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RIBBONS OF BLUE/WATERWATCH WA
WATERCHEMISTRY
SERIESThe chemistry behind water
quality testing
Prepared byStephanie Degens
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Editor/compiler Stephanie Degens. Community Monitoring and Environmental Officer
Ribbons of Blue/Waterwatch WA
Laboratory exercises author Allan Knight
Acknowledgments:
Ribbons of Blue/Waterwatch WA Regional Coordinators and their TEE chemistry
teachers: For reading draft documents and trialing them with students.
Staff from the Water and Rivers Commission. For comments on the draft documents.
Water and Rivers Commission, Swan-Canning Cleanup Program and Swan River Trust:
For support of the Ribbons of Blue/Waterwatch WA program.
For more information contact:
State Facilitator Ribbons of Blue/Waterwatch WA
on 9278 0300 or visit our web site www.wrc.wa.gov.au/ribbons
ISBN: 0-7309-7593-2
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CONTENTS
Electrical Conductivity of Aqueous Solutions Electrical Conductivity in Aquatic Systems 1
Measuring Conductivity Background 3
Experimental 5
Processing of results, and questions 5
Dissolved Oxygen Oxygen in Aquatic Systems 7
The analysis of Dissolved Oxygen in Water
Background 9
Experimental
Preparation of primary potassium bi-iodate solution 10 Preparation and standardisation of a sodium thiosulfate solution 11
Processing of results, and questions 12
Analysis of dissolved oxygen in a natural water sample 13
Processing of results, and questions 14
Biological Oxygen Demand 14
Considerations 15
pH and Aquatic Systems pH Buffering of Aquatic Systems 17
Factors Affecting pH of Aqueous Systems 18 pH Measurement
Background 19
Experimental 20
Processing of results, and questions 20
Plant Nutrients - Nitrogen Nitrogen in water 23
Nitrogen cycle 23
Analysis of nitrate/nitrite in a natural water sample
Background 26
Experimental 28
Processing of results, and questions 28
Plant Nutrients Phosphorus Phosphorus in water 30
Analysis of phosphate in a natural water sample
Background 32
Experimental 34
Processing of results, and questions 34
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ELECTRICAL CONDUCTIVITY OFAQUEOUS SOLUTIONS
TEE Year Subject Objectives fromChemistry Syllabus (2000-2001)11 1.38, 1.49, 1.L.6,
7.15
12 1.2, 2.L.1,
Electrical Conductivity in Aquatic Systems
The electrical conductivity of aqueous solutions is dependent upon the ion
concentration of the solution. Ions are free to move within the body of water and thus
are capable of carrying a current. The greater the ion concentration the higher the
conductivity. Note: Conductivity also increases with temperature.
Conductivity is used as a measure of salinity, the salt content of the water. Sodium
chloride is the main contributor to water salinity but other salts present may include
calcium carbonate (limestone) and other calcium and magnesium salts. The
conductivity/salinity can fluctuate naturally, especially in estuarine systems, however
it can be abnormally increased by human activity in a range of ways.
Overuse of fertilisers, causing leaching from the soil, can lead to an increase in the
concentration of phosphate, nitrate and ammonium ions (and their associated
cations/anions) thus leading to increased conductivity.
Excessive removal of native deep-rooted vegetation is of major concern in Western
Australia (as well as other parts of Australia). Removal of vegetation has caused a
rise in the level of underground water tables in many areas. As the water table rises it
dissolves salt that was previously held in the soil and brings it closer to the surface. If
the watertable rises to such an extent that it comes into contact with surface water the
salt can eventually enter rivers and streams and increase their salinity. That is reduce
the freshness of the water. This is termed as dryland salinity. This can have adverse
effects on the organisms depending upon this water for their well-being. The rising
water table can also cause an increase in bore water salinity.
The salinity of the Swan-Canning River system varies from fresh-to-brackish
conditions in the winter and spring to salty during the summer and autumn. When
river flow declines at the end of winter, sea water moves progressively up the estuary
reaching the upper Swan Maylands area and the Kent Street Weir in spring to early
summer. (In low rainfall years significant quantities of salty water often remains in
the system over winter.) The salt water flows in as a wedge beneath the less dense
fresh water (see figure 1). This layering process is called stratification.
The difference in densities reduces mixing of surface and bottom waters thus
preventing oxygen replenishment of the bottom waters. The oxygen concentration in
the bottom waters is further lowered as decomposition of organic matter consumesremaining oxygen. (This then causes the release of nutrients, such as nitrate and
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phosphate, stored in the sediments. This occurs because the solubilities of these
parameters tend to be higher in water with low oxygen concentration. These
dissolved nutrients accumulate in the stagnant salty waters at the bottom.)
Figure 1. Salt wedge in the Swan River (Swan-Canning Cleanup Program Action
Plan, 1999 11)
Every spring a dense wedge of salty water moves upstream while lighter fresh water
flows over the top as the summer progresses. The wedge moves upstream accordingto tidal and barometric influences, and freshwater flows reduce. The leading edge of
the salt wedge often has very low oxygen and high nutrient levels. (Swan-Canning
Cleanup Program Action Plan, 1999 11)
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Measuring Conductivity
Background
Conductivity can be measured relatively accurately with a conductivity meter. The
instruments measure the amount of electrical charge passing between two metalelectrodes 1 cm apart.
The units of salinity/conductivity can be expressed as microsiemens per centimetre
(S/cm), millisiemens per centimetre (mS/cm) or millisiemens per metre (mS/m). Be
aware of the units in which your instrument displays its readings. The total soluble
salts, expressed in milligrams per litre (mg/L), can be calculated from the following
standard conversion between different units:
Electrical conductivity (S/cm) 0.55 = total soluble salts (mg/L)
Electrical conductivity (mS/cm) 550 = total soluble salts (mg/L)
Electrical conductivity (mS/m) 5.5 = total soluble salts (mg/L)
As conductivity increases with temperature, it is necessary to compensate for
temperature variations between measurements.
Table 1 shows ranges of water electrical conductivity (EC) and restrictions that EC
places on water use.
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Table 1. Electrical conductivity ranges
Electrical
conductivity (EC)
Characteristics of the water EC ranges, restrictions for
usage and specific ECs.
0 - 800 S/cmor
0 - 440 mg/LFresh
Good drinking water for humans.
Generally good for irrigation, though above 300 S/cm somecare must be taken if using overhead spraying of salt sensitive
plants.Suitable for all livestock.
Distilled water 0 S/cm.
Rain water 66 S/cm.
800 2 500 S/cm
or
440 1 375 mg/L
Mod
erateBrackish Can be consumed by humans, though most would prefer water
in the lower half of this range.
When used for irrigation it requires special managementincluding suitable soils, good drainage and consideration of salttolerance of plants.Suitable for all livestock.
Tap water 400 S/cm.Maximum for hot water systems 1 600 S/cm.
Maximum for human drinking water 2 500 S/cm.
2 500 10 000
S/cm
or
1 375 5 500
mg/L
VeryBrackishTo
Salty
Not recommended for human consumption although water of up
to 3 000 S/cm could be drunk if nothing else was available.Not normally suitable for irrigation, though water of up to 6 000
S/cm can be used on very salt tolerant crops with special
management. Over 6 000 S/cm, occasional emergencyirrigation may be possible with care.
When used for drinking water by poultry and pigs, the electrical
conductivity should be limited to about 6 000 S/cm.
Most other livestock can drink up to 10 000 S/cm.
Over 10 000S/cm
or
over 5 500 mg/L
Salty
Not suitable for human consumption or irrigation.Not suitable for poultry, pigs or lactating animals, but beef
cattle can use water up to 17 000 S/cm and adult sheep on dry
feed can tolerate 23 000 S/cm. However chemical analysisshould be considered before using high electrical conductivitywater for stock.
Water up to 50 000 S/cm (the salinity of the sea) can be used(i) to flush toilets provided corrosion in the cistern can be
controlled and (ii) for making concrete, provided thereinforcement is well covered.
Pacific Ocean 58 000 S/cm.
Dead Sea 550 000 S/cm.
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Experimental
Equipment
Water sample(s)Thermometer (or temperature probe)
Conductivity meter
Procedure
Note: The conductivity can be measured in the field or water samples collected and
measurements made in the laboratory.
1. Your Ribbons of Blue/Waterwatch WA Regional Coordinator will support you
with information on water sampling techniques and help you to use the method
appropriate for your water-body to collect a sample(s).
2. Calibrate your conductivity meter as described in the instructions supplied with
your meter or by your Regional Coordinator.
3. Test and record the EC of the Mystery Solution supplied by Ribbons of Blue, to
ensure the equipment is working accurately.
4. Measure and record the conductivity of the water sample(s).
Processing of results, and questions
1. Calculate the total soluble salts (mg/L) using the conversion factors provided on
page 3.
2. Assess the water quality based on salinity using the information provided in
Table 1.
3. Suggest the ion species most likely to be responsible for the electrical conductivity
of your water sample.
4. At what time of year would you expect the salinity of your waterbody to belowest? Why?
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DISSOLVED OXYGEN
TEE Year Subject Objectives from
Chemistry Syllabus (2000-2001)
11 1.44,12 2.48, 2.60, 2.61, 2.62, 2.63, 2.72, 2.73,
5.2, 5.5, 5.6, 5.7, 5.35, 5.36, 5.37, 5.39,
Oxygen in Aquatic Systems
The solubility of gases in liquids can be studied in the context of oxygen in natural
and man made waterbodies (lakes, rivers, streams, drains etc.). Factors affecting the
solubility of oxygen in water include:
Temperature increasing solubility with decreasing temperature
Atmospheric pressure (altitude) the greater the pressure the higher the solubility
Salt concentration the lower the salt concentration the higher the oxygen
concentration
The sources and consumption of oxygen in water bodies also influence the amount of
dissolved oxygen. The sources are:
Absorption There is continuous exchange of oxygen between water and the
surrounding air. The greater the contact between the water and the air the more
oxygen that can dissolve, thus a turbulent stream will tend to have a higher oxygen
concentration than a still body of water.
Photosynthesis This redox process carried out by aquatic (and land) plants
results in oxygen directly entering the water. Those things that reduce the amountof sunlight able to penetrate the water (e.g. suspended solids) will lower the rate of
photosynthesis and hence lower oxygen concentration. As photosynthesis takes
place only during the day, the concentration of dissolved oxygen will vary over
the 24-hour daily cycle. Levels peak early afternoon and are lowest just before
sunrise.
Oxygen is consumed by:
Respiration all organisms (aquatic or terrestrial) consume oxygen during
respiration.
Decomposition the decomposition of plant and animal waste (whether from
living or dead organisms) is carried out by bacteria and other micro-organismsthat use oxygen to oxidise the organic matter.
The level of organic wastes present in a water sample can be estimated by measuring
oxygen consumption in the water. The biological oxygen demand (BOD) provides a
measure of the amount of oxygen consumed by a sample of water under standard
conditions. The oxygen concentration in the water is measured immediately on
collection and again after the sample has been incubated at 20C for 5 days. The
difference in oxygen concentrations is given in units of milligrams per litre (mg/L) or
parts per million (ppm) and is an indication of the quantity of organic wastes in the
sample.
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The low solubility of oxygen in water can be explained by the nature of the solvent
and solute. Water has a polar molecule whilst the oxygen molecule is non-polar. The
predominant type of molecular interactions between water molecules is hydrogen
bonding and between oxygen molecules are dispersion forces. Thus there will be
little interaction between water molecules and oxygen molecules, and so the solubility
of oxygen in water is low.
Oxygen is essential for the survival of aquatic organisms, and a shortage of dissolved
oxygen is not only a sign of pollution, it is harmful to fish. The sensitivity of aquatic
species to oxygen depletion varies, but some general guidelines to consider when
analysing test results are:
5 6 ppm Sufficient for most species
< 4 ppm Stressful to most aquatic species
< 2 ppm Fatal to most species
Work by the Waters and Rivers Commission to increase the oxygen concentration in
the Canning River has included an oxygenation technique involving the injection of
oxygen rich water into the colder deeper water (see Figure 2).
Figure 2. Oxygenation plant on the Canning River (Swan Canning Cleanup Program
Action Plan, 1999 53)
The purpose of oxygenation is to pump low-oxygen water from the bottom of the river
to a mixing plant to increase oxygen content and then return the treated water.
Oxygenation thus specifically targets the low-oxygen layer of the river and does not
interrupt natural layering caused by salinity and temperature differences. BOC gases
donated material and expertise for a trial conducted on the Canning River in 1997-
98. In the Thames River (London), authorities utilise a barge that injects oxygen
directly to the bottom of the river to solve the problem of poor water quality followingstorm and pollution events. (Swan Canning Cleanup Program Action Plan, 1999 53)
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The analysis of dissolved oxygen in water
Background
The analysis of natural water samples for dissolved oxygen can be done using a redox
titration known as the Winkler method. In this procedure, the water sample is firsttreated with an excess of manganese II, potassium iodide and sodium hydroxide. The
white manganese II hydroxide that initially precipitates is oxidised readily by
dissolved oxygen to give the brown manganese III hydroxide. The reactions are
Mn2+(aq) + 2OH-(aq) Mn(OH)2(S)
4Mn(OH)2(S) + O2 + 2H2O 4Mn(OH)3(S)
When acidified, the manganese III hydroxide dissolves and the freed manganese III
ion oxidises iodide to iodine.
Mn(OH)3(S) + 3H+
(aq) Mn3+
(aq) + 3H2O(l)
2Mn3+(aq) + 2I- I2 + 2Mn
2+(aq)
The liberated iodine can then be titrated with thiosulfate (S2O32-).
2 S2O32-
(aq) + I2(aq) S4O62-
(aq) + 2I-(aq)
The sequence of laboratory activities involved in the Winkler analysis is as follows:
I. Preparation of potassium bi-iodate solution as a primary standardII. Preparation and standardisation of a sodium thiosulfate solution
III. Analysis of the dissolved oxygen in a natural water sample
Table 2 gives an indication of what your dissolved oxygen results mean for the
waterbody sampled.
Table 2. Rating of dissolved oxygen ranges in natural waterbodies
DISSOLVED OXYGEN (percent saturation)
in standing or flowing water
Normal Some
Pollution
High
Pollution
80 120 120 135
or
55 80
> 135
or
< 55
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Experimental
Preparation of primary potassium bi-iodate solution
Equipment
Drying oven
Beaker (250 mL)
Desiccator
Balance
Volumetric flask (1.00 L)
Wash bottle
Storage bottle (1.00 L, dark glass)
Deionised water
Primary standard potassium bi-iodate [KH(IO3)2]
Procedure
1. Calculate the mass of anhydrous KH(IO3)2 needed to make up 1.00 L of
approximately 8.35 10-4 mol/L solution.
2. Place a little more than this calculated mass in an oven at 105 C for 1 hour. After
drying place the KH(IO3)2 in a desiccator to cool.
3. Accurately weigh out into a beaker a mass of KH(IO3)2 approximately equal to
that calculated in step 1.
4. Dissolve the solid in about 100 mL of water and then transfer the solution to the
volumetric flask and make up to the graduated mark. Stopper the flask and mix
the solution thoroughly.
5. Transfer the solution to the clean dark glass storage bottle and store in a dark
cupboard until required.
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Preparation and standardisation of a sodium thiosulfate solution
Equipment
Bottle (250 mL)
Beaker (250 mL)Volumetric flask (1.00 L)
Conical flask (250 mL)
Graduated cylinders (10 mL and 100 mL)
Burette and stand
Funnel
Pipette (10 mL)
Pipette filler
Deionized water
Alkaline-iodide-azide reagent (1.0 mL)
Concentrated sulfuric acid [H2SO4] (1.0 mL)
Starch indicator solution (usually VITEX in deionized water)Sodium thiosulfate pentahydrate [Na2S2O3.5H2O] (~5 g)
Sodium carbonate [Na2CO3] (0.1 g)
Manganese II sulfate solution (1.0 mL)
Alkaline-iodide-azide Reagent
The alkaline-iodide-azide reagent is prepared as follows (probably best done by the
laboratory technician prior to the laboratory activity):
1. Dissolve, in small increments, 400 g of sodium hydroxide pellets in 500 mL of
freshly boiled deionized water.
2. Add 900 g of sodium iodide while the solution is still hot.
3. Dissolve 10 g of sodium azide, NaN3, in 40 mL of deionized water.
4. Add the sodium azide solution to the first solution (step 2) and make up to 1.0 L.
Manganese II sulfate solution
The manganese II sulfate solution is prepared as follows (probably best done by the
laboratory technician prior to the laboratory activity):
1. Dissolve 365 g of manganese II sulfatemonohydrate, MnSO4.H2O, in freshly
boiled deionized water and dilute to 1.00 L.
Note: Manganese II sulfatedihydrate (400 g) or tetrahydrate (480 g) can be used.
Procedure
1. Dissolve approximately 4.8 g of Na2S2O3.5H2O and 0.1 g of Na2CO3 in about 100
mL of deionized water in a beaker.
2. Transfer to a 1.00 L volumetric flask and make up to the graduated mark. Mix thesolution thoroughly.
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3. Fill a 250 mL bottle with deionized water (NOT chlorinated tap water). To this
add, mixing after each addition,
1.0 mL alkaline-iodide-azide reagent,
1.0 mL concentrated sulfuric acid, and
1.0 mL manganese II sulfate solution.
4. Transfer 100 mL of this solution to a conical flask and pipette a 10 mL aliquot of
the potassium bi-iodate solution into the flask.
5. Allow the mixture to stand for 2 minutes in the dark and then titrate the liberated
iodine with thiosulfate solution until a very pale straw colour develops in the
solution. Add four drops of starch indicator solution and continue titrating until
the blue colour just disappears.
Processing of results, and questions
1. In the preparations of the alkaline-iodide-azide reagent and manganese II sulfate
solution why is it necessary to use boiled deionized water?
2. From the volume of thiosulfate solution used, calculate its concentration.
3. Why is it not acceptable to use the sodium thiosulfate as a primary standard?
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Analysis of dissolved oxygen in a natural water sample
Equipment
Biological Oxygen Demand Bottle or ordinary glass stoppered bottle (250 mL)
Conical flask (250 mL)Graduated cylinder (10 mL)
Burette and stand
Funnel
Pipette (20 mL)
Pipette filler
Deionized water
Manganese II sulfate solution (1.0 mL)
Alkaline-iodide-azide reagent (1.0 mL)
Concentrated sulfuric acid [H2SO4] (1.0 mL)
Sodium thiosulfate pentahydrate [Na2S2O3.5H2O] (~5 g)
Starch indicator solution (usually VITEX in deionized water)
Note: If you wish to carry out a BOD measurement collect sufficient water to obtain
two sets of results 5 days apart.
Procedure
Note:When adding solutions to the water sample place all pipettes below the water
surface to avoid agitation.
1. Collect your water sample(s). To avoid contamination, thoroughly rinse the
collection bottle with sample water. It is important to avoid trapping air bubbles
in the bottle. Tightly cap the bottle and submerge to the desired depth. Remove
the cap to fill the bottle to the top and then stopper or seal it whilst still
submerged. Once the sample is collected the bottle should only be opened when
the sample is to be analysed.
2. To the collected natural water sample, add 1.0 mL of manganese II sulfate
solution followed immediately by 1.0 mL of alkaline-iodide-azide reagent. (Some
overflow of water will occur.) Re-stopper the bottle at once, ensure no air is
trapped, and shake vigorously for at least 20 seconds or until the precipitatedmanganese II and manganese III hydroxide is evenly distributed.
3. Let the bottle stand for 2 3 minutes and then shake again.
4. Allow the precipitate in the bottle to settle by at least 1/3 (10 20 minutes).
5. Add 1.0 mL of concentrated, reagent grade H2SO4 well below the surface. (The
overflow will be alkaline so avoid contact with your skin.) Re-stopper the flask
and shake gently until dissolution of the precipitate.
6. Using a pipette, transfer 100 mL of the solution to a conical flask.
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7. Titrate at once with the previously standardised thiosulfate solution until a very
pale straw colour appears.
8. Add 4 drops of starch indicator solution, and continue the titration until the blue
colour just disappears.
Processing of results, and questions
1. When collecting the water sample why is it important to ensure that no air is
trapped in the bottle?
2. Which ion species in this process reacts directly with any oxygen dissolved in the
water?
3. Suggest a reason why it is necessary to shake the bottle in step 1 of the procedure
and then allow it to stand for 2 3 minutes as stated in step 2?
4. Calculate the number of moles of thiosulfate used in the titration?
5. What is the mole relationship between the thiosulfate and the oxygen?
6. How many moles of oxygen are present in 100 mL of your water sample?
7. What is the concentration of your dissolved oxygen in mg/L and ppm?
8. What volume of oxygen would this be at S.T.P.?
9. Rate the health of the water source from which your sample was taken.
10. Why would the accuracy of this technique for measuring dissolved oxygen be
reduced by the presence of iron II (Fe2+) in the water sample?
Biological Oxygen Demand (BOD)
To obtain an indication of the amount of organic waste in your water sample by the
BOD measurement, store your water sample at 20C for 5 days and then analyse for
oxygen as described above. The difference between your two measurements is the 5
day BOD of your water. The oxygen consumed during this 5 day period has beenused by bacteria and other micro-organisms in the oxidation of organic wastes in the
water.
Note: If you are unable to store the water sample at 20C, it can be kept at ambient
temperature and the BOD measured for this temperature.
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Considerations
This method is not applicable under the following conditions:
1. Samples containing more than 1 mg/L of iron II. This can be overcome by the
addition of potassium fluoride solution prior to acidification. Potassium fluoridesolution is made by dissolving 40g of KF.2H2O in distilled water and diluting to
100 mL.
2. Samples containing sulfite, thiosulfite, free chlorine or hypochlorite.
3. Samples with high concentrations of suspended solids. Filtration can remedy this.
4. Samples containing other oxidising or reducing ion species.
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pH AND AQUATIC SYSTEMS
TEE Year Subject Objectives from
Chemistry Syllabus (2000-2001)
11 4.1, 4.3, 4.4, 4.5, 4.12, 4.17,12 3.17, 3.18, 3.21, 3.L.2,
4.4, 4.5, 4.6, 4.9, 4.13, 4.19.
pH Buffering of Aquatic Systems
The pH of natural fresh waters is usually about 7 but values may vary from ~5.0 8.5
depending upon factors such as soil type, geology and vegetation. Marine waters
have a pH close to 8.2. Control of pH in many waters is achieved by the carbonate
hydrogencarbonate buffer system. A buffer solution is one that maintains the
approximate pH of a solution when a small amount of an acid or a base is added to the
solution. A buffer solutions function is an application of the equilibrium principles as
expressed in Le Chateliers Principle.
The buffering of natural waters by the carbonate hydrogencarbonate system initially
involves the dissolution of atmospheric carbon dioxide (or CO2 released by aquatic
organisms into the water) to give carbonic acid. The reaction for this can be
represented as
H2O(l) + CO2(aq) H2CO3(aq)
The carbonic acid will ionize to a small extent to give hydronium andhydrogencarbonate ions:
H2O(l) + H2CO3(aq) H3O+
(aq) + HCO3-(aq)
The hydrogencarbonate ion will then hydrolyse as represented below:
H2O(l) + HCO3-(aq) H3O
+(aq) + CO3
2-(aq)
Addition of acid to the water results in an increase in hydronium ion concentration
and as predicted by Le Chateliers principle the reaction favours the reverse direction.
Conversely, the addition of a base would result in a reduction of hydronium ionconcentration and so the reaction will shift to the products.
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Factors Affecting pH of Aqueous Systems
Many manufacturing and chemical industries release waste water into ground or
surface waters in the environment at some stage of their process. The pH of this
wastewater should, ideally, be the same as the pH of the natural water system into
which it is being released.
The presence of certain salts can influence the pH of aqueous systems. Some of these
are present naturally in water but others are pollutants. Salts that are naturally present
can have an unwanted effect on pH if their concentrations are increased by the actions
of people. Hydrolysis of ionic compounds to give either hydronium ions or hydroxide
ions can alter the normal pH of a natural aquatic system. People using inorganic
phosphate fertilisers or the release of phosphate containing detergents can increase
phosphate concentrations in waterbodies near to farming or residential areas. The
hydrolysis of phosphate ion can increase the pH of the water as illustrated below:
H2O(l) + PO43-(aq) HPO42-(aq) + OH-(aq)
The pH can also be increased by the presence of ammonia.
H2O(l) + NH3(aq) NH4+
(aq) + OH-(aq)
The application of ammonium based fertilisers can lead to a decrease in pH due to
hydrolysis of the ammonium ion.
H2O(l) + NH4+
(aq) NH3(aq) + H3O+
(aq)
Table 3 gives an indication of what your pH results mean for the waterbody sampled.
Table 3. Rating of pH ranges for natural waterbodies
pHin standing or flowing water
Normal May be
Polluted
Pollution
Problem
5.0 7.0 (no
limestone)
7.0-8.5 (limestone)
8.5 9.0
or
4.0 5.0
< 4
or
> 9
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pH Measurement
Background
The pH of an aqueous solution can be measured with the use of pH paper, acid-base
indicators, such as universal indicator, or pH meters or scans. In this activity, the pHof a natural water sample will be measured using a pH meter.
When using a pH meter it is first necessary to calibrate the meter. Calibration
involves placing the pH probe into a buffer solution of known pH to ensure the meter
is reading values accurately. Conventionally two buffer solutions are chosen such that
their pH values span the pH range of interest. For example, if the test sample were
thought to have a pH of about 6, typical buffer solutions used for calibration would
have values of, say, pH 4 and 9. The pH meter would be placed in the pH 4 buffer
solution and, if necessary, the reading on the meter would be adjusted to 4. The pH
probe would then be similarly placed in pH 9 buffer solution and the reading adjusted
if required.
A typical pH 4 buffer is a 0.05 mol/L solution of potassium hydrogenphthalate
[KHO2CC6H4CO2 ] (A salt of an aromatic carboxylic acid. See Figure 3 below.) A
typical pH 9 buffer is sodium tetraborate (borax) [Na2B4O7.10H20].
Figure 3: Potassium hydrogenphthalate
O
C OH
C OK
O
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Experimental
Equipment
pH meter
pH buffer solutions (pH 4 and 9)2 beakers (100 mL)
bottle (250 mL)
water samples from your local environment
Procedure
Note: Remember to rinse the pH electrodes with distilled water before each
measurement.
1. Calibrate your pH meter as follows:
I. Rinse a clean beaker with a small quantity of the pH 4 buffer solution.
II. Pour sufficient pH 4 buffer solution into the beaker to cover the glass electrode
of the meter to the correct depth. (Your teacher will show you this.)
III. Rinse the pH probe with pH 4 buffer solution and then place the electrode into
the buffer solution and check the reading on the meter. If necessary adjust
the reading to a value of 4.
IV. Repeat steps I III for the pH 9 buffer solution.
2. Rinse a clean beaker with a small quantity of your water sample and then pour in
sufficient quantity of your water sample to cover the glass electrode to the correct
depth. Rinse the pH probe with sample water and then place it into the water,
measure and record the pH value.
3. Repeat step 2 until you have measured the pH for all your water samples.
Processing of results, and questions
1. Write the hydrolysis equation for the phthalate ion, HO2CC6H4CO2-, in the pH 4
buffer solution (Hint: Remember the solution is acidic.) Using the appropriatechemical principles, explain why the addition of a small quantity of acid or base to
this solution should not greatly alter its pH.
2. Write the hydrolysis equation for the tetraborate ion, B4O72-, from the pH 9
solution. Using the appropriate chemical principles, explain why the addition of a
small quantity of acid or base to this solution should not greatly alter its pH.
Note: The following questions can probably only be answered if students have access
to any long term results of pH measurements made of waterbodies in their local
environment or they have made measurements of phosphate, ammonia or ammonium
levels in the water samples.
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3. Try to account for any variations in the pH of your water samples. Use
appropriate equations to explain any of your results.
4. Assess the quality of your water sample based on pH using the information
provided in Table 3.
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PLANT NUTRIENTS - NITROGEN
TEE Year Subject Objectives from
Chemistry Syllabus (2000-2001)
11 5.1, 5.2, 5.3, 5.4, 5.5, 5.6,12 5.1, 5.2, 5.3, 5.5, 5.6, 5.7,
6.5, 6.14
Nitrogen in water
Nitrogen is a part of living organisms in their protein, DNA and other components. It
is released naturally into water with the decay of dead organisms and their wastes and
through the fixation of atmospheric nitrogen. Humans introduce nitrogen into
waterbodies through a number of sources including:
fertilisers
plant and animal wastes
rural and urban run-off
sewage effluent
industrial discharges
The total nitrogen present in the water at any given time is composed of organic
nitrogen and inorganic nitrogen. Organic nitrogen is often associated with biological
material and is not soluble and consequently is less mobile that inorganic nitrogen.
Inorganic nitrogen is made up of dissolved or particulate nitrate, nitrite and
ammonium. The majority of nitrate occurs in dissolved form due to the high
solubility of nitrate compounds. This contrasts to the phosphates that have lowsolubility and so often occur in a particulate form.
Nitrogen cycle
A large amount of nitrogen enters the Swan-Canning system in dissolved form. This
provides nutrients for spring algal blooms that take up the nitrate and convert it to
organic nitrogen as algal cells. After spring these algae die and fall to the sediment on
the river floor. In the sediment, decomposition by microbes converts the nitrogen to
ammonia and then nitrate. This ammonia and nitrate is now available for uptake by
more algae in the summer, which subsequently die, fall to the bottom and decompose.
This cycle continues throughout summer.
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Figure 4. Nitrogen cycle in a freshwater system
Biodegradation of organic nitrogen compounds produces ammonium ion as the first
inorganic species. Nitrificationis a two-stage redoxprocess in which micro-
organisms convert ammonium ion to nitrate ion. In the first stage the ammonium is
oxidised to nitrite ion.
2NH4+
(aq) + 3O2(g) 2NO2-(aq) + 4H
+(aq) + 2H2O(l)
The nitrite ion is then oxidised to nitrate:
2NO2-(aq) + O2(g) 2NO3
-(aq)
This process uses considerable quantities of oxygen and contributes significantly to
reducing the amount of oxygen available to other aquatic organisms.
Nitrate concentration can also increase in summer. The incoming salt wedge reduces
the amount of dissolved oxygen in bottom waters. Microbes release nitrogencompounds under these conditions.
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Nitrogen can only be lost from a water system in the following ways:
denitrification, - where nitrate is reduced to N2 which then escapes into the
atmosphere
uptake by living organisms
burial in sediments
export to the ocean
Denitrification is when bacterial enzymes catalyse the reduction of nitrate to nitrogen
gas by a reducing agent such as carbohydrate (Sugars, such as glucose [C6H12O6], are
an example of carbohydrates.):
24NO3-(aq) + 5 C6H12O6(aq) + 24H
+(aq) 12N2(g) + 30CO2(g) + 42H2O(l)
The nitrogen gas can then escape to the atmosphere.
The processes of nitrification and denitrification as well as movement of nitrogen
through a freshwater system are illustrated in Figure 4.
Nitrogen is also found in groundwater. It percolates through the soil and collects in
aquifers, most commonly as nitrate. Some stream systems and wetlands are
groundwater fed. Groundwater nitrogen can therefore enter the surface water system.
Table 4 gives an indication of what your nitrate results mean for the waterbody
sampled.
Table 4. Rating of nitrate content in waterbodies
Nitrate Content in milligrams per litre (mg/L)Water Type
Low Medium HighSTANDING 0 0.025 0.025 0.25 > 0.25
FLOWING 0 0.05 0.05 0.4 > 0.4
TANNIN 0 0.25 0.25 1.5 > 1.5
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Analysis of nitrate/nitrite in a natural water sample
Background
The free (dissolved) nitrate tests used in water testing kits today are based on a
method developed in the late 1800s. The concentration of free nitrate in a watersample can be determined via a colorimetric technique. This technique involves
converting colourless nitrate to a pink coloured compound by reacting a series of
chemicals with a filtered water sample. The intensity of colour produced is directly
proportional to the amount of free nitrate in the water sample. The colour of the
tested water sample can be compared to a set of known colour standards and thus the
concentration of the nitrate estimated. Alternatively, the amount of nitrate in a sample
can be determine using a photometer and comparing light transmittance to a
calibration chart.
Free nitrate tests should be carried out on the filtrate from water samples passed
through 0.45 m filter paper. However, due to costs, Ribbons of Blue uses GFC filter
papers which filter out particles greater than 1.2 m
Note: The test described below actually measures both the nitrate and nitrite in the
sample but the nitrite is usually at such a low concentration that the results are fairly
accurate. The combination of the nitrate and nitrite is known as the total oxidised
nitrogen (TON).
The first step in the process is the reduction of nitrate ion to nitrite by the addition of
cadmium or zinc to the water sample. Below is the reaction for cadmium with nitrate.
NO3-(aq) + Cd(s) + 2H
+(aq) NO2
-(aq) + Cd
2+(aq) + H2O(l)
Palintest Kit uses Nitratest Powder containing 70% zinc as the reducing agent.
The LaMotte Kit uses Nitrate Reducing Agent that contains 7% Cadmium, which is added
after the Mixed Acid Reagent.
The nitrite produced thus is determined by a diazotisation reaction. In acid solution,
nitrite ions and aromatic amines react to form diazonium salts.
aromatic amine diazonium salt
In the Palintest Kit, sulfanilic acid, contained in the Nitricol Tablets acts as the aromatic amine.
The acid solution is provided by the ammonium chloride, from the Nitratest Tablets, whichdissociates to form a weak acid, and the sulfanilic acid.
The LaMotte Kit method has already acidified the water sample by adding the Mixed Acid
Reagent containing Citric and Acetic Acid. The aromatic amine is sulfanilamide contained in the
Nitrate Reducing Powder.
NH2HO3S + NO2 + 2H+ N2+HO3S + 2H2O
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The diazonium salts can then in turn react with diaromatic amines to form intensely
colouredazo compounds. (Aromatic compounds are those that contain a benzene ring
and amines are characterised by the functional group NH2.) For example:
diaromatic aromatic azo compound
amine (highly colouredred in this case)
The Palintest Kit uses the diaromatic amine N-1-napthlethylene diamine dihydrochloride
contained in the Nitricol Tablet to form the azo-compound.
The LaMotte Kit also uses the diaromatic amine N-1-napthlethylene diamine dihydrochloride
contained in the Nitrate Reducing Agent to form the azo-compound.
N2+HO3S
+ NH2R NH2RN2HO3S + H
+
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Experimental
Equipment
Water sample(s)
Palintest Kit or other Nitrate testing kit
Procedure
Note:If you wish to measure the nitrite and nitrate concentrations separately you can
do so by omitting the addition of the nitrate reducing reagent to a sample of water.
Add only the second reagent to give the colour. An estimate of the nitrite
concentration can then be made by obtaining a percent transmission reading and
reading from the table of values. The nitrate concentration can then be obtained by
carrying out the full test as outlined in the test kit and subtracting the nitrite
concentration from this second value.
1. Collect your water sample(s). Be sure to rinse your bottles with the water to be
collected before obtaining your sample. Record the site(s) from which your
sample is taken. If possible also record the depth and temperature from which you
obtained your water.
2. Read the instructions in the test kit to determine the nitrate content of your water
sample. If possible test a number of samples at different sites of the waterbody
(this may be done as a class) to ascertain a profile of the waterbody.
Processing of results, and questions
1. Using the outlined procedure in your kit, state the concentration of the nitrate in
your water sample(s). Rate the concentration(s) using information from Table 4.
2. Suggest likely sources of nitrate for the waterbody you have tested.
3. What is the oxidation state of the nitrogen before and after reaction with the zinc
(or cadmium)? (The Palintest uses zinc as the nitrate reducing agent.)
4. Write a balanced redox equation for the reaction between zinc and nitrate ion.
5. One of the ingredients used in the reagent tablets for this test is an acid. Suggest a
reason for its inclusion?
6. Based on the value for the nitrate concentration you determined, calculate the
mass of zinc (or cadmium, as appropriate to the test kit used) needed to react with
this quantity of nitrate. Is the quantity of zinc (or cadmium) added to your water
sample likely to be more or less than your calculated value? Why?
7. In the nitrification process, ammonium ion is converted to nitrite ion, which in
turn is converted to nitrate. What are the oxidation states of the nitrogen in these
three species? Explain why it is possible for nitrogen to have this range ofoxidation states. (Hint: Consider the electron configuration of nitrogen.)
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8. In the denitrification process, nitrate is converted to nitrogen gas. Write a
balanced redox equation for this conversion where sucrose, C12H22O11, is the
reducing agent. (Assume the sucrose is oxidised to carbon dioxide.)
9. Why does most of the nitrogen gas produced in denitrification escape to theatmosphere rather than dissolve in the water?
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PLANT NUTRIENTS - PHOSPHORUS
TEE Year Subject Objectives from
Chemistry Syllabus (2000-2001)
11 1.36, 1.49,5.2, 5.3, 5.4, 5.5, 5.6
12 3.19, 3.20, 3.21, 3.L.2,
5.1, 5.2, 5.3, 5.4, 5.5, 5.6,
7.5, 7.6, 7.7, 7.14
Phosphorus in water
The total phosphorus content of a waterbody consists of three components.
Phosphorus can occur in the form of phosphate bound up in the organic matter
(organic phosphate), as dissolved phosphate free in the water (known as reactive
phosphate) or as particulate phosphate bound to clay or insoluble metal phosphatecompounds (particularly iron and aluminium oxides).
Phosphorus is one of the principal plant nutrients entering the Swan-Canning river
system. In most natural systems phosphorus is present as either organic or inorganic
phosphate. As no simple method exists for measuring phosphorus content, and it is
rare in its pure form, the following method is for analysis of dissolved phosphate.
This procedure will identify the amount of phosphorus in the water sample that exists
as phosphate ions (PO43-). Compounds such as calcium phosphate (Ca3(PO4)2) will
have dissociated in the sample after the addition of acid so this phosphate will also be
detected. We commonly work out phosphorus content by working out how much
phosphate occurs in a sample and then calculating how much of this is phosphorus
(since the phosphorus content of phosphate is always constant).
Natural sources of phosphate in an undisturbed system include weathering of rocks
and decomposition of organic matter. These natural sources lead to a low phosphate
concentration in the water. Human activity significantly increases the concentration
of phosphate. Human activities that cause phosphate to enter waterbodies include:
application of inorganic fertilisers, both on home gardens and agricultural lands
use of detergents organic wastes (ie. from garden, agricultural wastes and animal effluent)
sewage effluent
industrial discharges
The major problem that can arise from an excess of phosphates (and other plant
nutrients) in water is eutrophication. This is the process whereby the excess plant
nutrients cause an explosion in the algal population in a waterbody. When the algae
die it is decomposed by micro-organisms which consume much of the oxygen from
the water. This reduces the oxygen concentration to a level where fish and otheraquatic organisms have insufficient for their survival.
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(Some of the algae species present can have adverse health effects on humans. The
extent of these species is not a major threat in the Swan-Canning system as yet but
there have been occasions during the summer when health authorities have advised
against swimming in parts of the river where blooms have been present.)
Phosphate applied as fertiliser is usually soluble and it is as a result of fertiliser usethat most phosphate enters the Swan River. In most soils the phosphate is
immobilised by reacting with iron, manganese, calcium and aluminium ions to form
less soluble compounds (refer to Solubility Rules) or by being adsorbed onto the
surfaces of clay and silt particles. It is then said to be bound and so much less
mobile. Bound phosphate only enters waterbodies as part of soil carried into the
waterbodies. However, the sandy soils of the Swan Coastal Plain contain few binding
metals and little clay or silt. This results in a high concentration of dissolved
phosphate entering ground and surface waters.
Phosphate concentration is also influenced by the quantity of dissolved oxygen.
Under oxygenated conditions, phosphorus can be retained in sediment by theformation of iron, manganese, aluminium and calcium compounds. The lower
oxygen concentrations during summer caused by the salt wedge moving up the Swan
River (and generally higher temperatures and reduced flow rates) can result in release
of phosphate into the water column.
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Analysis of phosphate in a natural water sample
Background
The measurement of total phosphate content in the water requires the conversion of
the particulate forms into the dissolved reactive form. The conversion can be carriedout by acid hydrolysis at boiling water temperature. Acid hydrolysis involves mixing
concentrated acid with the water sample to react with the particulate forms of
phosphorus. For example, if there were suspended particles of calcium phosphate in
the water, reaction with concentrated acid would produce the soluble calcium
hydrogenphosphate. This is an equilibrium process whereby the hydrogen ions
initially react with free phosphate to give hydrogenphosphate:
H+(aq) + PO43-
(aq) HPO42-
(aq)
This reaction results in a lowering of the free phosphate ion concentration thus units
of the particulate phosphates, such as calcium phosphate, will dissolve:
Ca3(PO4)2(s) 3Ca2+
(aq) + 2PO43-
(aq)
Note: When we say that an ionic compound is insoluble, as in the Solubility Rules,
this is not strictly accurate. Those that we say are insoluble should more accurately be
described as having very low solubility. For any ionic compound in water some
dissociation will take place. An equilibrium is established between the solid and the
dissociated ions. Hence by adding excess acid to an insoluble phosphate, the
concentration of phosphate ions is reduced and the equilibrium is altered in such a
way as to cause the solid to dissolve.
It is not practical or safe to measure the total phosphate content in a water sample out
in the field. Therefore, Ribbons of Blue/Waterwatch WA only analyses for dissolved
phosphate. Acid hydrolysis is carried out at room temperature.
The dissolved phosphate tests used to analysis water samples are based on the
Ascorbic Acid Method. The concentration of dissolved phosphate in a water sample
can be determined via a colorimetric technique. This technique involves converting
colourless phosphate to a blue coloured compound by reacting a series of chemicals
with a filtered water sample. The intensity of colour produced is directly proportional
to the amount of dissolved phosphate in the water sample. The colour of the testedwater sample can be compared to a set of known colour standards and thus the
concentration of the phosphate estimated. Alternatively, the amount of phosphate in a
sample can be determine using a photometer and comparing light transmittance to a
calibration chart.
Dissolved phosphate tests should be carried out on the filtrate from water samples
passed through 0.45 m filter paper. However, due to costs, Ribbons of
Blue/Waterwatch WA uses GFC filter papers which filter out particles greater than
1.2 m
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The two acid hydrolysis reactions can be represented in one step as:
Ca3(PO4)2(s) + 2H+
(aq) 3Ca2+
(aq) + 2HPO42-
(aq)
The Palintest Kit uses potassium hydrogen sulphate from the Palintest Phosphate No 1 LR
tablet to acidify the water sample.
The LaMotte Kit uses sulphuric acid from the Phosphate Acid Reagent to acidify the watersample.
The concentration of phosphate in a filtered water sample can be determined via a
colorimetric technique. The colorimetric technique involves the reaction of the
hydrogenphosphate ion, in the presence of acid, with ammonium molybdate,
(NH4)6Mo7O24, to form ammonium molybdophosphate, (NH4)3[P(Mo3O10)4]. This is
a complex reaction and may be represented as follows (This is not a balanced
equation.)
MoO42-
(aq) + HPO42-
(aq) [P(Mo3O10)4]3-
(aq)
The Palintest Kit uses Ammonium Molybdate from the Palintest Phosphate No 2 LR tablet tofacilitate the above reaction.
The LaMotte Kit uses Ammonium Molybdate from the Phosphate Acid Reagent to facilitatethe above reaction.
The molybdophosphate ion, [P(Mo3O10)4]3-, can then be reduced with a species such
as acidified tin II or ascorbic acid (also known as vitamin C), molecular formula
C6H8O6. This reduction reaction produces the intensely blue-coloured phosphorus
molybdenum blue, (MoO2.4MoO3)2. H3PO4. The exhibiting of colour is a property
typical of many transition metal compounds. The reduction using tin II is represented
by the following equation:
[P(Mo3O10)4]3- + 11H+ + 4Sn2+
(MoO2.4MoO3)2. H3PO4 + 2MoO2 + 4Sn4+ + 4H2O
The intensity of the blue colour is proportional to the concentration of phosphate in
the filtered water sample.
Palintest Kit reduces the molybdophosphate ion using sodium metabisulphate from the Palintest
Phosphate No 2 LR tablet.
LaMotte Kit reduces the molybdophosphate ion using Ascorbic acid from the Phosphate
Reducing Agent.
Table 5 gives an indication of what your phosphate results mean for the waterbody
sampled.
Table 5. Rating of phosphate content in waterbodies
Phosphate Content in milligrams per litre (mg/L)Water Type
Low Medium HighSTANDING 0 0.005 0.005 0.05 > 0.05
FLOWING 0 0.01 0.01 0.1 > 0.1
TANNIN 0 0.05 0.05 0.2 > 0.2
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Experimental
Equipment
Water sample(s)
Palintest Kit or other Phosphate testing kit
Procedure
Note: If you wish to test for reactive phosphate separately from total phosphorus your
water sample must first be filtered. Standard chemical methods require a 0.45 m
cellulose nitrate filter paper, however due to the expense of these filter papers
Ribbons of Blue groups use 1.2 m GFC filter papers.
1. Collect your water sample(s). Be sure to rinse your bottles with the water to be
collected before obtaining your sample. Record the site(s) from which your
sample is taken. If possible also record the depth and temperature from which you
obtained your water.
2. Read the instructions in the test kit on how to determine the phosphate content of
your water sample. If possible test a number of samples at different sites of the
water-body (this may be done as a class) to ascertain a profile of the waterbody.
Processing of results, and questions
1. Using the outlined procedure in your kit, state the concentration of the phosphorus
in your water sample(s). Rate the concentration(s) using information from Table
5.
2. Suggest likely sources of phosphate for the waterbody you have tested.
3. The formation of the blue colour depends upon a redox reaction. Give the
oxidation states of the molybdenum before and after reaction with the
hydrogenphosphate. Write a balanced half equation for the reduction of the
molybdophosphate ion to the phosphorus molybdenum blue.
4. Assuming the ascorbic acid is oxidised completely to carbon dioxide write the half
equation for its oxidation. Note: Some kits do not use ascorbic acid.
5. Combine the two half equations to give the redox equation for the reaction
between the molybdophosphate ion and the ascorbic acid solution (or other
reducing agent).
6. Which of the reagents in the reaction between the molybdophosphate ion and the
ascorbic acid would you expect to be the limiting reagent? Explain why this is
necessary to the success of the test.
7. Explain why most phosphate occurs in particulate form.
8. What effect(s) can the water temperature have on the concentration of phosphate?
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