Measuring the Concentration of Ozone in Solution, with Time

Table Contents

Abstract

1.0 Aims and objectives of the experiment

2.1 Theoretical and experimental basis:

2.2 Experimental Procedure:

2.2.1 Chemicals and Equipment

2.2.2 Iodometry Process

3.0 Table of Results

3.1 Titration values with calculations:

3.2 Calculated values

3.3 Graph Plot of mass Concentration of Ozone vs Time

3.4 Graph Plot of Molar Concentration of Ozone vs Time

3.5 Graph of Volume of Thiosulphate against Iodine concentration

4.0 Precautions Taken

5.0 Conclusion and Recommendations

Bibliography

Abstract

This was an analysis using iodometry to estimate the amount of ozone bubbled from an ozonator over time. 7 samples, varied by time were bubbled into a 50ml sample of 0.084336 M concentration of KI solution. The product was titrated with a 0.04074 M concentration of thiosulphate solution. Calculations for the mass of ozone were carried and the values also tabulated accordingly, using an excel sheet. Graphs were plotted for concentration values of ozone liberated over time and mass of iodine liberated against the volume of thiosulphate used.

1.0 Aims and objectives of the experiment

To employ the use of iodometry, in estimating the mount of ozone bubbled into solution over time.

  To be able to tell how much ozone remains in solution and also for how long. During this process, it is expedient that dosing be systematically done, so as to optimise cost and quality control.

To use this experiment as a practice session for the use of ozone in disinfecting wastewater effluent.

In the application of ozone as a disinfectant in wastewater treatment, it is very important to measure, the amount absorbed and unused by the system. They are various methods for doing obtaining the amount of ozone. These methods include the indigo, ampero-metric-membrane, stripping, gas phase detection, UV (for deionized high-purity water), ampero-metric bare-electrode and iodometry methods(Bollyky 2003, p. 1541). The last mentioned method, is the used in this experiment. It is said to be effective when ozone concentration is between 1 g/m³ to 200 g/m³   at 0⁰C or 273.15 K and 1.01325 x 105 Pa or 1 ATM((IOA) 1996).

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Iodometry is for oxidants like ozone and in this case, is carried out in the gas and/or liquid phase, with the bubbling of ozone into the water sample and KI solution(Gottschalk, Libra & Saupe 2009). During this process, ozone oxidises the iodide ion I^–, liberating I₂, which is titrated with thiosulphate. At turning to a pale yellow colour, starch is used as an indicator to increase the intensity of the end-point. The ozone concentration is then estimated as an equivalent value to the volume of thiosulphate titrated from the burette. This method is common mostly because it is cheap and affordable(Gottschalk, Libra & Saupe 2009).

2.1 Theoretical and experimental basis:

2KI + O₃ + H₂O        I₂+O₂+2KOH ………………. equation (1) (In flask)

I₂ + 2S₂O₃^2-           3I^– + S₄O₆^2- ………………….…… equation (2)

To calculate for required concentration of thiosulphate, we relate equivalent mole ratios as follows:

Therefore 1 mole of ozone would require 2 moles of thiosulphate.

1 m-mole of O₃ = $\frac{1}{2}$

C ^S₂^O₃^2-V^S₂^O₃^2-

Molar weight of O₃ is 48 and that of Na₂S₂O₃ is 158

∴1 m-mole O₃ $=\frac{\mathit{weight\; of\; ozone}}{48\frac{\mathit{mg}}{\mathit{mmole}}}$

$\frac{\mathit{weight\; of\; ozone}}{48\frac{\mathit{mg}}{\mathit{mmole}}}$

=  $\frac{1}{2}$

C ^S₂^O₃^2-V^S₂^O₃^2- ………………… equation (3)

weight of ozone = 0.5×48 $\frac{\mathit{mg}}{\mathit{mmole}}$

x C ^S₂^O₃²–x V^S₂^O₃^2- ……………………………..equation (4)

The above formula is used in obtaining the mass of ozone liberated. In similar fashion, mass of iodine is also calculated.

To prepare  a concentration of C S₂O₃^2-=0.04074 $\frac{\mathit{mmole}}{\mathit{ml}}$

, mass of anhydrous Na₂S₂O₃ required is calculated as:

1000mL×0.04074 $\frac{\mathit{mmole}}{\mathit{ml}}$

× $158\frac{\mathit{mg}}{\mathit{mmole}}$

= 6440 mg

Concentration of potassium iodide prepared= $\frac{\mathit{Mass}}{\mathit{ vol }}$

                                                                            = $\frac{14}{1000}$

                                                                            = 0.014 g/mL

= $\frac{\mathit{Mass}}{\mathrm{ }\mathit{mw}\mathrm{ }x\mathrm{ }\mathit{vol}\mathrm{ }}$

Molar Concentration is: $\frac{14000}{\mathrm{ }166.0028\mathrm{ }x\mathrm{ }1000\mathrm{ }}$

= 0.084M

2.2 Experimental Procedure:

2.2.1 Chemicals and Equipment

1. Anhydrous Sodium thiosulphate powder (Na₂S₂O₃)
2. Potassium Iodide (KI)
3. Potassium Iodate (KIO₃)
4. Sodium Carbonate (Na₂CO₃)
5. Sulphuric Acid (H₂SO₄)
6. Starch indicator
7. Deionised water
8. Ozonator machine
9. Burette
10. Pipettes
11. Erlenmeyer Flasks
12. Electronic balance
13. Funnel

2.2.2 Iodometry Process

1. A 0.04074 Na₂S₂O₃ stock (from the standardisation of a 0.04048 sample) was prepared by weighing 3.2 grams of anhydrous thiosulphate and topping up with deionised water in a 500ml flask. This is the titrant.
2. An acidified KI solution was prepared by weighing 7grams KI + 20 mL glacial acetic acid into a 500mL flask half-filled with deionised water. This was topped up to the line with more deionised water.
3. A 50ml aliquot of KI solution had ozone bubbled in it for 0.5, 1, 1.5, 2, 2.5, 3 and 3.5 minutes.
4. Each sample turned into a dark brown colour. This was titrated with thiosulphate.
5. On becoming pale yellow, few drops of starch indicator were added, this gave a dark violet colour to the solution.
6. Continue adding drops of thiosulphate, until solution turns colourless.
7. Read volume of thiosulphate used from burette.
8. Recalling that each sample was bubbled into 0.05L of KI solution, we obtain concentration values for ozone.

3.0 Table of Results

 3.1 Titration values with calculations:

Recall that 1mg of O₃≈10ml of thiosulphate, to obtain concentrations for ozone, we divide each mass by 0.05L, since this was the volume of the solution.

Molar Concentration of KI= 0.08434 M

Molar Concentration of thiosulphate = 0.04074 M

Weight of ozone, obtained from equation (4) = 0.5×48 mg/mmole x 0.04074 x V S2O32

weight of iodine  = 0.5×253.81 $\frac{\mathit{mg}}{\mathit{mmole}}$

x 0.04074 x V^S₂^O₃² ……. Equation (5)

3.2 Calculated values

Table ;++: Table of values

3.3 Graph Plot of mass Concentration of Ozone vs Time

Figure 1: Calibration graph with line of regression inserted.

From the equation in the graph above, I can calculate for concentration of ozone per time as:

Y = 9.714X + 1.95 with an R² value of 0.9921, which is not too bad for an experiment carried out under non-standard conditions.

To calculate for concentration of ozone, we substitute time as X.

3.4 Graph Plot of Molar Concentration of Ozone vs Time

3.5 Graph of Volume of Thiosulphate against Iodine concentration

Figure 2: Volume of Thiosulphate vs Mass of iodine

4.0         Precautions Taken

1. Appropriate laboratory wear was used, especially in the handling of ozone because of its toxicity.
2. Efforts were made to avoid the use of materials that might decompose in contact with ozone.
3. The KI crystals had to be properly grinded, to ensure a sharper endpoint detection. Otherwise, purple crystals were noticed in solution, while swirling. This made it difficult to obtain defined endpoints.
4. Tetra-Oxo sulphate six acid (10%) was used in the KI solution as a buffer, to ensure a pH condition of 2.
5. Sodium bi-carbonate was used to ensure that the pH of the solution was above 7, to prevent the dissociation of thiosulphate into sulphur-dioxide.
6. Rough runs were done, to ensure familiarity with procedure.
7. All chemicals used are toxic in nature. Care was taken to avoid contact with skin and also prevent a spill.
8. Appropriate tools were used to ensure quantities were measured accurately.
9. In order to reduce the amount of reagents, small quantities, appropriate for accuracy were used.
10. Time or duration of the experiment, was within a practicable range, to cater for ozone’s reaction profile in water. This is known to be unstable and also of a short span, especially with temperature increases.

5.0 Conclusion and Recommendations

The experiment was conducted under very non-standard conditions. There was a steady rise in the concentration of ozone. The temperature, pH and all other factors remained unchanged. Concentration for ozone were 8 mg/L and 36mg/L at 0.5 and 3.5 minutes respectively. Time range was left as it is to minimise the possibility of a reduction in ozone’s solubility and stability over time.

An essential recommendation, would be the performance of this experiment in an enclosed column, like a gas wash bottle, rather than the open flasks used. It would also be very paramount to recommend that ozone is applied at lower temperature, to see if obtained values for ozone would increase as stated by literature(Urbano et al. 2017, p. 227).

The results showed a gradual corresponding increase and then eventual decline in the ozone concentration over time, in solution. This could be due to:

• This could most probably be due to the instability of ozone, being a triatomic form of oxygen (Gordon et al. 1989, p. 76). It would usually lose an atom.
• Another possible reason might be the trend of ozone’s solubility, which is known to decreases as the temperature increases(Choi & Nielsen 2005, p. 15).
• The formation of ^·OH by OH^–, and the consequent secondary oxidation by ozone in solution(Sonntag 2012, pp. 19,187).
• Ozone is known to possess a very short half-life in water and this is largely affected by temperature increases as already mentioned above, 3 days at 20°C and 1.5 seconds at 250°C(Batakliev et al. 2014, p. 48).
• Very probably without air-tight coverage, the ozone could have eventually left the solution with the length of time and at uncontrolled conditions like pH and temperature(Peleg 1976, p. 362).
• Given that a 0.4mg O₃/L of residual ozone is required for efficient disinfection of water,

Bibliography

• (IOA), IOA 1996, Iodimetry method of ozone measurement – STANDARDISED PROCEDURE.
• Batakliev, T, Georgiev, V, Anachkov, M, Rakovsky, S & Zaikov, GE 2014, ‘Ozone decomposition’, Interdiscip Toxicolology, vol. 7, no. 2, pp. 47-59.
• Bollyky, LJ 2003, ‘Measuring Ozone in Water.pdf’, in BG Liptak (ed.), INSTRUMENT ENGINEERS’ HANDBOOK: Process Measurement and Analysis, Fourth Edition edn, CRC Press, vol. 1.
• Choi, LH & Nielsen, SS 2005, ‘THE EFFECTS OF THERMAL AND NONTHERMAL PROCESSING METHODS ON APPLE CIDER QUALITY AND CONSUMER ACCEPTABILITY’, Journal of Food Quality, vol. 28, no. 1, pp. 13-29.
• Gordon, G, Rakness, K, Vornehm, D & Wood, D 1989, ‘Limitations of the Iodometric Determination of Ozone’, Journal American Water Works Association, vol. 81, no. 6, pp. 72-6.
• Gottschalk, C, Libra, JA & Saupe, A 2009, Ozonation of Water and Waste Water: A Practical Guide to Understanding Ozone and its Applications, Technology & Engineering, John Wiley & Sons.
• Peleg, M 1976, ‘THE CHEMISTRY OF OZONE IN THE TREATMENT OF WATER’, Pergamom, vol. 10, pp. 361-5.
• Sonntag, Cv 2012, Chemistry of ozone in water and wastewater treatment : from basic principles to applications, London ; New York : IWA Pub., London ; New York.
• Urbano, VR, Maniero, MG, Perez-Moya, M & Guimaraes, JR 2017, ‘Influence of pH and ozone dose on sulfaquinoxaline ozonation’, J Environ Manage, vol. 195, no. Pt 2, pp. 224-31.