Sian ZhangCHM122BApparatus:Reagents:

Potassium Iodate (solid) Potassium Iodide (solid) Sodium Thiosulphate xH20 (solid) Sulphuric acid 1M (H2SO4(aq)) Starch solution (1%) Orange juice (Analyte) Deionized WaterEquipment: Burette; appropriate clamps and stands; Burrete clamp + Retort Stand Pipette; 10ml ; Pipette filler Volumetric flasks Conical Flask Measuring Cylinders (10ml, 25ml) Digital scales Hot Plate Bottles (5) Thermometer Spatula, Stirring rod Beakers Funnels Aluminium dishes Wash bottlesMethod (including revisions):A) Potassium Iodate Solution:

1. 4 grams of solid potassium iodate (KIO3) was measured using digital scales and transferred to a volumetric flask.

2. The volumetric flask was filled up to a graduation mark at 1L. Its contents were agitated until all the solid potassium iodate was dissolved. This created a 0.01869M solution.

B) Sodium Thiosulphate Solution:

1. 15.8 grams of solid sodium thiosulphate (Na2S2O3) was measured using digital scales and transferred to a volumetric flask.

2. The volumetric flask was filled up to a graduation mark at 1L. Its contents were agitated until all the solid sodium thiosulphate was dissolved. Due to the uncertainty regarding the water content of sodium thiosulphate, standardization of the solution was necessary to ascertain the molarity of the solution created.

C) Triiodine Solution:

1. 10ml of potassium iodate solution from method A was transferred into a conical flask using a pipette.

2. 0.5 grams of potassium iodide was measured out using digital scales and transferred to the volumetric flask. Note that any amount of potassium iodide can be added, given that it is in excess in the solution. The 8:1 stoichiometric ratio between potassium iodine and iodate, which indicates that eight moles of potassium iodide is required to react with one mole of potassium iodate to form three moles of triiodine.

3. 2.5ml of 1M sulphuric acid was measured in a measuring cylinder and transferred into the volumetric flask. Upon the addition of hydrochloric acid, the contents of the volumetric flask turned dark brown, indicating the formation of the dark-brown triiodine ion (I3-). The concentration of the triiodine solution is equal to three times the molarity of the potassium iodate solution used, as the potassium iodate solution acted as the limiting reagent in the reaction. The molarity of the triiodine solution was 0.05607M.Note that the triiodine solution is volatile in acidic environments, and is liable to be oxidized into iodide. For this reason, the triiodine solution must be used immediately after it is generated.Figure 2.1 Triiodine solution (Left pre-H2SO addition 4. Right post-addition of H2SO4)

D) Standardization of Sodium Thiosulphate Solution:

1. A burrete filled with the sodium thiosulphate solution from method B was positioned directly above a conical flask which contained the triiodine solution from method C.

2. Sodium thiosulphate was titrated into the flask until the solution changed from a dark-brown colour to a pale yellow.

3. 3 drops of 1% starch solution was added to the flask, causing its contents to turn deep black-brown. Sodium thiosulphate was added until the solution turned colourless, indicating the endpoint of the titration.

4. The volume of sodium thiosulphate used was noted. Steps 1-3 were repeated twice to produce three concurrent sets of data, and the average of the three values was calculated.

Thiosulphate ions react in a 2:1 ratio with the iodine present in equilibrium in the triiodine solution. Hence the amount of thiosulphate supplied by the noted volume of solution is equal to double the known moles of triiodine ions. From this the moles of sodium thiosulphate are inferred, and subsequently, the molarity of the sodium thiosulphate solution is determined. Figure 2.2 Standardization (Left pre-starch addition. Centre Starch Added. Right End-point)

E) Preparation of Juice Samples for testing variables; Exposure duration, Temperature

Exposure Duration:

1. Five separate aliquots of 200mls of the same orange juice was prepared and stored in separate bottles. Each bottle was left without its lid in order to expose the orange juice to atmospheric oxygen. Samples were stored in a refrigerator set at 5-6oC in order to prevent the growth of mold in the samples, which may interfere with results. Each of these aliquots represents a sample used for each varied length of exposure-duration.

Temperature:

2. 25 ml aliquots of analyte were separated into measuring cylinders, and heated to desired temperatures (temperature measured using thermometer) immediately prior to carrying out experimentation, using a hot plate.

F) Titration for Ascorbic acid content in juice

1. When testing for oxygen-exposure, 25ml of orange juice from the oxygen-exposure juice samples was measured out with a measuring cylinder and transferred into a conical flask containing triiodine solution from method C. When testing for temperature, the contents of the aliquot of juice prepared in E) 2. was used instead.

2. A burrete filled with the sodium thiosulphate solution from method B was positioned directly above the conical flask.

3. Sodium thiosulphate was titrated into the flask until the solution changed from a dark-brown colour to a pale green-yellow.

4. 3 drops of 1% starch solution was added to the flask, causing its contents to turn deep black-brown. Sodium thiosulphate was added all traces of black and brown disappeared, with the solution turning green. This indicated the endpoint of the titration.

5. The volume of sodium thiosulphate used was noted. Steps 1-4 were repeated twice to produce three sets of concurrent data, the average value being calculated. This was repeated for each different parameter of the variables (juice exposure time to oxygen; temperature) being tested. In total, 5 parameters were tested for each variable, for each of which three sets of data was obtained.

The known moles of sodium thiosulphate used correlates with the excess iodine in the triiodine solution, after it has reacted with ascorbic acid in the juice. The moles of sodium thiosulphate used are double the moles of excess iodine. The exact amount of excess iodine was calculated and subsequently, the ascorbic acid, which reacts in a 1:1 ratio with iodine, was determined, by finding the difference between the initial moles of iodine solution, (before it had been reacted with ascorbic acid) and the moles of excess iodine reacted with sodium thiosulphate.Figure 2.3 Titration (Left pre-starch addition. Centre Starch Added. Right End-point)

Revisions Made:In the initially proposed method, it was suggested 20mls of potassium iodate solution was to be used in creating the triiodine solutions used in the standardization of sodium thiosulphate solution and ascorbic acid titration. This was changed to 10mls, in order to reduce the amount of titrant required to reach the end-point in titrations, while maintaining accuracy. Originally, a 0.25 M thiosulphate solution was used; however this caused each drop of solution to contain more thiosulphate, therefore limiting the accuracy of the titrations carried out. Therefore, a dilute solution was used instead (0.06M) to increase the accuracy of results yielded. This also reduces the relative uncertainties in the results yielded, as a greater amount of titrant is used in each titration.When preparing the orange juice samples for titration testing the effect of temperature, a beaker of 200ml of analyte was heated and then separated into 25ml aliquots for testing. This caused errors, especially when the temperature being tested began to approach 100oC, at which point, the evaporation of water in the juice samples occurred. This resulted in the significant reduction in juice volume, causing the test sample to become much more concentrated for experiments testing higher temperatures, resulting in inconsistent initial ascorbic acid values for each different (20, 40, 60, 80, 100oC) temperature tested. In order to remedy this, aliquots of 25mls were separated prior to heating into separate conical flasks, and then heated. This meant that each aliquot of analyte began with consistent ascorbic acid contents, and any subsequent variations in ascorbic acid content was due to temperature, rather than being attributed to inconsistencies in initial content levels.

Results:Figure 2.4 Ascorbic acid content vs oxygen-exposure time

Duration of Exposure (nearest hour)Ascorbic Acid (mg)Uncertainties Mg

015.8302.665

4316.0952.661

14211.7623.432

16712.3812.715

21111.4962.905

Experimentation testing the duration of exposure, shown in figure 2.4, shows that the ascorbic acid content of the analyte decreased as the duration of the juices exposure to atmospheric oxygen increased. Hence the independent variable, duration of exposure, and the dependent variable, ascorbic acid content, share an inversely proportional relationship, as predicted in the hypothesis. The chosen trend-line fitted to the data points is one of exponential decay, which indicates that the rate of ascorbic acid deterioration is proportional to its the amount of ascorbic acid in the solution. Hence, when there is more ascorbic acid content in a solution exposed to oxygen, it will degrade relatively faster than a similar solution with less ascorbic acid content. There are a number of points that deviate from this trend line, however, the degree of deviation of these points lie is smaller than the uncertainties propagated, and hence, it is concluded that these discrepancies are due to experimental errors.

Figure 2.5 - Ascorbic Acid content vs Temperature

Temperature Ascorbic Acid (mg)Uncertainties Mg

2020.6932.593

4019.4552.611

6020.3402.598

8016.9792.648

10010.2582.923

Experimentation testing the temperature variable, shown in figure 2.5, shows that increasing the temperature of the analyte causes a corresponding decrease in the ascorbic acid content of the sample. Hence, the independent variable, temperature, and the dependent variable, ascorbic acid content, share an inversely proportional relationship, as predicted in the hypothesis. The chosen trend-line for the data points is a negative linear relation. This indicates that the rate of ascorbic acid decay, given by the gradient of the trend-line, is constant at all temperatures. There are a number of points that deviate from this trend line, however, the degree of deviation of these points lie is smaller than the uncertainties propagated, and hence, it is concluded that these discrepancies are due to experimental errors.In figures 2.4 and 2.5, exponential and linear trend-lines were chosen for their respective data points. Both trend-lines have the best R^2 value (excluding polynomial trends) for each of their sets of data. Although polynomial trends have greatest R^2 value (see appendix), extension of these graphs suggest that the ascorbic acid content of samples increases parallel to increase in the independent variables after a certain point, which should not happen, as shown by background research. Sample Calculations:Triiodine Solution: Firstly, the concentration of the potassium iodate solution used in producing triiodine is found using the equation c=m/v, where m is the moles of KIO3 and v is the volume of distilled water used. This concentration remains constant for all calculations.

There is an uncertainty of + 0.02g associated with the digital scales in measuring the mass of KIO3. A + 0.3ml uncertainty is associated with the volumetric flask used to measure 1L of distilled water. In order to find the total uncertainty in the concentration of the KIO3 solution, these absolute uncertainties must be converted into relative uncertainties. (Calculations involving multiplication or division require relative uncertainties)

The total uncertainty in the potassium iodate solution is the sum of these two values

However, each time KIO3 solution was utilized; it was used in 10ml portions. Therefore the moles of iodate used in each titration are as follows:

The pipette used to measure these moles has an associated error of 0.04ml.

Hence, the uncertainty of the moles of KIO3 in a 10ml sample is actually the sum of the KIO3 solution uncertainty and this uncertainty.

Here an excess of KI is added to the KIO3 solution calculated previously, as well as an excess of H2SO4. Hence, given that the values of the moles associated with these two substances do not dictate the molarity of the triiodine solution formed, the uncertainties associated with these two substances are irrelevant. Upon inspection, it is found that KIO3 and KI to form 3 moles of triiodine. Hence the triiodine solution (for the actual equation, refer to stage 1) has three time the number of KIO3 moles, and shares the same relative uncertainty.Moles of 3I3 formed:

Standardization Sample Calculation:The sodium thiosulphate is known to react with iodine (equivalent to triiodine) in a 2:1 ratio. Hence, the amount of thiosulphate moles required to react with the triiodine solution is:

The relative uncertainty of the moles of thiosulphate is the same as the uncertainty of the triiodine as these moles are supplied via titration, 0.93%. There is however, another uncertainty associated with the how these moles were supplied the error of the burrete used for titration of the solution. TrialStart (ml) EndDelta

1422.618.6

222.641.218.6

3018.618.6

AVG18.6

The total apparatus uncertainty (or individual uncertainty) is. All titrations yielded the same amount of thiosulphate solution used; consequently as there is no deviation from a mean, the error associated with the volume can be concluded to be

From this, it can be concluded that 18.6mls of the sodium thiosulphate used contains of thiosulphate. Hence the molarity of the thiosulphate solution is as follows:

Total uncertainty for the sodium thiosulphate concentration:

Note: all values calculated for thiosulfate and triiodine remain consistent for all ascorbic acid calculations

Ascorbic Acid Calculation:The triiodine and thiosulphate solution used in determining the ascorbic acid concentration are as the previous calculations specify. The moles of triiodine used are. This has an associated uncertainty of 0.68%.The concentration of sodium thiosulphate used is, with a 1.4676% associated uncertainty.It was observed that after these moles of triiodine had reacted, the following volumes of thiosulphate were required to react with the remaining moles of triiodine.TrialStartEnddelta

131.647.215.6

219.735.315.6

319.6535.315.65

AVG15.61666667

Here, the greatest deviation from the mean was given by 15.61666667-15.65 = 0.03333... This value is within the individual uncertainty of + 0.1 ml, hence the uncertainty of the average volume is concluded to be + 0.1ml.The moles of sodium thiosulphate used are given by:

The relative uncertainty of the moles of sodium thiosulphate used is the sum of the uncertainties of the concentration and volume of the sodium thiosulphate. The concentration uncertainty is stated above.

The moles of iodine with which the sodium thiosulphate is half the moles of the sodium thiosulphate used (due to 2:1 ratio), and its associated errors is the same as the relative uncertainty of the moles of sodium thiosulphate.

From this, the amount of iodine reacted with ascorbic acid can be determined, and consequently, the moles of ascorbic acid in the sample can be calculated.

As ascorbic acid reacts with iodine in a 1:1 ratio, the moles of ascorbic acid in the sample is:

The uncertainty of this value is equal to the sum of the absolute uncertainty of the total moles of iodine and the absolute uncertainty of the sodium thiosulphate (addition and subtraction calculation, therefore absolute uncertainties are used.)