study on the kinetics of degradation of ascorbic acid from different orange juices' (ib extended...

Study on the kinetics of degradation of ascorbic acid from different orange juices

Francisco Jess Moreno Velarde

WORDS: 3719

Extended Essay: Study on the kinetics of degradation of ascorbic acid

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Abstract

Kinetics of ascorbic acid degradation in three different types orange juice (natural, refrigerated and non-refrigerated) during a period of ten days was investigated. The samples were kept in refrigerated conditions and studied at ambient temperature and pressure during the month of July. Redox titration with iodine was used for ascorbic acid determination. Results showed vitamin C concentration decreased within the following days, following a zero-order kinetic model, a different result to the ones in bibliography. The reaction rate constants calculated were 0.0011 molL-1s-1 for natural juice, 0.0021 molL-1s-1 for refrigerated and 0.0026 molL-1s-1 for non-refrigerated. Difference between natural and non-natural have related to the addition of chemicals during manufacturing process. Words: 110


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INDEX

1. Introduction.3

2. Materials...5

3.1. Analytes.5

3.2. Apparatus.5

3.3. Reagents5

3. Method6

4.1. Reagent Manufacture.6

4.2. Titration Process.6

4. Results.6

5.1. Obtaining the concentration of vitamin C per sample.7

5.2. Measurement error calculation...9

5.3. Reaction rate constant calculation....10

5. Discussion.....11

6.1. Results.11

6.2. Conclusion...13

6.3. Evaluation and Improvements...13

6. References....14


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1. Introduction

Mothers often tell their children to drink their orange juice quickly if they do not want vitamins to go away. I found it interesting to investigate the rate of degradation of vitamin C in orange samples to check if this was true or simply popular belief.

Our objective in this investigation is to study the degradation of vitamin C in samples of

three different types of orange juice within a period of 10 days. We will determine the order of reaction followed by our samples and we will calculate the reaction rate constant k based on the results of the experiments.

Vitamin C is an essential nutrient for our organism. It is necessary for our bodys

development. Humans, as well as other primates and other species, cannot synthesize this nutrient due to the absence of the enzyme L-gulonolactone oxidase in their organisms. This enzyme is capable of catalyzing the conversion of glucose in vitamin C (without enzymatic action). Therefore, it is extremely important to maintain the intake of vitamin C in order to develop a healthy organism. Vitamin C is the L-enantiomer (or optic isomer) of L-ascorbate (an ion of ascorbic acid) ascorbic acid being the reduced form of vitamin C as hydrogen molecules add to L-ascorbate. Vitamin C is very unstable in aqueous solutions. Tests showed ascorbic acid to be a weak, monobasic acid and a strong reducing agent.1 Therefore, it oxidizes and the product of such oxidation involving the loss of two hydrogens (C6H6O6), was named dehydroascorbic acid.

2 This substance is easier to absorb and can therefore cross cellular membranes more easily. This way, organisms which cannot synthesize ascorbic acid absorb it through active transport and passive diffusion. This step is reversible, but it does not take long for this substance to change into 2,3-diketo-L-gulonic acid. This step is irreversible, therefore, it is necessary to avoid. Even though this substance is named acid it is a lactone; therefore, its acidity and ease of oxidation are due to the presence of an enediol group3(an alkene with a hydroxyl group on both sides of the double bond).

Figure 1: L-ascorbic acid molecule (left) and L-ascorbate ion (vitamin C, right)

The reactions in which vitamin C and its derivatives participate can be summarised in the

following scheme:

1 Davies, Michael B.; Austin, John and Partridge, David A. (1991): Vitamin C: Its Chemistry and Biochemistry p. 30



Enediolgroup


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Figure 2: Reaction routes in which vitamin C or derivatives participate

Vitamin C is mainly used for tissue growth and repair. It is necessary for the creation of collagen which is a protein molecule responsible for the formation of the skin, tendons, ligaments and blood vessels. Vitamin C is also very important for skin and bone maintenance. It is responsible for wound healing and osteogenesis, which is suggested to be an effective antiviral agent and its antioxidant properties are well known as it delays other molecules oxidation. Oxidation damages other cells when free radicals are created in the beginning of reactions. Free radicals are created when decomposing food or when exposed to radiation or even cigarette smoke. If free radicals are grouped, the ageing process may be stronger, and can lead to important diseases, such as cancer or heart attacks. Antioxidants eliminate these free radicals making them unable to take part in other reactions. Hence, vitamin C blocks the ageing of the body.

However, there are secondary effects when an excess amount of vitamin C is taken in. These

are, amongst others, gastric irritation, taste deterioration and renal problems due to the action of the vitamins metabolic action (oxalic acid). This can lead to inhibition of natural processes. Therefore, vitamin C is an essential nutrient for our organism. It is necessary for our bodys development. Humans, as well as other primates and other species, cannot synthesize this nutrient due to the absence of the enzyme L-gulonolactone oxidase.

This vitamin is soluble in water, making its consumption easier. The excess amount of vitamin C consumed can be eliminated by the body. Even still, the vitamin intake must be constant. Several food groups contain vitamin C. Citrus fruits , such as the orange, lemon, lime or grapefruit are known for their content in this vitamin, but there are foods with even more vitamin C in them. Our study subjects, oranges, contain 50 mg of ascorbic acid per 100 g4.

One of the common methods to determine vitamin C quantities in food is a redox titration determination. This method shows better results than a normal acid-base titration due to the presence of other acids, apart from ascorbic acid, which can interfere in the oxidation of the primary acid and hence change the results. Ascorbic acid is being oxidized by iodide.

Iodide is relatively insoluble, but this characteristic can be fixed as we add iodine to form a triiodide ion.

I2 + I- I3

-



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As we mentioned before, the triiodide ion oxidizes ascorbic acid so as to convert it into

dehydroascorbic acid, as shown in the reaction equation below:

C6H8O6 + I3- C6H6O6 + 3I

- + 2H+

It is necessary to also add a few drops of a prepared starch solution which acts as an iodine detector.

As long as there is vitamin, the triiodide ion will react to form iodine ion very quickly. However, when the vitamin is completely oxidized, there will be the same quantity of iodine as of triiodide. This will eventually react with the starch solution applied to the mixture and a deep blue stain will form. This shows the end of the titration.

2. Materials

2.1. Analytes:

- Don Simn non-refrigerated orange juice - Pascual refrigerated orange juice - Fresh juice (as of Wednesday 6th of july 2011)

2.2. Apparatus:

- 25 mL burette - 25 mL pipette - Clamp and stand for the burette - 250 mL erlenmeyer flasks - 500 mL volumetric flask - 100 mL beakers - Plastic Pasteur pipettes

2.3. Reagents:

- Potassium iodide (KI), Panreac - Potassium iodate (KIO3), Panreac - Concentrated (98%) sulfuric acid (H2SO4), Panreac - Ascorbic acid powder, Acofarma - Starch powder, Panreac - Distilled water


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3. Method

The method to manufacture reagents and titrate samples was taken from the following web page:

3.1. Reagent manufacture

- Iodine solution: We dissolved 5.005 grams of KI and 0.303 grams of KIO3 in 200 mL of

distilled water. Then we added 30 mL of 3 M H2SO4. Finally, we added up to 500 mL of distilled water. We noticed a change in colour from transparent to brown orange.

- Vitamin C standard solution: We had to dissolve 0.252 grams of vitamin C powder in distilled water up to a final solution volume of 250 mL.

- 1% starch indicator: We had to mix 0.50 grams of starch with 50 mL of hot water and wait until it cooled down

3.2. Titration process

- Before titrating the samples, the iodine solution must be titrated against the standard vitamin C solution by carrying out the following procedure. This titration should be repeated if a new iodine solution is made (because of finishing the first). CLARIFY

- Pour 20.00 mL (all of the samples were 20.00 mL) of the substance we want to titrate into the Erlenmeyer flask using the pipette.

- Fill the burette with the iodine solution. - Before starting the titration, 10 drops of starch solution have to be added. - Once everything is ready, open the tap and mix the solution with the substance that

contains vitamin C slowly until it changes colour (normally deep blue or grey). This would mean the ascorbic acid has been fully oxidized and now iodine is in excess

- Repeat the titration with a second sample in order to reduce random error and to establish a mean value for the calculations.

4. Results First of all, we titrated the vitamin C standard solution. This was the data obtained.

4th July 2011

Vitamin C

11.2 mL

11.3 mL

11.0 mL

Figure 3: Table showing the volume to neutralize the standard solution This data would help us in the future to determine the amount of vitamin C. This will be explained in the next section. The three juices were labeled and kept separately in the fridge. Measurements were taken every two days (except on weekends, when they were not taken) for a period of two weeks, and


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therefore, we had 5 measurements by the end of this period. Day 0 was established the 6th of July and Day 9 was the 15th of July. Throughout the entire 10-day process, the average temperature in the fridge was about 4 C. Measurements were made at about 30 C and 1 atmosphere. Firstly, we are going to show the raw data for the titrations in the following table:

Day 0 Day 2 Day 5 Day 7 Day 9

Don Simn NON-REFRIGERATED

12.9 mL 11.5 mL 10.8 mL 10.1 mL 9.2 mL

12.6 mL 11.8 mL 10.7 mL 10.1 mL 8.9 mL

Pascual REFRIGERATED

10.6 mL 10.1 mL 9.7 mL 8.8 mL 8.2 mL

10.5 mL 10.2 mL 9.5 mL 9.3 mL 8.1 mL

Natural 10.4 mL 10.5 mL 9.5 mL 9.2 mL 9.2 mL

10.8 mL 10.3 mL 9.3 mL 9.2 mL 9.2 mL

Figure 4: Table containing raw data collected from the experiment

4.1. Obtaining the concentration of vitamin C per sample This table shows the quantity of iodine necessary to neutralize the juice. We can figure out

how much vitamin C there was in the sample based on the amount of iodine used to neutralize it. We followed the following method in order to calculate the amount of vitamin C. This will be later explained using real data to maximize comprehension.

1) First of all, calculate the mean value of the volumes (the sum of both volumes divided by the

number of measurements, two in this case).



12.8 0.2 mL 11.7 0.2 mL 10.8 0.1 mL 10.1 0.1 mL 9.1 0.1 mL


10.6 0.1 mL 10.2 0.2 mL 9.6 0.1 mL 9.1 0.3 mL 8.2 0.1 mL

Natural 10.6 0.2 mL 10.4 0.1 mL 9.4 0.1 mL 9.2 0.1 mL 9.2 0.1 mL

Figure 5: Table containing mean volumes and their respective errors.

2) Knowing how much iodine was used to neutralize our standard vitamin C solution, we can calculate a value for the specific sample.

3) We can divide this data by the volume to obtain concentration by litres or divide it by its molecular mass to obtain the number of moles.

Using the raw data obtained from the experiments, we are going to show the calculations

involved to determine the final quantities of vitamin C. Let us use the data for the first titration of Don Simn non-refrigerated juice: Day 0. We are going to calculate concentration in grams and moles.

1) The volumes necessary to neutralize the sample were 12.9 mL and 12.6 mL. The mean

value is 12.8 mL


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2) Knowing how much iodine was necessary to neutralize our sample, we can use cross-

multiplication to solve for the mass of vitamin C in our sample. Our standard solution was made using 0.252 grams of vitamin C powder and the mean

volume to neutralize it was 11.2 g.

In this first table below, we are showing the amount of vitamin C in each sample in grams:



0.29 g 0.26 g 0.24 g 0.23 g 0.20 g


0.25 g 0.23 g 0.22 g 0.20 g 0.18 g

Natural 0.24 g 0.23 g 0.21 g 0.21 g 0.21 g

Figure 6: Table showing the amount of vitamin C in grams

3) Dividing by the volume of the sample (20 mL), we obtain the concentration by volume

However, the reaction rate constant kv is measured in units; therefore, we need to

calculate concentration in moles. To calculate the amount of moles in each sample, we simply divide the grams by the molecular mass of ascorbic acid, C6H8O6 which is 176 g/mol. The tables for moles and for concentration in moles per litre are going to be shown below:



0.0016 mol 0.0015 mol 0.0014 mol 0.0013 mol 0.0011 mol


0.0014 mol 0.0013 mol 0.0013 mol 0.0011 mol 0.0010 mol

Natural 0.0014 mol 0.0013 mol 0.0012 mol 0.0012 mol 0.0012 mol

Figure 7: Table showing the amount of vitamin C in moles


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Concentration in moles per litre (dividing by the volume of the sample):



0.080 mol/L 0.075 mol/L 0.070 mol/L 0.065 mol/L 0.055 mol/L


0.070 mol/L 0.065 mol/L 0.063 mol/L 0.055 mol/L 0.050 mol/L

Natural 0.070 mol/L 0.065 mol/L 0.060 mol/L 0.060 mol/L 0.060 mol/L

Figure 8: Table showing the concentration of vitamin C in moles/L

4.2. Measurement error calculation

Errors are very important to determine the precision of our results. Throughout the experiment, we have been calculating indirect magnitudes from the ones obtained directly from the lab. We have taken decimals and this could lead to error. We calculated the amount of vitamin C in grams and moles and the concentration in grams per litre and moles per litre. As the molecular mass does not have errors, we are going to calculate the error of the measurements in grams.

To calculate the mass of vitamin C in our sample, we had to multiply the quantity of vitamin

C in the standard solution by the volume used to neutralize the sample and divided by the volume taken to neutralize the standard solution.

Therefore, to calculate the error of the mass in the sample, we have to add the relative

errors of the mass of vitamin C in the standard solution, the relative error in the volume to titrate the sample and the relative error of the volume used to titrate the standard solution.

m and V being the mass and volume used in the samples and mst and Vst the mass and

volume in the standard solution. For example, for the first value which is 0,29 g, the total error is:

Once calculated, here is a table showing all the errors for mass:



0.29 0.04 g 0.26 0.04 g 0.24 0.03 g 0.23 0.03 g 0.20 0.03 g


0.25 0.03 g 0.23 0.04 g 0.22 0.03 g 0.20 0.05 g 0.18 0.03 g

Natural 0.24 0.04 g 0.23 0.03 g 0.21 0.03 g 0.21 0.03 g 0.21 0.03 g

Figure 9: Table showing the total error of the mass


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To calculate the error in the concentration, we simply add the relative error of the recently calculated mass and we add it to the relative error of the volume of the sample, which is 0.2 mL, as we are dividing mass per volume.

This is the final table after the calculations:



14.5 0.1g/L 13.0 0.2 g/L 12.0 0.1 g/L 11.5 0.1 g/L 10.0 0.2 g/L


12.5 0.1 g/L 11.5 0.2 g/L 11.0 0.2 g/L 10.0 0.3 g/L 9.0 0.2 g/L

Natural 12.0 0.2 g/L 11.5 0.1 g/L 10.5 0.2 g/L 10.5 0.2 g/L 10.5 0.2 g/L

Figure 10: Table showing the total error of the concentrations As was previously mentioned, molecular mass does not have errors, therefore, our absolute

error will be the result of multiplying the error of the concentration (this is our relative error) by the molarity data. Our final data, concentration in moles per litre (moles divided by the volume of the sample):



0.080 0.008 mol/L

0.08 0.02 mol/L

0.070 0.007mol/L

0.065 0.007 mol/L

0.06 0.01 mol/L


0.070 0.007 mol/L

0.07 0.01 mol/L

0.06 0.01 mol/L

0.06 0.02 mol/L

0.05 0.01 mol/L

Natural 0.07 0.01

mol/L 0.065 0.007

mol/L 0.06 0.01

mol/L 0.06 0.01

mol/L 0.06 0.01

mol/L

Figure 11: Table showing the concentration of vitamin C in moles/L

4.3. Reaction rate constant calculation

Before calculating the reaction rate constant, we have adjusted the data to the zero, first and second-order kinetic models in order to select the one that best fits the data.

1) Zero-order kinetic model

Figure 12: Graph showing the change in concentration with time

R = 0.9535

R = 0.958

R = 0.8179

0 0.01 0.02 0.03 0.04 0.05 0.06 0.07 0.08 0.09

0.1

0 2 4 6 8 10 Vit

amin

C c

on

cen

trat

ion

(m

ol/

L)

Time(days)

NON REFRIGERATED Don Simn

REFRIGERATED Pascual

NATURAL

Change in the concentration of vitamin C through time


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As we can observe by looking at the graph, the concentration of vitamin C decreases with time. While non-refrigerated juice (the blue series) or the refrigerated juice (the red series) both follow this trend, we can find an interesting feature, however, when we talk about natural juice. From Day 6 until Day 10, concentration is kept constant. We may relate this property later on with the fact of it being 100% natural juice unlike the other chemically-treated juices.

2) First-order kinetic model

Figure 13: Graph showing the change of the natural log in time

In this graph, we can observe a similar tendency as in the previous one, but in general terms,

the data is irregular and does not follow a linear model as well as the previous graph. The natural juice series (the green one) however, follows the same exact path and the three

last dots are constant, while the other two series decrease with time. 3) Second-order kinetic model

Figure 14: Graph showing the change of the inverse of the concentration in time

R = 0.9278

R = 0.9459

R = 0.8236

-3.1

-3

-2.9

-2.8

-2.7

-2.6

-2.5

-2.4

0 2 4 6 8 10

Nat

ura

l lo

g o

f vi

tam

in C

co

nce

ntr

atio

n

Time(days)



NATURAL

Change in natural log of the concentration of vitamin C through time

R = 0.8966

R = 0.9302

R = 0.8287

0

5

10

15

20

25

0 2 4 6 8 10 Inve

rsa

de

la c

on

cen

trac

in

(1

/co

nc.

)

Tiempo (das)



NATURAL

Variacin del inverso de la concentracin de Vitamina C a travs del tiempo


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In this graph, we observe three increasing series. In this particular case, the green series is

following the same trend, as the three last dots are constant. Both of the other two series have a positive slope. The inverse of the concentration increases with time as the quantity of juice decreases in time.

To obtain our reaction rate constant, we had to observe the graphs and choose the best

ones based on the coefficient of determination, R2, which is a measure of how well the dots of a graph fit in the trend line of the graph. Once having determined the best graph, the reaction rate constant for each juice will be the m coefficient (the slope) of the lineal function which represents the trend line.

The best graph was the first one, which corresponds to a zero-order kinetic model, as it has

two of the three highest R2 coefficients. To obtain the function of the trend line we can use Excel, select the trend line and click show equation in chart. This is the final data:

R2 kv

Non-refrigerated 0.9535 (3.00.3)10-8 Refrigerated 0.958 (2.40.3)10-8

Natural 0.8179 (1.30.3)10-8

Figure 15: Table showing the final calculation of the constant kv

5. Discussion

5.1. Conclusions

Ascorbic acid in the different juices decrease with time showing a zero-order decomposition kinetic model in all the types of juice studied. The highest reaction rate constant was achieved by Don Simn non-refrigerated juice while the lowest was achieved by the natural juice. In fact, when considering the error, the rate constants for the two industrial juices are the same and significantly differ from the one for the natural juice. This means the vitamin C in natural juice will last longer than in non-refrigerated juice and in refrigerated juice. If we observe the first graph showing the change of concentration in time we find out that natural juice does not contain the greatest amount of vitamin C, but non-refrigerated juice does. However, this non-refrigerated juice loses vitamin C faster than the natural juice. This can be due to chemically-added vitamin C during the production process, which may give the juice more vitamins, but only for a limited amount of time, as it decomposes faster than the other species. On the other hand, refrigerated juice shows similar a concentration of vitamin C than natural juice until Day 3. From Day 3, the constant rate of natural juice seems to decrease whereas the reaction rate in refrigerated juice is kept unvaried. We can deduce from these observations that refrigerated juice does not decompose as quickly as non-refrigerated juice, maybe due to the presence of chemicals to preserve the vitamin C whilst kept in the fridge.


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5.2. Evaluation and Improvements

To ensure the quality of my results, I compared them with similar studies carried out by doctors or students who allowed their results to be published. According to the results obtained by Drs. Burdurlu, Koca and Karadeniz, the decrease in the amount of citrus juice concentrates at all storage temperatures described a first-order kinetic model. Along the same lines, the article published by the EJEAFChe (Electronic Journal of Environmental, Agricultural and Food Chemistry) featuring the study by Drs. Abassi and Niakousari, described the loss of ascorbic acid in the natural juice at all storage temperatures as a first-order kinetic model while the concentrate followed a second-order kinetic model. These two papers show similar results however, they differ from mine. This may be due to the short period of time in which the measurements were taken (10 days), while Drs. Burdurlu, Koca and Karadeniz kept the juice stored for a period of 8 weeks and Drs. Abassi and Niakousari stored juice for a period of 12 weeks. It is possible that my results may be influenced by the fact that the time lapse I chose for the study was short in comparison with the time taken by the doctors in their studies. This may be a source of improvement for my study. Another way to improve my study would be to take more measurements of the titrations in order to avoid possible random errors.


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6. References

Bibliographical References: Abbasi, A. and Niakousari, M. (2008): Kinetics of Ascorbic Acid Degradation in Un-Pasteurized Iranian Lemon Juice during Regular Storage Conditions, Found in Pakistan Journal of Biological

Sciences, volume 11, Issue 10, January 2008(p. 13651369) (Print). Burdurlu, Hande S.; Koca, Nuray and Karadeniz, Feryal. (2006): Degradation of Vitamin C in Citrus Juice Concentrates during Storage, Found in Journal of Food Engineering, volume 74, Issue 2, May

2006 (p. 211216) (Print). Davies, Michael B.; Austin, John and Partridge, David A. (1991): Vitamin C: Its Chemistry and Biochemistry, Cambridge: The Royal Society of Chemistry Roberts, John D. and Caserio, Marjorie C. (1977), Basic Principles of Organic Chemistry, Menlo Park, CA: W.A. Benjamin Electric. Inc. Internet-based bibliography: Galiano Ramos, lvaro, VITAMINA C, IQB: MEDCICLOPEDIA, (in Spanish) Web, 03 Jan, 2012. . Helmenstine, Anne Marie, Vitamin C Determination by Iodine Titration, About.com Chemistry,The New York Times Company.Web. 5 July 2011. . Licata, Marcela, Vitamina C - cido Ascrbico. Zonadiet.com - Nutricin. Alimentacin. Salud. Deportes Y Vida Sana (in Spanish). Web. 30 Dec. 2011. .

study on the kinetics of degradation of ascorbic acid from different orange juices' (ib extended...

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