topic18

School Science Lessons
Topic 18 Environmental chemistry, pH tests, water hardness, air pollution, water pollution
Updated 2009-09-17
Please send comments to: J.Elfick@uq.edu.au
See: Interesting websites

18.1.0 pH tests
18.2.0 Total dissolved solids and suspended solids in water
18.3.0 Tests for air and dissolved oxygen in water
18.4.0 Ions in a water sample
18.5.0 Tests for standing water
18.6.0 Air pollution
18.7.0 Water pollution
18.7.0A Swimming pool chemistry

18.1.0 pH tests
18.1.0.1 Acidity and alkalinity
18.1.1 pH with universal indicator
18.1.2 pH of water in the laboratory
18.1.3 pH of rainwater
18.1.4 pH of standing water

12.13.0 Hardness in water
12.13.11 Tests for metal ions in water with EDTA, chelates
12.13.12 Measure water hardness using EDTA titration
16.4.4 | EDTA, ethylene diamine tetra acetic acid, (HOOC.CH₂)₂N(CH₂)₂N(CH₂.COOH)₂
18.2.0 Total dissolved solids and suspended solids in water
18.2.1 Insoluble solids in rainwater
18.2.2 Soluble solids in rainwater
18.2.2.1 Chlorides in groundwater
18.2.2.2 Iron in drinking water
18.2.2.3 Sulfates in groundwater
18.2.3 Extracted soluble solids from rainwater
18.2.4 Contamination of groundwater from refuse deposits
18.2.5 Salinity
18.2.6 Conductivity
18.2.7 Cations and anions in rain, rivers and seawater
35.22.7.1 Calcium carbonate dissolves in rain water

18.3.0 Tests for air and dissolved oxygen in water, dissolved oxygen, DO
2.25 Gases dissolved in a water sample
18.3.2 Dissolved oxygen in water (Winkler method)
18.3.3 Amount of dissolved oxygen, titration

18.4.0 Ions in a water sample
18.4.1 Phosphate ions in water
18.6.9 Colorimeter tests and test methods

18.5.0 Tests for standing water
18.5.1 Tests of river, lake or ocean
18.5.1a Eutrophication of waterways
18.5.2 Anions in sewage and tap water
18.5.3 Physicochemical test methods for water samples

18.6.0 Air pollution
18.6.0.1 Commonly occurring air pollutants, safe air and clean air
2.32.1 Composition of the atmosphere and greenhouse gasses
3.39 Carbon monoxide, CO
3.48 Acid rain, nitrogen oxides, (NOx) and acid rain,
12.6.0a Acid rain, Formation of acid rain (SOx) by burning sulfur or sulfur compounds
18.6.2 Air pollution from burning refuse
18.6.3 Danger of vehicle exhausts, tailpipe gases

18.7.0 Water pollution
5.41 Keep water clean (Primary)
18.6.1 Contamination of groundwater by refuse, hazardous wastes
18.6.4 Water tests
18.6.5 Smell of water, hydrogen sulfide
18.6.6 Colour of water
18.6.7 Hydrogen ion concentration of water
18.6.8 Temperature of water
7.2.2.33.2 Detergent phosphates

18.1.0 pH tests [H⁺] hydrogen ion concentration
See 12.10.7.0: Buffer solutions | See 7.7.0: Solutions, solubility, solubility rules| See 12.10.3.2: Hydrolysis of ammonium chloride
pH tests use an indicator which changes color with changes in the concentration of hydrogen ions, or the acidity of the solution. The pH value of the ocean changes very little when acids or alkalis are added or when diluted with water because oceans are buffered solutions. However, many rivers and lakes are weakly buffered so their pH may change rapidly if acids or bases are added. The oceans, and even some rivers and lakes, contain many different equilibrium reactions. For example, when carbon dioxide dissolves in the ocean all the following reaction can occur to produce about pH 8 that remains constant because of the "carbonate buffer". The pH of drinking water normally ranges from 5.5 to 9.0. At pH levels of less than 7.0, corrosion of water pipes may occur, releasing metals into the drinking water. This is undesirable and can cause other concerns if concentrations of such metals exceed recommended limits. The reactions below produce both H⁺ and OH^- ions, but [OH^-] > [ H⁺] so the pH value remains steady at about pH 8. Phosphates and silicates have chemical reactions that contribute to buffering capacity.
CO₂(g) <--> CO₂(aq)
H₂O + CO₂(aq) <--> H₂CO₃
H₂CO₃<--> H⁺ + HCO₃^-
HCO₃^-<---> H⁺ + CO₃^2-
HCO₃^-<---> CO₂(aq) + OH^-
CO₃^2- + H₂O <---> HCO₃^- + OH^-

18.1.0.1 Acidity and alkalinity
Instead of listing all the reactions that contribute to pH of oceans and rivers, you can call the capacity of chemical species together in the water to neutralize a strong acid acidity or alkalinity, as measured in a pH value of hydrogen ion concentration. pH log₁₀(1 / [H⁺]). For many lakes and rivers, if the water has values between pH 5.0 and pH 8.0, the water is often sufficiently buffered so that, if acids, bases or salts are added, the pH value does not change greatly. Some salts, e.g. ammonium sulfate, ammonium chloride and aluminium chloride, and some gases, e.g. carbon dioxide and sulfur dioxide, will increase the acidity of water. A low alkalinity river or lake may have sudden pH value changes if acid or acid industrial waste pollute the water. A sudden change in pH value, below pH 5.0 or above pH 8.0, kills many organisms, including fish. Alkalinity is a measure of the presence of bicarbonate, carbonate or hydroxide constituents. Concentrations less than 100 ppm are desirable for domestic water supplies. The recommended range for drinking water is 30 to 400 ppm. A minimum level of alkalinity is desirable because it is considered a buffer that prevents large variations in pH. Alkalinity is not detrimental to humans. Moderately alkaline water (less than 350 mg / l) in combination with hardness, forms a layer of calcium or magnesium carbonate that tends to inhibit corrosion of metal piping. Many public water utilities employ this practice to reduce pipe corrosion and to increase the useful life of the water distribution system. High alkalinity (above 500 mg / l) is usually associated with high pH values, hardness and high dissolved solids and has adverse effects on plumbing systems, especially on hot water systems where excessive scale reduces the transfer of heat to the water. Water with low alkalinity, < 75 mg /l). For example, some surface waters and rainfall, is subject to changes in pH because of dissolved gasses that may be corrosive to metallic fittings

18.1.1 pH with universal indicator
Use Universal Indicator with a constant technique. Construct a database of the pH of water at the same places, e.g. along a river bank, but at different times. Collect many readings and look for patterns.

18.1.2 pH of water in the laboratory
3.34.3: Carbon dioxide, Solubility of acetic oxide carbon dioxide in water, acidity of soda water
Use Universal Indicator to test the pH of deionized water or demineralized water, tap water, boiled tap water, tank water, bottled water, non-gaseous mineral water, gaseous mineral water, soda water, fizzy lemonade.

18.1.3 pH of rainwater
Collect rainwater in a clean container. Use Universal Indicator to test the pH of an isolated rain shower (a "rain incident") the first rain after a dry period, during continued periods of rain (a "rain episode") during different wind directions, In an urban area, in a rural area, in an area near an industrial plant, e.g. a powerhouse or steel works. Rain water is in equilibrium with the carbon dioxide in the atmosphere that forms carbonic acid, a weak acid (k_a = 1.75 x 10^-5), the solution under normal atmospheric conditions has pH 5.7. The average pH of rain is pH 5.0, with range of pH 5.6 to pH 4.5 but in eastern industrialized North America the average pH of rainwater is reported as pH 4.2.

18.1.4 pH of standing water
1. Use Universal Indicator to test the pH of temporary standing waters, puddles, tree boles, along rivers, shallow and deep water, main stream and tributary, lake or sea or ocean, including shallow and deep water, sewage water before and after treatment, rivers where they discharge sewage water, effluent discharged into rivers from factories and industrial plants.
2. Dip the pH meter into solution. Stir gently for a few seconds, until the readings stabilize. Record reading. Rinse pH detector tip with deionized water.

18.2.0 Total dissolved solids and suspended solids in water
Colorimetric tests based on Beer's Law are based on the observation that the higher the concentration of a substance in a sample, the darker the colour developed in the test, i.e. more light is absorbed by the sample.
The total dissolved solids test measures the total amount of dissolved minerals in water. The solids can be iron, chlorides, sulfates, calcium or other minerals found on the surface of the earth. The dissolved minerals can produce an unpleasant taste or appearance and can contribute to scale deposits on pipe walls. The following levels of total dissolved solids are expressed in mg / l: Less than 500 Satisfactory, 500 to 1 000 Less than desirable, 1 000 to 1 500 Undesirable, Over 1 500 Unsatisfactory. The only effective means of reducing total dissolved solids is by using reverse osmosis; however, removal is not economical

18.2.1 Insoluble solids in rainwater
Use a previously weighed filter paper kept in a desiccator. Collect the water in a very clean beaker. Swirl the water sample to keep the dirt suspended and filter 100 mL into a measuring cylinder.
Note the volume of the filtrate. Dry the filter paper in the dissector. Weigh the filter paper and insoluble particles.

18.2.2.1 Chlorides in groundwater
The chlorides in groundwater can occur naturally or be caused by pollution from sea water or industrial wastes. Chloride concentration above 250 mg / l can produce a distinct taste in drinking water. Where chloride content is known to be low, a noticeable increase in chloride concentrations may indicate pollution from sewage sources. The following levels of chlorides are expressed in mg / l: 0 to 250 Acceptable, 250 to 500 Less than desirable, 500 to 1 000 Undesirable, Over 1 000 Unsatisfactory.

18.2.2.2 Iron in drinking water
The iron in drinking water can give a rusty colour to laundered clothes and may affect taste. Frequently found in water because of large deposits in the surface of the earth. Iron can also be introduced into drinking water from iron pipes in the water distribution system. In the presence of hydrogen sulfide, iron causes a sediment to form that may give the water a blackish colour. Maximum concentration for iron in drinking water of 1.0 mg / l. The following levels of iron (Fe) are expressed in mg / l: 0 to 0.3 Acceptable, 0.3 to 1.0 Satisfactory (however, may cause staining and objectionable taste) Over 1.0 Unsatisfactory. Iron as it exists in natural groundwater is in the soluble (ferrous) state but, when exposed to oxygen, is converted into the insoluble (ferric) state with its characteristic reddish brown or rusty colour. If allowed to stand long enough, this rusty sediment will usually settle to the bottom of a container. However, it is difficult to use this type of settling to remove the iron.
The four options for removing iron from potable water are as follows:
1. For dissolved iron in concentrations up to 2.0 mg / litre, add food grade phosphate that sequesters the dissolved iron, i.e. keeps the iron in solution.
2. Zeolite softening can remove up to 10 mg / litre of dissolved iron.
3. Potassium permanganate can remove up to 10 mg / litre of iron and will remove dissolved as well as particulate iron. The permanganate provides oxygen to oxidize and precipitate any dissolved iron.
4. Liquid chlorine solution can be used for any quantity of iron, dissolved or not, and kill iron bacteria.

18.2.2.3 Sulfates in groundwater
The sulfates in groundwaterare caused by natural deposits of magnesium sulfate, calcium sulfate or sodium sulfate. Concentrations should be below 250 ppm. Higher concentrations cause laxative effects. Sulfates cannot be economically removed from drinking water. The following levels of sulfates are expressed in mg / l: 0 to 250 Acceptable, 250 to 500 Can be tolerated, 500 to 1 000 Undesirable, Over 1 000 Unsatisfactory.
Weigh a clean dry evaporating dish. Put the filtered rainwater in the evaporating dish and heat to dryness. If heated too rapidly, the solution "spits" and some solids may be lost. When you evaporate the solution to dryness, cool the evaporating dish and leave it in a desiccator for one day. Record the weight of the dissolved solids.

18.2.3 Extracted soluble solids from rainwater
Observe the dried filtrate under a microscope to identify its origin, e.g. sand, soot, organic matter or coal washing residues.

18.2.4 Contamination of groundwater from refuse deposits
The main problem of waste disposal today is the possible contamination of groundwater by direct seepage from refuse deposits or by substances leached from such deposits. The following experiments illustrate these processes:
1. Support a funnel with a top diameter of 100 mm with a support stand using a right angle clamp and a universal clamp. Put a small wad of cotton wool in the funnel and then fill it with soil to within a thumb width of the top. Put a beaker beneath it. Spread about 1 g of copper (II) sulfate on the soil and pour water over it. The water dissolves the copper (II) sulfate and seeps into the soil. After a few minutes the blue copper (II) sulfate solution drips from the funnel into the beaker placed beneath it.
2. Spread an equal amount of sodium sulfate, potassium sulfate or ammonium sulfate on the soil in the funnel, instead of the copper (II) sulfate. The liquid dripping into the beaker from the funnel will be colourless or slightly yellowish from the soil. Acidify it with hydrochloric acid and add 2% barium chloride solution, A thick white precipitate of barium sulfate forms (sulfate test).

18.2.5 Salinity
The salinity is the total solids, “salts”, in water after all carbonates have been converted to oxides, all bromide and iodide have been replaced by chloride, and all organic matter has been oxidized. So it is a measure of the total dissolved salts in a water sample, mainly Ca, Mg, Na, bicarbonate, Cl and sulfate.

18.2.6 Conductivity
The total dissolved solids in water can be measured by evaporating a water sample but it is more convenient to use methods based on water conductivity. The conductivity is a measure of how well a water sample transmits an electric current. It depends on the ionized substances dissolved in the water and temperature. Conductivity is expressed in umhos per cm. and is usually measured across one centimetre. deionized water is a poor conductor of electricity with conductivity 0.5 to 2 umhos / cm. However, water containing dissolved salts shows greater conductivity. Usually conductivity is directly proportional to the concentration of dissolved salts in the water. The conductivity of potable waters may range from 50 to 1500 umhos / cm. Dissolved ionic matter can be estimated from conductivity by multiplying by 0.54 to 0.96, depending on components in water and temperature. Total dissolved solids, TDS, is the concentration of minerals and salts impurities in the water, measured in parts per million, ppm. (1 ppm = 1 mg per litre). Specific conductivity is expressed as mhos per centimetre (M / cm) i.e. siemens per centimetre, S / cm. However, the mho (siemen) is a large unit, so usually the millimhos (millisiemen) (mS / cm) is used. A value of 0.67 is commonly used to convert conductivity as mS / cm into total dissolved solids as ppm. TDS ppm = Conductivity µS / cm x 0.67. Electronic conductivity instruments can automatically compensate for temperature and correct readings to 25°C. Many authorities think that potable water should contain less than 500 ppm of dissolved solids.
Remove the conductivity meter protective cap. Stir the sample with the sensor tip for a few seconds, until the readings stabilize. Record value as micromhos per centimetre. Estimate total dissolved solids by multiplying the conductivity by 0.67. Rinse the sensor tip with deionized water.

18.2.7 Cations and anions in rain, rivers and seawater
Average chemical compositionof natural waters, quoted by J. N. Butler, Carbon Dioxide Equilibria and Their Application, Addison-Wesley, Reading, MA, 1982, Chapter 5

Ion	Rain	Rivers	Seawater
Ca²⁺	.	0.53	10.6
Mg²⁺	.	0.21	54.6
Na⁺	0.009	0.39	479.0
K⁺	.	0.036	10.2
H⁺	0.072	.	.
NH4⁺	0.016	.	.
HCO₃^-	.	1.1	2.3
SO₄^2-	0.028	0.21	28.9
Cl^-	0.012	0.23	546.0
Br^-	.	.	0.85
F^-	.	0.008	0.07
NO₃^-	0.026	0.017	0.0001
H₂SiO₄^2-	.	0.15	.

18.3.0 Tests for air and dissolved oxygen in water
A good indicator of the health of water is how much air is dissolved. Low air levels usually mean high levels of water pollution. The mass of air that dissolves in water depends on the temperature of the water. As water is heated to near boiling point, the dissolved gases become less soluble.

18.3.1 Air dissolved in a water sample
Stand a beaker of water in sunlight. Bubbles of air appear. The taste of boiled water is different from tap water because boiled water has lost its dissolved oxygen. Note the temperature of a sample of water. Boil the water until no more bubbles appear. Collect the air from the water in an inverted measuring cylinder.

18.3.2 Oxygen content of water, dissolved oxygen, DO
1. The oxygen content of pure water is usually 7 to 10 mg per litre, depending on the temperature and atmospheric pressure. The following table shows the effect of temperature, at a constant pressure of 1013 mb (millibar) 760 torr. (1 mm mercury 133.3 Nm^-2)

Temperature ^oC	Oxygen saturation mg / L
0	14.16
5	12.37
10	10.92
15	09.76
20	08.84
25	08.11
30	07.53
35	07.04
40	06.59

2. In natural stretches of water, the oxygen content is also affected by oxygen consumption because of contamination and the breakdown process associated with it, and the production of oxygen because of the assimilation of underwater plants. The oxygen wasting processes predominate, i.e. if the loading because of insufficiently purified, effluent, for example, is too great, the stretch of water will gradually become a stinking, repulsive sewer in which life is no longer possible.
3. A sufficiently accurate test for determining the oxygen content of a water sample in schools is possible from the oxygen determination method devised by L. W. Winkler. This method uses the fact that when a manganous salt solution (e.g. manganous sulfate) is treated with caustic soda, a white precipitate of manganous hydroxide is produced.
MnSO₄ + 2 NaOH ---> Mn(OH)₂ + Na₂SO₄
In the presence of oxygen, the manganous hydroxide is oxidized to brown hydrated manganese oxide.
2Mn(OH)₂ + O₂ ---> 2MnO(OH)₂
The formation of hydrated manganese oxide is proportional to the amount of oxygen so that the intensity of the brown coloration is an indication of the oxygen content. The precipitate produced by the addition of manganous sulfate solution and caustic soda solution is coloured almost brown depending on the oxygen content of the water sample. The oxygen content may be estimated from the coloration. If the precipitate remains white or almost white, there is no oxygen or very little oxygen present. If the precipitate is light yellow, the water sample contains little oxygen. If the precipitate is coloured brown, it is rich in oxygen.

18.3.3 Amount of dissolved oxygen, titration
Manganese ions react with potassium iodide to produce iodine. Titrate the iodine with sodium thiosulfate using starch as an indicator.
Lower a sampler into the water from a convenient place, e.g. a bridge. Collect a water samples 1 meter below the surface by pulling a string attached to the stopper of the sampler. When bubbles no longer rise to the surface, note the water temperature and pull up the sampler. Check the sample for bubbles that can give a false, high reading. To the sample add 8 drops of manganous sulfate solution and 8 parts of alkaline potassium iodide azide. A precipitate forms. Add 8 parts of 1 to 1 sulfuric acid. Shake the sample until the reagent and the precipitate dissolve. The colour of the sample is now clear yellow if dissolved oxygen is low to brown-orange if dissolved oxygen is high. Add 8 parts of the starch indicator solution to the sample. It turns blue. Titrate the sample against known molarity sodium thiosulfate solution until the blue colour becomes colourless throughout the water sample. Record the results as ppm dissolved oxygen.

18.4.0 Ions in a water sample
See 12.11.3.2: Tests for metals with flame tests, metals and their compounds
Use these test solutions:

For Cl^-	AgNO₃ solution turns milky white
For SO₄^2-	BaCl₂ solution turns milky white
For Pb²⁺	Na₂S solution forms a black precipitate

Evaporate a water sample in a non-aluminium container. Heat slowly to avoid "spitting" when little water remains. Dissolve the crystalline mass after evaporation in 10 mL of deionized water. Add one drop of test solution and record the results.

18.4.1 Phosphate ions in water
Excess phosphate ions in water can cause eutrophication. Most modern detergents do not contain phosphate ions.
To check this, do the following experiment with 1 g of detergent dissolved in 1 g of water. Dissolve 1.5 g di-sodium hydrogen phosphate (Na₂HPO₄.12H₂0) in deionized water and make up to 1 litre. Make solutions with nine different concentrations by making up to 100 mL with deionized water the following volumes and label the containers with the concentrations.
Be careful! Use a burette or a pipette with rubber suction attachment. Do not suck by mouth!
Add 15 g of ammonium molybdate(VI)-4-water ([NH₄]₂Mo0₄] to 150 mL deionized water in a flask in crushed ice. Leave the solution to cool. Add 250 mL concentrated sulfuric acid to 250 mL deionized water. Stop adding the acid when the flask becomes too hot. Leave it to cool in ice. Slowly add the cold ammonium molybdate solution to the cold sulfuric acid solution. Add 10 mL of ammonium molybdate and acid solution to each of the nine phosphate solutions. Add 10 mL of ammonium molybdate and acid solution to a 100 mL sample of water. Add crystals of L-ascorbic acid and boil. Compare the colour with the colour of the standard solutions. This technique is called colorimetric analysis.

L phosphate (20 ug / L)	20 mL phosphate (4 ug / L)
75 mL phosphate (15 ug / L)	10 mL phosphate (2 ug / L)
50 mL phosphate (10 ug / L)	5 mL phosphate (1 ug / L)
40 mL phosphate (8 ug / L)	0 mL phosphate (0 ug / L)
30 mL phosphate (6 ug / L)	.

18.5.1 Tests of river, lake or ocean, record site data
1. Sample number, names of testers
2. Location: Distance from shore and location upstream or downstream from a marker
3. Date and time
4. Weather: Fine or cloudy or rainy, wind speed and direction
5. Air temperature
6. Water temperature
7. Current flow and depth: stagnant, calm, brisk, raging t.
8. Smell: no distinctive character, musty, fresh, putrescent, earthy, sewage-like, putrid, like liquid manure, peaty, chemical.
9. Colour and appearance: Compare the colour of the water against a white background. Appearance: scum, muddy, clear, brown, foamy, milky.
10. Visibility: Use a Secchi disc at depth of one metre
11. Turbidity: Observe settling on standing after 30 minutes. Precipitate any fine suspensions and colloids that pass through the filter paper by adding aluminium potassium sulfate (potassium alum, Al₂(SO₄)₃.K₂(SO₄).24H₂O) then filter. However, you can also use device called a nephelometer to measure turbidity, NTU
12 pH
13. Floating debris
14. Oil or petrol: Look for the rainbow effect of petrol or oil on water.
15. Detergents: Half fill a flask with river water, insert a stopper and shake for one minute, rate as "no detergent" if bubbles disappear in less than three seconds, rate as "slightly frothy" if the bubbles take up to ten seconds to break up, rate as "frothy" if the froth takes up to five minutes to disperse.
16. Microscopic examination: Note the size of suspended particles and any life forms, e.g. algae.
17. Dissolved oxygen, ppm
18. Conductivity, umhos / cm, conductivity x 0.67 = TDS, total dissolved solids

18.5.1a Eutrophication of waterways
Waterways can be have excess plant nutrients from leaching of land, especially agricultural land containing excess fertilizers, and discharge of effluents containing nitrogen and phosphorus. Decay of water weeds and other primary organic matter causes depletion of oxygen especially in the summer, e.g. decay of Eichhornia crassipes, water hyacinth, "the worst weed in the world", South America, Pontederiaceae.

18.5.2 Anions in sewage and tap water
See 12.11.5.0: Tests for anions in unknown solution, tests for acid radicals in solution
Compare concentrations by comparing the intensity of colour or amount of precipitate.

18.5.3 Physicochemical test methods for water samples
Many individual tests are necessary to detect with certainty the purity or contamination of a stretch of water including physico-chemical, bacteriological and biological test methods. You can study water samples taken from a river before and after passing through a community area. The physico-chemical examination of the water samples can include tests for the presence of pollutant indicators besides determining the smell, colour, temperature, and oxygen content. These pollution indicators are substances because of contamination with faecal matter in the water. However, many of these compounds are normally present in water in small amounts so precise guidelines and maximum permitted values have been decided. Occasionally, however, these values may differ considerably from the specified figure, because of geological conditions for example, although no contamination of a type that is harmful to health may be present in the water. The results of the individual physico-chemical tests must therefore only be used as a whole for assessing the quality of water. A selection of the most important physico-chemical tests for judging the purity or contamination of a stretch of water, which can also be carried out without great expense using facilities available in schools is given below.

18.6.0.1 Commonly occurring air pollutants, safe air and clean air

Pollutant	Source	Health	Safe air, level if harm not detected	Clean air
Sulfur dioxide	Sulfur in fuel	lung disease	2 pphm (annual mean)	0.01 ppm
Nitrogen oxides	High temperature combustion	lung disease bronchitis	16 pphm²(1 hour average)	< 0.01 ppm
Dust, soot particulates	Combustion, mining, clearing	lung disease if acid gases	90 micrograms / m³TSP, (annual mean)	10-20 g / m³
Ozone	Nitrogen oxides and carbon	lung irritation	12 pphm (1 hour average)	nil
Lead	Smelting, points, leaded petrol	cumulative poison	1.5 micrograms / m³ (3 month mean)	nil
Carbon monoxide	Motor vehicle combustion	body motor functions	9 ppm² (8 hour average)	< 1 ppm

pphm = parts per hundred million
ppm = parts per million
TSP = total suspendable particulates

18.6.1 Contamination of groundwater by refuse, hazardous wastes
The main problem of waste disposal is the possible contamination of groundwater by direct seepage from refuse deposits or by substances, leached from such deposits.
1. A 100 mm funnel is supported from a support stand using a right angle clamp and a universal clamp. Put a small wad of cotton wool in the funnel and fill it with soil to within a thumb width of the top. Put a beaker beneath. Spread l g copper (II) sulfate crystals on the soil. Pour water is poured over it. The water dissolves the copper (II) sulfate and seeps into the soil. After a few minutes the blue copper (II) sulfate solution drips from the funnel into the beaker beneath.
2. Spread an equal amount of sodium sulfate, potassium sulfate or ammonium sulfate on the soil in the funnel, instead of the copper (II) sulfate. The liquid dripping into the beaker from the funnel will be colourless or at any rate only slightly yellowish from the soil. On acidifying it with hydrochloric acid and adding some 2% barium chloride solution, a thick white precipitate of barium sulfate is forms.
3. Hazardous wastes must be disposed on in a geologically responsible way, usually by incineration in special furnaces. Hazardous wastes include pesticides, car batteries, petrol, oil, swimming pool chemicals, solvents, aerosol products, fire extinguishers, barbecue gas bottles, paint fluorescent lamps and tubes. Ask your local city council or town council how hazardous wastes should be disposed of and what facilities for disposal are available to individual householders.

18.6.2 Air pollution from burning refuse
Considerable pollution of the environment by harmful combustion products can be caused by the burning of rubbish that is often carried out without proper understanding. A typical example is the burning of polyvinyl chloride (PVC). This synthetic plastic is used to make many types of domestic objects but it is also used as a packaging material that is later scrapped becoming refuse. When it is burnt, the chlorine combined in its molecules is converted to hydrogen chloride among other products and this dissolves in the water produced to form dilute hydrochloric acid.
1. A few small PVC rods are laid on two small tablets of solid fuel (Esbit, Blitzem, firelighter, "meta" fuel, canned heat, snail bait (Metaldehyde) in a porcelain basin. A strip of moistened, blue litmus paper is hung 25 cm above the basin, using a support rod, a right angle clamp and a universal clamp. The fuel is set alight and the PVC is burnt. Within no more than a minute, the blue litmus paper is turned red by the hydrochloric acid formed by the burning PVC. The experiment should be carried out in a fume cupboard. Prove that the red coloration of the blue litmus paper is because of combustion of the PVC, two tablets of solid fuel without any PVC are burnt in a parallel experiment. The colour of the blue litmus paper is unchanged even after the two tablets have completely burnt away.
2. Put a porcelain dish containing a tablet of solid fuel and PVC rods on a glass plate. The solid fuel tablet is easier to ignite when it is broken in two and one piece is placed at an angle across the other. Use a 5 litre glass bell jar with a strip of moistened blue litmus paper secured by a rubber stopper hanging 10 cm inside the neck. Invert the glass bell jar over the porcelain dish when the solid fuel has been ignited. Although only a little PVC burns in the small quantity of air enclosed under the bell jar, the colour of the litmus paper turns strongly to red. These experiments show that rubbish should be deposited only on sites where the nature of the ground makes it unlikely that groundwater can become polluted by substances leached from the rubbish. Burning of the rubbish to reduce its bulk should be permitted only in appropriately designed refuse destructors in which the resulting combustion products can be rendered harmless.

18.6.3 Danger of vehicle exhausts, tailpipe gases
Exhaust gases from motor vehicles contain lead dust (5 to 30 mg / m³) nitric oxide (0.005 to 0.3% by volume) hydrocarbons (0.01-1%) and carbon monoxide (1-10% by volume). Carbon monoxide is especially dangerous to human beings because it cannot be detected by the senses because it has no colour, smell or taste. Its affinity to haemoglobin is 250 times that of oxygen. During carbon monoxide poisoning, a rapid breakdown in the supply of oxygen to the body occurs, leading to headaches and dizziness at low concentrations. At high concentrations (0.2% by volume) it can very rapidly lead to death.
To monitor the carbon monoxide content of the air you can use a carbon monoxide gas detector, e.g. the PHYWE gas detector, that allows detection of concentrations as low as 0.001% by volume. Make ten strokes of the pump and squeeze the bellows of the pump until they reach their limit so that the suction stroke admits 100 mL air. The colour of the preparation in the test-tube changes while the pumping is in progress. The white indicator layer, which contains iodine pentoxide (I₂O₅) as the effective reagent with fuming sulfuric acid (H₂S₂O₇) turns brown green during the reaction time, under the influence of selenium dioxide (SeO₂) as catalyst. The reaction can be described by the following equation:
SeO₂ + H₂S₂O₇
5CO + I₂O₅ ---> I₂+ 5CO₂
The length of the coloured zone depends on the carbon monoxide concentration. The carbon monoxide content of the air, in percentage volume can be read directly on the printed scale. MWC, maximum working concentration, is that concentration in the air of a workshop, measured in depth of breathing, at which no damage to health is to be expected, even with exposure during the day. Comparative measurements of the carbon monoxide content of the air should be repeated under different weather conditions because the carbon monoxide content of the air depends on rain, mist, wind, sunshine and smog conditions with an inversion layer.

CO content of the air in vol. %	CO concentration in ppm	CO concentration in the blood	Effect on humans
0.01	100 (MWC)	10 to 20%	No perceptible effect
0.025	250	30%	Headaches, slight fatigue
0.05	500	40 to 50%	Headaches, collapse and fainting on exertion
0.1	1000	60 to 70%	Unconsciousness, cessation of breathing on prolonged action
0.2	2000	>70%	Instant death

Nowadays, direct plug-in electrochemical 12-24V in-car carbon monoxide alarms are available to help detect dangerous levels of carbon monoxide in motor vehicles and workshops.

18.6.4 Water tests
Use individual tests to find the degree of purity or contamination of water. Tests include physicochemical, bacteriological and biological tests. You may do the tests at sites in a river before and after passing through a community area. You should test for the presence of "pollutant indicators" besides determining the smell, colour, temperature, oxygen content. These indicators are substances in the water because of contamination with faecal matter. Many of these compounds are normally present in water in small amounts so precise guidelines and maximum permitted values have been laid down. These values may differ considerably from the specified figure, because of geological conditions, although no contamination of a type that is harmful to health may be present in the water. So the results of the individual physico-chemical tests must be considered as a whole for assessing the quality of water. Some more important physico-chemical tests for judging the degree of purity or contamination of water are described below.

18.6.5 Smell of water, hydrogen sulfide
Water used for drinking must not smell unusual, repugnant, or revolting. Many substances can affect the smell of water. The bad smell of underground water may be caused by hydrogen sulfide that can be produced by the reduction of iron sulfide. You should make a preliminary test of smell to provide initial information when the sample is taken because some smells, like that of hydrogen sulfide, may rapidly disappear. You can more readily detect odours if the substance is heated slightly. Hydrogen sulfide, when dissolved in water, produces an offensive odour resembling that of rotten eggs. The presence of hydrogen sulfide in deep well-water is because of the reduction of sulfate. The acceptable level of hydrogen sulfide is 0.05 mg / L or less. Hydrogen sulfide can be removed through oxidation or by aeration or chlorination. The precipitated sulfur should be removed by filtration to prevent it from reverting back to hydrogen sulfide through the action of certain micro-organisms. The oxidation of hydrogen sulfide by chlorine may be advantageous in cases where it is otherwise unnecessary to repump the water, normally required with aeration, because chlorine can be applied directly into the system. Enough chlorine must be used to maintain a distinct chlorine residual
Put 100 mL of the water sample in a 250 mL wide mouth bottle. Close with a glass stopper and heat in a water bath to 40^oC. After shaking it vigorously, open the bottle and test the smell of the water immediately. Unbiased measurement of smell by means of numerical values is impossible but you can use the following terms to describe the smell: no distinctive character, musty, fresh, putrescent, earthy, sewage-like, putrid, like liquid manure, peaty, chemical. The general description "chemical" can be amplified, as smelling of hydrogen sulfide, chlorophenol (pharmacy shop smell) chlorine, tar, ammonia, mineral oil, phenol. You can use the terms "slight or "pronounced" to describe the intensity of the smell.

18.6.6 Colour of water
Pure, clean water is colourless or possibly just slightly bluish in colour. A colour other than this may be because of the most different kinds of foreign or contaminating matter. Thus humus materials generally produce a yellow brown coloration, iron a yellowish reddish one, micro-organisms, e.g. plankton organisms may give a greenish, yellowish or brownish colour to the water. Water that is distinctly contaminated has a greyish yellowish to grey black colour. Different substances that may cause differing amounts of harm may produce similar discoloration so an unbiased measurement using numerical values is impossible. It follows that the colour of water used for assessing its quality must be used only with all the other test results. Devices called tintometers can be used to record water colour. The name of a colour may be based on the "Munsell colour system".
Use two tall colourless glass beakers. Put 200 mL of the water sample in one beaker. Put the same amount of deionized water in the other beaker. Put both beakers on a white background, e.g. a sheet of writing paper or a white tile. The colour of the water sample, as viewed from above, is compared with that of the distilled water. It can be described as follows: colourless brown, slightly yellowish, yellowish green, yellowish, greenish, yellow, green, yellow brown, grey yellow, brownish, grey black. The room must be well-lighted by daylight.

18.6.7 Hydrogen ion concentration of water
Pure water is very slightly dissociated, i.e. split into its ions (H⁺ and OH^-). One litre of pure water contains 1 / 10 000 000 (10^-7) g of hydrogen ions (H⁺). Since the same amount of hydroxyl ions (OH^-) is also present, pure water has a neutral reaction. For the sake of simplicity, instead of the value 10^-7, the pH value is used which is the logarithm of the reciprocal of the hydrogen ion concentration. Water with a pH value of 7 is neutral; below it is acid and above 7 it is alkaline. Natural stretches of water usually have a pH value approximating to that of the neutral point. Extreme values, e.g., because of the nature of the soil from which the water originates, are pH 3 in the acid range and pH 12 in the alkaline. The organisms living in water thrive best of all at a pH between 6.8 and 7.8. Changes of hydrogen ion concentration, e.g. because of the introduction of insufficiently neutralized industrial effluent, may cause gross disturbances in the ecological equilibrium. Also, the changes may cause the direct poisoning of underwater life because of the materials introduced.
The simplest way to find the pH value is to use universal indicator paper, pH 1 to 10, for the whole pH range and as special indicator paper for various pH ranges, e.g. indicator rods: pH 2.5 to 4.5, pH 4.0 -7.0, pH 6.5 to 10.0. First wet a strip of universal indicator paper the water sample. After one minute find the pH value by comparing the colour produced with that of the colour scale provided. The colour scale of the universal indicator paper is subdivided into complete pH units. By estimating the intermediate stages, half pH units can also be read off. The test is repeated in the same way, using a strip of special indicator paper of the corresponding pH range. The graduation of the colour scale of the special indicator paper is so fine that 2 / 10 of a pH unit can be read off. By this means, determining the hydrogen ion concentration of a water sample with an accuracy normally sufficient for school purposes is possible. For more accurate measurements use a battery operated electronic pH meter.

18.6.8 Temperature of water
Drinking water should be neither too hot nor too cold and have a temperature between 8^oC and 12^oC. Colder water is generally regarded as unacceptable. At temperatures below 5^oC, stomach or intestinal troubles may even occur. Water at a temperature of more than 15^oC has no longer a refreshing effect. The temperature is important for the ecological equilibrium of a stretch of water since the oxygen content of water is very closely related to it. The introduction of large amounts of insufficiently cooled cooling water into a river by an industrial undertaking may have catastrophic consequences and, e.g. contribute to the mass death of fish observed recently.
Tie a rope, e.g. a Perlon cord 0.5 mm in diameter, just above the spherical bulb at the lower end of an aquarium thermometer. Also, a piece of metal, e.g. an old key, is attached to make it sink in water. Tie pieces of red string at intervals of 25 m. This device allows temperatures to be taken at various depths in water. The thermometer is let down, the depth of water is recorded, and after two minutes it is pulled up quickly and the temperature immediately read. Temperature measurements on water samples are done in the workroom using a chemical thermometer.

18.6.9 Colorimeter tests, and test methods
Tests for to Test method
Aluminum to Eriochrome, Cyanine R
Ammonia to Nitrogen, Salicylate
Ammonia to Nitrogen, Nesslerization,
Bromine to DPD, Tablets
Chlorine to DPI, Tablets
Iodine to DPI, Tablets
Chlorine Dioxide to DPI, Tablet / Glycine
Chromium (Hexavalent) to Diphenylcarbazide, -
Chromium (Total, Hexa and Trivalent) to Diphenylcarbazide,-
COD, Low to Digestion (mercury free)
COD, Low to Digestion contains mercury
COD, Standard to Digestion (mercury free)
COD, Standard to Digestion (contains mercury)
COD, High to Digestion (contains mercury)
Copper to Diethyldithiocarbamate
Copper to Bicinchoninic Acid
Cyanide to Pyridine to Barbituric Acid
Cyanuric Acid to Melamine
Fluoride to SPADNS
Hydrazine to P to dimethylaminobenzaldehyde
Hydrogen Peroxide to DPI Tablets
Iron to Bipyridyl
Iron to 1,10 Phenanthroline
Manganese (LR) to PAN
Manganese (HR) to Periodate
Molybdenum to Thioglycolate
Nickel to Dimethy1glyoxime
Nitrate Nitrogen to Cadmium Reduction
Nitrite Nitrogen to Diazotization /Coupling
Oxygen, Dissolved to Winkler Colorimetric
Ozone to Indigo Trisulfonate
pH (5.0 to 9.6) to Colorimetric
Phenols 4 to Aminoantipyrine
Phosphate (HR) to Molybdovanadate
Phosphate (LR) to Ascorbic Acid Reduction
Potassium to Tetraphenylboron
Silica (LR) to Heteropoly Blue
Silica (HR) to Silicomolybdate
Sulfate to Barium Chloride
Sulfide to Methylene Blue
Tannin to Tungsto to Molybdophosphoric Acid
Turbidity (0 to 400 FTU) to Absorptimetric
Zinc to Zincon