Inorganic Chemistry

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Jeeten
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12832
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Inorganic Chemistry
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2010-05-30 14:49:45
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Inorganic Chemistry
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Inorganic Chemistry
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  1. What does OIL RIG stand for?
    • Oxidation is losing
    • Reduction is gaining
  2. A loss of electrons is called...
    Oxidation
  3. A gain in electrons is called...
    Reduction
  4. When reduction and oxidation happen simultaneously, it is called a...
    Redox reaction
  5. An oxidising agent...
    Accepts electrons and gets reduced
  6. A reducing agent...
    Donates electrons and gets oxidised.
  7. All atoms are treated as ____ when talking about oxidation numbers, even if they are covalently bonded.
    Ions
  8. Uncombined elements have an oxidation number of...
    0
  9. Elements just bonded to identical atoms, like O2 and H2 have an oxidation number of...
    0
  10. The oxidation number of simple monatomic ion such as Na+ is...
    The same as it's charge. (+1 for for Na example)
  11. In compounds/compound ions, the overall oxidation number is...
    the ion charge.
  12. The sum of the oxidation numbers for a neutral compound is...
    0
  13. Combined oxygen is nearly always ___, except in peroxides where it's ___.
    And in fluorides OF2 where its ___ and O2F2 where it's ___
    • -2
    • -1
    • +2
    • +1
  14. Combined hydrogen is ___ except in metal hydrides where it is ___.
    • +1
    • -1
  15. If you see Roman numerals in a chemical name, it's an...
    Oxidation number
  16. The oxidation number for an atom will _____ by 1 for each electron lost.
    Increase
  17. The oxidation number for an atom will _____ by 1 for each electron gained.
    Decrease
  18. Elements which are oxidised and reduced by the same time is called..
    Disproportionation.
  19. Ionisation energy ____ down group 2. What are the 3 reasons for this?
    • 1 - Each element down group 2 has an extra electron shell compared to the one above.
    • 2 - The extra inner shells shield the electrons from the attraction of the nucleus.
    • 3 - Also, the extra shell means that the other electrons are further away from the nucleus, which greatly reduces the nucleus's attraction.
  20. When group 2 elements react, they're oxidised from a state of ___ to ___ forming M__ ions. This is because Group 2 atoms contain...
    • 0
    • +2
    • M2+
    • 2 electrons in their outer shell.
  21. The Group 2 metals react with water to give a _____ ______ and _____.

    They get increasingly reactive down the group because the ionisation energies ____.
    • Metal hydride
    • Hydrogen
    • Decrease
  22. Group two metals burn in oxygen with...
    Characteristic flame colours.
  23. Group 2 metals react with Cl to form...
    White solid chlorides.
  24. The oxides and hydroxides are ____.
    Bases.
  25. The oxides of Group 2 metals react readily with ____ to form ___ _____, which dissolve. The hydroxide ions make these solutions...
    • Water
    • Metal hydroxides
    • Strongly alkaline

    MgO is an exception - it only reacts slowly and the hydroxide isn't very soluble.
  26. The oxides form more _____ alkaline solutions as you go down the group, because the hydroxides get more ____.
    • Strongly
    • Soluble.
  27. Because oxides are bases, both the oxides and hydroxides will _____ dilute acids, forming solutions of the corresponding salts.
    Neutralise.
  28. What is the general equation of when oxides react with water?
    MO (s) + H20 (l) → M(OH)2 (aq)
  29. What is the general equation of when oxides react with an acid? (Cl in this case)
    MO (s) + 2HCl (aq) → MCl2 (aq) + H20 (l)
  30. What is the general equation of when hydroxides react with water?
    • +H20
    • M(OH)2 (s) → M2+ (aq) + 2OH- (aq)
  31. What is the general equation of when hydroxides react with acid?
    M(OH)2 (aq) + HCl (aq) → MCl2 (aq) + 2H20 (l)
  32. Solubility trends depend on the _____ _____.
    Compound Anion.
  33. Generally, compounds of Group 2 elements that contain singly charged negative ions (e.g OH-) ______ in solubility down the group.
    Increase
  34. Compounds that contain _____ charged negative ions (e.g SO4-2) decrease insolubility down the group.
    Doubly
  35. Group 2 element Hydroxide (OH-) Sulfate (SO42-)
    Mg
    Ca
    Sr
    Ba
    • Group 2 element Hydroxide (OH-) Sulfate (SO42-)
    • Mg
    • Ca ↓ ↑
    • Sr ↓ ↑
    • Ba
  36. Most sulfates are soluble in water but _____ sulfate is insoluble.

    Compounds like MgOH that have very low solubilities are said to be ______ soluble.
    Barium

    Sparingly
  37. Thermal stability of carbonates and nitrates ____ down the group.
    Changes.
  38. Thermal decomposition is when a substance _____ ____ (decomposes) when ____. The more thermally stable a substance is, the more heat it will take to break it down.
    • Broken down
    • Heated
  39. Thermal stability _____ down a group.
    Increases.
  40. The carbonate and nitrate ions are large and can be made unstable by the presence of a ______ ______ ____. (A cation) The cation polarises the anion, distorting it. The greater the distortion, the less stable the anion.

    Large cations cause ____ distortion than small cations. So the further down the group, the larger the cations, the ___ distortion caused and the ____ stable the carbonate/nitrate ion.
    Positively charged ion.

    • Less
    • Less
    • More
  41. Group 2 compounds are ____ thermally stable than Group 1 compounds.
    Less
  42. The greater the ___ on the cation, the greater the _____ and the ___ stable the carbonate/nitrate ion becomes.

    Group 2 cations have a __ charge, compared to a __ charge Group 1 cations.
    • Charge
    • Distortion
    • Less

    • 2+
    • 1+
  43. Group 1 carbonates are ______ ____ - you can't heat them enough with a bunsen burner to make them decompose (though they do decompose at higher temperatures.
    Thermally stable

    Except Li2CO3 which decomposes to Li2O and CO2
  44. Group 2 carbonates decompose to form the ____ and _____ _____. Give an example:
    Oxide and Carbon Dioxide

    MCO3 (s) → MO (s) + CO2 (g)
  45. Group 1 nitrates decompose to from the _____ and _____.

    Give an example.
    Nitrate and oxygen.

    2MNO3 (s) → 2MNO2 (s) + O2 (g)

    Except LiNO3 which decomposes to form LiNO, NO2 and O2
  46. Group 2 nitrates decompose to form the ____, ____ _____ and _____.

    Give an example:
    • Oxide
    • Nitrogen dioxide
    • Oxygen

    2M(NO3)2 (s) → 2MO (S) + 4NO2 (g) + O2 (g)
  47. How easily nitrates decompose and be tested by measuring...
    Hw long it takes until oxygen is produced. (i.e to relight a glowing split)

    OR

    How long it takes until a brown gas (NO2) is produced. this needs to be done in a fume cupboard because NO2 is toxic.
  48. How easily carbonates decompose can be tested by measuring...
    How long it takes for carbon dioxides to be produce. You test for carbon dioxide using lime water - which is a saturated solution of calcium hydroxide. This turns cloudy with carbon dioxide.
  49. What is the characteristic flame colour associated with Li?
    Red
  50. What is the characteristic flame colour associated with Na?
    Orange/Yellow
  51. What is the characteristic flame colour associated with K?
    Lilac
  52. What is the characteristic flame colour associated with Rb?
    Red
  53. What is the characteristic flame colour associated with Cs?
    Blue
  54. What is the characteristic flame colour associated with Ca?
    Brick red
  55. What is the characteristic flame colour associated with Sr?
    Crimson
  56. What is the characteristic flame colour associated with Ba?
    Green
  57. How do you carry out a flame test?
    Mix a small amount of the compound you are testing with a few drops of hydrochloric acid

    Heat a piece of platinum or nichrome wire in a hot Bunsen flame to clean it

    Dop the wire into the compound/acid mixture. Hold it in a very hot flame and note the colour produced.
  58. What is the origin of a flame test?
    The energy absorbed from the flame causes electrons to move to higher energy levels. The colours are seen as the electrons gall back down to lower energy levels, releasing energy in the form of light, The difference in energy between the higher and lower levels determines the wavelength of the light released which determines the colour of the light.
  59. The movement of electrons between energy levels is called ______ ______.
    Electron Transition.
  60. Halogens are the highly reactive non-metals of...
    Group 7
  61. The word halogen should be used when describing the atom __ or the molecule ___, but the word halide is used to describe the ______ ___.
    • X
    • X2
    • Negative Ion (X-)
  62. What are the main properties (Formula, colour, physical state and electronic structure) of Fluorine.
    • F2
    • Pale yellow
    • Gas
    • 1s2 2s2 2p5
  63. What are the main properties (Formula, colour, physical state and electronic structure) of Chlorine.
    • Cl2
    • Green
    • Gas
    • 1s2 2s2 2p6 3s2 3p5
  64. What are the main properties (Formula, colour, physical state and electronic structure) of Bromine.
    • Br2
    • Red-Brown
    • Liquid
    • 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p5
  65. What are the main properties (Formula, colour, physical state and electronic structure) of Iodine
    • I2
    • Grey
    • Solid
    • 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 5s2 5p5
  66. Halogens in their natural state exist as covalent diatomic molecules. Because they are covalent, they have ___ solubility in water.
    Low
  67. Halogens easily dissolve in ______ compounds like hexane. Some of these resulting solutions have distinctive colours that can be used to identify them.
    Organic.
  68. What is the colour of Cl2 in water and in hexane?
    • In water - Virtually Colourless
    • Hexane - Virtually Colourless
  69. What is the colour of Br2 in water and in hexane?
    • In water - Yellow/Orange
    • In hexane - Orange/Red
  70. What is the colour of I2 in water and in hexane?
    • In water - Brown
    • In hexane - Pink/Violent
  71. Halogens get ___ reactive down the group.
    Less
  72. Halogen atoms react by ____ an electron in their outer p sub-shell. This means they're ____ and they ____ another substance (it's a redox reaction) - so they're _____ agents.
    • Gaining
    • Reduced
    • Oxidise
    • Oxidising
  73. As you down down Group 7, the atoms become _____ so the outer electrons are further from the nucleus. The outer electrons are also _____ more from the attraction of the positive nucleus, because there are more inner electrons. This makes it harder for larger atoms to attract the electron needed to form an ion, so larger atoms are less _____.
    • Larger
    • Shielded
    • Reactive
  74. The smallest halogen, _____ is the most reactive non-metal element.
    Fluorine
  75. The melting and boiling points of Group 7 _____ down the group. So you can predict that fluorine would be a ___ as room temperature, like chlorine below it.
    • Increase
    • Gas
  76. Halogens undergo _______ with alkalis.
    Disproportionation.
  77. Give the equation and ionic equation of the reaction between halogen X2 and the cold alkali NaOH.
    Equation : X2 + 2NaOH → NaXO + NaX +H20

    Ionic Equation : X2 + 2OH- → XO- + X- + H20
  78. Give the equation between the halogen X2 and the hot alkali NaOH.
    Equation : 3X2 + 6NaOH → NaXO3 + 5NaX + 3H2O

    Ionic Equation : 3X2 + 6OH- → XO3- + 5X- + 3H2O
  79. All the halogen -ate ions have a single halogen atom and a charge of __.
    -1
  80. When halogens react they are reduced - and they ____ other substances.
    Oxidise.
  81. Bromine is a ____ oxidising agent so you get a mixture of Iron (II) and iron (III) bromide. WIth iodine you only get iron (II) iodide - no Fe3+ ions form.
    Weaker
  82. Halogens oxidise metals. Give the reaction between the metal Fe and the halogen Cl
    2Fe(s) + 3Cl2 (g) → 2FeCl3 (s)

    • Fe is oxidised: 2Fe → 2Fe3+ + 6e-
    • Cl is reduced: 3Cl2 + 6e- → 6Cl-

  83. Halogens react with non metals. Give the equation of the reaction between Cl and S.
    S8(s) + 4Cl2(g) → 4S2Cl2(l)

    • S is oxidised to +1
    • Cl is reduced to -1
  84. Halogens also react with some ions. All the halogens except iodine (which is less strongly oxidising than the others) will oxidise iron (II) ions to _________ ions in solution. The solution will change colour from ____ to _____.
    • Iron (III)
    • Green
    • Orange
  85. A halide ion can react as a _____ agent by losing an electron from it's outer shell. How easy this is depends on the attraction between the halides _____ and the outer _____. As you go down the group, the attraction gets ____ because:
    • Reducing
    • Nucleus
    • Electrons
    • Weaker
    • 1) The ions get bigger, so the electrons are further away from the positive nucleus
    • 2) There are extra inner electron shells, so there's a greater shielding effect
  86. This explains their reactions with Sulphuric Acid. All the halides react with conc. sulphuric acid to give a ____ halide as a product to start with. But what happens next depends on which halide you've got.
    Hydrogen
  87. What is the reaction between KF or KCl with H2SO4 ? And explain how it would happen and what you would see at each step of the process and what type of process it is.
    KF (s) + H2SO4 (l) → KHSO4 (s) + HF (g)

    KCl (s) + H2SO4 (l) → KHSO4 (s) + HCl (g)

    • 1) HF or HCl is formed. You'll see misty fumes as the gas comes in contact with moisture in the air.
    • 2) But HF and HCl aren't strong enough reducing agents to reduce the sulphuric acid so the reaction stops there.
    • 3) It is not a redox reaction - the oxidation states of the halide and sulphur stay the same (-1 and +6)
  88. What is the reaction between KBr and H2SO4 (2 equations), what happens at each step and what you would notice, and the type of reaction?
    • KBr (s) + H2SO4 (l) → KHSO4 (s) → + HBr (g)
    • 2HBr (aq) + H2SO4 (l) → Br2 (g) + SO2 (g) + 2H2O (l)

    Ox state of S : +6 → +4 (Reduction)

    Ox state of Br : -1 → 0 (Oxidation)

    • 1) The first reaction gives misty fumes of HBr
    • 2) But the HBr is a stronger reducing agent than HCl and reacts with the H2SO4 in a redox reaction.
    • 3) The reaction produces choking fumes of SO2 and orange fumes of Br2
  89. What is the reaction of KI with H2SO4?
    • KBr (s) + H2SO4 (l) → KHSO4 (s) → + HBr (g)
    • 2HBr (aq) + H2SO4 (l) → Br2 (g) + SO2 (g) + 2H2O (l)

    Ox state of S : +6 → +4 (Reduction)

    Ox state of Br : -1 → 0 (Oxidation)

    • 1) The first reaction gives misty fumes of HBr
    • 2) But the HBr is a stronger reducing agent than HCl and reacts with the H2SO4 in a redox reaction.
    • 3) The reaction produces choking fumes of SO2 and orange fumes of Br2
  90. What is the reaction between KI and H2SO4?
    KI (s) + H2SO4 (l) → KHSO4 (s) + HI (g)

    2HI (g) + H2SO4 (l) → I2 (s) + SO2 (g) + 2H20 (l)

    • Ox state of S : +6 → +4 (Reduction)
    • Ox state of I : -1 → 0 (Oxidation)


    6HI (g) + SO2 (g) → H2S (g) +3I2 (s) + 2H20 (l)

    • Ox state of S : +4 → -2 (Reduction)
    • Ox state of I : -1 → 0 (Oxidation)

    • 1) Same initial reaction giving HI gas
    • 2) The HI then reduces H2SO4 as before.
    • 3) But HI keeps going and reduces the SO2 to H2S. The H2S gas is toxic and smells of bad eggs.
  91. The hydrogen halides are _____ gases.
    Colourless
  92. They're very ___, dissolving in water to make strong ____. They turn ___ litmus paper ___.
    • Soluble
    • Acids
    • Blue
    • Red
    • HCl (g) → H+ (aq) + Cl- (aq)
  93. Hydrogen chloride forms _____ acid, hydrogen bromide forms _____ acid and hydrogen iodide gives _____ acid.
    • Hydrochloric
    • Hydrobromic
    • Hydroiodic
  94. They react with ammonia gas to give _____ _____.
    White fumes.

    • For example:
    • Hydrogen chloride gives ammonium chloride.

    NH3 (g) + HCl (g) → NH4Cl (s)
  95. Halide ions are ____ from solution by more _____ halogens.
    • Displaced
    • Reactive
  96. The halogens' relative ____ strengths can be seen in their displacement reactions with halide ions.

    For example, if you mix bromine water (Br2 (aq)) with KI solution, the bromine displaces the iodide ions (it oxidises them) giving the I2 (aq) and KBr (aq). You can see what happens by following the _______ changes.
    • Oxidising
    • Color changes
  97. What are the colour changes between Cl2 (aq) with KCl (aq), KBr (aq) and KI (aq)
    KCl (aq) - No reaction

    KBr (aq) - Orange solution (Br2) formed.

    KI (aq) - Brown solution (I2) formed.
  98. What are the colour changes between Br2 (aq) with KCl (aq), KBr (aq) and KI (aq)
    KCl (aq) - No reaction

    KBr (aq) - No reaction

    KI (aq) - Brown solution (I2) formed.
  99. What are the colour changes between I2 (aq) with KCl (aq), KBr (aq) and KI (aq)
    KCl (aq) - No reaction

    KBr (aq) - No reaction.

    KI (aq) - No reaction.
  100. You can make the changes easier to see by shaking the reaction mixture with an ______ ______ like hexane. The halogen that's present will dissolve readily in the ______ _____, which settles out as a distinct layer above the _____ solution.
    • Organic solvent
    • Organic solvent
    • Aqueous
  101. What are the ionic equations between Cl2, Br- and I- (Displacement reaction)
    Cl2 (aq) + 2Br- (aq) → 2Cl- (aq) + Br2 (aq)

    Cl2 (aq) + 2I- (aq) → 2Cl- (aq) + I2 (aq)
  102. What is the ionic equation between Br2 and I- (Displacement reaction)
    Br2 (aq) + 2I- (aq) → 2Br- (aq) + I2 (aq)
  103. What is the reaction between I and F-, Cl- and Br- ?
    No reaction
  104. A halogen will displace a halide from solution if the halide is _____ it in the periodic table.
    Below
  105. You can also say a halogen will oxidise a halide if the halide is below it in the periodic table.
    Show this using Cl and Br.
    Cl2 (aq) + 2Br- (aq)2Cl- (aq) + Br2 (aq)

    • 0 -1
    • -10
  106. Halides give coloured precipitates with ______ _____ solution.
    Silver Nitrate
  107. What is the test to help you find out which halide ion you are dealing with?
    • 1) First you add dilute nitric acid to remove ions that might interfere with the test.
    • 2) Then you add silver nitrate solution (AgNO3 (aq))
    • A precipitate is formed (of the silver halide)

    Ag+ (aq) + X- (aq) → AgX(s)

    Where X is F, Cl, Br or I
  108. The colour of the precipitate identifies the halide. You can test these results by adding ammonia solution because...
    Each silver halide has a different solubility in ammonia.
  109. What precipitate is formed and it's solubility in ammonia of F-, Cl-, Br- and I-?
    Fluoride F- : No precipitate

    Chloride Cl- : White precipitate, dissolves in dilute NH3 (aq)

    Bromide Br- : Cream precipitate, dissolves in conc. NH3 (aq)

    Iodide I- : Yellow precipitate, insoluble in conc. NH3 (aq)
  110. Silver halides react with ____.
    Sunlight
  111. Silver halides ____ when light shines on them producing silver and the halogen.

    Give the example of Ag and Br in this case.
    Decomposes

    2AgBr → 2Ag + Br2
  112. Titrations are used to find out the ______ of acid or alkali solutions. They're also handy wen you're making soluble salts of soluble bases.
    Concentration.
  113. Titrations need to be done accurately. They allow you to find out exactly how much acid is needed to ____ the alkali.
    Neutralise.
  114. How do you carry out a titration?
    1) You measure out some alkali using a pipette and put it in a flask, along with some indicator (such as phenolphthalein)

    2) Then you do a rough titration to get an idea where the end point is (the point where the alkali is exactly neutralised and the indicator changes colour). Add the acid to the alkali using a burette - giving the flask a regular swirl.

    3) Now you would carry out a accurate titration. Tun the acid into within 2cm3 of the end point, then add the acid dropwise. If you don't notice exactly when the solution changed colour then you've overshot and your result won't be accurate.

    4) Record the amount of acid used to neutralise the alkali. It's best the repeat this process a few times making sure you get the same answer each time.
  115. What would you use to carry out a titration?
    A burette
  116. Indicators show you when the reaction's just ____.
    Finished
  117. Indicators change _____. In titrations, indications that change colour quickly over a very ____ pH range are used so you know exactly when the reaction has ended.
    • Colour
    • Small
  118. The two main indicators for acid/alkali reactions are:
    Methyl orange - Turns yellow to red when adding acid to alkali

    Phenolphthalein - Turns red to colourless when adding acid to alkali
  119. 25 cm2 of 0.5 M HCl was used to neutralise 35 cm3 of NaOH solution. Calculate the concentration of the NaOH solution.
    First write a balanced equation and decide what you know and what you need to know:

    HCl (25cm3 0.5M) + NaOH (35cm3 xM) → NaCl + H20

    • No of moles HCl → (Conc. x Volume) / 1000
    • → (0.5 x 25) / 1000 = 0.0125

    • From the equation you know 1 mole of HCl neutralises 1 mole of NaOH.
    • So 0.0125 HCl must neutralise 0.0125 moles of NaOH

    • Conc. NaOH → (Moles NaOH x 1000) / Volume
    • → (0.0125 x 1000) / 35 = 0.36 mol dm-3
  120. Uncertainty is the amount of ____ your measurement might have.
    Error
  121. The _____ possible error is a useful measure of uncertainty.
    Maximum
  122. The uncertainty in yours measurement varies for different ____.
    Equipment.
  123. For example, the scale on a 50 cm3 burette has marks every 0.1 cm3. You should be able to tell which mark the level's closest to, so any reading you take won't be more than ____ cm3 out (as long as you don't make a mistake). The uncertainty of a reading from the burette is the maximum error you could have - so that's ___cm3
    • 0.05
    • 0.05
  124. There's uncertainty even when you weigh stuff too. Even electronic scales don't give an ____ mass. If the mass is measured to the nearest 0.01g the real mass could be up to ± _____ g
    • Exact
    • 0.005
  125. You can minimise some uncertainties. One obvious way is to ______ _____ in your measurements is to buy more _____ _____ available. In real life there's not much you can do about this, you're stuck with whatever you have.
    • Reduce Errors.
    • Precise Equipment.
  126. Planning can also improve your results. Think about the readings from a burette. You take two readings to work out a titre (the volume of liquid delivered from the burette) - the ____ volume and the ____ volume. Each reading has an uncertainty of ___ cm3. The titre is the second reading minus the first reading so the titre will have a total uncertainty of ___ cm3.
    • Initial
    • Final
    • 0.05
    • 0.1
  127. For any reading or measurement you can calculate the percentage uncertainty using the equation:
    Percentage Uncertainty = (Uncertainty / Reading) x 100
  128. Using the equation, what has the larger uncertainty, 10 cm3 or 20 cm3 ?
    (0.1 / 10) x 100 = 1%

    (0.1 / 20) x 100 = 0.5%

    • This shows that the larger the volume used, the uncertainty percentage decreases.
    • The same principle can be applied to other measurements such as weighing solids.
  129. Errors can either be ____ or ____.
    • Systematic
    • Random
  130. What are systematic errors?
    They are the same every time you repeat an experiment. They may be caused by the set-up or equiptment you're using. If the 10.00 cm3 pipette you're using to measure out a sample for titration actually only measures 9.95 cm3, your sample will be about 0.05 cm3 too small every time you repeat the experiment.
  131. What are random errors?
    Random errors vary. They're what make the results a big different each time you repeat an experiment. The errors when you make a reading from a burette are random. You have to estimate or round the level when it's between two marks - so sometimes your figure will be above the real one, and sometimes it will be below.
  132. Repeating an experiment and finding the mean of your results helps to deal with ____ errors. the results that are a bit high will be cancelled out by the new ones that are a bit low. (Your results will be more ____) But repeating your results won't get rid of any _____ errors. (Your results won't get more ____)
    • Random
    • Reliable
    • Systematic
    • Accurate.
  133. How do you find the total uncertainty in a final result?
    1) Find the percentage uncertainty for each bit of equiptment

    2) Add the individual percentage uncertainty together. This gives the percentage uncertainty in the final result.

    3) Use this to work out the actual total uncertainty in the final result.
  134. 10.00 cm3 of KOH solution is neutralised by 27.3 cm3 of HCl of known concentration.
    The volume of KOH has an uncertainty of 0.06 cm3. The volume of HCl has an uncertainty of 0.1 cm3.
    The concentration of the KOH is calculated to be 1.356 mol dm-3.
    What is the uncertainty in this concentration?
    First work out the percentage uncertainty for each volume measurement.

    The KOH volume of 10.00 cm3 has an uncertainty of 0.06 cm3 :

    (0.06 / 10.00) x 100 = 0.60 %

    The HCl volume of 27.3 cm3 has an uncertainty of 0.1 cm3

    (0.1 / 27.3) x 100 = 0.37%

    Find the percentage uncertainty in the final result:

    0.6 + 0.37 = 0.97%

    Uncertainty of the final answer:

    0.97% of 1.365 mol dm-3 = 0.013 mol dm-3
  135. Iodine-sodium thiosulfate titrations are a way of finding the concentration of an ____ agent. THe more ____ the oxidising agent is, the more ___ will be oxidised by a certain volume of it.
    • Oxidising
    • Concentrated
    • Ions
  136. What is the first stage of how you can find out the conc. of a solution of the oxidising agent potassium iodide (V) ?
    Stage 1 - Use a sample of oxidising agent to oxidise as much iodide as possible.

    Measure out a certain volume of potassium iodide (V) (the oxidising agent) - say 25 cm3

    Add this to an excess of acidic potassium iodide solution. The iodate (V) ions in the potassium iodide (V) solution oxidise some of the iodide ions to iodide.
  137. What is the second stage of how you can find out the conc. of a solution of the oxidising agent potassium iodide (V) ?
    Stage 2 - Find out how many moles of iodine have been produced.

    You do this by titrating the resulting solution with sodium thiosulfate. (You need to know the conc. of the sodium thiosulfate solution)

    The iodine in the solution tracts with the thiosulfate ions like this:

    I2 + 2S2O32- → SI- + S4O62-

    Titration of Iodine with Sodium Thiosulfate.

    • 1) Put all the solution from stage 1 in a flask.
    • 2) From the burette, add sodium thiosulfate solution to the solution in the flask
    • 3) It's had to see the end point, so when the iodine colour fades to pale yellow, add 2 cm3 of starch solution (to detect the presence of iodine). The solution in the conical flask will go dark blue showing theres still some iodine there.
    • 4) Add sodium thiosulfate one drop at a time until the blue colour disappears
    • 5) When this happens, it means that all the iodine has just been reacted
    • 6) Now you can calculate the number of moles of iodine in the solution.
    • You would then calculate the number of moles of iodine produced in stage one. In this case it equals 6.66 x 10-4 moles.
  138. What is the third stage of how you can find out the conc. of a solution of the oxidising agent potassium iodide (V) ?
    Calculate the concentration of the oxidising agent.

    You look back at your original equation:

    IO3- (aq) + 5I- (aq) + 6H+ (aq) → 3I2 (aq) + 3H2O (l)

    The equation shows that 1 mole of iodate (V) ions produces 3 moles of iodine. 25 cm3 of potassium iodate (V) solution produced 6.66 x 10-4 moles of iodine. So there must have been 6.66 x 10-4 / 3 = 2.22 x 10-4 moles of iodate (V) ions.

    There would be the same number of moles of potassium iodate (V) in the solution. So now it's straightforward to find the conc. of potassium iodate (V) solution:

    N = (CV/1000)

    2.22 x 10-4 = (Conc. x 25) / 1000

    Conc. of potassium iodate (V) solution = 0.00888 mole dm-3
  139. Using _____ apparatus could make your results inaccurate - so make sure the burette is very ___ and ____ it out with sodium thiosulfate before you can start because traces of water will dilute the solution.
    • Contaminated
    • Clean
    • Rinse
  140. It's important to ____ the burette correctly (from the _____ of the meniscus, with your eyes level to the liquid.
    • Read
    • Bottom
  141. To reduce the effect of random errors, ___ the experiment and take an ___.
    • Repeat
    • Average
  142. Remember to ___ the flask between experiments or use ____, ____ one.
    • Wash
    • New
    • Clean

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