Dissolved oxygen (microscopic bubbles of oxygen gas in aqueous solution, DO) is essential for the healthy functioning of all freshwater ecosystems. If more oxygen is consumed than is produced, DO levels decline and some sensitive animals may move away, weaken and even die. Therefore, DO presence in an ecosystem is a positive sign whereas low DO levels are an indication of severe pollution.
Gases are usually more soluble at colder temperatures. For example, oxygen is more soluble in cold water than in hot water. The decrease in oxygen solubility with increased temperature has serious consequences for aquatic life. Power plants that discharge hot water into rivers can kill fish by decreasing the dissolved oxygen concentration. With global industrialization, the conditions of the world have been altered drastically in a short time. Pollutants have been introduced, rivers have been diverted, and forests have been cut down. In general, the higher the DO levels, the more stable the freshwater ecosystem and thus the greater the capacity for this ecosystem to sustain life. Much of the DO in water comes from the atmosphere. Streams with a high kinetic energy and a tumbling water action promote the mixing of atmosphere oxygen, a non-polar compound, with water, a polar compound.
A high DO level in a community water supply is good because it makes drinking water taste better. However, high DO levels speed up corrosion in water pipes. For this reason, industries use water with the least possible amount of dissolved oxygen.
DO levels in streams fluctuate significantly during the day, especially if the freshwater ecosystem supports extensive plant life. DO levels are at their lowest during the early morning, tend to rise during the day and peak in the afternoon. Temperature has a significant influence on DO levels. The concentration of DO in natural water and wastewater is a function of the temperature of the air and water, the degree of hardness of the water, and the demand for oxygen in the body of water. The solubility of oxygen increases with decreasing water temperature (oxygen solubility in water is inversely proportional to temperature).
If you were to put the presence of DO into an everyday layman context then it would prove useful to observe a pot of water being heated. One can notice that bubbles form on the walls of the pot prior to reaching the boiling point. These cannot be filled with only water vapour because liquid water will not begin to vaporize until it has reached its boiling point. One can assume that this gas is oxygen, or at least a mixture of gases from the air, because bubbles of this sort form in water from virtually every source. When these bubbles form, they eventually grow to a sufficient size to leave the surface of the pot and escape to the air and therefore the dissolved gas in the liquid has decreased. This seems to support my hypothesis that dissolved oxygen will decrease when temperature is increased.
When testing for dissolved oxygen, the concerns for safety involve water hazards and exposure to chemicals.
Precautions to follow include:
1. Cover all abrasions and if possible, wear good quality latex gloves.
2. Wash hands frequently and always wash hands prior to eating.
3. Wear a protective smock, apron or lab coat, and surgical or rubber gloves when working in the laboratory to protect clothes and skin.
4. Read all labels carefully and know what to do in case of a spill.
The WINKLER TEST is used to determine the concentration of dissolved oxygen in water samples. An excess of manganese (II) salt, iodide (I-) and hydroxide (OH-) ions is added to a water sample causing a white precipitate of Mn(OH)2 to form. This precipitate is then oxidized by the dissolved oxygen in the water sample into a brown manganese precipitate. In the next step, sulphuric acid is added to acidify the solution resulting in Mn(SO4)2, (a brown precipitate). The Mn(SO4)2 then converts the iodide ion (I-) to iodine (I2). Under controlled conditions, the amount of sodium thiosulfate titrated, is equivalent to the amount of dissolved oxygen present in the sample.
1. Manganese sulfate solution *
2. Alkaline potassium iodidesodium solution *
3. Sulfuric acid (H2SO4), concentrated *
4. Starch indicator solution
5. Sodium thiosulfate (Na2S2O3 5H2O), 0.025 N *
6. Potassium iodate (KH(IO3)2), 0.025 N
*These reagents are poisonous and/or corrosive and should be handled with caution.
1. 7 conical flasks
2. Large bucket
4. Water bath
6. 3 Droppers
8. Burette holder
10. Safety glasses
1. Put your safety glasses on and secure your burette into your burette holder.
2. Put a spare beaker underneath the burette and run a small amount of the thiosulfate solution through the burette to clean it. Make sure all the thiosulfate has run out.
3. Close the cockstop, fill the burette with the thiosulfate solution, take note of where your start volume is, and take the funnel out.
4. Fill a clean, large bucket with tap water. Run the water at full pressure before starting to fill the bucket. Fill the bucket and do not change the pressure for the duration of the fill.
5. Fill 7 conical flasks with 200ml of water from the bucket.
6. Place one flask with a thermometer in it into a water bath and allow to heat until it reaches 35oC. (Depending on the temperature of the water straight from the tap, you will need to fill a secondary bucket and place ice into the water, creating a second water bath for the lower temperatures, ie. 5oC and 10oC. Place your conical flask and thermometer into the water bath and continue adding ice until you reach your desired temperature.)
7. Take the flask out of the water bath and immediately insert stopper.
8. Using a dropper inserted well below the surface, immediately add 1ml manganese (II) sulfate solution followed by 1ml of alkali-iodide solution to the heat treated water. Be sure to use separate droppers for each solution so you don’t contaminate anything.
9. Immediately re-stopper to reduce mixing with air.
10. Invert the flask a few times. The manganese (II) will be oxidized to form a brown precipitate. Allow the precipitate to settle to half the volume of water.
11. Using a separate dropper, CAREFULLY add about 1ml of concentrated sulfuric acid will below the surface. Re-stopper and mix thoroughly.
12. Pour the sample into a 500ml wide-mouth flask to be titrated.
13. Titrate the water sample against the thiosulfate solution to a pale straw colour.
14. Add a few drops of starch indicator and continue titrating to the first disappearance of the blue colour.
15. Record the volume of thiosulfate used.
16. Re-fill the burette and note the starting volume.
17. Repeat steps 6-16 twice more to obtain an average titre.
18. Repeat steps 6-17 for your other temperatures.
19. Repeat the entire experiment again to make sure your results are accurate. Be sure to note if you have made any changes to the method and reasons why.
RESULTS FOR FIRST TITRATION OF THIOSULFATE AGAINST HEAT-TREATED WATER SAMPLE
AVG dissolved oxygen concentration
The graph in figure 1 presents the results of the experiments comparing temperature to dissolved oxygen. Each blue diamond represents the average result of three tests performed for each temperature.
RESULTS FOR SECOND TITRATION OF THIOSUFATE AGAINST HEAT-TREATED WATER SAMPLE
AVG dissolved oxygen concentration
The graph in figure 2 presents the results of the experiments comparing temperature to dissolved oxygen. Each blue diamond represents the average result of three tests performed for each temperature.
Set 1: 5oC
c(S2O32-)= 0.250 mol/L
v(S2O32-)= 25.4 ml or 0.0254 L
n(S2O32-)= c x v
= 0.000635 mol
n(O2)= 1/4 n(S2O32-)
= 0.25 x 0.000635 mol
= 0.00015875 mol
v(O2)= 0.20 L
= 0.00015875 mol/0.2 L
= 0.00079375 mol/L
c(g/L)= c(O2) x M(O2)
c(g/L)= c(O2) x 32 g/mol
= 0.02254 g/mol
c(mg/L)=0.02254 g/mol x 1000
= 25.40 mg/L
In general, the trend was as hypothesized. The dissolved oxygen concentration decreased with increasing temperature. The numerical trend is a downward one, as indicated by the linear equation in the above graphs.
ACCURACY AND VALIDITY:
The Winkler method is a rather difficult procedure to complete with the least amount of error. Therefore if the experimenter lacks experience it is possible to collect incorrect results. However in spite of this, the results from both trails were very similar, showing a downward linear trend. Taking extra care during the heating process and prior to the “fixing” agents being added will help to keep your sample from mixing with air. Also performing a ‘rough’ trial for each temperature first, is ideal because it gives the experimenter a measurement of how much thiosulfate solution was needed at each temperature. This will make it easy to first run a large amount and then finish off by closing the stopcock and dripping the next few millilitres in, when the volume of thiosulfate was close to the rough value.
Furthermore, not only did I perform a rough trial, I also duplicated my entire experiment, keeping the variables the same throughout. This allowed me to check the accuracy of my first set of results. The second experiment provided me with concordant results and ensured the accuracy of my procedure and samples.
This downward trend entirely supports my hypothesis and background research and therefore, one can regard my results as accurate. Regarding my specific trend line, the data in figure 1 show that 97.74% of the variance in X (Temperature) can be explained by variation in Y (Amount of DO) . Likewise the data in figure 2 show that 98.06% of the variance in X (Temperature) can be explained by variation in Y (DO).
However regarding the validity, the specific linear equations (Figure 1: -0.2457x + 26.914, Figure 2: -0.2457x + 26.871), will only be valid if the same variables were applied, ie., same temperatures, same thermometer, same solutions. Also, if the experimenter were to gather the water sample from a different tap, they would have differing results (although, from researching DO, the results should create a similar downward trend). Therefore, even though these results seem to be accurate, they are only relative to the specific tap I gathered my sample from and under the same conditions I applied. The only applicable piece of information that could be drawn from my results, is that there is a definite trend; the higher the temperature of the water, the less dissolved oxygen it contains.
Very slight variations between the first and second set of data collected, could be due to the different time of day that the procedure was conducted. The ‘tap temperature’ for the first set of data collected was exactly 10oC. Whereas, the ‘tap temperature’ for the second set of data, (the next day), was slightly less, at 8oC. This made it a little harder to test for 10oC in the second sample (because of the small interval between 8oC and 10oC), and caused me to heat it too much at first. I took another sample and tried to be more careful and therefore obtained concordant results; varying only 0.1 ml when averaged.
THE RELATIONSHIP BETWEEN THE BACKGROUND INFORMATION AND THE OUTCOME:
The hypothesis on the correlation between temperature and dissolved oxygen was supported by the data: the calculations showed that they were indeed inversely related, with dissolved oxygen falling as the temperature rose. Also, these results entirely supported the background research. The trend in each of the performed trials, (except for results at 25oC – see below, “Possible Sources of Error”), followed a similar pattern and showed similar relationships between temperature and dissolved oxygen. The general trend discovered here is completely universal: as a liquid is heated, gas dissolved in it will evaporate and leave the liquid.
POSSIBLE SOURCES OF ERROR:
All of the data collected in this experiment seems to agree with and support the hypothesis. There was a piece of data, at temperature 25oC, that fit the hypothesis but did not fit perfectly with the trend shown by my results. This piece of data from both sets of results is only at a 0.5ml interval whereas the rest of the data is at, at least 1.0ml intervals. It is likely that this came about due to human error and inexperience with the Winkler method. A result of this nature probably occurred because I overshot the amount of thiosulfate titrated. However, the sample may have been over exposed to the air or the temperature may have changed before the fixing procedure was finished. Especially considering that the rest of the data indicates that the hypothesis is correct and follows a specific trend, it seems that this one anomaly is the cause of some sort of error in the execution of the Winkler method. To prevent, or at least remedy this sort of problem, the experimenters should, in the future, test each sample multiple times to determine if the results are consistent or if the results are merely inaccurate.
ADVICE GIVEN FOR FUTURE EXPERIMENTS:
Were the experiment to be performed again, the experimenter should calculate the concentration of dissolved oxygen multiple times to make sure that consistent and precise data is found, but more importantly, all of the water samples would be taken from the same source. In my case, I used the same tap, same thermometers and the same bottles of reagents. Also, other than repeating the entire procedure, another method of verifying the dissolved oxygen concentration should be used, like a “Hach testing kit” or similar apparatus.
Further investigations could involve the effects of other variables on dissolved oxygen, such as the presence of pollutants or how certain aquatic plants contribute to the dissolved oxygen level of water. It would be interesting to research whether or not, all tap water is safe for drinking. Factors that you could prove that it is unsafe, could be heavy industrial and environmental pollution. Toxic bacteria, chemicals and heavy metals can penetrate and pollute natural water sources making people sick while exposing them to long term health consequences such as liver damage, cancer and other serious conditions. The world has reached the point where all sources of drinking water, including municipal water systems, wells, lakes, rivers, and even glaciers, contain some level of contamination. Even some brands of bottled water have been found to contain high levels of contaminants in addition to plastics chemical leaching from the bottle.
Also, you could follow this experiment by conducting the same procedure and just change a few variables. Such as: “Does seawater hold as much dissolved oxygen as freshwater at the same temperature?” If the ocean is too far away, make your own saltwater by adding between 30 and 35 g of table salt for each litre of water. Use a water bath, (same one you used for this experiment), for cooling the saltwater sample down without diluting it.
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