School Science Lessons
Chemistry
2025-10-03
Environmental chemistry
Contents
18.1.0 Tests for air pollution
18.2.0 Tests for pH in the environment
18.3.0 Tests for water pollution
18.1.0 Air pollution
18.1.1 Catalytic converter in a motor vehicle
18.1.2 Commonly occurring air pollutants, safe air and clean air
18.1.3 Danger of vehicle exhausts, tailpipe gases
18.1.4 Indoor air pollution, formaldehyde pollution
18.1.5 Tests for air pollution from burning refuse
18.1.6 Tests for carbon monoxide
18.2.0 pH tests
18.2.1 pH tests
18.2.2 Tests for pH of water in the laboratory
18.2.3 Tests for pH of soil samples
18.2.4 Tests for pH of standing water
18.3.0 Water pollution
18.3.1 Acidity and alkalinity of water
18.3.1a Drinking water test media, (Agar)
18.3.2 Cations and anions in rain, rivers and sea water
18.3.3 Chlorides in groundwater
18.3.4 Conductivity, TDS and electrical conductivity
18.3.5 Effect of household detergents on freshwater organisms
18.3.6 Eutrophication of waterways
18.3.7 Iron in drinking water
18.3.8 Salinity
18.3.9 Tests for air in a water sample
18.3.10 Tests for colour of water
18.3.11 Tests for contamination of groundwater from refuse deposits
18.3.12 Tests for dissolved oxygen, titration
18.3.13 Test for dissolved oxygen, DO (Winkler method)
18.3.14 Tests for electrical conductivity with a conductivity meter
18.3.15 Tests for environment of river, lake or ocean
Detergents
Smell
18.3.16 Tests for hydrogen ion concentration of water
18.3.17 Tests for insoluble solids in rain water
18.3.18 Tests for ions in a water sample
18.3.19 Tests for phosphate ions in water
18.3.20 Tests for smell of water, hydrogen sulfide
18.3.21 Test for soluble solids in rain water
18.3.22 Tests for sulfates in groundwater
18.3.23 Tests for temperature of water
18.3.24 Tests for turbidity
18.3.25 Tests for water pollution
18.3.26 Total dissolved solids and suspended solids in water
18.3.27 Using rivers for water testing
18.1.1 Catalytic converter in a motor vehicle
A catalytic converter reduces emissions of carbon monoxide, nitrogen oxides and hydrocarbons from motor vehicles It contains a 3-way catalyst of Platinum (Pt,), Palladium, (46Pd) and Rhodium, (45Rh) on a ceramic base.
Carbon monoxide is oxidized to carbon dioxide.
Hydrocarbons are oxidized to carbon dioxide and water.
Nitrogen oxides are reduced to nitrogen gas.
If the catalytic converter uses an oxidation catalyst, it cannot reduce the oxides of nitrogen, which must be lessened by engine design.
From "Ford Australia" (edited):
"The catalytic converter is located in the exhaust system and is designed for exhaust gases to pass through before being released to the atmosphere.
The catalytic converter turns the harmful pollutants in vehicle exhaust systems into harmless gases such as steam or water vapour.
When the exhaust gas comes into contact with the precious metals (or catalyst) found in the catalytic converter, a chemical reaction takes place.
It weakens the bonds of the polluting chemicals and allows them to easily convert into more desirable by-products of combustion.
The heart of the catalytic converter is a "ceramic monolith".
It is a honeycombed structure that has many small channels through which the exhaust gases flow.
The honeycomb structure means the gases touch a bigger area of catalyst at once, so they are converted more quickly and efficiently.
The entire surface of the monolith is coated with a washcoat of aluminium oxide.
This enlarged surface is then coated again with the precious metals platinum and rhodium or palladium and rhodium in a ratio of 5:1.
This creates reaction surfaces for oxidation, i.e. reduction, of the exhaust pollutants.
There may be several monoliths in the catalytic converter to compensate for expansion due to the high exhaust gas temperatures.
The pressure-sensitive monoliths are provided with an elastic intermediate layer in a two-shell, thermally insulated stainless steel casing.
A specific operating temperature is necessary for the catalytic converter to function correctly.
18.1.2 Commonly occurring air pollutants, safe air and clean air
World Health Organization, WHO, report:
In new estimates released on 25 March, WHO reports that in 2012 around 7 million people died (1 in 8 of total global deaths), as a result of air pollution exposure.
This finding more than doubles previous estimates and confirms that air pollution is now the world's largest single environmental health risk.
There is a stronger link between both indoor and outdoor air pollution exposure and cardiovascular diseases, e.g. strokes, ischaemic heart disease, and cancer.
Much of the health risk results from burning of fossil fuels.
The new figures underline the opportunity for more sustainable development choices to simultaneously save lives and mitigate climate change.
Table 18.1.0.1 Air pollutants
Pollutant
 Source
 Health
 Safe air
 Clean air
sulfur
dioxide
 sulfur in fuel
 lung disease
 2 pphm
(annual mean)
 0.01
ppm
nitrogen
oxides
 high temperature
combustion
 lung disease
 16 pphm2
(1 hour average)
 < 0.01
ppm
dust, soot
 combustion, mining, land clearing
 lung disease
if acid gases
 90 mg / m3TSP, (annual mean)
 10-20
g / m3
ozone
 nitrogen oxides
and carbon
 lung irritation
 12 pphm
(1 hour average)
 nil
lead
 smelting, leaded petrol
 cumulative
poison
 1.5 mg / m3
(3 month mean)
 nil
carbon
monoxide
 motor vehicle
combustion
 body motor
functions
 9 ppm2
(8 hour average)
 < 1
ppm

mg = micrograms, pphm = parts per hundred million, ppm = parts per million, TSP = total suspendable particulate
18.1.3 Danger of vehicle exhausts, tailpipe gases
Exhaust gases from motor vehicles contain lead dust (5 to 30 mg / m3), NO (0.005 to 0.3% by volume), hydrocarbons (0.01-1%), and CO (1-10% by volume).
Carbon monoxide is very dangerous to humans, 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, use a carbon monoxide gas detector that allows detection of concentrations as low as 0.001% by volume.
18.1.4 Indoor air pollution, formaldehyde pollution
The pollution in an indoor environment includes the following:
Suspended particulates, smoke nuisance, volatile and semi-volatile organic compounds, formaldehyde, and burning gases.
Formaldehyde pollution is caused by the following:
1. pressed wood products (hardwood plywood wall panelling, particle board, fibre board) and furniture made with these pressed wood products,
2. urea-formaldehyde foam insulation (UFFI),
3. combustion sources and environmental tobacco smoke,
4. durable press drapes and other textiles, and
5. glues.
Formaldehyde gas can cause watery eyes, burning sensations in the eyes and throat, nausea.
Also, it can cause difficulty in breathing in some humans exposed at elevated levels (above 0.1 parts per million).
Health effects include eye, nose, and throat irritation; wheezing and coughing; fatigue; skin rash; severe allergic reactions.
Average concentrations in older homes are usually well below 0.1 ppm.
In homes with many new pressed wood products, levels can be greater than 0.3 ppm.
18.1.5 Tests for 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 done without proper understanding.
A typical example is the burning of polyvinyl chloride (PVC), (C2H3Cl)n.
This synthetic plastic is used to make many types of domestic objects and packaging material later scrapped to become refuse.
When burnt, chlorine in its molecules is converted to hydrogen chloride dissolving in water to form dilute hydrochloric acid.
WHO advises that the acceptable level of fine particles in the air measuring less than 2.5 microns, known as PM2.5, should >25 μg/m3, (micrograms per cubic metre).
Above 300, and the United States Environmental Protection Agency warns outdoor activity becomes hazardous, even for healthy adults.
On January 12, 2013 in Beijing, the PM2.5 pollution hit 886, officially off the charts for dangerous air quality.
The problem is not limited to Beijing, because many other cities in China have also been suffering from air pollution.
Experiments
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 done 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 amount of air enclosed under the bell jar, the litmus paper colour turns strongly to red.
Rubbish should be deposited only on sites where groundwater cannot 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, where combustion products can be rendered harmless.
18.1.6 Tests for carbon monoxide
There is no specific chemical test for carbon monoxide gas.
Table 18.1.3 Carbon monoxide effects
CO content of the
air in vol.
per cent		 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

Direct plug-in electrochemical 12-24V in-car carbon monoxide alarms are now available to help detect dangerous levels of CO in motor vehicles and workshops.
18.2.1 pH tests
The pH tests use an indicator that changes colour 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.
When carbon dioxide dissolves in the ocean all the following reactions 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+ + CO32-
HCO3-<---> CO2 (aq) + OH-
CO32- + H2O <---> HCO3- + OH-.
18.2.2 Tests for pH of water in the laboratory
Use Universal Indicator to test the pH of the following:
Deionized water, (demineralized water), tap water, boiled tap water, tank water, bottled water, non-gaseous and gaseous mineral water, soda water, fizzy lemonade.
18.2.3 Tests for pH of soil samples.
1. Take soil samples with a teaspoon.
Add an equal amount of water to each sample, enough to cover the soil.
Shake the samples thoroughly then drain off the liquid or use a filter.
Test the pH of the collected liquid.
2. Use a commercial pH testing kit.
Put a level teaspoon of soil on a white tile.
Add 5 drops of pH due indicator liquid and stir with the rod provided.
Dust the paste with the white powder, barium sulfate.
Wait for one minute.
Read from the colour card the pH value of the colour nearest to that of the sample.
3. Most soils are either slightly acid or slightly alkaline.
Acid soils have pH values of less than 7.
Alkaline soils have pH values of more than 7.
Plant growth is affected by soil pH.
Few plants grow well in soils with pH values below 5.5.
Most other plants grow best in soils with pH values 6 to 7.4.
Plants adapted to acid soils may not get enough iron and manganese from alkaline soils.
Their young leaves show yellowing, chlorosis, and growth is poor.
Severe deficiency leads to death.
Plants adapted to alkaline and slightly acid soils can be harmed by dissolved aluminium and manganese in very acid soils and may not get enough calcium.
You can raise soil pH by adding agricultural lime or dolomite.
5. Check the pH of mix in pots, because most fertilizers produce acidity.
Raise pH with a suspension of 5 g of hydrated lime (builders' lime), in a litre of water.
Lower pH with 2 g of iron sulfate in a litre of water.
Within two minutes, heavily water the pot to remove excess salt.
Soil test kit
1. The test kit contains one bottle of pH dye indicator and one bottle of barium sulfate solution.
Place a level teaspoon of mixed soil or potting mix on the test plate.
Add 3 to 5 drops of indicator liquid and stir with the rod provided.
Dust the paste with the white powder provided.
Wait one minute.
Read from the colour card the pH value of the colour nearest to that of the sample.
For a garden bed, take at least five samples from holes dug in different parts of the bed.
Each sample is to extend from the surface to a depth of 10 cm.
Test each sample separately.
For farm paddocks, take at least 20 samples from each area.
Mix samples together thoroughly and test as one sample.
For bought and home made potting mix, thoroughly mix the bulk lot.
For mix in a pot, first knock the root ball from the pot.
Remove a wedge of mix representing the whole depth of the root ball.
Mix thoroughly.
For a mix in large tubs, dig down the side of the root ball as deeply as is possible.
Thoroughly mix the sample removed.
2. Show how to use a soil pH test kit
Plants cannot absorb plant nutrients from the soil if the soil is too acid or too alkaline.
Soils that are not well-drained are too acid.
Soils made from coral rocks are too alkaline.
You can test the soils using a colour test.
If the colour of soil in the test turns yellow orange.
The soil is too acid.
If colour if soil in the test turns blue purple.
The soil is too alkaline.
If colour if soil in the test turns dark green.
The soil is not too acid nor too alkaline.
Collect just enough soil from just under the surface of the soil to cover your little finger nail, and place on a white plate.
Shake two drops of the indicator on the soil and mix to a paste with the stick.
Sprinkle some special white power on the paste.
Wait a few minutes then match the colour of the powder with the colour chart.
Do this for swampy soil, coral soil, dark well drained soil.
3. Acids have a sharp sour taste and can dissolve substances, e.g. in a car battery.
Alkalis have a slippery feel and can dissolve substances, e.g. soap.
Plants cannot absorb plant nutrients from the soil if the soil is too acid or too alkaline.
Soils in swampy ground are too acid for most plants.
Soils made from coral sand are too alkaline for most plants.
You can make soil less acid by adding burnt shells hammered to a powder, and by draining the soil.
You can make soil less alkaline by adding rotten plants from a compost heap.
Good soil is dark in colour from the rotten plants and is well drained.
To make sure that the soil is not too acid and not too alkaline the agriculture field officer can do a soil test.
Test soil pH, acid soils and alkaline soils
Over years of cropping, the pH of the soil can become too low (acidification),.
When the soil has a pH of below 5.5, it is called an acid soil.
One of the reasons why soils become more acid, is that under continual cropping, many commonly used fertilizers have an acidifying effect.
Care should be taken when fertilizers such as ammonium sulfate or mono ammonium phosphate are used.
Loss of some basic nutrients from the soil, e.g. calcium, also causes acidification.
Soils high in clay and organic matter are more resistant to acidification, conversely, once a soil has been acidified, they are more resistant to rehabilitation.
Adding organic matter is preventative, and adding lime is rehabilitative.
Acid soils and alkaline soils
This topic is very important, because you can improve the fertility of your garden soils by treating them so that they are not too acid or too alkaline.
The pH scale measures whether substances are acid or alkaline:
pH 1 very strong acid that can burn you, e.g. battery acid,
pH 6 weakly acid, e.g. soda water,
pH 7 neutral, neither acid nor alkaline, e.g. water,
pH 8 weakly alkaline, e.g. soap, and
pH 14 very strong alkali that can also burn you.
Acids have a sharp raw taste, e.g. unripe oranges or bush limes.
Alkalis have a slippery feel, e.g. soap, saliva.
Plants can absorb plant nutrients best when pH is 6 to 7. In soils formed from coral rock the pH will be too high.
In swampy land the pH will be too low.
To lower the pH add rotten compost.
To raise the pH add lime.
18.2.4 Tests for pH of standing water
1. Use Universal Indicator to test the pH of the following:
Temporary standing waters, puddles, tree boles, along rivers, shallow and deep water, main stream and tributary, lake or sea or ocean.
Also, test 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.
18.3.1a Drinking water test media, (Agar)
Bismuth Sulfide Agar, Brilliant Green Agar, Brilliant Green Bile Lactose Broth, Cetrimide Agar, DCLS Agar, ENDO Agar (Base), Formate Ricinoleate Broth, MacConkey Agar No 1, Malachite Green
Broth, Methyl Red Voges Proskauer Broth, Nutrient Broth No 3, Peptone Water, Phenylalanine Agar, Plate Count Agar, Selenite Broth (Base), Simmons Citrate Agar, SS-Agar, Tetrathionate
Broth, Tetrathionate Broth (Base), Triple Sugar Iron Agar, Tryptic Soy Broth, Tryptone Glucose Extract Agar, Tryptone water, Urea Broth, Violet Red Bile Agar.
18.3.1 Acidity and alkalinity of water
The capacity of chemicals 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+]).
In lakes and rivers, if water pH is pH 5.0 to pH 8.0, the water is sufficiently buffered, so, if acids, bases or salts are added, the pH value does not change greatly.
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.
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.
For example, in hot water systems excessive scale may reduce the transfer of heat to the water.
Water with low alkalinity, (< 75 mg / l), e.g. some surface waters and rainfall, is subject to changes in pH, because of dissolved gasses.
These dissolved gasses may be corrosive to metallic fittings.
Experiment
Tests for pH with universal indicator
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.
18.3.3 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 / Litre 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 / Litre:
0 to 250 Acceptable,
250 to 500 Less than desirable,
500 to 1 000 Undesirable, and
Over 1 000 Unsatisfactory.
18.3.4 Conductivity, TDS and electrical conductivity
Total solids, mg / L, Conductivity, umhos / cm, conductivity × 0.67 = TDS, total dissolved solids, Floating debris.
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.
(conductivity: mho (ohm backwards) = 1 siemens (s) = ohm-1 = amps / volts).
18.3.5 Effect of household detergents on freshwater organisms
Most detergents are "biodegradable", but degradation takes time, so detergents in rivers may be harmful before being degraded.
Detergents may affect the permeability of biological membranes.
Investigate the tolerance of certain freshwater organisms to different concentrations of detergents.
Use different organisms for testing, e.g. Daphnia, freshwater snails, small fish, the monocotyledon duckweed Lemna, and tadpoles.
18.3.6 Eutrophication of waterways
Waterways can 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 Water hyacinth (EicHornia crassipes), "the worst weed in the world", South America, Pontederiaceae.
18.3.7 Iron in drinking water
The iron in drinking water can give a rusty colour to laundered clothes and may affect taste.
It is 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 /Litre.
The following levels of iron (Fe) are expressed in mg / Litre:
* 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
The ferric state has a 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.
18.3.8 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.
It is a measure of the total dissolved salts in a water sample, mainly Ca, Mg, Na, bicarbonate, Cl and sulfate.
18.3.9 Tests for air 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.
18.3.10 Tests for colour of water
1. 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.
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 chart".
2. 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.
Compare the colour of the water sample, as viewed from above, with the colour of the distilled water.
Colour can be described as: 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.
18.3.11 Tests for contamination of groundwater from refuse deposits
A problem of waste disposal is the possible contamination of groundwater by direct seepage from refuse deposits or by substances leached from such deposits.
Experiments
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 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).
3. Hazardous wastes must be disposed on in a geologically responsible way, usually by incineration in special furnaces.
Hazardous wastes include the following:
pesticides, car batteries, petrol, oil, swimming pool chemicals, solvents, aerosol products, fire extinguishers, barbecue gas bottles, paint fluorescent lamps and tubes.
Experiment
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.
18.3.12 Tests for 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 metre 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 so that it turns blue.
Titrate the sample against known molarity sodium thiosulfate solution until the blue colour becomes colourless through the water sample.
Record the results as ppm dissolved oxygen.
18.3.13 Test for dissolved oxygen, DO (Winkler method)
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.
When water is heated to near boiling point, the dissolved gases become less soluble.
Experiments
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)
The solubility of dissolved oxygen decreases as salinity increases.
Table 18.3.2 Dissolved oxygen, DO
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.
It was devised by L. W. Winkler, 1863 - 1939, Hungary.
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 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.
18.3.14 Tests for electrical conductivity with a conductivity meter
See diagram: 18.2.6: Electrical conductivity (EC) and Total Dissolved Solids (TDS).
1. Electrical conductivity (EC) measures the capacity of water to conduct electrical current.
It is directly related to the concentration of salts dissolved in water, Total Dissolved Solids (TDS).
Salts dissolve into positively charged ions and negatively charged ions, which conduct electricity.
Portable meters are used in the field to measure electrical conductivity where it is difficult to measure TDS.
Distilled water does not contain dissolved salts to conduct electricity, so has an electrical conductivity of zero.
However, at high salt concentrations electrical conductivity is not directly related to salts concentration.
TDS (ppm) = 0.64 × EC (μS / cm) (microSiemens / cm).
2. Conductivity is a measure of the specific conductance of an electrolyte solution.
By using the AC resistance between electrodes, it measures the ionic content in a solution.
Also it indirectly measures total dissolved solids, TDS.
Conductivity SI unit: siemens per meter, S / m
3. The total dissolved solids in water can be measured by evaporating a water sample.
However, 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 0.67 S / cm × 0.67.
Electronic conductivity instruments can automatically compensate for temperature and correct readings to 25 oC.
Many authorities think that potable water should contain less than 500 ppm of dissolved solids.
18.3.15 Tests for environment of river, lake or ocean
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, daily range, oC
6. BOD, Biochemical Oxygen Demand, mg / L
7. Colour and appearance.
Compare the colour of the water against a white background.
Appearance: scum, muddy, clear, brown, foamy, milky.
8. Current flow and depth: stagnant, calm, brisk, raging.
9. Detergents
Detergents
Experiment
Half fill a flask with river water, insert a stopper and shake for one minute,(br) 1. Rate as "no detergent" if bubbles disappear in less than three seconds, 2. Rate as "slightly frothy" if the bubbles take up to ten seconds to break up, 3. Rate as "frothy" if the froth takes up to five minutes to disperse.
10. Dissolved oxygen, ppm, % saturation
11. Faecal coliform bacteria colonies / 100 mL
12. Microscopic examination: Note the size of suspended particles and any life forms, e.g. algae.
13. Oil or petrol: Look for the rainbow effect of petrol or oil on water.
14 pH
15. Smell: Smell
No distinctive character | musty | fresh | putrescence | earthy | sewage-like | putrid | like liquid manure | peaty | chemical|.
16. Turbidity
18.3.24 Tests for turbidity
17. Total nitrates, mg / L
18. Total phosphorus, mg / L
19. Visibility: Use a Secchi disc at depth of one metre
20. Water temperature, daily range, oC
Each item can be given a comparative rating so that an overall rating can be calculated for comparison with other sites.
18.3.16 Tests for hydrogen ion concentration of water
1. 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.
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 occur, because soil from which the water originates, are pH 3 in the acid range and pH 12 in the alkaline range.
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.
Experiment
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, you can find the hydrogen ion concentration of a water sample with an accuracy sufficient for school purposes.
For more accurate measurements use a battery-operated electronic pH meter.
18.3.17 Tests for insoluble solids in rain water
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.
18.3.18 Tests for ions in a water sample
Use the following test solutions:
Table 18.4.0 Test solutions
For Cl- AgNO3 solution turns milky white
For SO42- 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.
18.3.19 Tests for phosphate ions in water
Excess phosphate ions in water can cause eutrophication.
Nowadays, most 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]2MoO4] 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, then 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.
Table 18.4.1 Phosphate ions in water
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) .

18.3.20 Tests for smell of water, hydrogen sulfide
1. 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.
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.
2. 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.
Experiment
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 use the following terms to describe the smell:
1. no distinctive character, 2. musty, 3. fresh, 4. putrescence, 5. earthy, 6. sewage-like, 7. putrid, 8. like liquid manure, 9. peaty, 10. chemical.
The general description "chemical" can be amplified, as smelling of hydrogen sulfide, chlorophenol (pharmacy shop smell), chlorine, tar, ammonia, mineral oil, phenol.
Use the terms "slight or "pronounced" to describe the intensity of the smell.
18.3.21 Test for soluble solids in rain water
Rain water is in equilibrium with the carbon dioxide in the atmosphere that forms carbonic acid, a weak acid (ka = 1.75 × 10-5).
The rain water 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 rain water is reported as pH 4.2.
Experiments
1. Test for soluble solids in rain water
Observe the dried filtrate under a microscope to identify its origin, e.g. sand, soot, organic matter or coal washing residues.
2. Collect rain water in a clean container.
Use Universal Indicator to test the pH of the following:
An isolated rain shower ( "rain incident"), first rain after dry period, during continued periods of rain ("rain episode") during different wind directions.
Also test in an urban area, a rural area, in an area near an industrial plant, e.g. a powerhouse or steel works.
18.3.22 Tests for sulfates in groundwater
The sulfates in groundwater are 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:
1. 0 to 250 Acceptable, 2. 250 to 500 Can be tolerated, 3. 500 to 1 000 Undesirable, 4. Over 1 000 Unsatisfactory.
Experiment
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.
18.3.23 Tests for temperature of water
1. 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 water, because 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 severe consequences and contribute to the mass death of fish Experiment
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.
18.3.24 Tests for turbidity
Turbidity, mg / L.
Observe settling on standing after 30 minutes.
Precipitate any fine suspensions and colloids that pass through the filter paper by adding aluminium potassium sulfate, (Al2(SO4)3.K2(SO4).24H2O), then filter.
However, you can also use device called a nephelometer to measure turbidity, NTU.
18.3.25 Tests for water pollution
Use individual tests to find the degree of purity or contamination of water.
Tests include physicochemical, bacteriological and biological tests.
Do the tests at sites in a river before and after passing through a community area.
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 must be 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.
List of tests
1. Salinity: meter and refractometer
2. Turbidity: Tube and meter
3. Nitrite: Hach test strips
4. Nitrate: Hach test strips and spectrophotometer
5. Phosphate: Hach test strips and spectrophotometer
6. pH: Narrow range Neutralist pH strips and meter
7. Dissolved oxygen: Meter and spectrophotometer
8. Conductivity: Meter.
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.
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 done without great expense using facilities available in schools is given below.
18.3.26 Total dissolved solids and suspended solids in water
Colorimetric tests based on the Beer-Lambert law assume the higher the concentration of a substance in a sample, the darker the colour in the test.
More light is absorbed by the sample.
The intensity of light that passes through a sample diminishes exponentially with the concentration of substances and thickness of the sample.
Absorption of light depends on the pathway length and concentration of the material passed through.
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 / Litre:
Less than 500 Satisfactory,
500 to 1 000 Less than desirable,
1 000 to 1 500 Undesirable, and
Over 1 500 Unsatisfactory.
The only effective means of reducing total dissolved solids is by using reverse osmosis, however, such removal is not economical.
18.3.27 Using rivers for water testing
Two experiment methods:
1. Manipulate one variable to see the effect on the other, keeping other variables constant For example: What is the effect of temperature (the manipulated or independent variable) on the concentration of Vitamin C in orange juice (the dependent variable)?
2. Concomitant variation.
A naturally occurring variation (Variable 1) is correlated against another variation (Variable 2).
The correlation method.
For example: How does temperature in the natural environment affect stoma opening?
In this experiment you do not need you to control the environmental temperature, but you do need to measure the dependent variable at different temperatures.
However, there may be other potentially influential variables, e.g. humidity.
That does not preclude the variables in relationship being considered.
One way to address the other variables is to collect data on the other variable as well and run the statistics on each pair of variables separately.
Where the data are collected from a local river, the method used is the second approach.
Students are relying on some naturally occurring change in water quality.
For example: BOD, turbidity, temperature as one set of variables, then trying to correlate this with the riverside development at sites along the river.
However, there are so many confounding variables and the differences between two sites may be many).
A river may be just used as a source of field data for collecting samples for water quality testing.
18.3.2 Cations and anions in rain, rivers and sea water
Average chemical composition of natural waters, quoted by J. N. Butler, "Carbon Dioxide Equilibria and Their Application", Addison-Wesley, Reading, MA, 1982, Chapter 5
Table 18.2.7 Cations and anions in rain, rivers and sea water
Ion
 Rain
 Rivers
 Sea water
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
SO42-
 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
H2SiO42-
 .
 0.15
 .