group 7

• what are the group 7 elements called?

  THE HALOGENS
• formula of halogens

  they are diatomic:

  F2

  Cl2

  Br2

  I2
• colour and physical state of halogens

  fluorine - yellow gas

  chlorine - green gas

  bromine - red/brown liquid

  iodine - grey solid
• what is the trend in ELECTRONEGATIVITY down group 7 (halogens)

  • electronegativity DECREASES
  • as you go down the group, the atomic radius increases
  • larger atoms attract electrons in bonding pairs less
  • this is because the e-s are further away from the nucleus and are shielded by other e-s
• how does BOILING POINT change down group 7?

  • BPs INCREASE
  • this is due to increasing strength of VAN DER WAALS forces between molecules, as the size and Mr of the molecules increase down the group
• what is the trend in REACTIVITY down group 7?

  • REACTIVITY DECREASES
  • halogens must GAIN an electron when they react
  • as you go down group 7, the atoms get bigger and the outer shell is further from the nucleus
  • therefore the electrons are less strongly attracted to it
• what happens to the halogens' relative oxidising strengths down the group

  THEY DECREASE - the halogens become less oxidising down group 7
• A halogen will displace a halide from solution if the halide is below it on the periodic table
• what colour solution is formed when Br- is displaced

  an orange solution of Br2
• what colour solution is formed when I- is displaced

  brown solution of I2
• CHLORINE GAS + SODIUM HYDROXIDE

  • if you mix chlorine gas with cold, dilute, aqueous sodium hydroxide you get sodium chlorate (I) solution (NaClO), sodium chloride (NaCl) and WATER!!!
  • in this reaction, chlorine is oxidised and reduced (disproportionation)
  • sodium chlorate (I) solution is household bleach that can kill bacteria
• USES OF SODIUM CHLORATE (I) SOLUTION (BLEACH)

  • water treatment
  • bleaching paper and textiles
  • cleaning
• REACTION BETWEEN CHLORINE AND WATER

  NORMALLY

  • it undergoes disproportionation (Cl is both oxidised and reduced)
  • the products are a mixture of chloride, chlorate (I), and hydrogen ions
  • the reaction is REVERSIBLE

  IN SUNLIGHT

  • chlorine can decompose water
  • forms chloride ions, hydrogen ions and OXYGEN!!
  • the reaction is REVERSIBLE
• WHY DO WE ADD CHLORINE TO DRINKING/SWIMMING WATER

  • chlorate ions kill bacteria
  • so adding chlorine or a chlorate (I) containing compound to water can make it safe to drink/swim in
  • HOWEVER...chlorine is TOXIC
• why is chlorine used in UK water treatment

  • chlorine kills disease causing micro-organisms
  • some chlorine persists in the water and prevents reinfection further down the supply
  • chlorine prevents the growth of algae (eliminating bad tastes/smells)
  • chlorine removes discolouration caused by organic compounds
• risks of using chlorine in UK water treatment

  • chlorine gas can irritate the respiratory system if breathed in
  • chlorine liquid on skin or eyes causes chemical burns
  • chlorine can react with organic compounds to form chlorinated hydrocarbons
  • many chlorinated hydrocarbons are carcinogenic
  • HOWEVER the cancer risk is smaller than the risks from untreated water
• what happens to REDUCING POWER of HALIDES down the group

  • IT INCREASES
  • to reduce another species, the halides must lose an electron
  • this depends on how strongly the electron is attracted to the nucleus
  • the attraction gets weaker down the group because:
  • the ionic radius increases, increasing the distance
  • there are extra shells and electrons that shield the electron
• F- AND Cl- REACTIONS WITH CONC. H2SO4

  • they are not strong enough reducing agents and cannot reduce S in H2SO4
  • there is no redox reaction, but there is acid + base reaction where nothing happens to charges or oxidation states
  • HF and HCl gases that are produced produce WHITE STEAMY FUMES that are ACIDIC and TOXIC
• Br- REACTION WITH CONC. H2SO4

  • stronger reducing agent
  • after the initial acid + base reaction the Br- ions reduce S from +6 to +4
  • HBr produced in acid + base reaction produces WHITE STEAMY FUMES that are ACIDIC and TOXIC

  REDOX:

  • Br is oxidised and S is reduced
  • the reaction produces CHOKING FUMES of SO2, ORANGE FUMES OF Br2 and water
• IODIDE REACTION WITH CONC. H2SO4

  • strongest reducing agent
  • after acid + base reaction and reducing S once, HI reduces S again from +4 to -2
  • HI also produces white, steamy, acidic and toxic fumes

  2ND REDOX:

  • I is oxidised and S is reduced
  • the reaction produces H2S gas which is TOXIC and smells like ROTTEN EGGS, SOLID IODINE which is a dark grey solid that sublimes with purple fumes, and water
• TEST FOR PRESENCE OF HALIDE IONS

  • initially add dilute nitric acid to remove unwanted ions
  • then add 2/3 drops of silver nitrate solution (AgNO3)
  • a precipitate is formed of the silver halide

  the colour of the precipitate determines the halide:

  F- = NO precipitate

  Cl- = white

  Br- = cream

  I- = yellow

  the yellow precipitate forms fastest
• TEST TO DETERMINE WHICH HALIDE ION IS PRESENT

  you can test whether the halide present is I-, Br- or Cl- using ammonia solution - each silver halide has a different solubility in ammonia

  Cl- = soluble in dilute ammonia

  Br- = soluble in conc. ammonia

  I- = insoluble in conc. ammonia (least soluble)
• formula of chlorate (I) ION

  ClO-
• SO2

  • sulfur dioxide
  • acidic (therefore can be removed from flue gases through wet scrubbing)
  • colourless and choking gas
  • S has oxidation state of +4
• I2

  • molecular lattice structure due to van der waals forces between molecules
  • dark grey solid at room temperature that sublimes with purple fumes
  • when formed after I- is displaced, I2 forms a brown solution
• Br2

  • orange fumes when a gas
  • red/brown liquid
  • forms orange solution when formed after Br- is displaced
• Sulfur

  • yellow solid
• H2S

  • hydrogen sulphide
  • toxic gas
  • smells of rotten eggs
• MOCK EXAM CORRECTION

  half equation of iodine from iodide

  2I- -----> I2 +2e-
• MOCK EXAM CORRECTION

  half equation for the formation of H2S from H2SO4

  8e- + 8H+ + H2SO4 ----> H2S + 4H20