Water Chemistry Book

8/22/2019 Water Chemistry Book

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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


Analysis of dissolved oxygen in a natural water sample 13


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


Plant Nutrients - Nitrogen Nitrogen in water 23

Nitrogen cycle 23

Analysis of nitrate/nitrite in a natural water sample

Background 26

Experimental 28


Plant Nutrients Phosphorus Phosphorus in water 30

Analysis of phosphate in a natural water sample

Background 32

Experimental 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.


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.


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



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.


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



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.


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



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.


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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Water Chemistry Book