Water is essential for the support of all life forms and is a vital requirement in a technical civilization. The chemical reactions that sustain life all take place in aqueous solution, with water itself frequently involved as a reactant, as in the process of photosynthesis by which green plants convert carbon dioxide and water into complex sugars, liberating oxygen as a by-product:
6 CO2(g) + 6 H2O(l) C6H12O6(aq) + 6 O2(g)
Equally important are the processes in the lung whereby oxygen is transferred to hemoglobin for transport throughout the circulatory system, where water plays a vital role--dehydration of the lungs is a fatal condition.
Besides the normal requirements of each individual for drinking water to maintain the 70% water content of the body, agricultural and industrial processes require enormous amounts of water. Thus adequate rainfall or irrigation is needed for crop production; 600 kg of water are used in the preparation of 1 kg of fertilizer, 300 kg in the production of 1 kg of steel and 200 kg in the production of a single copy of this laboratory manual from the original tree. A plentiful and reliable supply of water is thus essential to civilization, and the development of all societies has been closely linked to availability and control of water supply.
Although about 80% of the planet's surface is covered with water, surprisingly little of the total estimated mass of water, some 1.3 x 1024 kg or about 5% of the mass of the earth, is available for ready use. Over 97% of that total, for example, is in the seas and oceans as a dilute saline solution not directly useful either for drinking water or for agricultural purposes. This huge volume of water, however, plays an important physical role, acting as a heat sink to absorb most of the tremendous amount of solar energy that falls on the earth daily. Liberation of this energy by radiation during hours of darkness helps to stabilize the surface temperature of the earth. Without this water buffer the earth's surface temperature would fluctuate between -150oC and +120oC every 24 hours.
The ability of water to sustain life is closely linked to its ability to dissolve molecular oxygen from the atmosphere. Dissolved oxygen also results from photosynthetic processes taking place in aquatic plants in the ocean. Over three-quarters of all the oxygen produced by photosynthesis is in the oceans. It is this aspect of the ecocycle, more than any other, that prompts concern over the effects of oceanic pollution. By killing the plant life of the sea we cut off a major source of the oxygen needed to sustain human life. The problem is a very real one and demands immediate action. We have traditionally used the oceans as dumping grounds for many of our wastes and the process continues at an alarming rate; there is already a very real prospect of the Mediterranean becoming a dead sea in our lifetime.
The minimum concentration of dissolved oxygen needed to support life is only 4 mg/L or four parts of dissolved oxygen per million parts of water (4 ppm). The equilibrium concentration of dissolved oxygen in water, although it varies with temperature, is normally higher than that: 14.6 ppm at 0oC, 11.3 ppm at 10 C, 9.2 ppm at 20oC and 7.6 ppm at 30oC. It might be thought, therefore, that a sufficiency of dissolved oxygen would always be present. Unhappily, the process of solution is a very slow one and is ultimately dependent on the surface-to-volume ratio of any body of water.
A raging glacial torrent in a mountain stream, where water is constantly being broken up and tossed into the air will quickly and naturally become saturated with oxygen; a large, still deep lake may only achieve total saturation over centuries, if at all. In public drinking water the dissolved oxygen content typically averages 4 ppm or more. The desirable level is somewhere between 8-10 ppm.
In this experiment we will quantitatively determine the dissolved oxygen content of your water sample by Winkler's method. The sample is first treated with excess manganese(II) sulfate solution and then with an alkaline solution of potassium iodide. The following reactions are important:
Mn2+(aq) + 2 OH-(aq) Mn(OH)2(s)
4 Mn(OH)2(s) + O2(aq) + 2 H2O 4 Mn(OH)3(s)
The Mn(OH)2 initially formed reacts with the dissolved oxygen. The process is a heterogeneous reaction, involving the combination of a gas with a colloidal solid, and is slow.
The amount of Mn(OH)3 formed is determined by reaction with iodide ion, which is inert in basic solution, but in acidic solution reacts with Mn(OH)3 to form Mn2+ and iodine. Therefore the sample is acidified with sulfuric acid and the following reactions take place:
Mn(OH)2(s) + 2 H+(aq) Mn2+(aq) + 2 H2O
2 Mn(OH)3 + 2 I-(aq) + 6 H+(aq) 2 Mn2+(aq) + I2(aq) + 6 H2O
The iodine formed may be titrated against standard thiosulfate solution, using starch as an indicator, to determine the amount of iodine formed:
I2(aq) + 2 S2O32-(aq) 2 I-(aq) + S4O62-(aq)
From the stoichiometries of the various equations, the amount of dissolved oxygen can be readily calculated.
Natural water supplies contain varying amounts of dissolved cations and anions of various kinds reflecting the nature of the ground over which the water has passed or is stored. Small amounts of Na+ or K+ occur in fresh water and cause no serious problems. Trouble is caused in some applications by the presence of Ca2+, Mg2+, and Fe3+ ions because of their tendency to precipitate as insoluble salts, particularly carbonates, phosphates, and hydroxides. When this precipitation takes place in pipes or boilers, flow is often critically reduced and dangerous situations may arise.
More commonly these ions present a problem because they combine with the anionic component of many soaps and detergents to form insoluble precipitates-- the familiar scum in the sink or washing machine. Soaps are sodium or potassium salts of aliphatic fatty acids. These binary salts are soluble, but the ternary calcium or magnesium salts are not.
Water with a high concentration of Mg2+ and Ca2+ is referred to as hard water, and the degree of hardness is linked to the concentrations of these two cations. [typical public drinking water: 60-150 mg CaCO3/L] Water with a low concentration of these ions is referred to as being soft.
The hardness of your water sample will be determined by titration of the alkaline earth cations with ethylenediaminetetracetic acid (EDTA), a complexing agent. The titration is monitored by using another complexing agent, eriochrome black T (EBT), as an indicator. EDTA is a more powerful complexing agent than EBT and displaces EBT from its calcium complex. Ca(EBT) is a magenta shade whereas EBT itself is a deep blue color; Ca(EDTA) is colorless. As the titration proceeds, added EDTA removes Ca2+ from Ca(EBT), and when all the calcium has been attached to EDTA at the end point, only the blue color of EBT is present. The color change is progressive, with magenta giving way to blue, so that the end point may not be too obvious in an initial run. A small amount of Mg2+ ion is added to the standard EDTA solution to make this change sharper. Thus a blank titration is required.
EDTA is a powerful complexing agent for nearly all metal ions forming exclusively 1:1 complexes. Discrimination between different metal ions is achieved by careful control of the pH of the solution. Ca2+ and Mg2+ are bound in a tight complex only in strongly basic solution, optimally at pH 10. It is necessary to use an appropriate buffer solution to ensure the proper pH of the sample during the EDTA titration.
In the actual experiment you will standardize the stock solution of EDTA against a CaCO3 solution of known concentration (1.000 mg/mL ) and then use the EDTA to determine the unknown Ca2+ content of your sample.
A certain amount of chloride ion is almost always present in domestic water supplies as the result of chlorination procedures adopted in the course of purification. Chloride ion is an essential nutrient and one danger of the current antisalt fad is that chloride sufficient to sustain the stomach's need to synthesize hydrochloric acid may not be ingested. Excessive amounts of chloride are likely to be the result of industrial or agricultural pollution of the water. Hydrochloric acid is used extensively in the production of steel, and wastes are commonly discharged into many lakes and rivers. Chloride ion also results from discharge of wastes by paper and pulp mills which use chlorine as a bleaching agent.
The chloride ion content of the water sample can be determined by precipitation with Ag+ ions. The formation of a cloudy white precipitate of AgCl can be used to estimate the Cl- content. But a more accurate--and more interesting--method involves a titration using an electrochemical cell, known as a potentiometric titration. You will construct an electrode consisting of the half cell Cu|Cu2+. The electrode will be immersed in the water sample and standard silver nitrate solution added from a buret. During the titration, the voltage is monitored between the copper wire in the electrode and a silver wire in the water sample. The following reaction occurs in the complete cell as silver ion is added:
2 Ag+(aq) + Cu(s) 2 Ag(s) + Cu2+(aq)
As predicted by LeChâtelier's principle, the addition of silver ion propels the reaction forward and the changing concentration of silver and copper ions causes a change in the voltage (as the Nernst equation would show).
Initially the changes are small and gradual. But as the ion product of [Ag+] and [Cl-] approaches the Ksp of AgCl and the amount of free Cl- ion in the sample diminishes, addition of more
silver ion causes a sudden increase in the cell voltage, similar to the sudden change in pH at the end-point in an acid/base titration. A plot of voltage vs. mL of AgNO3 in fact yields a typical s-curve and the end-point can be determined graphically from a derivative plot. [typical public drinking water contains about 250 mg Cl-/L]
Natural water supplies in contact with air also absorb carbon dioxide, with production of the weak acid carbonic acid. This acid, found only in aqueous solution, interacts with water to form hydrated protons and hydrogen carbonate anions:
CO2(g) + H2O H2CO3(aq) H2CO3(aq) H+(aq) + HCO3-(aq) Ka = 4.3 x 10-7
Most natural water sources are therefore slightly acidic from this cause. [typical drinking water: pH 6.0-8.5] To sustain life water must remain fairly close to neutral in character.
A gradual control is being gained over the direct discharge of strongly acid or base waste effluents into rivers and lakes, though these controls are often evaded by the socially irresponsible or through accident. A serious problem is arising in certain parts of the world with the steady increase in so-called acid rain, which is leading to a gradual increase in the acidity of many natural waters and the destruction of many plant species in affected areas.
Ironically, this situation has arisen in part from attempts to solve the local pollution problems. The construction of extremely tall chimneys to discharge sulfur dioxide-containing gases, and the removal of alkaline fly ash from these gases has led to a cleaner Cleveland but a less healthy Maine.
You will test the pH of your water sample by a universal indicator paper, i.e., a paper designed to produce different colors for different pH throughout the range 1-14.
YOU WILL NEED ABOUT 1 LITRE OF WATER FOR THIS EXPERIMENT. BRING SOME FROM HOME OR YOUR FAVORITE LOCAL SUPPLY. USE THE 1 LITRE PLASTIC BOTTLE FROM YOUR LOCKER FOR THIS PURPOSE.
COLLECT YOUR SAMPLE IN SUCH A WAY AS TO MINIMIZE THE AMOUNT OF AIR CONTACT. IF YOU HAVE A FAUCET AERATOR, REMOVE IT FIRST BEFORE FILLING THE BOTTLE, AND THEN LET THE WATER COME OUT ONLY IN A SLOW STREAM, NOT A GUSHER.
PROCEDURE: Determination of acidity
Clean and dry a glass stirring rod and use it to remove a drop of water. Touch the drop of water to a 3 cm length of universal pH paper. Match the color produced to the pH scale provided and record your observation.
[always touch a drop of solution to a small piece of paper; NEVER dip the paper in the solution]
PROCEDURE: Determination of Cl- content
To construct the electrode, obtain a short length of glass tubing and a copper/silver wire pair. The bare copper wire is inserted into the glass tube and the silver wire remains outside (you can fasten the silver wire to the glass tube with a small rubber band if you like). Insert the wire as far into the tube as it will go and secure it with a piece of electrical tape.
At the other end of the tube, insert a threaded rubber septum cap (wetting the stopper end will help). You can thread the cap with a needle and the thread provided. A double strand of thread works well. The thread will function as the salt bridge in the cell.
Fold the collar of the septum cap over the body of the electrode to seal the end. The final assembly should look something like the diagram below:
[you can also inspect the sample electrode]
When you are ready to use the electrode, place enough 0.1 M CuSO4 inside the glass tube to immerse the bare copper wire. The wire should not be touching the thread.
Measure out 100 mL of your water sample with a graduated cylinder and pour it into a tall 200 or 250 mL beaker wrapped with black paper. Fill a buret with standard 0.004 M AgNO3 solution (note exact concentration on bottle!) and prepare to titrate the water sample. Place the electrode in the water sample but keep the water from entering the top of the tube and diluting the copper solution. Be sure the silver wire is actually in the water before you start. Attach the wire leads to a voltmeter (one of the wires has a white line on the insulation--this is connected to the silver wire), attaching the copper wire to the negative (-) terminal. Suspend your electrode in the water so that the threaded end is not sitting on the bottom of the beaker but is about half-way into the water sample.
Chloride ion contents vary. Therefore you can save some time by taking an initial voltage reading of your sample. It is not necessary to begin recording data until the voltage is around 0.13 volts. If your cell gives less than this, add silver nitrate from the buret in a slow stream (while stirring) until the cell voltage approaches 0.13 v. At that point record the volume of AgNO3 added and the voltage.
Now add AgNO3 in 0.5 mL increments, recording the voltage after each addition. You should notice a gradual increase, then a marked increase and then a gradual flattening out, just as in an acid/base titration, but perhaps not so dramatic. Be sure to record data at least 2 mL past the end-point.
Disassemble your electrode and clean the parts, returning them to the instructor. Be sure to clean the buret thoroughly.
PROCEDURE: Determination of dissolved oxygen
Take a conical funnel and attach its stem to a piece of rubber tubing about 25 cm in length. Introduce the tube into a clean, dry 500 mL flask and gently pour water from your sample into the container. Avoid any agitation of the system so as to minimize the amount of oxygen entering the water from the air during transfer.
Allow the flask to overflow, remove the tubing, and place a stopper in the container. Water should overflow as the stopper is inserted. DO NOT FORCE THE STOPPER INTO THE NECK OF THE FLASK!!!!! DO NOT PRESS YOUR WEIGHT DOWN ON THE FLASK TO SEAT THE STOPPER!!!! Such actions have led to serious cuts in the past as the flasks have given way under pressure. The stopper does not have to be seated tightly. Simply make sure no air is trapped between the stopper and the water. Hold the flask by the upper neck only as you insert the stopper.
Remove the stopper from the flask and gently remove 4 eyedroppers full of water (do not introduce air bubbles into the sample). With the same dropper, add 2 droppers full of MnSO4 solution by placing the tip of the dropper below the surface of the water and gently squeezing on the rubber bulb. DO NOT expel excess air into the sample.
Clean the dropper and use it to add 2 droppers full of KI/NaOH solution, following the same cautions. Top off the flask, if needed with a little more of your water and then re-insert the stopper, following the same cautions as before, taking care that no air is trapped.
Grasp the container firmly with a good grip on the stopper and invert it several times to ensure complete reaction. Allow the container to stand until the precipitate has settled. This would be a good time to set up the titration for this part of the experiment. Fill a buret with 0.005 M Na2S2O3 solution. You must complete the titration within 8 hours at this point.
Remove the stopper from your sample once the precipitate has settled and use an eyedropper to remove 2 full droppers of water. Now add, as before, 2 full droppers of concentrated sulfuric acid [CAUTION!]. Replace the stopper carefully---add a little more water if you must. Holding the stopper firmly in place, invert the flask several times, mixing the contents. The precipitate should dissolve within a few minutes. If not, add another dropper of H2SO4 and repeat. You must complete the titration within 45 minutes at this point.
Using a graduated cylinder, measure 200 mL of the acidified sample into a 500 mL flask. Titrate the sample until the color becomes a pale yellow. Add 5 mL of the starch solution to the flask. The deep blue color of the starch-iodine complex should develop immediately. Continue the titration until the blue color just disappears and record the volume of thiosulfate used. Repeat the titration with a second 200 mL sample only if problems are encountered with the first sample.
PROCEDURE: Determination of water hardness
Standardization of the EDTA solution
Pipet exactly 25 mL of the CaCO3 solution into a clean, dry 250 mL Erlenmeyer flask. Add 10 drops of EBT indicator and 7 mL of the buffer solution [CAUTION: the buffer contains a high concentration of ammonia]. Swirl the contents of the flask to mix the reactants.
Carefully rinse the buret used in the last step first with distilled water and then twice with small amounts of the EDTA solution. Be sure to allow the EDTA solution to drain through the buret tip. Ready the buret with EDTA and titrate the CaCO3 sample until the color changes from magenta to blue. This color change is very gradual; titrate to a distinct blue color and keep the flask and its contents as a reference for subsequent titrations with EBT in this step. Record the volume of EDTA required and repeat the titration for a second sample of CaCO3. The two titrations should agree to within 0.5 mL. Finally, titrate a "blank", using 25 mL of distilled water instead of the CaCO3 solution.
EDTA Titration of the Water Sample
Using a graduated cylinder, measure 100 mL of your water sample into a 250 mL Erlenmeyer flask. Add 40 drops EBT and 7 mL of buffer solution. Mix well and use a drop of the mixture to check that the pH is not less than 9. Add more buffer if needed to achieve a suitable pH. Titrate against EDTA to the same blue end point as before and record the volume of EDTA used. Carry out a
duplicate titration if time permits.
1. moles thiosulfate used in dissolved O2 titration [may be average or a single value]
2. equivalent moles of I2, Mn(OH)3, and O2
3. mass of O2 in mg
4. total dissolved oxygen content expressed as parts per million by mass (mg/L)
5. the volume of EDTA solution needed to titrate 1 mg CaCO3 [the standard solution is 1 mg CaCO3/mL]
6. the hardness of the water expressed as mg CaCO3/L
7. graph of your potentiometric titration data (voltage or E vs. mL AgNO3)
8. a table of E, V, E/V and Vav (average mL AgNO3) for your titration (near the endpoint region only)
9. graph of E/V vs. Vav (first derivative plot)
10. Cl- concentration for your water sample expressed in mg/L
You should include a summary table for all of the tests done as well as any explanations for repeat titrations with unusual results. Be sure to include the source of your water sample.
Why are 40 drops of EBT used in the titration of the water sample when only 10 drops were used in the standardization titration of the CaCO3? Be specific!