Inorganic chem (2) Chem unit 2 Pt 4

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Inorganic chem (2) Chem unit 2 Pt 4
2012-04-04 08:10:55
chemistry unit inorganic halogens

Starting with halogens!
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  1. Give the colours of solution when chlorine, bromine and iodine is in water.
    • Chlorine - virtually colourless
    • Bromine - Yellow/orange
    • Iodine - Brown
  2. Give the colours of solution when chlorine, bromine and iodine dissolved in hydrocarbon solvents (eg. hexane).
    • Chlorine - virtually colourless
    • Bromine - Orange/red
    • Iodine - Pink/violet
  3. What is the physical state (in standard conditions) and colour of chlorine, bromine and iodine?
    • Chlorine gas is green
    • Bromine liquid is red-brown
    • Iodine solid is grey
    • (Also fluorine atom before chlorine is gas and is pale yellow)
  4. Explain the trend in reactivity down the halogen group.
    • Decrease in reactivity down the group. (Becomes less oxidising down the group).
    • Because atom becomes larger as you go down, so outer electrons further from nucleus, more shielding from the nucleus. Thus makes it harder for larger atoms to attract the electron needed.
  5. Halogens undergo _______ with alkali solution. Give the products for cold and hot solutions.
    • Disproportionation
    • COLD - halate (I), halide and water
    • HOT - halate (V), halide and water
    • eg. of Bromine with Sodium hydroxide
    • COLD - NaBrO + NaBr + H2O
    • HOT - NaBrO3 + 5NaBr + 3H2O
  6. When halogens react with metals, what does it do to the metal? Give example of how chlorine, bromine and iodine react with hot iron.
    • It oxidises metals. (gets reduced itself)
    • As halogens become less oxidising down the group;
    • Chlorine and bromine react with iron to form iron (III) halides.
    • Iodine react with iron to form iron (II) - it is less oxidised.
  7. Chlorine reacts with Phosphorus (non-metal) to produce what? In excess chlorine, what is produced?
    • PCl3
    • PCl5 (in excess chlorine)
  8. When a halogen (except iodine which is less oxidising) is added to Fe2+ ions (Fe(II)) what happens?
    • Fe2+ is oxidised to Fe3+ ions in solution.
    • The solution will change colour from green to orange.
  9. The reducing power of halides ______ down the group. Why?
    • Increases down the group.
    • By losing an electron from outer shell.
    • Because ions get bigger, electrons further away, also more shielding from inner electrons.
  10. Describe the reaction of KCl with H2SO4. What is produced and why?
    • HCl gas is produced - misty fumes
    • But HCl is not a strong enough reducing agent (remember reducing power increases down group) so it won't reduce sulfuric acid. Reaction stops there.
    • It is not a redox reaction (oxidation states have not changed)
    • Products - KHSO4 + HCl
  11. Describe reaction of KBr with H2SO4. What is produced and why?
    • Br2 fumes - orange - is produced. (along with choking fumes of SO2)
    • Because HBr is a stronger reducing agent and can reduce H2SO4 in a redox reaction.
    • Br2, SO2, H2O produced.
  12. Describe reaction of KI with sulfuric acid.
    • Same initial reaction as KBr and KCl, giving HI gas. This gas then reduces H2SO4 (like KBr).
    • But HI further reduces SO2 to H2S.
    • H2S gas is toxic and smells of rotten eggs.
    • 2HI + H2SO4 ---- I2 + SO2 + 2H2O
    • 6HI + SO2 ------ H2S + 3I2 + 2H2O
  13. Give features of hydrogen halides.
    • Colourless gas
    • Very soluble - dissolveing in water to make strong acids. (turn blue litmus red)
    • React with ammonia gas to give white fumes (NH4Cl)
  14. In a displacement reaction of halide ions, what can displace what?
    • Chlorine can displace bromide and iodide ions. (most oxidising). - Br2 (orange solution) formed, or I2 soluton (Brown solution) formed.
    • Bromine can displace iodide ions.
    • Iodine cannot displace any of the above (weak oxidising agent)
    • eg of ionic equation - Cl2 + 2Br- --- 2Cl- + Br2
    • You can make changes easier to see by mixing with organic solvent like hexane (halogen will dissolve, but halides won't - 2 layers formed)
    • A halogen will displace halide from solution if halide is below it in periodic table.
  15. How can you test whether it's Cl-, Br- or I- ? Results?
    • Add dilute nitric acid (remove ions that might interfere), then add silver nitrate solution. A precipitate will form.
    • Chloride - white precipitate, dissolves in dilute NH3
    • Bromide - cream precipitate, dissolves in conc NH3
    • Iodide - yellow precipitate, insoluble in NH3
  16. How do silver halides (eg. AgBr) react with sunlight?
    • Silver halides decompose when light shines on them.
    • Producing silver and the halogen.
    • eg. 2AgBr ----- 2Ag + Br2
  17. Halide ions are?
    • Reducing agents
    • More strongly reducing as you go down group.
    • eg. 2Fe3+ + 2I- ----- 2Fe2+ + I2
    • (Iodide can reduce, but chlorine can't - more reducing as you go down group - may need to predict)
  18. What equipment would you use to measure out the known solution in a flask during preparation of titration? And what would you use to add the unknown solution during the titration?
    • Pipette - can only measure one volume of solution (therefore in preparation)
    • Burette - can measure different volumes, and let you add solution drop by drop. (therefore during titration)
  19. What are the 2 main indicators used for acid/alkali titrations?
    • Methyl orange - yellow to red when adding acid to alkali.
    • Phenolphthalein - red to colourless when adding acid to alkali.
    • (Used because colour change occurs very quickly over a very small pH range)
  20. How do you work out the concentration from titrations?
    • Write out balanced equation
    • Work out moles for known solution
    • Use molar ratio to work out moles of unknown solution
    • Use volume and moles to work out concentration.
    • (Remember, concentration is in mol dm-3 )
  21. What is a standard solution?
    One whose concentration is known and does not change over time.
  22. What do you need to do to convert mol dm-3 into g dm-3 ?
    • Use formula - moles = mass / molar mass
    • So, multiply the moles by molar mass to get grams.
    • eg. 0.36 mol dm-3 of NaOH will be 0.36 x 40 = 14.4 g dm-3
  23. What are some uncertainties found when measuring substances during titrations? What is a good measure of uncertainty?
    • Good measure of uncertainty is maximum possible error (eg. uncertainty from a burette that marks every 0.1cm3 is maximum error of 0.05cm3)
    • Uncertainty of weighing substances (eg. nearest 0.01g - real mass could be 0.005g smaller or larger)
    • Pieces of equipment measuring liquid such as fixed-volume pipettes and volumetric flasks. (Manufacturer provide these uncertainty values).
  24. Outline some methods to minimise some uncertainties.
    • Buy most precise equipment available (though this is not easily done)
    • Check accuracy of pipette by transferring its contents to a weighed beaker and find mass and density to work out exact volume delivered.
    • For any reading/measurement you can calculate percentage uncertainty using equation - (uncertainty/reading) x 100.
    • As this shows, the larger the volume being measured, the less percentage uncertainty there will be - so plan titration with larger volume.
    • Same principle can be applied to other measurements like weighing solids.
  25. Outline the two types of errors.
    • Systematic errors: are the same every time experiment is repeated. May be caused by set-up or equipment used. (eg. 10cm3 pipette might actually be measuring 9.95cm3 leading to inaccuracies every time).
    • Random errors: different every repeat - will sometimes be above the real value, or sometimes it could be below. (eg. random errors measuring burette reading).
    • Repeating experiment will deal with errors (high values cancel out low values) but not systematic errors. Results get more reliable, but not more accurate.
  26. How do you calculate the total uncertainty in the final result?
    • Find percentage uncertainty fro each part of experiment (mainly volume measurement)
    • Add individual percentage uncertainties together. This gives percentage uncertainty in final result.
    • Use this to work out actual total uncertainty in the final result (this will be in mol dm3 if final result is a concentration)
  27. Suggest some causes of errors in a titration.
    • Air bubbles in burette
    • Contaminated equipment
    • Parallax when reading meniscus line
    • Impurities in solution (no substance will be 100%)
    • Errors in transferring substances (some may be left in container etc)
    • All equipment calibrated to be used at 20oC - if lab is warmer or colder - burette may not be accurate.
    • Balance/equipment will only show to a certain decimal place.
    • Judging end point of titration
  28. Judgement of the end point of an Iodine-Sodium Thiosulfate titration by?
    Adding a few drops of 1% starch solution when solution colour has become very pale. (Dark blue colour will suddenly go colourless at the end point)
  29. What can the iodine-sodium thiosulfate titration be used to calculate?
    • The concentration of an oxidising agent (such as potassium iodate (V) )
    • Percentage purity of potassium iodate (V)
  30. Outline how you could find out conc of oxidising agent from sodium-iodine thiosulfate titration. HOWEVER CHECK WITH TEACHER!
    • React a certain volume of unknown conc of potassium iodate (V) (oxiding agent) with excess potassium iodide solution. [Iodate(V) ions oxidise some of the iodide ions to iodine]
    • Titrate resulting solution with sodium thiosulfate (known conc). To "sharpen" result - add starch solution at end.
    • Use equation to work out moles of iodine in the solution.
    • Use original balanced equation and moles of iodine to find out concentration of potassium iodate (V) solution.
  31. Outline what you need to find out in each step of the iodine/thisulfate titration to find out percentage purity of potassium iodate (V).
    • Moles of Sodium thiosulfate used
    • From this, moles of iodine that reacted with this.
    • From this, moles of Iodate (V) ions involved in producing this iodine.
    • From this, mass of potassium iodate used.
    • Percentage purity = (mass KIO3 calculated / mass of crude KIO3) x 100
  32. Give the ionic equation for potassium iodate (V) reacting with acidified potassium iodide solution. And also the ionic equation of the product of this reaction reacting with thiosulfate.
    • IO3-(aq) + 5I-(aq) + 6H+(aq) ---- 3I2(aq) + 3H2O(l) .
    • I2 + 2S2O32- ----- 2I- + S4O62-