The Chemistry of Natural Waters

Purpose

The purpose of this experiment is to determine the concentrations of specific constituents in various water sources, to test the theory of electroneutrality (where the sum of the cations is equivalent to the sum of the anions – neutralizing their respective charges), and to explore the sources of the constituents, and why the concentrations may vary according to time, location, and water source.

Methods

Nine Tests were performed on twelve water samples for a variety of chemical characteristics. Tests for nitrate, sulfate, chloride, sodium, potassium, conductivity, alkalinity, pH, and hardness were performed. (Hardness includes total hardness, calcium hardness, and magnesium hardness. Calcium and magnesium cation concentrations were derived from these tests.)

Collection

Samples from freshwater bodies and the ocean were collected near the edge of the water body using a plastic bottle. Snow samples were collected and allowed to melt before being placed into the sample bottle.

Filtration

Filtered samples were created from a part of the original sample; filtration was required prior to atomic absorption and the spectrophotometric processes used to determine the sulfate, nitrate, and major cations. The spectrophotometric processes especially are dependent on turbidity and/or color of the sample and foreign particulate matter was not acceptable if accurate results were to be obtained.

All samples were filtered using a vacuum-assisted filter funnel/flask with a GF/C filter. Some samples (especially the urban snow samples) required pre-filtering with coarser filter paper to remove large particles.

Dilutions

Some samples (especially the sea water sample) required precision dilutions to bring their ion concentrations down to levels usable by the instrumentation. Then the value on the instrument readout is simply multiplied by the dilution factor when it comes time to report the results.

Precision dilutions were performed by pipetting an aliquot (such as 1 mL) into a volumetric flask. The flask is then filled with deionized water. If a 1/10000 dilution is required, 1 mL of sample is diluted in a 100 mL flask. 1 mL of this dilution is then diluted in another 100 mL flask, 100 x 100 = 10,000.

Conductivity

By itself, water is a poor conductor of electricity. The presence of various ions in water allows it to become a conductor. The amount and type of ions in the solution govern how conductive the water will be.

A Hach SensIon 5 conductivity meter was used to test the samples’ conductivity. The meter is first calibrated using a standard potassium chloride solution; the probe is then rinsed with deionized water and then is placed in the sample. The electrical resistance between two electrodes in the water is used to determine the conductivity. Because conductivity varies with temperature, the meter also measures the sample temperature and corrects the conductivity measurement at 25°C (the standard).

A 1003 microsiemen/cm standard solution was used for calibration of the conductivity meter.

pH

pH is a measure of the hydrogen ion concentration in a solution. It is logarithmic scale; therefore a pH of 7 is 10^-7 g/L H⁺ (pH is the negative log of the H⁺ concentration).

A Hach SensIon 1 portable pH meter was used to test the samples for pH. Calibration of the meter involved the use of two buffer solutions, one of pH 4.0 and another of pH 7.0. The probe is rinsed with deionized water and the placed into the 4.0 buffer solution, removed, rinsed, and placed in the 7.0 buffer. 4.0 is entered into the meter, and 7.0 for the second buffer. These pH values are used as a reference.

pH electrodes such as the one on the SensIon 1, use two parts – a glass half cell and a reference half-cell. The glass half-cell contains a solution of a known pH. The sample under test is outside the glass half-cell. A constant electrical potential is developed inside of this half-cell. The reference half-cell is in contact with the sample and has a constant potential. Changes in potential are due to changes in pH of the sample, and this varies linearly when the sample is at a constant temperature.

The probe is then rinsed and the sample and the probe placed in the sample. The READ button is pressed and the meter will stabilize. When it has stabilized, the pH value will be displayed.

Atomic Absorption Spectrometry – Basic Principals

Atomic absorption uses the excitation of electrons to create a spectral signature which is measured. A typical atomic absorption spectrometer contains a lamp, whose cathode is coated with the metal to be measured by the spectrometer. An electrical potential is placed between the anode and cathode of the lamp, causing a “sputtering” effect on the cathode which throws off radiation corresponding with the metal’s spectral signature. This radiation is directed out of the lamp to the sampling area.

Sodium and potassium cations were measured using a Perkin Elmer Analyst spectrometer. In the sampling area, water sample is drawn into a nebulizer which converts it into a fine mist. This mist is drawn into an air-acetylene flame. The flame provides energy to the atoms and ions in the sample, but not enough to bring their electrons up from ground state. The radiation from the lamp is passed into the flame, and this extra “burst” of energy brings the electrons up from ground state and their return to ground state throws off radiation which is picked up by a detector on the other side of the sampling chamber where its intensity is measured. Only atoms of the same element as that in the lamp are excited by its radiation, therefore other substances present in the sample do not emit their spectra.

Titrations

Titration is the process of reacting a substance in question with another substance of known concentration in order to determine the concentration of the substance in question. A known volume of sample is reacted with a known volume of a standard solution. An indicator is used to determine when the reaction reaches equilibrium. For acid-base reactions, pH indicators are common. The volume and concentration of standard solution is used to determine the number of moles of standard used. The stoichiometric ratio of the standard and the substance under test (sample) is compared to find the number of moles of sample used. From this and the sample volume, a concentration can be established.

A Hach digital titrator was used to perform all titrations. This titrator uses a syringe filled with reagent of a known concentration. A knob is used to very gradually push the plunger of the syringe to dispense the reagent. A counter records the number of times the knob has been turned. A calculation is provided to determine how much reagent was dispensed based upon the number of turns.

Chloride Titration

The mercuric nitrate method was used to test for chloride in the water samples.

The above reactions show that the mercuric nitrate (in the syringe) reacts with chloride ions in the sample, producing mercuric chloride and nitrate ions. Once all of the chloride ions have been used, mercuric ions (now in excess) combine with the diphenylcarbazone (included in the powder pillow that is added) to produce a purple-colored compound which appears showing that the titration is complete.

The above table is used to perform a chloride titration. The first two rows are used for the low-range reagent syringe. This is used on samples with a chloride concentration up to 160 mg/L. The remaining rows are for the high-range cartridge. The Sample Volume corresponds with the volume of sample being titrated. The Digit Multiplier is multiplied by the number of turns of the titrator dial. The product is the chloride concentration in mg/L.

For samples 3, 4 and 12, the 100 mL/Low Range, 50 mL/High Range (with 1/100 dilution), and 100 mL/Low range titrations were performed respectively.

Alkalinity Titration

There are three types of alkalinity that this test can measure; hydroxide, carbonate, and bicarbonate. The first step is to add a phenolphthalein indicator to the sample. If the sample remains colorless, there is no carbonate or hydroxide alkalinity – the total alkalinity is bicarbonate alkalinity. The three forms are dependent upon pH – at low pH, CO₂ is dissolved in the solution. As the pH gets higher, this becomes bicarbonate, then carbonate, then hydroxide. Each form is more basic than the other. All twelve samples had only bicarbonate alkalinity.

To titrate a sample without PPTH alkalinity, a Bromcresol Green/Methyl Red indicator powder pillow is added. The sample is titrated with sulfuric acid until a light-pink color is reached. The number of turns on the titrator is multiplied by the appropriate digit multiplier to give alkalinity in mg/L CaCO₃. Because all of the samples have bicarbonate alkalinity, the carbonate alkalinity is multiplied by 122/100 to give the appropriate bicarbonate alkalinity.

The alkalinity titration was a simple acid-base titration with sulfuric acid.

The above table is used for alkalinity titration and is very similar to the one used for chloride. For samples 3, 4, and 12, the low range, high-range, and low-range titrations were used respectively.

Hardness Titration

The hardness test contains two parts from which three values are derived – the total hardness (THA), calcium hardness, and magnesium hardness. The results of the total hardness test and the calcium hardness test are subtracted to obtain the magnesium hardness value.

The above table shows the range of the titration, titration cartridge concentration, sample volume, and digit multiplier. For samples 3, 4, and 12, the 50 mL/High Range, 25 mL/Low Range (with 1/100 dilution), and 50 mL/High Range titrations were used respectively.

The above table is used for both the total hardness and calcium hardness procedures.

In the total hardness test, one mL (or 2 mL is hardness is high) of Hardness Buffer solution is added to sample, and contents of one ManVer 2 hardness indicator powder pillow is added as well. The sample is titrated with EDTA (in the titration cartridge) from red to blue.

In the calcium hardness test, 2 mL of KOH standard solution is added to the sample, and then a CalVer 2 calcium indicator powder pillow is added to the sample before titration. The sample is titrated with EDTA from a pink color to blue.

Once the hardness values are obtained, they are in the form of CaCO₃ equivalent. They can be converted to find the concentrations of their respective cations. Calcium hardness (CaCO₃) is divided by 2.5 to give the calcium cation concentration. This number is derived from the molecular weight of calcium carbonate (~100) divided by the atomic weight of calcium (~40). A similar procedure is used for magnesium hardness, where the magnesium hardness (CaCO₃) value can be divided by 4.11 to give the magnesium cation concentration.

Spectrophotometry

Sulfate and Nitrate concentrations were measured using spectrophotometry. In this process, a chemical reaction(s) between the sulfate/nitrate anions and reagents added to the sample change the color and/or opacity of the sample in the reaction cuvette. A specific wavelength of light is passed through the sample to a detector. The differences between the light source and detector are used to determine the anion concentration.

A Hach DREL 2000 spectrophotometer is used for each of these tests.

Nitrate

The Hach Cadmium reduction method was used to determine nitrate. The test procedure is as follows:

1. Enter the program number, 351, then press READ/ENTER

2. Set the wavelength dial to 507 nm.

3. Press READ/ENTER

4. Transfer 30 mL of sample to a small Erlenmeyer flask.

5. Add one NitraVer 6 to the flask. Cover with parafilm or stopper.

6. Press SHIFT TIMER. A 3-minute reaction period will begin. Shake the sample

flask during the entire 3-minute period.

7. After the timer is done, press SHIFT TIMER again and a two minute timer

begins so that the cadmium can settle.

8. After the timer is done, pour 35 mL of sample into a cuvette.

9. Add one NitriVer 3 powder pillow to the cuvette. Cover with parafilm.

* NOTE – The sample will turn a pink color if nitrate is present. If it turns a dark orange color, then the sample is likely too concentrated with nitrate and you will need to perform a dilution. Try a 1/10 dilution. Waters which may be heavily polluted with nitrate such as those from agricultural runoff may require even higher, such as 1/20.

10. Press SHIFT TIMER. A 10 minute reaction period will begin.

11. When the timer is done, fill another cuvette with sample and this will be the

“blank” to zero the instrument.

12. Press ZERO

13. Within 10 minutes after the reaction period timer is done, remove the parafilm

and place it into the spectrophotometer. Close the lid.

14. Press READ/ENTER. The nitrate concentration will be shown as mg/L.

This test works from 0-0.4 mg/L of nitrate. After the test, the contents of the cuvette to which the powder pillow were added should be disposed of properly, as it contains cadmium.

The reactions which take place in the nitrate test are as follows:

(1) Nitrate is reduced to nitrite by action of cadmium.

(2) Nitrite plus added sulfanilic acid yields a diazonium salt

(4) The salt plus added chromotropic acid forms a red-orange compound.

The cadmium is in the NitraVer powder pillow. The two acids are in the NitriVer powder pillow.

The final result of the nitrate test is in the form of nitrate nitrogen. Each nitrate ion has one nitrogen and three oxygens. To obtain the value as a mass of nitrate, multiply the result of the test by 62/14 – the molecular mass of nitrate over the atomic mass of nitrogen.

Sulfate

The Sulfate test involved adding excess barium chloride to the sample. This barium chloride reacts with sulfate anions in the sample and produces a white, insoluble barium sulfate precipitate. The precipitate concentration in the sample (and therefore the sulfate ion concentration) is determined by the turbidity (opacity) of the sample in the cuvette.

The Hach SulfaVer 4 test was used. The test procedure is as follows:

1. Enter the program number, 680, and then press READ/ENTER

2. Set the wavelength to 450 nm with wavelength dial.

3. Press READ/ENTER

4. Fill a cuvette with 25 mL of sample.

5. Add a full SulfaVer 4 powder pillow to the cuvette, swirl to dissolve.

6. Press SHIFT TIMER, five minute timer for reaction will start.

7. After the timer is done, fill a second cuvette with 25 mL of sample. This will be the “blank”.

8. Place blank into spectrophotometer. Close the lid.

9. Press ZERO, display will show “0 mg/L SO₄^2-.”

10. Place sample (with powder pillow added) into machine, close lid.

11. Press READ/ENTER.

12. The concentration of sulfate anion will be shown as mg/L.

The test works from 0-70 mg/L of sulfate. After the test, the contents of the cuvette to which the powder pillow were added should be disposed of properly, as it contains barium.

Data

The attached spreadsheets 1 and 2 contain the original source data and the ion balance respectively.

The atomic absorption data from the sodium and potassium measurements (along with the dilution factors of each) is listed below. The “absorbance” value is converted to a concentration by first calibrating the AA spectrometer with standard solutions. The absorbance varies linearly with concentration.

Data Discussion and Analysis

Adjustments

The ion balance protocol requires that values be in the form of parts per million (ppm). The values obtained in the lab were all in the form of mg/L. Fresh water has a density of approximately 1000 g/L, therefore one milligram in one liter of water would be equivalent to one part per million. In seawater, the density is higher – around 1020 g/L. The ppm values for seawater are found by dividing each mg/L concentration in the seawater sample by 1.02. Once ppm is obtained, it can be converted to epm – equivalents per million. This is the weight of a quantity of an element or compound that will react with another element or compound in a chemical reaction.

Analysis

Most of the ion balance calculations fell within rather reasonable levels; however there were some irregularities. The highest deviation is found on sample #2, with an ion balance of -20.4. The nitrate value was very high for sample 2, and there is a possibility that this value is not correct. There may have been a problem with dilution or an error in the analysis.

Deviations from the expected values can also arise due to factors other than poor analysis. The pH of sample waters may change due to loss/gain of carbon dioxide, although small pH changes will not greatly affect ion balance. pH values of about 6 and higher are virtually insignificant in affecting the ion balance.

Adsorption of calcium into the walls of the sample container may reduce the calcium cation concentration, especially since some samples were stored for several weeks prior to testing (Thomas).

Ions that were not tested for may also affect ion balance. Silicate, fluoride, phosphate, iron, and other substances not tested may be present. For example, if calcium phosphate is present, the calcium will be detected but its phosphate counterpart will not be included. Most of the ions which were not analyzed are much rarer or are not very soluble in water; therefore their presence is likely little or none.

Issues with the analysis method itself may cause deviations as well, but most of the time they will not place an effect on the ion balance, just on the concentration of the individual ions. The chloride test registers bromide and iodide, therefore any bromide or iodide present will be counted with the chloride. This will cause the chloride concentration to be higher than actual, but since the bromide and iodide anions will likely be combined with a measured cation such as sodium or potassium, the ion balance will not be affected. The same condition applies to the nitrate test, which registers both nitrate nitrogen and nitrite nitrogen.

Comparison with Other test results

Several samples have been selected from the USGS National Water Information System’s database. This database contains results of analyses done by USGS or state departments of natural resources. The selected samples are (mostly) close to the areas where the actual sampling for this project has taken place. It is interesting to note that the pHs are about the only measurement that is very close to the value that we obtained. The other values seem to vary by quite a bit. These samples were taken in the summer time, when weather is warmer and biological activity is much higher. Dissolution of substances into water can take place at a faster rate in warmer weather, and increased biological activity changes the rate at which materials are added/removed from the water body. Even so, there are few outstanding differences (other than the nitrate value of the Ashland/Schuylkill County groundwater – which is very high)

Reasons for Chemical Composition of Waters

Precipitation

The chemistry of precipitation is often thought to be only water with few external constituents. This belief comes from the fact that precipitation is essentially distilled water – water is evaporated and all minerals and foreign matter are left behind. The water then condenses and falls back to the surface. While the water is in the atmosphere, it can pick up a number of substances from the air. Snow and ice samples can also be contaminated while they are still in solid form on the ground.

Precipitation from urban locations generally contains more dissolved solids than that found in rural areas. In urban areas, there are many more opportunities for the water to pick up substances as it moves through the air. Aerosols and gases produced by automobiles, construction and demolition, and industry add to substances in urban air which can end up in precipitation. Natural forces such as wind can also bring particles into the air, which further become dissolved in atmospheric water.

Water is naturally acidified by dissolution of carbon dioxide in the air, creating a pH of about 5. Man-made pollutants such as nitric oxide and sulfur dioxide can acidify the water further by means of acid-rain production pathways.

The urban snow samples tested in the lab show definite contamination with chloride; only seawater has a higher chloride concentration than sample #11. The most likely source of this chloride is from salt used as a deicer on the road surfaces. Sodium is also concentrated more highly in the urban snow, and its concentration correlates fairly well with the chloride concentration in both samples (ratios of chloride/sodium in both samples is fairly consistent). Rural snow had very low chloride, less than 1 ppm.

Calcium and magnesium constituents can be picked up from dust and other particulates. In an urban environment, erosion of concrete sidewalks and building materials creates dust which can be kicked up by wind and action of vehicles. Concrete is made up primarily of calcium and magnesium compounds. In Wilkes-Barre, dust from the demolition of the Hotel Sterling near the snow collection sites may be a source of some of the calcium and magnesium components. Concrete is not terribly soluble in water (if it were, nearly all of the world’s structures would be in trouble!), but when the water is acidified (as in acid precipitation) it becomes more soluble and in dust form the surface area available for reaction is very high.

Calcium in snow samples. Sample #5 is the rural snow – with a very low calcium (0.02 ppm – negative on the logarithmic scale). The urban samples (10, 11) had concentrations of about 14 and 1 ppm respectively. Aerosols from building materials and other activities contribute to the calcium content of these snows. Some ice melt mixtures (such as the kind bought in stores for use by homeowners) also contain calcium compounds – yet another possible source.

Surface Water and ground waters

The chemistry of surface and ground waters is more dependent upon the geology of the ground beneath the water body and the biological activity within it. Nitrates are common in surface waters, especially those which pass near agricultural areas when animal manure and fertilizers are used. In rivers and streams, there is little plant life in fast-moving water, therefore nitrate is not used up. In more stagnant bodies of water such as in lakes and ponds, high nitrate levels encourage the growth of phytoplankton and algae. These ‘algal blooms’ can block sunlight from penetrating the water, and when the algae die, its decomposition robs the water of oxygen which is needed for organisms such as fish. Excessive nitrate levels usually come from man-made sources, such as overuse of fertilizers on farms, septic systems, and wastewater treatment plants. Contamination of groundwater with nitrate is a major worldwide environmental issue. Nitrate in drinking water can lead to “blue baby” syndrome if consumed by infants. This dangerous condition results from the nitrate being reduced to nitrite in the stomach, which bonds with hemoglobin in the blood and eliminates its ability to carry oxygen. Some studies also suggest that nitrate is a precursor to carcinogens such as nitrosamines (van Maanen).

Calcium and magnesium levels in surface water and ground water are greatly dependent upon the geology of the region. Limestone is calcium carbonate, so naturally waters passing through or over limestone formations will be harder and have elevated calcium levels. The same holds true for magnesium-bearing rock. As surface waters erode rock formations, small amounts of sodium, potassium, and other cations which have been trapped in the crystalline structure of the rock are removed as well. These are not found in high concentrations in surface waters, but they can be concentrated over time in isolated lakes and in the oceans. The alkali metals especially are highly concentrated in the oceans because there is little or no biological demand for them and all of their compounds are very soluble in water.

The river samples (8, 9) and the Ashland groundwater sample (2) have higher alkalinity – a sign that they have contacted limestone geology. The effect is also defined (though not as vividly) when looking at pH (below). Sample 3 has a high pH – it is the Ashland Stream (Deep Creek).

Waters passing over limestone formations will have a greater capacity to deal with acidic precipitation than those which are passing over silicate rocks, because the calcium carbonate buffers the acidity, producing carbon dioxide, water, and the calcium salt of the acid (such as calcium sulfate).

The Geologic Map of Pennsylvania shows that waters in Northeast Pennsylvania will be encountering rock from the Devonian period – mostly shale, sandstone, and limestone. Sandstone will contribute little alkalinity to the water. Shale is made up of compressed clay and mud, and when these materials weather they release all kinds of elements. Clays contain much sodium, potassium, aluminum, and magnesium.

The Susquehanna River begins in a region of New York State with limestone geology. This would affect pH, alkalinity, bicarbonate, and calcium levels positively. We took two river water samples from the Susquehanna at different intervals, one about five days after the other. The most interesting aspect of this is that the second sample has elevated levels of nearly everything compared with the first sample. The best explanation is probably due to the rapid snow melt which occurred between the two sampling periods. This snow melt carried substances into the river, including road salt (increased chloride and sodium levels). The pH has increased; the snow melt water running over the land introduces carbonate minerals into the water. Nitrate levels were lower than before the snow melt period, but during winter runoff will contain little nitrate. Urban runoff is probably the greatest variable in these samples, as the samples were taken close to shore near the city, where storm water discharge enters the river.

Another possible explanation is that surface water inputs to the river dropped off while groundwater inputs remained constant. The general increase in pH, bicarbonate, and alkalinity would come from the groundwater which has been flowing through the carbonate rock structure while dissolving some calcium carbonate itself along with the other cations which are present in the rock structure. Nitrate levels in groundwater would be less likely to be elevated (as nitrate tends to originate at the surface) and this can explain the lower nitrate values.

Seawater

Seawater chemistry is one of the most dynamic, and it is affected by all factors – geochemical and biological.

The source of ions in seawater is mostly a product of chemical weathering of rock on land. Acidic precipitation dissolves rock and mineral deposits and solutions are created which end up in the oceans. Rock contains trace amounts of elements such as sodium and potassium which are released as the rock weathers. Nitrate, phosphate, and other nutrients are also renewed in the oceans by upwelling from the deep.

Substances in seawater which are in high biological demand are of lower concentration than those substances which are not. Nitrate and calcium carbonate are two examples. Nitrate is a nutrient used by photosynthesizing organisms, and calcium carbonate is used by many marine animals to build shells and other structures. This is a reason why the oceans contain a much greater magnesium hardness (and magnesium ion concentration) than calcium hardness, as opposed to surface waters which are the opposite. Organisms and chemical reactions in the oceans can create precipitates which fall to the ocean floor. Over time these will be recirculated into the mantle and eventually incorporated into new land.

Sodium and chloride are both in very low biological demand, and are both the ions of highest concentration of seawater. The balance of these ions in seawater is kept due to the fact that sodium chloride is deposited by evaporation in desert regions.

Conclusion

The chemistry of water can be used as an indicator not only of the suitability of the water for a particular purpose, but an indicator of its origins, biological activity within it, geology of the land on which it flows, and the impact that humans have upon it. Measuring the concentrations of these various ions gives a picture of the health of a water system.

Bibliography

Barica, Jan and F. A. J. Armstrong. “Contribution by Snow to the Nutrient Budget of

some small Northwest Ontario Lakes.” Limnology and Oceanography 16.6 (1971): 891-899. JSTOR. Wilkes U, Wilkes Barre, Farley Lib. 11 April 2007 <http://www.jstor.org>.

van Maanen, Jan M. S., Danielle M. F. A. Pachen, Jan W. Dallinga, Jos C. S. Kleinjans.

“Formation of Nitrosamines during Consumption of Nitrate- and Amine-Rich Foods, and the Influence of the Use of Mouthwashes.” Cancer Detection & Prevention 22.3 (1998): 204-212. Abstract. 28 April 2007 <http://www.blackwell-synergy.com/links/doi/10.1046/j.1525-1500.1998.0oa26.x>.

Mann, K.H. “Annual fluctuations in Sulphate and Bicarbonate Hardness in Ponds.”

Limnology and Oceanography 3.4 (1958): 418-422. JSTOR. Wilkes U. Wilkes

Barre, Farley Lib. 11 April 2007 <http://www.jstor.org>.

The Susquehanna River Hydrologic Observing System. 2004. Pennsylvania State

University. 30 April 2007 <http://srbhos.psu.edu/srb_testbeds/OstegoLake.asp>.

Thomas, J.F.J. Industrial Water Resources of Canada: Scope, Procedure, and

Interpretation of Studies. Canada Department of Mines and Technical Surveys, 1946.

United States. Dept. of the Interior. U.S. Geological Survey. National Water Information

System: Web Interface. 28 April 2007 <http://nwis.waterdata.usgs.gov/usa/nwis/qwdata>.