Experiment A Quantitative Study of Electrolysis filler0

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					            S.6 Chemistry Laboratory Menu (2007 – 2008)
 Date      Experiment   Title
10/9/07        1        An acid-base titration

18/9/07        2        Standardization of hydrochloric acid by sodium carbonate

27/9/07        3        Calcium carbonate in eggshell

8/10/07        4        Analysis of two commercial brands of bleaching solution

16/10/07       5        An analysis of aspirin tablets

8/11/07        6        An analysis of commercial vitamin C tablets

                        Investigation the amount of vitamin C in different kinds of
16/11/07       7
                        vegetable or fruit juice
                        Application of Hess's law to determine the enthalpy change of
27/11/07       8
                        hydration of magnesium sulphate(VI)

5/12/07        9        Determination of dissolved oxygen in a water sample

                        Order of reaction between hydrogen peroxide and potassium
14/1/08       10
                        iodide in Acidic Medium
                        Investigating the order of reaction between iodine and
22/1/08        11
                        propanone
                        To determine the activation energy of the reaction between
30/1/08       12
                        bromide ions and bromate(V) ion in acid solution

20/2/08       13        Determination of the equilibrium constant of esterification

                        Determination of the partition coefficient of ethanoic acid
28/2/08       14
                        between water and 2-methylpropan-1-ol

11/3/08       15        Hydrogen bonding

24/4/08       16        Investigation of Some of the Properties of a Pair of Cis-Trans


                                Chem – is – try
                    S.6 Chemistry Experiment grouping list

Group                          Name/Class number/                    locker number
        A. Cheung Hung Man, Mandy(3)                (Titration)
  1                                                                       2
        B. Ip Sze Lok, Mandy(11)              (Electrochem)

        A. Lam Pui Pui, 佩佩(15)                (Redox reaction)
  2                                                                       4
        B. Yung Ping Yeung, Ayumi(33)         (Esterification)


        A. Ho Tsui Yan, Snoopy(7)             (Calculation)
  3                                                                       10
        B. Law Hiu Tung, Karen(18)       (Homologous series)


        A. Huang Lau Shan, 柳飽(9)         (detergent/hydrocarbons)
 4.                                                                       12
        B. Liu Wai Han, Agnes(20)        (Esterification)

        A. Yau Fung Lam, Candy(29)            (Chemical cell)
  5                                                                       16
        B. Lau Yan Yan, Yen(16)               (Atoms)

        A. Chloe Yip(31)                      (Plastics)
  6                                                                       22
        B. Priscilla Tso(24)                  (Balancing equation)

        A. 葉珮雅, Kitty(30)                     (Hydrocarbons)
  7                                                                       28
        B. Wong Ching Yee, Polly(25)          (Alkanol)
                                 FORMAT OF CHEMISTRY REPORT
Your report should include some or all of the following components as instructed by teachers
A typical report is divided into sections:

1.         Title of the experiment

2.         Objective(s) : clear statement(s) of the purpose(s) of the experiment

3.         Theory / Introduction / Background

4.         Chemicals and apparatus used

5.         Diagram for the set-up

6.         Procedures and Observations

7.         Results

8.         Treatment of data

9.         Questions / Discussions

10.        Conclusions

Notes on writing reports:
1.    In the section of introduction, students should write down the methods used to achieve the objectives,
      the principles involved and other relevant information BRIEFLY.
2.    You are not required to copy the procedures of the experiment from laboratory manual. Just cut
      and glue the laboratory manual to the pages of report book. If special apparatus have been used in an
      experiment (especially in organic practical), it is required to draw the apparatus used.
3.    In the section of RESULT, students should record accurately the observations made during an
      experiment. The result should be presented in a logical sequence. It is advisable to record results in
      table form.
4.    The section of TREATMENT OF RESULT is an important part of a report. This section may include
      calculations, drawing graphs and derivation of equations etc.
5.    DISUSSION AND CONCLUSION section is the „spirit‟ of a report. In this section students have to
      draw conclusions from observations made during an experiment, comment on unexpected results,
      suggest modification of experimental set-up and sequence to overcome difficulties and safety (hazard
      may arise or the use of harmful chemicals) aspects. It is also requires to discuss the accuracy of the
      results and the precautions, assumptions used and limitation.
6.    Reports should be written in „PASSIVE FORM‟. Avoid using the subject „I‟ and „We‟.
7.    Don‟t try to cheat the teacher by „making‟ results or use other students‟ results without the teacher‟s
      permission. If you fail to do an experiment, it is wise to discuss why you failed and how the
      difficulties may be overcome.
Experiment 1 --- An acid-base titration
 Objective: To determine the concentration of a solution of sodium hydroxide by titration against a
            standard solution of hydrochloric acid.
 Theory: Sodium hydroxide can be neutralized by hydrochloric acid according to the following
            equation:
                                    NaOH(aq) + HCl(aq)  NaCl(aq) + H2O(l)
            As the number of mole of sodium hydroxide reacted equals the number of hydrochloric
            acid, the amount of sodium hydroxide can be found by the known amount of hydrochloric
            acid used for complete neutralization.
                 To show you when the reaction is completed, i.e. equivalence point is reached, you
            can use an indicator called phenolphthalein, which is colourless in acid and pink in
            alkaline.
                 The volume of titrant added to cause the mixture changing colour is called the
             end-point. Finding the end-point by titrate standard hydrochloric acid with sodium
             hydroxide solution, the concentration of sodium can be found.
Apparatus: safety spectacles, filter funnel, 50 cm3 burette, stand and clamp,
           100 cm3 beakers x 2,
           250 cm3 conical flasks x 4, 25 cm3 pipette, pipette filler
           white tile, washing bottle of distilled water
 Reagents: sodium hydroxide solution (unknown concentration),
           standard (        M) hydrochloric acid,
           phenolphthalein indicator solution
Procedures:
 1. Using the funnel, rinse the burette with the sodium hydroxide solution and fill it with the same
    solution. Do not forget to rinse and fill the tip.

 2. With the help of a pipette filler, draw up some of the standard hydrochloric acid and rinse the
    pipette.
 3. Using pipette, transfer 25.00 cm3 of the standard hydrochloric acid solution to a clean 250 cm3
    conical flask.
 4. Add 2 to 3 drops of the phenolphthalein indicator solution to the solution in conical flask.
 5. Record the initial burette reading in the „trial‟ column of results table given.
    Run sodium hydroxide solution from the burette into the flask, with swirling, until the solution just
    turn pink. This is the trial run.
    Record the final burette reading.
 6. Refill the burette with the sodium hydroxide solution, and again record the initial burette reading to
    the nearest 0.02 cm3 (i.e. corrected to the one-fifth).
 7. Repeat the step from 3 to 5 at least three more times.
 8. Empty the burette and wash it carefully immediately after the titration, especially if it has a ground
    glass tap.
Results:
    Pipette solution:                              Burette solution:

     Indicator solution:

     Result table
                                       Trial           1               2            3             4
       Final burette reading/cm3

      Initial burette reading/cm3

       Volume used(titrant)/cm3


     Mean titre/cm3 =



Calculation:
    Calculate the concentration of the sodium hydroxide solution.




Questions:
1. What effect would each of the errors described below have on the calculated value of the
    concentration of sodium hydroxide?
    (a) The burette is not rinsed with the sodium hydroxide solution.
    (b) The pipette is not rinsed with the standard hydrochloric solution.
     (c) The tip of the burette is not filled before titration begins.
     (d) The conical flask contains some distilled water before the addition of
           standard hydrochloric acid solution.
2.   In using phenolphthalein as an indicator, we prefer to titrate from a colourless to pink solution rather
     than from pink to colourless. Suggest a reason for this.
3.   Why is it advisable to remove sodium hydroxide from the burette as soon as possible after the
     titration?
Suggested solutions for experiments            Experiment 1                 Date:       10/9/07
1.   An acid-base titration
2. Objectives:         Determine the concentration of sodium hydroxide solution
3. – 6.          Try to remember…in writing procedures and observation, you should
     Procedures:       The volume measured by pipette, burette should be have 2 decimal place.
                 Writing “ 25 cm3 “ instead of “ 25.00 cm3 “ will cause mark deduction.
     Observations: The colour change at the end-point,” from colourless to pink” should be stated clear.
7. Results: Try to remind yourself that there must have “2 d.p.” in recording the burette reading.
     Sample results:       Burette solution = NaOH (with unknown concentrations)
                           Conical flask solution = 25.00 cm3 of 0.1 M HCl
                                               Trial          1         2           3           4
                 Final burette reading         43.98       37.92     24.48      37.60       33.82
                 Initial burette reading       20.60       15.02     1.86       14.94       11.20
                 Volume used / cm3             23.38       22.90     22.62      22.66       22.62
     (     If data aren’t 2 decimal placed or 23.38 included in calculation will cause mark deduction.)
                 Mean titre = 22.70 cm3                (i.e. vol. of NaOH used = 22.70 cm3)
8. Treatment of data:
     (           Equation must be shown or presentation mark will be deducted.              )

                 NaOH + HCl  H2O + NaCl
Calculations: Vol. of HCl used = 25.00 cm3 ; [HCl] = 0.10 M ; Vol. of NaOH used = 22.70 cm3

                 no. of mol of HCl used  no. of mol of NaOH                 [NaOH]
                 0.025 x 0.1 = 0.0025                  0.0025 mol           0.0025/0.0227 = 0.11M
     Full steps:

         no. of mol of HCl present = 0.025 x 0.1 = 0.0025 mol

         no. of mol of NaOH used for titration = 0.0025 mol (HCl + NaOH  NaCl + H2O
         Concentration of NaOH used = 0.0025 / 0.0227 = 0.11 M
9.   Questions and Discussions
1.   What effect would each of the errors described below have on the calculated value of the
     concentration of sodium hydroxide.
a)   burette is not rinsed:    NaOH may be diluted and the calculated value of [NaOH] is smaller.
b)   pipette is not rinsed:    HCl may be diluted. The volume of NaOH used will be smaller. In our
                               calculation, vol. of NaOH used decreases will make calculated value of
     [NaOH] to                      be increased.
c)   the tip of burette is not filled before titration:
          NaOH used will be increased. In our calculation, vol. of NaOH used increased will make
     the calculated value of [NaOH] to be decreased.
d)   the conical flask contains some distilled water before the addition of standard hydrochloric acid
     solution:
          It will not affect the no. of mole of HCl inside the conical flask. In our calculation, it
     makes no effect on the calculated value of [NaOH], i.e. the value remains the same.
2.   Alkaline solution in the conical flask will absorb CO2 from the air. Therefore, conical flask
     solution should be acidic solution that will change colour from colourless to pink when using
     phenolphthalein.
3.   It is advisable to remove sodium hydroxide from the burette as soon as possible after the titration.
     It is because NaOH will attack glass and NaOH remained inside the burette will form Na2CO3
     with CO2 in the air. That crystal of Na2CO3 will block the tip of burette.

10. Conclusions: the concentration of NaOH was found to be 0.11          M
NOTES for dealing with “Bad” data

Example 1
There are 5 data: 25.19, 26.25, 26.01, 26.17 and 26.32
Can 25.19 be rejected?
    Step 1: Remaining data: 26.25, 26.01, 26.17, 26.32
    Step 2: Average: (26.25 + 26.01 + 26.17 + 26.32) / 4 = 26.19
    Step 3: Deviations: 0.06, 0.18, 0.02, 0.13
    Step 4: Average deviation: (0.06 + 0.18 + 0.02 + 0.13) / 4 = 0.098
    Step 5: Deviation of 25.19 from 26.19: 1.00
    Step 6: As 1.00 > 4 (0.098), 25.19 is rejected and the average value = 26.19

Example 2
There are 3 data: 25.79, 26.01 and 26.17
Can 25.79 be rejected?
    Step 1: Remaining data: 26.01, 26.17
    Step 2: Average: (26.01 + 26.17) / 2 = 26.09
    Step 3: Deviations: 0.08, 0.08
    Step 4: Average deviation: (0.08 + 0.08) / 2 = 0.08
    Step 5: Deviation of 25.79 from 26.09: 0.30
    Step 6: As 0.3 < 4 (0.08), 25.79 can‟t be rejected and the average value = 25.99
Experiment 2 --- Standardization of hydrochloric acid by sodium carbonate

Procedures in preparing a standard solution of sodium carbonate
1.   To weigh about 1.3 g of sodium carbonate.     Then pour the weighed sodium carbonate into a clean
     and dry 100 cm3 beaker.

2. Wash the weighing bottle with a small amount of distilled water. Pour the rinse into the beaker.
   Repeat this step twice.

3. Add about 40 cm3 of distilled water into the beaker. Stir the mixture with a
   clean glass rod until all heated sodium carbonate has dissolved.

4. Pour the sodium carbonate solution into a 250.0 cm3
     volumetric flask through a filter funnel.


     Rinse the beaker with a small amount of distilled
     water and transfer the rinse into the volumetric flask.
     Repeat the step of rinsing twice.



     Rinse the funnel carefully with distilled water.




5.   Add distilled water to the volumetric flask until the liquid level is about 1 cm below the graduated
     mark on the flask.


     Place the flask on bench and allow the solution to settle.
     Using a dropper, add water drop by drop to the flask until the
     meniscus just touches the graduated mark.


6.   Stopper the volumetric flask tightly. Turn it upside down and
     shake several times in order to mix the sodium carbonate with
     water thoroughly.
Procedures in titration(Hydrochloric acid is titrated with sodium carbonate)
1.   Construct the following titration set-up as shown in the diagram.




2.   Close the stopcock of the burette. Add about 10 cm3 of
     distilled water in the burette through the funnel.


     Remove the funnel. Rotate the burette horizontally to rinse
     its inner wall.



     Open the stopcock and let the water run out into the sink. Repeat this step once.

3.   Follow step 2 to rinse the burette with several cm3 of the given hydrochloric acid (but not with
     water). Repeat this step once.

4.   Close the stopcock of the burette. Pour the given hydrochloric acid through the funnel into the
     burette until the liquid level is near zero. Open the stopcock of burette to allow the titrant to
     fill up the tip and then adjust the liquid level near zero (but not above zero).

5.   Rinse the conical flasks provided with distilled water.


6.   Take a pipette filler and a 25.00 cm3 pipette. Press the
     pipette filler and then put its tip into the top end of the
     pipette tightly. Keep the pipette and the pipette filler
     connected together. Insert the tip of the pipette into a
     beaker of distilled water. Release the air pressure in the
     pipette filler to suck several cm3 of water.


     Remove the pipette filler and close the top end of the
     pipette with a forefinger immediately.
     Remove the tip of the pipette from the water. Rotate the pipette




     horizontally to rinse its inner side. Then, let the water run out to the sink.


     Repeat this step once. Finally wipe the outside wall of the pipette with tissue. (Caution:
     Clamp the pipette or place it safely on the bench; otherwise, it may roll down the bench.)

7.   Pour about 40 cm3 of the standard sodium carbonate solution into another clean 100 cm3 beaker.


8. Follow step 6, rinse the pipette with the standard sodium carbonate solution in the 100 cm3 beaker
     (but not with water). Repeat this step once.


9.   Use the pipette filler to fill the pipette with the standard sodium carbonate solution until the solution
     is several cm above the graduated mark,


10. Remove the pipette filler and close the top end of the pipette with a
    forefinger immediately. Remove the tip of pipette from the standard
    sodium carbonate solution (better to wipe the tip with tissue quickly).
    Adjust the liquid level with the forefinger carefully until it just touches
    the graduated mark.



11. Insert the pipette into a conical flask. Remove the forefinger to allow the
    standard sodium carbonate to drain into the flask. When the drainage stops,
    touch the tip of pipette onto the inner wall of conical flask to allow the
    remaining droplet to flow out.


12. Add 3 drops of methyl orange indicator to the conical flask.


13. Record the initial burette reading in the table provided. The reading should be
    accurate to 2 decimal places.


14. Insert the tip of burette into the mouth of conical flask containing the
    standard sodium carbonate. Start titration by opening the stopcock of the
    burette to allow hydrochloric acid to drain into the flask. During the
    process of titration, swirl the conical flask continuously to mix the
    solutions. When the solution in the conical flask just changes from yellow
    to orange, close the stopcock of the burette immediately.
15. Record the final burette reading in the table provided. This is a trial titration to estimate the
    volume of hydrochloric acid required. Calculate the volume of hydrochloric acid added in
    titration.


16. Add the given hydrochloric acid to the burette through a filter funnel if the volume remained
    is not enough to carry out another titration.


17. Repeat steps 9-16 to obtain 2-3 sets of consistent results (difference of results within 0.10 cm3).
    However, stop draining hydrochloric acid at about 3 cm3 less than the estimated value. Then add
    hydrochloric acid drop by drop until the reaction mixture in conical flask just changes from yellow
    to orange.


Results:
    Pipette solution:                               Burette solution:

      Indicator solution:

      Result table
                                        Trial           1               2             3             4
        Final burette reading/cm3

       Initial burette reading/cm3

        Volume used(titrant)/cm3


      Mean titre/cm3 =


Treatment of the data
1. Calculate the concentration of sodium carbonate solution.
2. Calculate the concentration of hydrochloric acid solution.

Write a full report with the following items:
1.         Title of the experiment
2.         Objective(s) : clear statement(s) of the purpose(s) of the experiment
3.         Theory / Introduction / Background
4.         Chemicals and apparatus used
5.         Diagram for the set-up
6.         Procedures and Observations
7.         Results
8.         Treatment of data
9.         Discussions
10.        Conclusions
Experiment 2          Standardization of hydrochloric acid by sodium carbonate              Report sheet
S.6 Chemistry                   Name:________________________(              ) Locker no.(            )
1.   Preparation of standard sodium carbonate solution

       Mass of sodium carbonate used =

       Molar mass sodium carbonate =

       Volume of solution prepared =

       Concentration of sodium carbonate solution =

2.   Standardization of hydrochloric acid by sodium carbonate

     Conical flask solution =           Concentration =

       Volume used =

       Indicator used =

     Burette solution =

     Titration results:


                                Trial        1st   titration    2nd titration      3rd titration
        Final burette
        reading /cm3
        Initial
        burette
        reading /cm3
        Titre / cm3

     Mean titre =

Determine the concentration of hydrochloric acid:
Experiment 3 --- Calcium Carbonate in Eggshell

Introduction

In the past, pesticides such as dichlorophenyltrichloroethane (DDT) have been used extensively. Their

harmful effects on biological systems are gradually revealed. One effect on wild bird life is the weakening

or thinning of the eggshell. The consequence is the breaking of the eggshell before hatching. Eggshell

mainly consists of calcium carbonate. In the present investigation, we will determine the percentage

of calcium carbonate in eggshell         by acid-base titration.


Chemicals
0.10 M NaOH solution, 0.20 M HCl solution, ethanol, phenolphthalein


Procedures

1. Boil an egg and allow it to cool down.

2. Take off the eggshell and remove the attached membrane.

3. Dry the eggshell in an oven at 110℃ for overnight.

4. Remove the eggshell from the oven and allow to cool down.

5. Ground the eggshell into a fine powder with mortar and pestle.

Suggestion for the following steps :

6. Weigh 0.2 g of powder accurately and put into a conical flask.

7. Pipette 25.00 cm3 of 0.20 M HCl solution into the flask and add 5 cm3 of ethanol.

8. Boil the solution for 5-10 minutes and allow it to cool down. Add a few drops of indicator solution

   and titrate the solution with NaOH.

9. Repeat the titration twice. Record your results.




References

1. http://chem..lapeer.org/Chem1Docs/EggshellTitration.html

2. http://www.accessexcellence.org/AE/AEC/AEF/1996/tucker_eggshell.html
Experiment 4 --- Analysis of Two commercial Brands of Bleaching Solution
Introduction
      Sodium chlorate(I) (sodium hypochlorite) forms the basis of most commercial bleaches. The amount
of this active ingredient available can be estimated by the following method.
      In this analysis, the sodium chlorate(I) is allowed to react with an excess of potassium iodide
solution in the presence of acid, liberating iodine, which is then titrated against standard sodium
thiosulphate solution.
      The reaction involved are :

                    ClO-(aq)      + 2I-(aq) +   2H+(aq)                     I2(aq) + H2O(l) + Cl-(aq)

                         I2(aq)    + 2S2O32-(aq)                  2I-(aq)      + S4O62-(aq)
Chemicals
Commercial bleach(2 brands), 1M KI, dilute H2SO4, 0.050M Na2S2O3, starch indicator (freshly prepared)

Apparatus
Titration apparatus, measuring cylinder


Procedure
1. Pipette 25.0 cm3 of the bleach into a clean 250 cm3volumetric flask. Mark up to the mark using
    deionized water.
2. Pipette 25.0 cm3 of this solution into a conical flask, add to it 10 cm3 of 1M potassium iodide
    solution and 10 cm3 of dilute sulphuric (VI) acid.
3. Titrate this against the 0.050M sodium thiosulphate solution. Add a few drops of freshly prepared
     starch indicator when the reaction mixture turns pale yellow and continue to titrate to the end-point.
4.   Repeat steps (1) to (3) with another brand of bleach.


Discussion
1. For each brand, work out
    (a) the amount of the active ingredient available in g dm-3.
    (b) the cost per gram of this compound.
2. According to your results, which of the two brands of bleach is better buy ?
3. Why should potassium iodide be present in excess ?
4. What is the function of the dilute sulphuric(VI) acid ?
5.   Bleaching solution may deteriorate for two main reasons. One is the attack by carbon dioxide in air
     according to the equation :

            2ClO-(aq) + 2CO2(aq)                          CO32-(aq)     +      Cl2(aq)

     What is the other possible reason?
6.   The starch indicator should not be added too early. Why?
                                                    Results
Relevant equation(s):




Titration Results:                                            Set-up of the titration
Trial Titration

                        Final Burette Reading / cm3


                        Initial Burette Reading /cm3


                          Volume of titrant / cm3


                                         Mean titre =                   cm3
Calculations:
Experiment 5 --- An analysis of aspirin tablet

  Objective: To determine the percentage of 2-ethanolyydroxybenzoic
             acid (acetyl-salicylic acid) in aspirin tablets.




            CH3COOC6H4COOH + NaOH  CH3COONa + HOC6H4COONa + H2O
Equation(s):
                                   NaOH + HCl  NaCl + H2O
    Theory:     A known amount of standard sodium hydroxide solution is used in excess to
            hydrolyse a known mass of aspirin tablets:
                CH3COOC6H4COOH + NaOH  CH3COONa + HOC6H4COONa + H2O


                  The unused sodium hydroxide which remains is then titrated with standard acid.
                                    NaOH + HCl  NaCl + H2O


                 The amount of alkali required for the hydrolysis can now be calculated.
                 Next, from the equation for the hydrolysis of acetyl-salicylic acid, the number of
            moles of acetylsalicylic acid which have been hydrolysed can be found.
 Apparatus: safety spectacles, filter funnel, 50 cm3 burette, stand and clamp,
            100 cm3 beakers x 2, 250.00 cm3 volumetric flask,
            250 cm3 conical flasks x 4, 25 cm3 pipette, pipette filler
            white tile, washing bottle of distilled water, Bunser burner, tripod and gauze.
  Reagents: sodium hydroxide solution (approximately 1.0 M),
            standard hydrochloric acid (       M),
            phenolphthalein indicator solution
            aspirin tablets (about 5)


Part I:   Standardization of sodium hydroxide used in hydrolysis
 Procedure:
    Partner No. 1 should standardize the approximately 1.0 M NaOH used for the hydrolysis as
    follow:
 1. Using a safety filler, pipette exactly 25.00 cm3 of the approximately 1.0 M NaOH solution into a
    250.00 cm3 volumetric flask and make up to the mark.
 2. Pipette 25.00 cm3 of the diluted NaOH solution to a conical flask.
 3. Titrate the diluted NaOH solution against 0.10 M hydrochloric acid using phenolphthalein
    indicator.
 4. Carry out one rough and two accurate titrations.
 5. Record your results in a table given on the “data sheet”.
 6. Calculate the accurate molarity of the approximately 1.0 M NaOH.
Part II: Determination of acetyl-salicylic acid in aspirin tablet
Procedures:
   Partner No. 2 should hydrolyse the aspirin as follow:
 1. Weigh accurately between 1.3 g and 1.7 g of the aspirin tablets into a clean conical flask. (This
    will be about 5 tablets.)
 2. Using a pipette filler, pipette 25.00 cm3 of the approximately 1.0 M NaOH on to the tablets,
    followed by about the same volume of distilled water.
 3. Simmer the mixture gently for ten minutes to hydrolyse the acetyl-salicylic acid.
    CARE Eye protection must be worn.
 4. Now, cool the mixture and transfer with washings to 250 cm3 volumetric flask and make up to
    the mark with distilled water.
 5. Record the weight of aspirin taken.


    Both partners should now estimate the quantity of unused NaOH after the hydrolysis as follows:
 6. Pipette 25.00 cm3 of the diluted and hydrolysed solution into conical flask.
 7. Titrate this against 0.10 M hydrochloric acid using phenolphthalein indicator.
 8. Record your titration results in a table on the “data sheet”.


Questions:
1. Why is the mixture simmered gently and carefully during hydrolysis?         Why is it unwise to boil it
    vigorously?


2.    How the result be affected if the washings be transferred NOT carefully to the 250 cm3 volumetric
      flask?


3.    What might be the remainder of the tablets be made of?

4.    What are the uses of aspirin and what are their side effects?


Write a full report with the following items:
1.         Title of the experiment
2.         Objective(s) : clear statement(s) of the purpose(s) of the experiment
3.         Theory / Introduction / Background
4.         Chemicals and apparatus used
5.         Diagram for the set-up
6.         Procedures and Observations
7.         Results
8.         Treatment of data
9.         Questions / Discussions
10.        Conclusions
                                              Data Sheet
                    Part I: Standardization of sodium hydroxide used in hydrolysis
Results:
    Pipette solution:                            Burette solution:

    Indicator solution:

    Result table
                                    Trial       1          2         3          4
                                3
     Final burette reading/cm
     Initial burette reading/cm3
     Volume used (titrant)/cm3


    Mean titrant/cm3 =


Calculation:
    Calculate the concentration of the sodium hydroxide solution.


                    Part II: Determination of acetyl-salicylic acid in aspirin tablet

Results: Mass of aspirin used
                    Weight of paper         =                   g

         Weight of paper + weight of aspirins    =                   g

                             Weight of aspirins =                    g


Results: The amount of NaOH left after aspirin is reacted with 25cm3 of standard NaOH
         solution
    Pipette solution:                   Burette solution:
    Indicator solution:
                                    Trial       1          2         3          4
                                3
     Final burette reading/cm
     Initial burette reading/cm3
     Volume used(titrant)/cm3

    Mean titrant/cm3 =

Calculation:
1. How many moles of NaOH are added to the flask before hydrolysis of the aspirin?
2. How many moles of NaOH are remain after hydrolysis of the aspirin?
3. How many moles of NaOH are used in the hydrolysis of the aspirin?
4. How many moles of acetyl-salicylic acid have been hydrolysed?
5. What percentage of the aspirin tablets is acetyl-salicylic acid?
Experiment 6 --- An analysis of commercial vitamin C tablets
Introduction
1.   In this experiment, the content of vitamin C tablets are investigated and compared with
     manufacturer‟s specifications.

2.   Vitamin C(C6H8O6) is ascorbic acid. This is rapidly and quantitatively oxidized by iodine in an acid
     medium according to the following equation :




3.   The standard method for determining ascorbic acid present in a sample is to titrate against a standard
     iodine solution. But the low solubility of iodine makes this procedure less than ideal.
     In this experiment, a known excess of iodine is created in situ by the reaction with the ascorbic acid is
     titrated against standard sodium thiosulphate solution.

Chemicals
Potassium iodate, 1M potassium iodide, 0.5M sulphuric acid, sodium thiosulphate solution, freshly
prepared starch solution, vitamin C tablet

Procedures
1.   Accurately weigh out 0.6 – 0.7 g of potassium iodate. Dissolve in distilled water and make up to the
     mark in a 250 cm3 volumetric flask.
2.   Use this iodate solution to standardise the sodium thiosulphate solution. Put 25.0 cm3 portions of the
     iodate solution, 5 cm3 1M KI and 5 – 10 cm3 0.5 M sulphuric acid and titrating against the sodium
     thiosulphate solution. Add starch solution when the solution becomes pale yellow.
3.   Dissolve a weighed vitamin C tablet in about 150 cm3 0.5 M sulphuric acid in a beaker.
4.   Transfer this solution to a 250 cm3 volumetric flask and make up to the mark using distilled water.
5.   Pipette 25.0 cm3 of the vitamin C solution into a clean conical flask and add 5 cm3 1M KI solution.
6.   Pipette 25.0 cm3 of the standard potassium iodate solution into the flask containing the vitamin C
     and KI solution.
7.   Back titrate the excess iodine against the stardardised sodium thiosulphate solution.
8.   Calculate the mass of ascorbic acid in the tablet.

Task:     Write a short report with the items: Title, Objectives, Results, Treatment of data, Discussions
     and Conclusion.
Experiment 8 --- Application of Hess’s Law to determine the Enthalpy Change
                  of Hydration of Magnesium sulphate(VI)
Chemicals
Anhydrous MgSO4 , MgSO4‧7H20

Apparatus
Balance, 100 cm3 measuring cylinder, polystyrene foam cup, -10-110℃ thermometer, watch glass

Procedure
A. To determine the enthalpy change of solution of anhydrous magnesium sulphate(VI)
    1. Weigh an empty polystyrene foam cup.
    2. Pour 50 cm3 of distilled water from a measuring cylinder into the polystyrene foam cup and record
       the temperature.
     3. Weigh accurately 0.025 mole of anhydrous magnesium sulphate(VI) and add it to the water in the
        foam cup. Stir to dissolve it as quickly as possible. Record the highest temperature of the solution.
     4. Calculate the molar enthalpy change of solution of anhydrous magnesium sulphate(VI).
       (Assume that the specific heat capacity of the solution in the foam cup is 4.2 kJ kg-1 K-1, and that
       of the foam cup is 1.3 kJ kg-1 K-1.)

B.   To determine the enthalpy change of solution of anhydrous magnesium sulphate(VI)-7-water
     1. Repeat the experiment as described in part A above, but use 0.025 mole of magnesium
        sulphate(VI)-7-water instead of the anhydrous salt.
     2. Calculate the molar enthalpy change of solution of magnesium sulphate(VI)-7-water.
        (Assume that the specific heat capacity of the solution in the foam cup is 4.2 kJ kg-1 K-1, and that
       of the foam cup is 1.3 kJ kg-1 K-1.)

Discussion
1. Write equations for the following processes:
     (a) Dissolution of anhydrous magnesium sulphate(VI) into water.
     (b) Dissolution of magnesium sulphate(VI)-7-water into water.
     (c) Hydration of anhydrous magnesium sulphate(VI) to form magnesium sulphate(VI)-7-water.

2.   Draw an energy cycle to link the corresponding enthalpy changes of the reactions together.

3.   Using Hess‟s law, calculate the molar enthalpy change of hydration of anhydrous magnesium
     sulphate(VI) to form magnesium sulphate(VI)-7-water.

4.   What assumptions have you made in your calculations?

5.   What are the sources of error in this experiment?

6.   Why cannot the molar enthalpy change of hydration of magnesium sulphate(VI) be measured
     directly in the laboratory ?
Experiment 9 – Determination of Dissolved Oxygen in a Water Sample
Introduction
In an alkaline solution, dissolved oxygen will oxidize manganese(II) to the trivalent state.

     8OH-(aq) + 4Mn2+(aq)       + O2(aq) + 2H2O(l)                4Mn(OH)3(s)

The analysis is completed by titrating the iodine produced from potassium iodide by manganese(III)
hydroxide.

     2Mn(OH)3(s)      + 2I-(aq) + 6H+(aq)                  2Mn2+(aq) +     I2(aq)   + 6H2O(l)

Sodium thiosulphate is used as the titrant.

Success of the method is critically dependent upon the manner in which the sample is manipulated. At all
stages, every method must be made to assure that oxygen is neither introduced to nor lost from the sample.
Furthermore, the sample must be free of any solutes that will oxidize iodide or reduce iodine.

Chemicals
Manganese(II) sulphate solution, alkaline potassium iodide solution, concentrated sulphuric acid,
0.0125M sodium thiosulphate solution, starch solution

Apparatus / Materials
250 cm3 volumetric flask, 250 cm3 conical flask, measuring cylinders, titration apparatus

Procedures
1.   Use a 250 cm3 volumetric flask to collect a water sample. Fill the flask completely with water
     without trapping any air bubbles.

2.   Add 1 cm3 of manganese(II) sulphate solution to the sample using a pipette. Discharge the solution
     well below the surface (some overflow will occur).

3.   Similarly introduce 1 cm3 of alkaline potassium iodide solution. Be sure that no air becomes
     entrapped. Invert the bottle to distribute the precipitate uniformly.
     [Hazard Warming: Care should be taken to avoid exposure to any overflow, as the solution is quite
     alkaline.]

4.   When the precipitate has settled at least 3 cm below the stopper, introduce 1 cm3 of concentrated
     sulphuric(VI) acid well below the surface. Replace the stopper and carefully mix until the precipitate
     disappears.

5.   Allow the mixture to stand for 5 minutes and then withdraw 100 cm3 of the acidified sample into a
     250 cm3 conical flask.

6.   Titrate with 0.0125M sodium thiosulphate solution until the iodine colour becomes faint. Then add 1
     cm3of starch solution, and continue adding the thiosulphate solution until the blue colour disappears.

7.   Record the volume of thiosulphate solution used and calculate the dissolved oxygen content in the
     sample in mg dm-3.

Remark
If the water sample has a low DO value. It is recommended to withdraw 200 cm3 of the acidified sample
into a 500 cm3 flask for the titration described in step(5).
Experiment 10 Order of Reaction between H2O2 and KI(aq) in Acidic Medium
Introduction
     In acidic medium, hydrogen peroxide reacts with iodide ions as shown in the following chemical
equation:       H2O2(aq) + 2H+(aq) + 2I-(aq)               I2(aq) + 2H2O(l)
     In this experiment, the orders of reaction with respect to hydrogen peroxide, hydrogen ions and
iodide ions are to be determined.
     The initial concentration of one of the reactants is varied to carry out several sets of experiments
(while keeping the concentration of the two other reactants constant). The iodine produced will first react
with a small amount of sodium thiosulphate as shown in the equation below:
                I2(aq) + 2S2O32-(aq)                       2I-(aq) + S4O62-(aq)
     When all sodium thiosulphate is reacted, any further iodine produced will form a dark blue complex
with starch indicator.
     By graphical method, the order of reaction with respect to the reactant of varying concentrations can
then be determined. The oorders of reaction of the two other reactants can also be determined similary by
carrying out two sets of experiments.

Procedures
A.   Determination of the order of reaction with respect to hydrogen peroxide
1.   According to the following table, measured 0.10M hydrogen peroxide and distilled water using clean
     measuring cylinders. Pour each combination of hydrogen peroxide and water into a clean and dry
     100 cm3 beaker.
     Volume of 0.10M hydrogen peroxide (cm3)           2.0        4.0       8.0     12.0        16.0
     Volume of water (cm3)                            23.0       21.0      17.0     13.0         9.0
                      3                                    3
     Start with 2.0 cm of hydrogen peroxide and 23.0 cm of water in the clean and dry beaker.

2.   Measure 1.0 cm3 of 0.010M sodium thiosulphate solution using a measuring cylinder. Pour it into the
     beaker containing 2.0 cm3 of hydrogen peroxide and 23.0 cm3 of water.

3.   Add 3 drops of freshly prepared starch solution to the beaker.

4.   Measure 10.0 cm3 of 0.50 M sulphuric acid and 10.0 cm3 of 0.10M potassium iodide solution using
     clean and dry 10.0 cm3 measuring cylinders. Pour them into the beaker.

5.   Start the stop watch immediately.

6.   Stir the reaction mixture quickly with a clean and dry glass rod.

7.   Observe any colour change against a white tile.    Try to record the time for dark blue colour appears.

8.   Wash the beaker. Repeat steps 1-7 for the other combinations mentioned in step 1.

B.   Determination of the order of reaction with respect to potassium iodide
1.   According to the following table, measured 0.10M potassium iodide and distilled water using clean
     measuring cylinders. Pour each combination of potassium iodide and water into a clean and dry
     100 cm3 beaker.
     Volume of 0.10M potassium iodide (cm3)            2.0         4.0       8.0     12.0       16.0
                           3
     Volume of water (cm )                            23.0        21.0      17.0     13.0        9.0

     Start with 2.0 cm3 of potassium iodide and 23.0 cm3 of water in the clean and dry beaker.
2.   Measure 1.0 cm3 of 0.010M sodium thiosulphate solution using a measuring cylinder. Pour it into the
     beaker containing 2.0 cm3 of hydrogen peroxide and 23.0 cm3 of water.

3.   Add 3 drops of freshly prepared starch solution to the beaker.

4.   Measure 10.0 cm3 of 0.50 M sulphuric acid and 10.0 cm3 of 0.10M hydrogen peroxide using clean
     and dry 10.0 cm3 measuring cylinders. Pour them into the beaker.

5.   Start the stop watch immediately.

6.   Stir the reaction mixture quickly with a clean and dry glass rod.

7.   Observe any colour change against a white tile. Record the time for the dark blue colour to appear.

8.   Wash the beaker. Repeat steps 1-7 for the other combinations mentioned in step 1.

C.   Determination of the order of reaction with respect to sulphuric acid

1.   According to the following table, measured 0.50M sulphuric acid and distilled water using clean
     measuring cylinders. Pour each combination of hydrogen peroxide and water into a clean and dry
     100 cm3 beaker.
     Volume of 0.50M sulphuric acid (cm3)             2.0       4.0         8.0       12.0        16.0
     Volume of water (cm3)                          23.0       21.0        17.0       13.0         9.0
                      3                             3
     Start with 2.0 cm of sulphuric acid and 23.0 cm of water in the clean and dry beaker.

2.   Measure 1.0 cm3 of 0.010M sodium thiosulphate solution using a measuring cylinder. Pour it into the
     beaker containing 2.0 cm3 of hydrogen peroxide and 23.0 cm3 of water.

3.   Add 3 drops of freshly prepared starch solution to the beaker.

4.   Measure 10.0 cm3 of 0.10 M hydrogen peroxide and 10.0 cm3 of 0.10M potassium iodide solution
     using clean and dry 10.0 cm3 measuring cylinders. Pour them into the beaker.

5.   Start the stop watch immediately.

6.   Stir the reaction mixture quickly with a clean and dry glass rod.

7.   Observe any colour change against a white tile. Record the time for the dark blue colour to appear.

8.   Wash the beaker. Repeat steps 1-7 for the other combinations mentioned in step 1.

Discussions

1.   Explain whether dilute hydrochloric acid can be replaced sulphuric acid in this experiment.

2.   Explain whether sodium thiosulphate can be omitted in this experiment.

3.   Explain why the reciprocal of time, 1/t, can be used to express the initial rate.

4.   Derive equations relating log(1/t) to log [H2O2], log(1/t) to log [I-]   and    log(1/t) to log [H+].

5.   Plot suitable graphs to determine the order of reactions with respect to hydrogen peroxide, potassium

     iodide and hydrogen ions.
Experiment 11 – Investigating the order of Reaction between Iodine and
     Propanone
Introduction

      Iodine reacts with propanone in acidic medium. Its rate of reaction is slow enough to be measured

with simple apparatus in the laboratory.

      In this experiment, iodine is allowed to react with excess propanone. At constant time intervals, a

fixed volume of the reaction mixture is withdrawn. The amount of remaining iodine is then titrated by

sodium thiosulphate as shown in the following chemical equation:

           I2(aq) + 2Na2S2O3(aq)                           2NaI(aq) + Na2S4O6(aq)

      By graphical method, the orders of reaction with respect to iodine and propanone can be determined.

Procedures
A.    Preparation of sodium thiosulphate solution for titration
1.    Rinse a 25.0 cm3 pipette with distilled water and then with 0.1M sodium thiosulphate solution.

2.    Pipette 25.0 cm3 of the sodium thiosulphate solution into a clean 250.0 cm3 volumetric flask.

3.    Make it up to 250.0 cm3 using distilled water and shake it well.

4.   Rinse a burette with distilled water and then with the diluted sodium thiosulphate solution.

5.   Pour the diluted sodium thiosulphate solution into the burette through a filter funnel and make sure
     that tip of the burette is filled with solution.

6.   Adjust the liquid level near the zero mark of the burette.

B.    Preparation of reaction mixture

1.    Rinse a 50.0 cm3 pipette with distilled water and then with 0.02M iodine solution.

2.    Pipette 50.0 cm3 of the iodine solution into a clean conical flask.

3.    After consulting your teacher, choose one of the combinations of solutions, A to E, indicated in the
      following table.
       Group                                     A           B           C            D            E
                                 3
       Volume of propanone (cm )               25.0        20.0        15.0         10.0          5.0
                                    3
       Volume of distilled waer (cm )           0.0         5.0        10.0         15.0         20.0

4.    According to the combination , measure the specified volumes of propanone and distilled water
      using measuring cylinders. Pour them into the conical flask containing iodine solution.

5.    Measure 25.0 cm3 of 1.0 M sulphuric acid using a measuring cylinder. Pour it into the clean conical
      flask quickly.

6.    Start the stop water when half of the liquid in the pipette has run into the flask.

7.    Swirl the conical flask to mix the contents thoroughly.
C.   Titration of reaction mixture
1.   Measure 10.0 cm3 of 0.5M sodium hydrogencarbonate solution using a measuring cylinder.

2.   When the reading of the stop watch is approaching 5 minutes, deliver 10.0 cm3 of the reaction
     mixture into the conical flask using a clean 10.0 cm3 pipette.

3.   Record the exact time at which all reaction mixture in pipette runs into the flask.

4.   Titrate the reaction mixture in the conical flask with diluted sodium thiosulphate solution using
     starch as indicator.

5.   When the stop watch reading is approaching 10, 15, 20, 25 and 30 minutes, 10.0 cm3 of the reaction
     mixture is withdrawn from the flask. Repeat steps 2-7 for each portion.

6.   Obtain and record the results of the other groups.

Discussion
1.   Write a balanced chemical equation to represent the reaction between iodine and propanone in acidic
     medium.

2.   Explain why it is necessary to add the 10.0 cm3 sample of the reaction mixture into sodium
     hydrogencarbonate before titration.

3.   Plot a graph with the volume of sodium thiosulphate added against the time at which the 10.0 cm3
     sample of the reaction mixture was added to the sodium hydrogencarbonate solution.

4.   Determine the concentration of sodium thiosulphate from the graph you plotted.

5.   Based on the results obtained, determine the order of reaction with respect to iodine for its reaction
     with propanone in acidic medium.

6.   Does iodine take part in the rate determining step of the reaction between iodine and propanone ?

7.   Collect the slopes of the graphs for the other combinations from A to E. Plot a suitable graph to
     determine the order of reaction with respect to propanone.
Experiment 12 – To determine the Activation Energy of the Reaction between Bromide
ion and Bromate(V) ion in Acid solution

Introduction
The reaction can be represented by

                         5Br (aq) + BrO3(aq) + 6H+(aq)  3Br2(aq) + 3H2O(l)
The progress of the reaction may be followed by adding a fixed amount of phenol together with some
methyl red indicator. The bromine produced during the reaction reacts very rapidly with phenol. Once all
the phenol is consumed, any further bromine bleaches the indicator immediately. So, the time for the
reaction to proceed to a given point may be determined.

Chemicals
Solution A (0.083 M with respect to KBr and 0.017 M with respect to KBrO3), 0.5M H2SO4, 0.01M
phenol solution, methyl red indicator

Apparatus

Beaker, boiling tube, stop watch, 10110C thermometer, burette

Procedure
[Hazard Warning: 0.5M sulphuric(VI) acid is irritant, and phenol is toxic.]
                     3                              3
1.                 of 0.01M                             solution A and 10 drops of methyl red indicator into
     a boiling tube.
                 3
2.                of 0.5 M sulphuric(VI) acid into a second boiling tube.
3.   Place both boiling tubes into a beaker of water which is maintained at about 30C. Allow the contents
     of the tubes to reach the temperature of the water bath.
4.   Pour the sulphuric(VI) acid into the first boiling tube, and at the same time start the stop watch.
     Swirl gently.
5.   Keep the first boiling tube in the water bath throughout the experiment. Record the time (t) taken, to
     the nearest second, for the complete disappearance of the red colour.
6.   Record also the temperature (T ), to the nearest degree, of the reaction mixture at the end of the
     experiment.
7.   Repeat steps (1) to (6), maintaining the reaction temperature at about 35C, 40C, 45C and 50C.
8.   Record your result in the table below.

           Experiment       Time                    1       Temperature              1
                                               ln                                      / K 1
                             t/s                    t                                T

                                                         T / C         T/K

               1

               2

               3

               4

               5


Discussion
1.   Give an equation for the reaction between phenol and bromine.

2.   What is the use of methyl red in this experiment?

3.   Based on your results, is it advisable to perform the experiment at high temperatures such as 80C?

4.   Why is it not necessary to know how far the reaction has proceeded at the point where the methyl
     red is decolourized?

5.   The Arrhenius equation can be represented as:
                                   E   / RT
                         k = Ae a
     (a)   Can 1/t substitute k in this equation? Why?
     (b)   Derive an equation relating ln k and 1/T.
     (c)   Determine Ea by plotting a suitable graph.
           (Given: R = 8.314J mol1K1)

6.   Explain why the reaction rate can be affected by temperature.
Experiment 13 – Determination of the equilibrium constant of esterification
Introduction
In the presence of concentrated sulphuric(VI) acid, ethanoic acid reacts reversibly with propan-1-ol when
heated, forming propyl ethanoate and water.

Chemicals
Glacial ethanoic acid, propan-1-ol, 0.05M NaOH, concentrated H2SO4, phenolphthalein indicator

Apparatus / Materials
Quickfit set, balance, titration apparatus, 1.0 cm3 graduated pipette, measuring cylinder, ice bath,
anti-bumping granules

Procedures
[Hazard Warning : Glacial ethanoic acid and concentrated sulphur(VI) acid are corrosive, propan-1-ol is
                   flammable, and 0.5M sodium hydroxide is irritant.]

1.   Put 0.25 mole of glacial ethanoic acid (density = 1.05 g cm-3) and 0.25 mole of propan-1-ol (density
     = 0.8 g cm-3) into a clean, dry pear-shaped flask. Mix thoroughly.

2.   Transfer 1.0 cm3 of the mixture by pipette to a 250 cm3 conical flask containing about 25 cm3
     deionized water and 2 drops of phenolphthalein indicator. Titrate to end point with 0.50M sodium
     hydroxide solution. Record the titre (V1 cm3).

3.   Add 8 drops of concentrated sulphuric(VI) acid to the remainder of the acid-alcohol solutions while
     continuously swirling the flask. Another 1.0 cm3 sample is titrated immediately.

     Record the titre (V2 cm3).

     The difference between V1 and V2 represents the volume to be subtracted from subsequent titrations
     to correct for the amount of sulphuric(VI) acid present.

4.   Add a few anti-bumping granules to the flask, and attach it to a water-cooled reflux condenser.
     Reflux for 50 minutes. Cool the flask and its contents in an ice bath. Remove 1.0 cm3 sample from
     the flask for titration with the 0.050M sodium hydroxide as before. Record the titre needed and
     correct it for the sulphuric(VI) acid.

5.   Continue refluxing for an additional 20 minutes, cool, and titrate another 1.0 cm3 sample. The two
     titres should agree to within 0.2 cm3. Otherwise, repeat this step.

Discussion
1. What is the purpose of adding a small amount of concentrated sulphuric(VI) acid to the reaction
    mixture at the beginning of the experiment ?

2.   Why should anti-bumping granules be added to the reaction mixture before refluxing?

3.   Why should the refluxing be continued in step (5) until the titre of sodium hydroxide used
     approaching constant?

4.   Write the equation for the esterification reaction between ethanoic acid and propan-1-ol.

5.   Calculate the concentration of ethanoic acid remaining at the end of the reflux.

6.   Calculate the concentrations of the other species present at equilibrium.
7.   Write an equilibrium expression for the esterification reaction.

8.   Calculate the equilibrium constant for the esterification reaction.

9.   If the concentration of the sodium hydroxide solution is not known exactly, explain what effect, if
     any, it would have on the determination of the equilibrium constant for the esterification reaction.
Experiment 14 –           Determination of the Partition Coefficient of ethanoic acid
                          between water and 2-methylpropan-1-ol
Chemicals

2-methylpropan-1-ol (density = 0.805 g dm-3), 0.2 M ethanoic acid, 0.1 M NaOH, phenolphthalein

Apparatus / Materials

100 cm3 separating funnel, titration apparatus, 10.0 cm3 pipette, 50 cm3 measuring cylinder, thermometer

Procedures
1.   Record the room temperature.

2.   Using suitable apparatus, pour 25 cm3 of the given aqueous ethanoic acid and 25 cm3 of
     2-methylpropan-1-ol into a 100 cm3 separating funnel. Stopper the funnel and shake vigorously for 1
     to 2 minutes. (Release pressure in the funnel by occasionally opening the tap.)

3.   Separate approximately 20 cm3 of each layer, discard the fraction near the junction of the two layers.

4.   Pipette 10.0 cm3 of the aqueous layer into a conical flask and titrate it with 0.1M sodium hydroxide
     solution using phenolphthalein.

5.   Using another pipette, deliver 10.0 cm3 of alcohol layer into a conical flask and titrate it with 0.1M
     sodium hydroxide solution using phenolphthalein.

6.   Repeat procedures (2) to (5) with another separating funnel using either one of the following
     volumes:
     (a) 35 cm3 of aqueous ethanoic acid and 25 cm3 of 2-methylpropan-1-ol,
     (b) 45 cm3 of aqueous ethanoic acid and 25 cm3 of 2-methylpropan-1-ol.

7.   For each experiment, calculate the ratio of the concentration of ethanoic acid in the aqueous layer to
     that in the 2-methylpropan-1-ol layer.

Discussion

1.   Why is shaking necessary in procedure (2) ?

2.   Would you expect the partition coefficient to vary with temperature ? Explain briefly.

3.   Explain why the amounts of aqueous ethanoic acid and 2-methylpropan-1-ol placed in the funnel
     need not be measured out accurately, whereas the volumes if the aqueous and alcohol solution used in
     the titration must be known as accurately as possible.

4.   What assumptions are made in the above experiment ? Based on experimental evidence, are these
     assumptions valid ? Explain your answer.

5.   Why is it more efficient to extract a solute with two 25 cm3 potions of solvent rather than with a
     single 50 cm3 extraction ?

6.   Give two applications of the partition law.
Experiment 15      Determine the strength of hydrogen bonding between some molecules

Introduction
     Breaking or formation of intermolecular hydrogen bonds between molecules in liquids would cause
an enthalpy change when the liquids are mixed. This experiment is to investigate such enthalpy changes
and to measure approximate strengths of hydrogen bonds formed between molecules of ethanol and those
between molecules of trichloromethane and ethyl ethanoate using simple calorimetric methods.

Chemicals
Ethanol, cyclohexane, ethy1 ethanoate, trichloromethane, tetrachoromethane

Apparatus
10 cm 3 and 25 cm 3 measuring cylinders, 50 cm 3 , -10-110°C thermometer

Procedure
[Hazard Warning: Ethanol, cyclohexane and ethyl ethanoate are flammable, trichloromethane is harmful,
                and tetrachloromethane is toxic.]

A. To discover the existence of hydrogen bonds between ethanol molecules
1.  Try to use a measuring cylinder, add 10 cm3 of ethanol into an insulated 50 cm3 beaker.   Measure
   the temperature of the liquid.

2.    Then add 10 cm3 of cyclohexane to the ethanol in the beaker, mix well and record the lowest
     temperature attained.

     (a)    Why should the beaker be insulated?

     (b)    Is the mixing process endothermic or exothermic?

     (c)    Account for the temperature change.

B.   To measure the strength of hydrogen bond formed between ethanol molecules

      Repeat steps (1) and (2) in Part A above using the same volume of ethanol but
      20cm 3 of cyclohexane. From the temperature drop estimate the hydrogen bond
      strength (in KJ mol 1 ) in ethanol.

     (a) Explain why cyclohexane has to be used in excess in this experiment?

     (b) Comment on the reliability of the hydrogen bond strength obtained.
C.     To invest/gate the format/on of hydrogen bonds between molecules of ethyl ethanoate and
       trichloromethane
     [Hazard Warning:      Mixtures of trichloromethane and propanone have been known to explode on
                           standing. The solvent residues from this experiment should not be disposed of
                           into a container in which propanone is present.]

1.      Measure 10cm 3 of ethyl ethanoate into an insulated beaker. Record its temperature.

2.      Add to this 10 cm 3 of trichlororrsethane and mix well. Record the highest temperature attained.
        (a) Is the mixing process exothermic or endothermic?
        (b) Account for the temperature change.

D.     To measure the strength of hydrogen bonds formed between molecules of ethyl ethanoate and
       trichloromethane

       Repeat steps (1) and (2) in part C above using either one liquid in excess. From the temperature
       change estimate the strength of the hydrogen bond formed between molecules of ethyl ethanoate and
       trichloromethane.

       Explain why it does not matter which liquid is used in excess.

E.     To discover which atoms of the molecules of ethyl ethanoate and trichloromethane are essential for
       the formation of hydrogen bonds

           Measure 10cm 3 of ethyl ethanoate into an insulated beaker and record its temperature. Add to it
       10 cm 3 of tetrachloromethane, with stirring, and record the new temperature attained.
       (Hazard Warning: Tetrachloromethane is toxic. Avoid inhaling vapour and skin contact.]

        (a) Is there any significant change in temperature on mixing the two liquids?

        (b) Compare this with the mixing of ethyl ethanoate and trichloromethane in Part C above and
            suggest how the hydrogen bond is formed between them.

       The following physical data may be useful:

                             Specific heat capacity of glass = 0.78 kJ kg -1 K-1
                                                   Relative              Density        Specific heat capacity
         Liquid             Formula
                                             Molecular mass              /kg dm 3          / kJ kg 1 K 1

        Ethanol            CH3CH2OH                   46                  0.81                   2.44

      Cyclohexane            C 6 H 12                 84                  0.78                   1.83

Trichloromethane             CHCl 3                 119.5                 1.48                   0.98

     Ethyl ethanoate    CH3CO2CH2CH3                  88                  0.90                   1.92
Experiment 16        Investigation of Some of the Properties of a Pair of Cis-Trans Isomers

Introduction
Maleic acid and fumaric acid are geometrical isomers of butenedioic acid. Each of these isomers
has its own distinctive properties such as melting point, solubility, density and stability.

      In this experiment some maleic acid is converted to fumaric acid by heating an aqueous
solution of maleic acid in the presence of hydrochloric acid. The hydrochloric acid serves merely
as a catalyst of the reaction. The properties of these two isomers are then compared.

Chemicals
Maleic acid, magnesium ribbon, Na2CO3, concentrated HCl, bromine water, pH paper

Apparatus

100 cm3 and 250 cm3 beakers, watch glass, apparatus for suction filtration, melting point
apparatus, 25 cm3 measuring cylinder

Procedure

[Hazard Warning:     Maleic acid is irritant, concentrated hydrochloric acid is corrosive, and
                     bromine water is harmful.]
A.   Conversion of maleic acid to fumaric acid

1.   Weigh out about 1 g of maleic acid in a clean dry 100 cm3 beaker. Add 10 cm3 of deionized
     water and warm slightly to dissolve the acid.

2.   Add 10 cm3 of concentrated hydrochloric acid, and cover the beaker with a watch glass.
     Place the beaker inside a 250 cm3 beaker which is about one third full of water. Heat this
     water bath to boiling for about 5 minutes or until a solid material forms in the small beaker.

3.   Cool the solution to room temperature by placing the small beaker with its contents in a cold
     water bath or in an ice bath.

4.   Filter the reaction mixture by suction using the following set-up:

                                                 Buchner funnel
                                                 clamp

                                                 to filter pump



                                                 filtrate
5.   Stop suction, either by lifting the funnel or by disconnecting the tubing, and soak the residue
     in about 1 cm3 of cold water. (If you turn off the tap, you may get a "suck-back" of water.)

6.   Resume suction and dry the crystals by drawing air through them for a few minutes.
7.   Transfer the crystals into a weighed watch glass and dry in an oven at about 120C for 10
     minutes.

8.   Weigh the dried crystals of fumaric acid.
B.   Comparison of properties of the two isomers

1.   Solubility in water  place about 0.1 g of each isomer into 10 cm3 of water in separate test tubes,
     shake to help dissolving. See which isomer is more soluble.

2.   Melting point  using the electrical melting point apparatus, measure the melting points of the two
     isomers.

3.   Acid strength  for each of the two isomers, prepare a solution by dissolving about 0.1g of the
     compound in about 20cm3
     following tests:

     (a)   Measure the pH of the solution.

     (b)   Add a 3 cm strip of magnesium ribbon.

     (c)   Add a small amount of solid sodium carbonate.

4.   Reaction with bromine water  suspends about 0.1g of the acid in about 5 cm3 of water. Add 3
     drops of bromine water to the resulting solution/suspension. Shake and observe.

Results
A.   Conversion of maleic acid to fumaric acid
           Mass of fumaric acid + watch glass                          g

           Mass of watch glass                                         g

           Mass of fumaric acid                                        g

           Percentage yield                                            %

B.   Comparison of properties of the two isomers

     1.    Water solubility :                              acid is more soluble.

     2.    Melting point of maleic acid :                  C

           Melting point of fumaric acid :                   C
     3.      Acid strength
                    Test                   Maleic acid         Fumaric acid

          pH of solution

          Reaction with Mg

          Reaction with Na2CO3


     4.      Reaction with bromine water



             Maleic acid :



             Fumaric acid :


Discussions
1.   Assuming that equilibrium concentration was achieved in procedure A, which isomer would
     you classify as the more stable with respect to transformation of one into the other?

2.   What do each of the following tests contribute to your knowledge of the structure of each
     isomer?

             (a) The reaction with magnesium and sodium carbonate.

             (b) The pH value.

             (c) The melting point.

3.   Considering the structures of the two isomers, try to account for the observed differences in
     solubility and melting point.

4.   One of the isomers can lose a molecule of water from each molecule of acid when its two
     carboxyl groups react to form an anhydride. Which geometrical isomer, cis- or trans-, do
     you predict it is?
Experiment 17
Identification of a Carbonyl Compound by Preparing its Derivative

Introduction
Crystalline derivatives of many carbonyl compounds can be formed by condensation reactions
with compounds such as phenylhydrazine and 2,4-dinitrophenylhydrazine. These derivatives can
usually be isolated in relatively pure forms which have well defined melting points.
Phenylhydrazine forms derivatives (phenylhydrazones) readily with aromatic aldehydes, but in
general      2,4-dinitrophenylhydrazine      is     preferred   because       its    derivatives
(2,4-dinitrophenylhydrazones) have higher melting points and are less soluble. These derivatives
are useful in identification of carbonyl compounds.


Chemicals
2,4-dinitrophenylhydrazine in methanol, ethanol, methanol, carbonyl compound labelled X
(different groups of students may work on different carbonyl compounds), dilute H 2SO4


Apparatus

Beaker, 100m3 measuring cylinder, apparatus for melting point determination, apparatus for
suction filtration, rubber rings (cut from a rubber tubing of appropriate diameter), ice bath


Procedure

[Hazard Warning: Methanol solution of 2,4-nitrophenylhydrazine is flammable and toxic,
methanol, ethanol and many carbonyl compounds are flammable, and bench dilute sulphuric(VI)
acid is corrosive.]
A.   Preparation of 2,4-dinitrophenylhydrazone of compound X

1.   Add 10 drops of compound X to 5m3 of 2,4-dinitrophenylhydrazine solution in a 50m3
     beaker. If crystals are not formed, add about 1m3 of dilute sulphuric(VI) acid. If they are still
     not formed, warm the mixture gently, then cool with scratching in ice water.
2.   Filter off the crystals by suction filtration. While still on suction, wash the crystals with 1m3 of
     methanol.
3.   Recrystallize the crystals from hot ethanol as follows:
     (a)  Transfer the crystals to a 100m3 beaker standing on a steam bath (or in a 250m3
          beaker of hot water).
     (b) Dissolve the crystals in the minimum amount of hot ethanol.
     (c) When the crystals have dissolved, cool the solution in an ice-water mixture until
          crystals reappear.
     (d) Filter the crystals by suction. If necessary, rinse the beaker with the filtrate (not the
          extra solvent) to complete the transfer. Finally, wash the crystals with a few drops of
          cold ethanol. Dry the crystals by drawing air through them for a few minutes.
4.   Spread the crystals on a dry watch glass and leave overnight in a desiccator for drying.
B.   Determination of the melting point of the 2,4-dinitrophenylhydrazone of compound X

1.   Introduce a small amount of the crystals into a melting point tube until a total length of about
     0.5m is compacted at the bottom of the tube.
2.   Attach the prepared melting point tube to the thermometer, as shown in Figure 1.
3.   Half-fill a boiling tube with paraffin oil, and position the thermometer with attached tube and
     the stirrer as shown in Figure 2.
4.   Position the apparatus over a low Bunsen flame and gauze and gently heat the apparatus,
     stirring the paraffin oil all the time by moving the stirrer up and down.
5.   Note the temperatures when the crystals start to melt and when the melting is completed.



                             rubber ring

                             melting point                              clamp
                             tube
                                                                        stirrer

                              sample                                    paraffin oil
                              thermometer                               melting point tube
                              bulb                                      with sample


                                                       heat

                                                                  Figure 2
                          Figure 1


6.   Compare the melting point of the crystals with the values given in the following table and
     thus identify compound X.


                                                              Boiling         Melting point of
                   Name                    Formula            point          2,4-dinitrophenyl-
                                                               / C           hydrazone / C
        Aldehydes
         methanal                 HCHO                         21                     167
         ethanal                  CH3CHO                        21          146, 164 (2 forms)
         propanal                 CH3CH2CHO                     48                     156
         butanal                  CH3CH2CH2CHO                  75                     123
         2-methylpropanal         (CH3)2CHCHO                   64                     187
         benzaldehyde             C6H5CHO                      178                     237
        Ketones
         propanone                CH3COCH3                      56                     128
         butanone                 CH3CH2COCH3                   80                     115
         pentan-2-one             CH3CH2CH2COCH3               102                     141
         pentan-3-one             CH3CH2COCH2CH3               102                     156
         hexan-2-one              CH3CH2CH2CH2COCH3          128               107
         4-methylpentan-2-one     (CH3)2CHCH2COCH3           117                95
         cyclohexanone                                        156              162
                                           O

Discussion
1.   What soluble impurities may be present in your product before recrystallization? How can
     they be removed in the recrystallization process?
2.   What factors should be considered in selecting a suitable solvent in the recrystallization
     step?
3.   In the recrystallization procedure, why were the crystals dissolved in only the minimum
     amount of hot ethanol?
4.   If the sample contains insoluble impurities such as pieces of filter paper, cork, etc., suggest
     how they can be removed.
5.   If the melting point of the 2,4-dinitrophenylhydrazone is 156C, suggest how you can
     confirm whether compound X is propanal or pentan-3-one.

				
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Description: Experiment A Quantitative Study of Electrolysis filler0