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, 10110C 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 30C. 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 35C, 40C, 45C and 50C. 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 80C? 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 mol1K1) 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 120C 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 156C, suggest how you can confirm whether compound X is propanal or pentan-3-one.