pull from nucleus on bonding electrons
• HF is more polar

  Can form hydrogen bonds
• Predicted boiling temperature of HF without hydrogen bonds is 150-180 K

  
• Propanone

  Can dissolve a wide range of substances due to both polar and non-polar characteristics
• Propanone and water

  Hydrogen bonds form
• Propanone and octane

  London forces form
• HI has higher boiling temperature than HBr

  Stronger London forces in HI
• HF has higher boiling temperature than HCl

  HF forms hydrogen bonds, HCl does not

  
• H2O has higher boiling temperature than HF

  Water can form up to two hydrogen bonds per molecule, HF only one

  
• Phosphoric(V) acid is used instead of concentrated sulfuric acid to prepare hydrogen iodide
• When hydrogen iodide gas is inverted in water, water rises in the test tube
• Reaction of hydrogen iodide gas and ammonia

  NH3(g) + HI(g) → NH4I(s)
• 1. iodobutane reacts with hot aqueous silver nitrate to form a yellow precipitate
• Test for sulfate ions

  Add HCl, then add BaCl - white precipitate forms

  
• Boiling temperatures increase from chlorine to iodine

  Number of electrons increases, so strength of London forces increases

  
• Chlorine content in swimming pools

  Lower than 1 ppm - ineffective as disinfectant, higher than 3 ppm - increased toxicity/acidity
• Confirmatory tests for anion in cream precipitate

  Add dilute ammonia - precipitate insoluble, add concentrated ammonia - precipitate soluble
• Dilute hydrochloric acid is needed in Test 3 to remove any carbonate
• Disproportionation reaction of chlorine with hot aqueous sodium hydroxide

  Cl2 + 2OH- → ClO- + Cl- + H2O

  
• Oxidation numbers in disproportionation reaction

  Cl changes from 0 to -1 (reduced) and 0 to +5 (oxidised)

  
• Iodide ions are the strongest reducing agents, reducing sulfur from +6 to 0 in H2SO4 and -2 in S

  
• Reaction of concentrated hydrochloric acid and concentrated ammonia

  HCl(g) + NH3(g) → NH4Cl(s)

  
• Sodium carbonate dissolves in hydrochloric acid, forming a colourless solution with effervescence
• Ammonium and bromide are hazardous in school labs, chlorine is toxic
• Green precipitate turns brown due to oxidation of Fe2+ to Fe3+ by oxygen
• Dilute hydrochloric acid is needed in Test 2 to remove any carbonate

  
• Ionic half-equation for oxidation of chlorine to chlorate(I) ions

  Cl2 + 2OH- → ClO- + Cl- + H2O

  
• Disproportionation is simultaneous oxidation and reduction of an element in the same species
• Formation of London forces between halogen molecules

  Uneven distribution of electrons, induces temporary dipole in first molecule, which induces dipole in second molecule
• Bromine has higher boiling temperature than chlorine


  Bromine has more electrons, so stronger London forces

  
• Astatine is a solid due to increasing strength of London forces down the group
• Test for iodide ions

  Add AgNO3 (and HNO3) - yellow precipitate forms

  
• Test for carbonate ions

  Add HCl - fizzing

  
• Test for sulfate ions

  Add BaCl - white precipitate forms

  
• Reaction of chlorine with cold, dilute sodium hydroxide is disproportionation

  
• Molecular geometry of ammonia

  • Trigonal pyramidal, 107° angle, 3 bonding electron pairs and 1 lone pair
• Reaction is not redox as oxidation number does not change
• Hydrogen halides as reducing agents

  Hydrogen iodide reduces sulfur in sulfuric acid further than hydrogen bromide