2.2 G2 and G7

2.2 G2 and G7

G7: the halogens
1. properties
  1. highly reactive
  2. non metals
  3. fluorine (f2)
    1. pale yellow gas
  4. chlorine
    1. green gas
  5. bromine
    1. red-brown liquid
  6. iodine
    1. grey solid
  7. boiling point increase down group
    1. size of atoms increases so van der waals forces also increase
  8. electronegativity decreases down group
    1. atoms increase in size so electrons in bonding pair are further from the nucleus and are more shielded so less attraction
2. displacement of halide ions
  1. down the group the oxidising power of the halides decreases
    1. the electrons are further from the nucleus so less attracted to it
  2. a halogen will displace a halide if the halide is below it in the periodic table
    1. chlorine will displace bromide and iodide
      1. forms orange solution of bromine
        Cl2 + 2NaBr --> 2NaCl + Br2
    2. bromine will displace iodide
      1. produces brown solution of iodine
        Br2 + 2NaI ---> 2NaBr + I2
3. uses of chlorinw
  1. bleach
    1. mix chlrine gas with cold, dilute sodium hydroxide and you get sodium chlorate solutiion
      1. 2NaOH + Cl ---> NaClO + NaCl + H2O
    2. kill bacteria
    3. used in water treatment
    4. used to bleach paper and textiles
  2. water treatment
    1. when mixed with water chlorine undegoes disproportionation to produce chloride and chlorate ions
      1. Cl2 + 2H2O <--> 2H+ + Cl_ +ClO
    2. n sunlight chlorine decomposes water to form chloride ions and oxygen
      1. Cl2 + 2H2O <--> 4H+ +"Cl- = O2
    3. benefits
      1. kill disease carrying bacteria
      2. some remains in water to prevent reinfection
      3. prevents algae growth
      4. prevents bad smells
      5. removes dicolouration
    4. risks
      1. chlorine gas is harmful if breathed in
      2. liquid chlorine can cause chemical burns
      3. chlorine reacts with organic compounds and forms chlorinated hydrocarbons which are carcinogenic
    5. used because benefits outweigh risks (in Q be specific)
4. halide ions
  1. test for halide ions
    1. add silver nitrate and a precipitate forms
      1. chloride - white precipitate forms
        dissolves in dilute ammonia
      2. bromide- cream precipitate forms
        dissolves in concentrated ammonia
      3. iodide- yellow precipitate forms
        doesnt't dissolve in ammonia
      4. fluoride - no precipitate forms
  2. reducing power of halides
    1. reducing power decrreases down the group
      1. to reduce something the halide must lose and electron form its outer shell
        becomes easier down the group because the outer electrons are further from the nucleus and has more shielding do the attraction is less
    2. all halides react with H2SO4 to form NaXHSO4 and HX
      1. fluoride and chloride
        only do the acid based reaction
        produce whit misty fumes
      2. NaX + H2SO4 --> NaXHSO4 + HX
        this is an acid based reaction
        not redox as oxidation states stay the same
      3. iodide
        undergoes same reaction as bromide
        reduces SO2 to H2S
        6HI + SO2 --> H2S +3I2 + 2H2O
        H2S is toxic and smells like bad eggs
        can also produce sulfu
        yellow solid
      4. bromide
        HBr acts as a reducing agent and further reduces the H2SO4
        2HBr + H2SO4 --> Br2 + SO2 + 2H2O
        produces invisible choing gas aand orange Br2 fumes
Group 2: Alkalaine earth metals
1. Trends
  1. atomic radius increases down group
    1. electron shells are added as you go down the group so the outer electrons are further from the nucleus
  2. first ionisation energy decreases down group
    1. outer electrons are further from the nucleus and are more shielded so less energy is requred to remove the outer elctron
  3. reactivity increases down group
    1. when group 2 elements react they lose two electrons
      1. as you go down the group these electrons are further away from the nucleus and have more shielding so are more easily removed
  4. melting points decrease down the group
    1. down the group the metal ions increase in sze
      1. down the group the atomic radius increases so the electrons are furtther from the nuclei so have less of an attraction
        the attraction is more easily overcome so themelting point is lower
    2. there are 2 delocalised electrons per metal ion
2. reactions with water
  1. form a metal hydroxide and hydrogen
    1. they react more readily down the group berylium doesnt react whereas barioum reacts very rapidly
    2. solubility of hydroxides increases down the group
  2. magnesium reacts with steam rapidly but slowly with water
  3. water acts as an oxidising agent
  4. M + 2H2O --> M(OH)2 + H2O
  5. uses
    1. Magnesium hydroxide is an antacid- used in tablet form to neutralise stomach acid
    2. calcium hydroxide is used in agriculture to neutralise acidic soil
3. reactions with sulfate
  1. uses
    1. Barium sulfte is used in "barium meals" which are used to highlight soft tissue problems on x-rays
  2. reactions get faster down the group
  3. metal sulfate solubility decreases down the group (barium is insoluble)
4. uses
  1. magnesium is used in the extraction of titanium from its ore
    1. titanium ore (TiO2) is heated with carbon in a stream of chlorine gas this produce TiCl4
      1. this is then reduced by magnesium at 1000 degrees celcius to produce titanium and magnesium chloride
        TiCl4 +2Mg --> Ti + 2MgCl2
  2. calcium oxide or calcium carbonate are used to remove sulfur dioxide from flu gases
    1. this is known as wet scrubbing
    2. stops sulfur dioxide being released into the atmosphere and causing acid rain
    3. a slurry is produced by combining the calcium compound with wter and it is sprayed onto the flue gases
      1. produces calcium sulfite which can be sold on an used in plaster board
      2. CaO + 2H2O +SO2 --> CaSO3 + 2H2O
        OR: CaCO3 + 2H2O + SO2 --> CaSO3 +2H2O + CO2

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	Created by Lauren Cullinan over 5 years ago