Senior Biology Projects (Grades 9 - 12)
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• The effect of sound on plants
• Plants in different environments (light intensity, colour)
• The effect of nicotine, air, yeast on mold growth
• Factors affecting the strength of hair, the growth of bacteria, molds or yeast
• Experiment with Hydroponics
• Use seedlings started from seed with three types of soil and >different rates of fertilizer
• The effectiveness of Antiseptics and soaps on household bacteria
• The effect of air pollution on algae, protozoa, fish, insects or mosses and lichens
• Comparing types of artificial light on plant growth
• Conditions necessary for the life of a brine shrimp
• The commercial uses of algae methods of production
• Producing mutations in bacteria, yeast, protozoa or molds
• Best conditions for mushroom production, growth of ferns
• The effects of ultrasonic antibiotics temperature changes on bacteria count
• Microbial antagonism
• Reaction of paramecia, planaria to pH, light and temperature conditions
• Plant tropisms and growth hormones
• Transpiration rates for different plants and conditions
• Sugar level in plant sap at different times and dates
• Using radioisotopes to study uptake of plant nutrients
• A study of territoriality in mice
• A study of the cleaning habits of mice
• Observation of conditioned responses in different animals
• A study of animal phosphorescience and other biolumincescences
• Learning and perception in animals and humans
• Studies of memory span and memory retention
• Age versus learning ability
• A study of the relation between physical exercise and learning ability
• Is audio or visual information better remembered
• The effect of bleaching and dyeing on hair
• A study of the percentage of DNA (by weight) in different species
• Factors affecting the enzyme's reaction rates
• Genetic variations across a Sansevieria leaf
• Factors affecting seed germination (e.g. soil temperature, pH)
• Root formation in cuttings versus lighting conditions
• Factors affecting flowering
• Study of sterility in plant hybridgs (F1 and F2)
• Comparison of different plant's ability to add humus to the soil
• Factors affecting Nodule Formation in Legumes
• Can household compounds (e.g. tea) be used to promote good health in plants
• Effects of cigarette smoke on the growth of plants
• The effects of water impurities on plant growth
• The effects of phosphates on aquatic plants
• Effect of mineral deficiencies on protein content in soybeans
• The effect of excess salinity on plants
• A study of the tumours produced in plants by agrobacterium tumifacieus
• The effect of polarized light on plant growth directica
• The effects of solar activity on plant growth
• Tracing solar activity cycles in tree growth rings
• The effects of electric fields on plants
• The effects of magnetic fields on plant growth
• Effects of magnetism on the size and frequency of blooms and fruits
• Does magnetizing seeds before planting affect growth
• The effects of XRay and other radiation on plants
• The effect of music of varying types and duration on plants
• Organic fertilizer versus chemical Fertilizer
• Study of population fluctuations in insects
• A study of toxicity of insecticides versus temperature
• Is polarizes light the guidance system for Foraging ants
• A study of stimuli that attract mosquitos
• The factors affecting the rate at which a cricket chirps
• Study of insect of animal behaviour versus population density
• A study of diffusion through cell membranes
• Growth of plant and animal cells by cloning
• Regeneration in sponges, Paramecia, Planaria, etc.
• Manipulation of Vegetative reproduction in plants
• Search for nearvacuum environment tolerant plants
Senior Environmental Science Projects (Grades 9 - 12)
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• The study of flora in a given region
• Observations of urban wildlife
• Study of adaptations of city flora to smog
• An ecological study of the animal and plant populations occupying the same tree
• The effects of crowding (with the same or other species) on a certain plant
• Annual variations in the ecology of a body of water
• A study of a shoreline
• Observations of the spread of Dutch Elm disease
• A study of the relation between soil type and vegetation
• A study of the relation between vegetation and insects
• Monitoring the changes in wildlife caused by human encroachment
• The study of the impact of pollution on an ecosystem
• A study of water pollution from feed lot farms
• Tracing chemical(e.g. DDT) concentrations in successive food chain levels
• Ozone destruction experiments
• A study of air purification methods
• Efficient methods of breaking down crude oil in seawater
• Experimenting with microbial degradation of petroleum
• Experimenting with biodegradability
• Finding efficient methods of harvesting and using plankton
• Find and ink that would decompose for recycling paper
Environmental chemistry
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18.1.0 pH surveys
18.2.0 Measure 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 Survey standing water
18.6.0 Pollution
18.7.0 Swimming pool chemistry
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18.1.0 pH surveys
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 different types of standing water
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18.2.0 Measure 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
12.13.0 Hardness in water
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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
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18.4.0 Ions in a water sample
18.4.1 Phosphate ions in water
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18.5.0 Survey standing water
18.5.1 Site survey of river, lake or ocean
18.5.2 Anions in sewage and tap water
18.5.3 Physicochemical test methods for water samples
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18.6.0 Pollution
18.6.1 Contamination of groundwater by refuse, hazardous wastes
2.32.1 Composition of the atmosphere and greenhouse gasses
18.6.2 Air pollution from burning refuse
18.6.3 Danger of vehicle exhausts, tailpipe gases
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
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18.1.0 pH Surveys [H+] hydrogen ion concentration
See: 12.10.7: Buffer solutions | See: 7.7.0: Solutions, solubility equilibrium, solubility rules | See 12.10.3.2: Hydrolysis of ammonium chloride
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+ + CO32-
HCO3- <---> CO2(aq) + OH-
CO32- + H2O <---> HCO3- + OH-
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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
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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.
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18.1.2 pH of water in the laboratory
Use Universal Indicator to test the pH of distilled water or demineralized water, tap water, boiled tap water, tank water, bottled water, non-gaseous mineral water, gaseous mineral water, soda water, fizzy lemonade.
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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.
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18.1.4 Measuring pH of different types 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 distilled water.
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18.2.0 Measure total dissolved solids and suspended solids in water
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 - 1,000 Less than desirable, 1,000 - 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
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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.
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18.2.2.1 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 - 250 Acceptable, 250 - 500 Less than desirable, 500 - 1,000 Undesirable, Over 1,000 Unsatisfactory.
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18.2.2.2 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 -0.3 Acceptable, 0.3 - 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. Four options for removing iron from potable water: (a) For dissolved iron in concentrations up to 2.0 mg / l, add food grade phosphate that sequesters the dissolved iron, i.e. keeps the iron in solution. (b) Zeolite softening can remove up to 10 mg / l of dissolved iron. (c) Potassium permanganate can remove up to 10 mg / l of iron and will remove dissolved as well as particulate iron. The permanganate provides oxygen to oxidize and precipitate any dissolved iron. (d) Liquid chlorine solution can be used for any quantity of iron, dissolved or not, and kill iron bacteria.
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18.2.2.3 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: 0 - 250 Acceptable, 250 - 500 Can be tolerated, 500 -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.
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18.2.3 Extracted soluble solids from rainwater
Examine the dried filtrate under a microscope to identify its origin, e.g. sand, soot, organic matter or coal washing residues.
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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).
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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.
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18.2.6 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. Distilled 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. 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 be less than 500 ppm.
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 distilled water.
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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.
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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.
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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 break down 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.
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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.
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18.4.0 Ions in a water sample
See also 12.11.3.2: Flame tests to identify metals and their compounds
Use these 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 distilled water. Add one drop of test solution and record the results.
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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 distilled water and make up to 1 litre. Prepare solutions with nine different concentrations by making up to 100 mL with distilled 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 distilled water in a flask in crushed ice. Leave the solution to cool. Add 250 mL concentrated sulfuric acid to 250 mL distilled 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) -
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18.5.1 Site survey 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
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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.
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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.
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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.
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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", metaldehyde, ethanal [CH3CHO]4, acetaldehyde "meta" fuel, canned heat, snail bait) 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. If no fume cupboard is available, the experiment can be carried out as in 5.14.2.2.
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. Prepare 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.
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18.6.3 Danger of vehicle exhausts, tailpipe gases
Exhaust gases from motor vehicles contain lead dust (5 - 30 mg / m3) nitric oxide (0.005 - 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 break down 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 - 20% No perceptible effect
0.025 250 30% Headaches, slight fatigue
0.05 500 40 - 50% Headaches, collapse and fainting on exertion
0.1 1000 60 - 70% Unconsciousness, cessation of breathing on prolonged action
0.2 2000 >70% Instant death
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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.
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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 40oC. 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.
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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 distilled 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.
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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 - 4.5, pH 4.0 -7.0, pH 6.5 - 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.
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18.6.8 Temperature of water
Drinking water should be neither too hot nor too cold and have a temperature between 8oC and 12oC. Colder water is generally regarded as unacceptable. At temperatures below 5oC, stomach or intestinal troubles may even occur. Water at a temperature of more than 15oC 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.
MICROBIOLOGY
1. What happens to the way plants grow if there are no microorganisms in the soil?
Take a sample of fertile soil from a field or garden and divide it into two portions. Bake one in an oven (to destroy the microorganisms). Leave the other portion alone as a control. Plant the same number of seeds in each soil sample. Remember to treat both samples the same while the plants are growing. Make sure all the plants receive the same amounts of water and light, and are kept at the same temperature. How do the plants differ as they grow?
2. Are different plants affected in different ways by specific microorganisms?
Some microorganisms and plants form mutually beneficial partnerships. For example, certain bacteria make a natural nitrogen fertilizer for plants in the family called legumes. Peas, alfalfa and soybeans are legumes. The nitrogen-fixing bacteria are available from garden supply stores and by mail order. Grow both legumes and non-legume plants with and without the bacteria. Are there differences in how well the plants grow?
SCIENCE PROJECT WORK WITH PROCEDURES
1. Do different varieties of the same fruit have the same level of vitamin C?
What about different brands of orange juice? Or fresh juice compared to juice from frozen concentrate? Does the way a fruit is stored or how long it is stored change the level of vitamin C?
Background Info: Most birds and animals make their own vitamin C. But a few species, like people and guinea pigs, must get it from their food. Good sources of vitamin C are citrus fruits like oranges and grapefruit, strawberries, green peppers, broccoli and potatoes. Vitamin C is required for the body to make and maintain collagen, a protein. Collagen forms the base for all connective tissue in the body. If you don't have enough vitamin C in your diet, you might get the disease scurvy. Symptoms include loss of appetite, bleeding gums, loose teeth, swollen ankles and tiny hemorrhages (bleeding spots) in the skin.
Procedure to test for vitamin C content. (With this method, you can compare relative vitamin C content and rank foods from highest to lowest, but you won't be able to get exact concentrations.)
You'll need some 2% iodine solution (find it at your local pharmacy) to prepare the vitamin C indicator solution described in steps 1 to 4.
1. Mix 1 tablespoon of cornstach into enough water to make paste.
2. To this paste, add 250 milliliters of water and boil for 5 minutes.
3. Add 10 drops of the starch solution to 75 milliliters of water (use an eyedropper).
4. Add enough iodine to produce a dark purple-blue color. Now your indicator solution is ready.
5. Put 5 milliliters of indicator solution (about 1 teaspoon) in a 15-milliliter test tube (one for each sample).
6. To the test tube, use a clean eyedropper to add 10 drops of juice from the fruit or beverage (for solids, pulp them in a blender and strain the juice). Re-clean the eyedropper for each sample.
7. Hold the test tube against a white background. Line up the tubes from lightest to darkest purple. The lighter the solution, the higher the vitamin C content. That's because vitamin C causes the purple indicator solution to lose its color
2. Are there different amounts of iron in different breakfast cereals?
The iron in ready-to-eat breakfast cereals is in the form called elemental, not in combination with any other chemical compound. Iron is sprayed on the outside of cereal flakes. You can separate the iron with a strong magnet.
Background Info: Iron is essential in a healthy diet to build blood. Iron is easiest to absorb from meat, fish and poultry.
Procedure:
You'll need a fairly sensitive scale for this procedure. A bathroom scale won't cut it!
1. Crush 1/2 cup of cereal in a baggie, until the flakes are half their original size. Pour into a bowl.
2. Add 1 cup of hot water and mix with a wooden spoon.
3. Get a strong, 3-inch bar magnet that is not grey or black (so the iron filings will show up). Don't use a horseshoe magnet.
4. Put the magnet into the cereal mix and stir gently in a circle for a fixed amount of time, say 5 minutes. Try not to bump the bottom or sides of the bowl.
5. Take out the magnet. Remove the iron filings that it pulled from the cereal, and weigh them on a laboratory scale.
3. Are all apples equally sweet?
As apples ripen, the starch in the fruit changes to sugar, making the fruit sweet. What kind of sweet differences are there between apple varieties or individual apples of the same type?
Background Info: Starch levels in apples can be measured by dipping a portion of the apple into an iodine solution. The starch reacts with the iodine solution to produce a blue-black color in a pattern that is characteristic for each variety of apple. For example, Red Delicious apples lose starch in a fairly even ring, while Golden Delicious apples have an uneven pattern.
Recipe for the Iodine Solution:
Always make up a fresh iodine solution. Keep this solution in a dark-colored (or foil-wrapped) bottle and away from light. Since iodine is poisonous, treated apples should also be considered poisonous and should not be eaten by people or animals or used in composting. Do not allow pets to lick the fruit after testing.
Dissolve 10 grams (about 1/3 ounce) of potassium iodide in 10 ml (approximately 1/8 cup) of water. When it is dissolved, add 2.5 grams (about 1/12 ounce) of iodine crystals. Shake the mixture until the crystals are thoroughly dissolved. Dilute this mixture with water to make 1.14 liters (1 quart).
Make sure the solution is completely mixed by shaking the bottle everyday for several days.
Procedure:
1. It is best to test fresh apples that have not been stored, so this experiment is best done in the fall. Another way to use this test is to track apple ripening from a single tree over the harvest season to pinpoint the best time to harvest that tree's apples.
2. Pour the iodine solution into a shallow glass container to a depth of 5 to 7.5 mm (approx. 1/4 inch). Cut each apple in half horizontally across the core and put the exposed surface of one of the halves in the iodine solution. The apple stem can serve as a convenient handle, if the top half is used.
3. Wait 1 minute before removing the apple half. Repeat with the next test apple.
4. Compare the patterns of black spots, which indicate the presence of starch.
5. When there is no reaction and no color change, all of the starch has changed to sugar.
Which color light is best for plant growth?
Build a box with three separate compartments and different colored light bulbs--white, red and green--then plant seeds in each compartment. Water the seeds and watch to see which plant grows better.
Does acid rain affect the cell structure of Spirogyra?
Fill six fishbowls with two liters of distilled water each and Spirogyra cultures. Allow the cultures to grow for 10 days. Leave two of the bowls with distilled water only. Add 3cc of acid and 6.0pH to two bowls, and 12cc of acid and 3.0pH to the remaining two bowls. Take samples of Spriogyra after 24 and 48 hours. Record the differences in cell structure.