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Post by Robot on Jul 26, 2004 21:46:51 GMT -5
Water is Essential for Life
It covers 71% of the earth's surface and makes up 65 % of our bodies. Everyone wants clean water-- to drink, for recreation, and just to enjoy looking at. If water becomes polluted, its loses its value to us economically and aesthetically, and can become a threat to our health and to the survival of the fish living in it and the wildlife that depends on it.
How does water pollution occur?
Although some kinds of water pollution can occur through natural processes, it is mostly a result of human activities. We use water daily in our homes and industries, about 150 gallons per day per person in the United States. The water we use is taken from lakes and rivers, and from underground (groundwater); and after we have used it-- and contaminated it-- most of it returns to these locations.
The used water of a community is called wastewater, or sewage. If it is not treated before being discharged into waterways, serious pollution is the result. Historically, it has taken humanity quite a bit of time to come to grips with this problem. Water pollution also occurs when rain water runoff from urban and industrial areas and from agricultural land and mining operations makes its way back to receiving waters (river, lake or ocean) and into the ground.
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Post by Robot on Jul 26, 2004 21:48:33 GMT -5
What are some different types of water pollution?
Microbiological Disease-causing (pathogenic) microorganisms, like bacteria, viruses and protozoa can cause swimmers to get sick. Fish and shellfish can become contaminated and people who eat them can become ill. Some serious diseases like polio and cholera are waterborne.
Chemical A whole variety of chemicals from industry, such as metals and solvents, and even chemicals which are formed from the breakdown of natural wastes (ammonia, for instance) are poisonous to fish and other aquatic life. Pesticides used in agriculture and around the home-- insecticides for controlling insects and herbicides for controlling weeds-- are another type of toxic chemical. Some of these can accumulate in fish and shellfish and poison people, animals, and birds that eat them. Materials like detergents and oils float and spoil the appearance of a water body, as well as being toxic; and many chemical pollutants have unpleasant odors. The Niagara River, between the US and Canada, even caught fire at one time because of flammable chemical wastes discharged into the water.
Oxygen-depleting Substances Many wastes are biodegradable, that is, they can be broken down and used as food by microorganisms like bacteria. We tend to think of biodegradable wastes as being preferable to non-biodegradable ones, because they will be broken down and not remain in the environment for very long times. Too much biodegradable material, though, can cause the serious problem of oxygen depletion in receiving waters.
Like fish, aerobic bacteria that live in water use oxygen gas which is dissolved in the water when they consume their "food". (The oxygen in the compound H2O, water, is chemically bound, and is not available for respiration (breathing)). But, oxygen is not very soluble in water. Even when the water is saturated with dissolved oxygen, it contains only about 1/25 the concentration that is present in air. So if there is too much "food" in the water, the bacteria that are consuming it can easily use up all of the dissolved oxygen, leaving none for the fish, which will die of suffocation.
Once the oxygen is gone (depleted), other bacteria that do not need dissolved oxygen take over. But while aerobic microorganisms-- those which use dissolved oxygen-- convert the nitrogen, sulfur, and carbon compounds that are present in the wastewater into odorless-- and relatively harmless-- oxygenated forms like nitrates, sulfates and carbonates, these anaerobic microorganisms produce toxic and smelly ammonia, amines, and sulfides, and flammable methane (swamp gas). Add in the dead fish, and you see why we don't want large amounts of biodegradable materials entering lakes and streams.
Nutrients The elements phosphorus and nitrogen are necessary for plant growth, and are plentiful in untreated wasthingyer. Added to lakes and streams, they cause nuisance growth of aquatic weeds, as well as "blooms" of algae, which are microscopic plants. This can cause several problems. Weeds can make a lake unsuitable for swimming and boating. Algae and weeds die and become biodegrable material, which can cause the problems mentioned above (and below). If the water is used as a drinking water source, algae can clog filters and impart unpleasant tastes and odors to the finished water.
Suspended matter Some pollutants are dissolved in wastewater, meaning that the individual molecules or ions (electrically charged atoms or molecules) of the substance are mixed directly in between the molecules of water. Other pollutants, referred to as particulate matter, consist of much larger-- but still very small-- particles which are just suspended in the water. Although they may be kept in suspension by turbulence, once in the receiving water, they will eventually settle out and form silt or mud at the bottom. These sediments can decrease the depth of the body of water. If there is a lot of biodegradable organic material in the sediment, it will become anaerobic and contribute to problems mentioned above. Toxic materials can also accumulate in the sediment and affect the organisms which live there and can build up in fish that feed on them, and so be passed up the food chain, causing problems all along the way . Also, some of the particulate matter may be grease-- or be coated with grease, which is lighter than water, and float to the top, creating an aesthetic nuisance.
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Post by Robot on Jul 26, 2004 21:50:14 GMT -5
How do we prevent water pollution?To keep our used water from spoiling our water resources, we have to remove the pollutants before the water gets back into the environment. In urban areas in most developed countries, the wastewater from homes, businesses and factories is collected by a system of underground pipes-- sewers-- which carry it to one or more central treatment facilites. Most of these are located near bodies of water into which the treated wastewater is discharged. In the U.S., all such facilites must have a permit issued by the federal and/or state government, describing limits on the amounts of various pollutants which may be discharged. The U.S. EPA (Environmental Protection Agency) refers to these as NPDES (National Pollutant Discharge Elimination System) permits. Industries located in areas where they are not connected to a sewer can discharge directly into a waterway, but will need a permit, and will probably have to have their own treatment plants. Even industries which are connected to sewers may have to pretreat their wastewaters before discharging them into the sewers, because they may contain materials which will harm the sewers or the treatment plants-- or may be a danger to the people who work in maintaining the sewer system. [If you are some one who works in an industrial pretreatment program, you may be aware that the USEPA has recently (July 22, 1999) proposed a series of changes to the current regulations. I have prepared a summary of the proposals which you may find helpful before reading the entire document.] www.geocities.com/RainForest/5161/strmregw.htmHomes in non-urban areas that are not connected to a sewer are usually required by their town to have on-site treatment systems. Most common for single homes are septic systems, which consist of a buried tank connected to a set of perforated pipes, embedded in gravel, which distribute the water into the soil. (The "Links" page has several references on this subject.) Larger housing complexes may have treatment systems based on the principles used in full-scale sewage treatment plants. A Word About Sewers: Besides having a set of sewer pipes-- called, strangely enough, sanitary sewers-- which carry wastewater to a treatment plant, cities and towns also need pipes to collect stormwater. These are needed to prevent street flooding and usually lead directly to a waterway without any treatment. The runoff of pollutants from streets and yards into these storm sewers contain oil and other automotive wastes, which may contain toxic metals and organic compounds-- as well as pesticides and nutrient-containing fertilizers from lawns and gardens, and pathogenic microorganisms from animal wastes. The problem of pollution from storm sewers is currently being addressed by the USEPA. Further complicating the situation is the fact that while some cities and towns have completely separate sanitary and storm sewer systems, many others have combined systems. During rainy periods, combined sewers cause two problems: overloading of the treatment plant with extra water and contaminating waterways with untreated sewage from overflows. Even in cities with separate sewer systems, the flows to the treatment plants often increase greatly when it rains because of cracks or separations in the pipes, which allow groundwater or stormwater from broken storm sewer pipes to infiltrate into the sanitary sewer-- or from direct inflow of stormwater into manholes and from illegal connections of roof drains and sump pumps in buildings.*
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Post by Robot on Jul 26, 2004 21:56:25 GMT -5
Water Analysis
Why do we need to analyze water?
If water is badly polluted-- like raw sewage--- it might be obvious from its appearance or odor.
It might be colored or turbid (cloudy), or have solids, oil or foam floating on it.
It might have a rotten odor, or smell like industrial chemicals.
A lot of dead fish floating on the surface of a lake would be a clear sign that something was wrong. But many harmful-- and beneficial-- materials in water are invisible and odorless. In order to go beyond the obvious, to determine what materials are in the water, and how much, we need to be able to conduct chemical or microbiological analyses.
Analysis of a natural body of water will tell us how clean or polluted it is. If there is damage to wildlife, the measurements will help pinpoint the cause-- and the source. In a wastewater treatment plant, analyses are necessary for monitoring the effectiveness of the treatment processes. In the United States, the Clean Water Act requires wastewater dischargers to have permits. These permits set limits on the amounts of specific pollutants which can be discharged, as well as a schedule for monitoring and reporting the results. Usually, the reports must be filed monthly, while the measurement frequency for a particular parameter (measurable property) can run anywhere from "continuously" to just once a year. Only standard analytical procedures specified in the "Code of Federal Regulations" may be used, so that the government agencies can feel reasonably confident that results from different laboratories are comparable.
Similar considerations apply to drinking water. The purity of the water we drink is of more concern to the average person than the quality of the wastewater discharged by the sewage plant. But we should not forget that in many places, especially along a river, one town's wastewater discharge may be part of the next town's water supply...
There are two aspects to water analysis that we need to consider:
what substances or organisms are we interested in testing for-- and why?
what procedures and equipment do we use to make the measurements, and how do they work?
Let's look at the "procedures and equipment" first:
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Post by Robot on Jul 26, 2004 21:58:30 GMT -5
Analytical Methods
Water analyses are done by several methods. The most common types of measurements are gravimetric (weighing), electrochemical (using meters with electrodes) and optical (including visual). Instrumental methods are becoming increasingly popular, and instrumentation is getting "smarter" and easier to use with the inclusion of microprocessors. In the simplest case, a sample may just be placed in an instrument and a result read directly on a display. More often some physical separation technique or chemical procedure is needed before a measurement is made, in order to remove interferences and transform the analyte-- the target of the analysis-- into a form which can be detected by the instrument.
Since even raw sewage is generally more than 99.9% water, most environmental analyses are measuring very low concentrations of materials. The results of these measurements are usually expressed in the units "milligrams per liter," abbreviated as mg/L. Since a milligram is one thousandth of a gram, and a liter of water weighs about a thousand grams, a mg/L is approximately equal to one part per million by weight. A part per million ("ppm") is only one ten thousandth of one percent. For toxic metals and organic compounds of industrial origin, measurements are now routinely made in the part per billion (microgram per liter) range or even lower. At such low levels, sensitive equipment and careful technique are clearly necessary for accurate results. Avoiding contamination of the sample and using methods which prevent interferences from other substances in the water are crucial requirements for successful analyses.
Separation Techniques:
Some measurements require separating the analyte from other substances in the water which may interfere with the measurement. Some measurements even require separating the analyte from the water entirely. Separation techniques include:
Filtration: The water is passed through a fine-pore filter which can be made of paper, glass fibers, a cellulose acetate membrane, etc. Filtration through a filter of some agreed-upon standard pore size can be used to separate "suspended" from "dissolved" portions of the analyte. The analyte may be the suspended matter which is captured on the filter-- or the filter may be used to clarify the water for analysis of a dissolved material. Often, the filtration is assisted by applying a vacuum below the filter, which is supported on a porous holder in some type of funnel.
Distillation: If the analyte can be boiled out of the water, or along with the water, then the vapors can be cooled and re-condensed or trapped in a liquid form in a different container. This way the analyte can be removed from the interfering substances in the original water sample. Often the sample is made acidic or alkaline, or treated chemically in some other way before distillation, to convert the analyte into a volatile (easily evaporated) form, and to immobilize or neutralize interfering substances.
Extraction: Some analytes may be much more soluble in an organic solvent than in water. If the solvent does not mix with water, and the sample is shaken with portions of the solvent, almost all of the analyte may be transferred from the water into the solvent, leaving interfering substances behind. This is known as a "liquid-liquid" extraction. The analysis may be completed using the organic portion. There are also continuous versions of this process for use with liquid or with dry samples.
Another type of extraction is called "solid-phase extraction." In this kind of procedure, the sample is passed through a column or filter containing a powdered or granulated material which retains (adsorbs) the substances of interest and allows other types of dissolved materials to pass through. Then a solvent, or an acid or alkaline solution, can be passed through to de-sorb and redissolve the analytes, a process known as elution.
Either type of extraction can also be used to concentrate the analyte into a smaller total volume, which increases the sensitivity of the analysis. This can be true for distillation or filtration, as well.
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Post by Robot on Jul 26, 2004 22:02:10 GMT -5
Measurement Techniques: Gravimetric analysis or, simply, weighing: Analytical balances routinely used for gravimetric analysis are sensitive to one tenth of a milligram, or one ten-thousandth of a gram. Most laboratories use electronic balances with direct digital readouts. For a measurement of the milligrams per liter of solids in the water, a measured volume of sample can be dried in a tared (pre-weighed) dish; the dish plus solids are weighed after the water has evaporated off; the weight of solids is calculated by subtraction, and the concentration figured by dividing the weight of solids by the volume of the sample. For a filtered sample, the tared filter itself is dried along with the solids it captured, and the suspended solids (those captured on the filter) calculated in the same way. In some chemical analyses, a precipitate is formed by reacting the analyte of interest with another chemical reagent (reacting chemical); then the precipitate can be filtered, dried, and weighed as a suspended solid. This type of analysis is more common with water solutions that are more concentrated than environmental samples, though, such as chemicals purchased for use in water or wastewater treatment. Electrochemical: The outer portions of all atoms and molecules consist of "shells" of electrons, and all chemical reactions involve interactions with these outer electrons-- sharing or transfer, or something in between. It is not surprising,, then, that electricity and chemistry are interrelated (just think of batteries), and that electrical measurements can be used to detect and determine some substances of interest. The procedures involve placing electrodes in a water sample and measuring either an electrical potential (voltage), in millivolts, or a current, in milliamperes, which is related to the concentration of analyte. Depending on what they are designed to measure, electrodes can be simple pieces of metals such as gold, silver, platinum, copper, etc.; or they may be elaborate systems with semi-permiable membranes and several internal electrodes and filling solutions. The instrumentation may be capable of reading out directly in concentration units. Usually some sort of calibration procedure is necessary, using one or more standard solutions of known concentration. Colorimetry or spectrophotometry: This method involves measuring the intensity of a color in a solution and relating it to the concentration of the analyte. While some materials of interest are already colored, most of these analyses require the analyst to add some chemical reagents (reacting chemicals) to a sample to produce a characteristic color. The simplest type of measurement is visual comparison of the intensity of the color to a set of color standards which represent various concentrations of the analyte. While this is method does not require any expensive equipment, color perception is rather subjective-- and many people have some degree of color-blindness. A more precise measurement can be made using a colorimeter. A colorimeter is a device consisting of 1) a light source, which can be as simple as tungsten-filament light bulb; 2) some optics for focusing the light 3) a colored filter, which passes light of the color which is absorbed by the treated sample; 4) a sample compartment to hold a transparent tube or cell containing the sample, 5) a light-sensitive detector, like the light meter on a camera, which converts the light intensity into an electric current, and 6) electronics for measuring and displaying the output of the detector. Some colorimeters may be designed to read out directly in concentration units, while others may show the results in units of light absorbance which need to be compared to a calibration curve. (An interesting point is that the filter is not the same color as the solution being tested, but rather the complementary color. We want to use a filter which transmits light of the color which the solution absorbs. A yellow solution looks yellow because it absorbs blue light, so a blue filter would be used.) If we want to get more precise and more interference-free measurements, we can use a spectrophotometer. This is very similar to a colorimeter, except that instead of using a filter to select the color of light to pass through the sample, we instead break the white light up into a rainbow (spectrum) of colors using a prism or a diffraction grating. The light is passed through a narrow opening (slit) before reaching the sample. By rotating the prism or grating, the color {"wavelength") of light can be selected more precisely and we can better match the color with that absorbed by the sample. The principle is shown in the diagram below. Needless to say, spectrophotometers cost more than colorimeters, and are likely to be more delicate and less portable, as well. While many tests are done using visible light, some analyses also make use of the invisible ultraviolet or infrared portions of the spectrum. Scanning spectro photometers can also be used to identify some types of analytes by the wavelengths or colors of the light they absorb. There is a variation of this type of testing, usually referred to as atomic spectroscopy, which is used mostly for trace metal analysis. The sample is converted to a gas by one of several methods-- usually involving heating. Then the light from a lamp containing the same metal is passed though the gas and the absorbance measured just as with a liquid sample (atomic absorption spectrophotometry). Alternatively, the intensity of the light emitted from the heated atoms of the metal in the gas can be used as a way of measuring the concentration (atomic emission spectrophotometry). A very popular atomic emission method in use today is called inductively coupled plasma spectrometry. The sample is carried in a stream of argon gas surrounded by coils which emit radio frequency energy that converts some of the gas into a very hot, ionized (electrically charged) form. An advantage of this method is that many elements can be measured simultaneously, or in rapid succession.
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Post by Robot on Jul 26, 2004 22:04:14 GMT -5
Titration: Titration depends on using a well-defined chemical reaction to measure the amount of a standard solution needed to react with certain amount of the sample. A known volume, such as 100 mL, of sample is placed into a flask or beaker. The standard reagent is dispensed from a graduated tube called a burette so the volume used can be measured. The "end point" of the reaction is usually determined by observing a color change in an indicator solution, which is added to the flask before the start of the titration. End points are also often determined using electrochemical equipment. Once we know how much of the standard reagent was needed, we can calculate the amount of the analyte that is in the sample, because the reaction will always use the same proportion of the two materials. A common example is measuring the concentration of an acid by titrating with a standard base, such as sodium hydroxide.
Chromatography: This technique got its name, which means "color picture", because it was first used to separate colored pigments from a single spot on a piece of paper. A solvent, such as alcohol, is allowed to move slowly across the paper, and the different components of the pigment travel at different rates. The result is a series of separated spots of different colors. They move at different rates because of differences in the pigments' relative attraction to the paper (the "stationary phase") and their solubility in the solvent (the "mobile phase"). This principle is used in modern instrumentation to separate mixtures of organic chemicals or inorganic ions. The components can be identified by their retention times,-- i.e., how long it takes them to pass through the instrument-- and detectors can be used to measure the amount of each component.
In gas chromatography, (or, simply, "GC") the mixture of substances is injected into a narrow, coiled column, several feet long, made of an inert material like glass, silica or stainless steel. The sample has usually been extracted into an organic solvent and concentrated by evaporation as a pretreatment step. The column may be filled with an oil-coated, powdered mineral, which forms the stationary phase. In the narrower capillary columns, the stationary phase is bonded directly to the wall of the tubing. The columns are usually contained in an oven, which may be programmable to raise the temperature at a controlled rate over time. Heating the column allows analysts to use this technique on many substances which are not gases at room temperature, including solvents and toxic chemicals like pesticides and PCB's. A continuous flow of an inert gas, such as argon, helium, or sometimes nitrogen, carries the evaporated mixture through the column. The substances are detected as they exit the column, usually by a technique that converts them into ions (electrically charged atoms or molecules), although one method uses heat conduction. The ions are produced by means such as flames, ultraviolet light, or radioactive materials. They are detected by being attracted to charged plates, where they produce an electrical current proportional to the amount present. The output of the detector usually is shown as a chart of "peaks" vs. time, called a chromatogram, often with the retention time and the intensity of the peak printed out. The retention time is used to identify the substance, while the height or area of the peak is used to quantify its concentration. A more positive identification is possible using a mass spectrometer (see below) as the detector.
For substances which cannot easily be vaporized because of high boiling point or instability at higher temperatures, there is a liquid version of this technique know as HPLC (High pressure or high performance liquid chromatography). Organic solvents are used as the mobile phase. Ultraviolet (UV) light absorption is often used for detection. Herbicides and pharmaceuticals are common types of substances analyzed by this technique. Another variation of LC is ion chromatography, (IC), where the target analytes are charged inorganic or organic substances. The mobile phase is an aqueous (water-based) solution, and the stationary phase is made up of an ion exchange resin. The detectors usually measure electrical conductivity, although UV absorption can also be used. This technique can be used to measure the concentrations of several important inorganic anions, such as fluoride, sulfate, phosphate, and nitrate all in one analysis
Mass Spectrometry: In a mass spectrometer, an ionized vapor is passed between magnets or radio frequency coils which separate the ions by mass (actually by charge to mass ratio). The pattern produced is characteristic of the particular substance, which can be identified by comparison with computerized "libraries" of mass spectra. While the instrumentation can be used alone, for environmental analyses it is usually used in tandem with another technique. Used as a "detector" for gas chromatography ("GC-MS"), it can positively identified components which have already been separated from a mixture. There is some use with liquid chromatography, as well (LC-MS). As a detector for metal ions produced in an ICP (see above), it provides very high sensitivity and is being used to determine very low levels of metal in drinking water, and may soon be approved for wastewater effluents and receiving waters.
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Post by Robot on Jul 26, 2004 22:05:38 GMT -5
Parameters / Analytes
The previous part described some of the techniques and equipment used in water and wastewater analysis. This part discusses what things we want to measure, what their significance is, and what methods are used for each of them. Click on the parameter in the left frame to view information here.
CONTENTS
pH, acidity, and alkalinity
Dissolved Oxygen (D.O.)
Oxygen Demand
Solids
Nutrients
Chlorine
Oil and Grease
Metals
Cyanide
Toxic Organic Compounds
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Post by Robot on Jul 26, 2004 22:09:37 GMT -5
pH is a measure of how acidic or alkaline a solution is. In pure water at room temperature, a small fraction (about two out of every billion) of the water molecules (H2O, or really, H-O-H) splits, or dissociates, spontaneously, into one positively charged hydrogen ion (H+) and one negatively charged hydroxide ion (OH-) each. There is an equal number of each ion, so the water is said to be "neutral".
Some materials, when dissolved in water, will produce an excess of (H+), either because they contain these ions and release them when they dissolve, or because they react with the water and cause it to produce the extra hydrogen ions. Substances which do this are called acids. Likewise, some chemicals, called bases or alkalis, produce an excess of hydroxide ions.
The scale which is used to describe the concentration of acid or base is known as pH, for power or potential of the Hydrogen ion. A pH of 7 is neutral. pH's above 7 are alkaline (basic); below 7, acidic. The scale runs from about zero, which is very acidic, to fourteen, which is highly alkaline. The scale is logarithmic, meaning that each change of one unit of pH represents a factor of 10 change in concentration of hydrogen ion. So a solution which has a pH of 3 contains 10 times as many (H+) ions as the same volume of a solution with a pH of 4, 100 times as many as one with a pH of 5, a thousand times as many as one of pH6, and so on. Some common materials and their approximate pH's are: Acids--- carbonated beverages, 2 to 4; lemon juice, about 2.3; vinegar,about 3; Bases: baking soda, 8.4; milk of magnesia.10.5; ammonia,11.7;lye,14 to 15. (Some of these figures are from the Handbook of Chemistry and Physics, 56th ed., CRC Press,1976) While the pH measures the concentration of hydrogen or hydroxide ions, it may not measure the total amount of acid or base in the solution. This is because most acids and bases do not dissociate completely in water. That is, they only release a portion of their hydrogen or hydroxide ions.
A strong acid, like hydrochloric acid, HCl, releases essentially all of its H+ in water. The concentration of H+ is the same as the total concentration of the acid. A weak acid, like acetic acid (the acid in vinegar), may release only a few percent of the hydrogen that it has available.
If you are trying to neutralize an acid by adding a base, like sodium hydroxide, the amount you would need to neutralize a strong acid could be calculated directly from the pH of the acid solution. But for a weak acid, the pH does not tell the whole story; the total amount of base needed would be a lot more. This is because as the OH- from the base reacts with the H+ in solution to form water, more H+ will break loose from the undissociated portion of the acid to take its place. The neutralization will not be complete until all of the weak acid has dissociated. To measure the total acidity, also called base-neutralizing capacity (BNC) of a water sample, it has to be titrated with base. That is, a solution of a base whose concentration is known must be added to the water sample slowly until the neutralization is complete. By measuring the volume of the base added, you can figure out the original concentration of acid.
In a similar way, the acid-neutralizing capacity (ANC), or alkalinity of a water sample has to be determined by titrating it with a solution of a strong acid of known concentration.
For a more technical explanation of pH and alkalinity, look at this "mini-tutorial", which includes formulas, reactions, examples, and titration curves.
Significance: Although there are some microorganisms which can function at extreme pH's, most living things require pH's close to neutrality. Many enzymes and other proteins are denatured by pH's which differ much from pH7, which disrupts the functioning of the organism and may kill it. Besides the harm to aquatic life in natural waters, pH imbalances can inhibit-- or completely wipe out-- biological processes in wastewater treatment plants, resulting in incomplete treatment and pollution of the receiving waters. Low (acidic) pH's also cause corrosion in sewers systems and increase the release of toxic and foul-smelling hydrogen sulfide gas. (This gas has been responsible for the deaths of numerous sewer workers.) Low pH's also increase the release of metals, some toxic, from soils and sediments. Alkalinity is an important parameter because it measures the water's ability to resist acidification, for instance, to acid rain. In wastewater treatment, some processes produce acidity. If there is not enough alkalinity to neutralize it, the pH of the process can drop and cause it to become inhibited. Alkalinity can be supplemented by chemical addition to avoid this.
Measurement: There are indicator solutions which change color in different pH ranges, and these can be used for approximate estimation of pH in solutions which contain high enough concentrations of pH-determining ions. "pH paper", impregnated with such indicators, are commonplace in testing laboratories. For accurate measurements and use in dilute solutions, electrochemical measurement (a "pH meter") is required. Alkalinity and acidity are determined by titration with strong base or acid, respectively, using either indicators or a pH meter to mark the endpoint.
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Post by Robot on Jul 27, 2004 9:56:22 GMT -5
More about pH and alkalinity...
pH is one of the most commonly made measurements in water testing, but one of the least understood. Here is an attempt at an explanation:
In order to understand the pH scale, we have to discuss the ideas of moles and of logarithms.
Moles: When chemists talk about the amount of a substance, they often like to use the unit of moles, rather than grams. A mole of a substance is, simply, the number of grams of that substance equal to its molecular weight. A mole of water, weighs about 18 grams, because water has a molecular weight of about 18. A mole of calcium carbonate, CaCO3, weighs about 100 grams; a mole of methyl alcohol, CH3OH, 32 grams, etc. The advantage for chemists of using moles is that an equal number of moles of any substance contains the same number of molecules, so it is easier to calculate the amounts of substances which react with one another.
In a liter of pure water at room temperature the number of moles of hydrogen ions is about 0.0000001. (For hydrogen, with an atomic weight of 1, this is also about equal to the number of grams of hydrogen ions.) In scientific notation, this is written as 1 x 10-7, where the superscript, -7, is known as a power, an exponent, or a logarithm. (All three terms mean the same thing. The seven indicates the number of places to the right of the decimal point that the "one" is located.) It turns out, when measuring hydrogen ion concentration electrochemically, that the electrical potential (voltage) generated at the measuring electrode is directly related not to the H+ concentration, but to the logarithm of the H+ concentration. This makes it more convenient to refer to the H+ concentration in terms of its logarithm. And since H+ concentrations in water solutions are almost always less than one mole per liter, the exponent is almost always going to be negative, because that is the way scientific notation expresses numbers less than one. So, the negative of the logarithm of the hydrogen ion concentration in a solution is given a special name. It is called the pH, which stands for the potential of the hydrogen ion.
Of course, in the pure water, the concentration of hydroxide ions is also 1 x 10-7 moles per liter, since each water molecule that dissociates produces one ion of each type. The water is said to be neutral. It has a pH of 7 and also a pOH of 7, where the term pOH refers to the negative logarithm of the hydroxide ion concentration. There are substances which, when dissolved in water, will upset that balance, and produce an excess of either H+ or OH-. They may contain those ions and release them (dissociate) when they dissolve, or they may react with the water (hydrolyze) and produce them that way. Those substances which increase the concentration of H+ are called acids; those which decrease it (and increase the OH- ) are bases or alkalis. For instance, if a strong acid solution increases the H+ concentration to 0.1 moles per liter (1 x 10-1), which has a million times as many H+ ions as a neutral solution, then the pH is equal to 1. Similarly, if a strong base solution contains 0.1 moles per liter of OH- ions, it has a pOH of 1. According to the laws of chemical equilibrium, the pH and the pOH always add up to 14 (at about room temperature), so the solution with the pOH of 1 has a pH of 13. Most solutions have a pH between 0 and 14, and 7 is the neutral point. pH's below 7 are increasingly acidic as the number decreases; pH's above 7 are increasingly alkaline. And since the scale is logarithmic, each unit change in pH represents 10 times as many ions in solution.
A strong acid or base is one which dissociates completely when it dissolves in water. The amount of it in solution can be estimated from the pH. Most acids and bases, however, are weak; they dissociate or hydrolyze only partially. Many solutions also contain mixtures of several acidic or basic substances. In these cases, it is difficult to estimate the total amount of acid or base by measuring the pH, so this must be done by titration. As every high school chemistry student knows, acids react with bases to form water and salts. Therefore, an acid is titrated using a standard base, and visa versa. In water and wastewater analysis, the amount of acid needed to titrate a solution to a particular pH is a measure of the acid neutralizing capacity of that solution, and is referred to as the solution's alkalinity. In natural waters, the pH is most often controlled by the concentrations of carbonate, bicarbonate, and carbon dioxide, since these are products of respiration and fermentation. Because of this, alkalinity is usually measured in terms of the amount of acid needed to reach the pH of a pure solution of one or another of these substances. Similarly, acidity is defined as base neutralizing capacity, and is measured by titration against a standard base.
(While a chemist might prefer to measure these quantities in moles per liter, engineers seem more comfortable with standard weight units. So acidity and alkalinity are usually expressed in units of milligrams per liter of calcium carbonate. Calcium carbonate, or limestone, is a weakly alkaline material, 50 grams of which react with one mole of hydrogen ions.)
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Post by Robot on Jul 27, 2004 10:00:13 GMT -5
Titrations and Buffer Solutions:A solution of an pure acid will have a pH determined by its concentration and by how strong an acid it is-- that is, how easily it releases a proton (hydrogen ion) when dissolved in water. As we titrate a solution of an acid with strong base, hydrogen ions are consumed by reacting with the added hydroxide ions to form water, which leads to an increase in the pH of the solution. For an acid with one proton which can dissociate ("monoprotic" acid), the titration will be complete when the number of moles of hydroxide added equals the number of moles of acid originally present. If we call the fraction of acid which has been neutralized f, then the titration is complete when f = 1. (For an acid with two replaceable hydrogens ("diprotic" acid), the titration is complete at f = 2, and so forth.) If more base is added after the acid is all neutralized, then the pH of the solution will be determined by the concentration of hydroxide-- essentially as if it were being added to plain water. The initial, final, and intermediate pH's will be a function of the acid's concentration, strength, and the value of f. The calculations are different for strong and weak acids, so let's consider them separately. For a strong acid, essentially all of the acid dissociates, so that the concentration of hydrogen ions (H+) is equal to the concentration of the acid. Therefore, the initial pH of the acid solution is equal to the negative logarithm of the concentration of the acid in moles per liter (by the definition of pH). When 90% of the acid has been neutralized (f= 0.9), the concentration of H+ is only one-tenth of its what it was originally-- so the pH will be one unit higher, since -log(0.1) = 1. When 99% has been neutralized (f= 0.99), the pH is 2 units higher, and so on. When f = 1, the pH should equal 7--- and any further addition will raise the pH to a value equal to [14 minus pOH], just as though it were being added to pure water. (Note that we have made the simplifying assumption here of ignoring the increase in the volume of the solution due to adding the base-- but this could easily be accounted for. We also assumed that the original acid concentration was a lot higher than 10-7 molar, so that we could ignore the H+ contributed by the dissociation of water.) For a weak acid, an approximate formula can be derived for the pH of a solution of the pure acid which states that pH = 1/2( pKa + pC ) The pKa is the negative logarithm of the "acid dissociation constant", and pC is the negative logarithm of the concentration of the acid in moles per liter. (The pKa is a property of each particular acid, and is a number which can be looked up in reference books.) So, for example, a 0.1 molar solution (pC = 1) of an acid which has a pKa of 5, would have a pH of about 1/2(5+1), or 3. For the hypothetical acid with the formula, HA, the reaction which occurs as the titration with strong base proceeds can be written as: HA + OH- ===> A- + H2O The major chemical species in the solution are the remaining acid, HA, which has not been neutralized, and the anion (negative ion), A-, which is called the "conjugate base." It is the ratio of these which determines the pH, according to the formula: pH = pKa + log [(A- )/(HA)] where (A-) means the molar concentration of A- and (HA) is the molar concentration of the remaining HA. Note that this formula can also be expressed as pH = pKa + log [f/(1 - f < )] When the concentration of the two species is equal, the ratio [(A-)/(HA)] equals 1-- and since the logarithm of 1 equals zero, the pH is equal to the pKa. At an earlier point in the titration, when, say, one-tenth of of the acid had been neutralized, the pH would be equal to pKa + log (0.1/0.9). This works out to about 0.95 pH units below the pKa. When 90% of the titration is complete, the pH should be about equal to pKa + log (0.9/0.1), or about 0.95 units above the pKa. So the pH change during the middle 80% of the titration will vary less than one unit below or above the value of the pKa. Likewise, you can easily show that between the 1% and 99% points of the titration, the pH will vary between 2 units below and two units above the pKa. (Note that the same assumptions are made as for the strong acid case discussed above.) For a monoprotic acid (also called a "monobasic" acid-- how's that for a confusing term) at the end of the titration (f = 1), there is another approximate formula for the pH: pH = 7 +1/2( pKa - pC ) For the previous example of a 0.1 molar solution (pC = 1) of an acid which has a pKa of 5, the endpoint pH would be about 7 + 1/2(5-1), or 9. "Diprotic" (also called "dibasic") acids can be thought of of dissociating in two steps. For a generic dibasic acid H2Z, loss of one proton can be written as H2Z <===> H+ + HZ- for which pKa is called pK1 and loss of the second proton is written as HZ- <===> H+ + Z= for which pKa is called pK2 Since HZ- is negatively charged, and positive charges are attracted to negative charges, it is harder for the second proton to break away. Because of this, the value of pK2 is usually several units higher than pK1. Often, a solution of a dibasic acid behaves essentially like a mixture of two independent acids, one being a much weaker acid than the other. The titration curve runs from f = 0 to f = 2, and looks like one monobasic titration curve followed by another one at a higher pK. The pH's at f = 0.5 and f = 1.5 correspond to the values of pK1 and pK2, respectively. If the acid is concentrated enough that the contribution due to the dissociation of water can be ignored, the pH at f = 1 is about equal to the average of pK1 and pK2. (In cases where the pK's are fairly close the simple model does not work so well.) The pH range near the pKa value of a particular weak acid is sometimes referred to as the buffer region. As we have seen, the pH does not change much in this region when strong acid or base is added-- even in amounts which are a significant fraction of the amount of the weak acid/base mixture itself. This property is made use of in chemical, biological and pharmaceutical work-- and in nature-- to keep solutions at a near-constant pH. To make a buffer solution, you do not actually need to titrate a weak acid or base. For instance, to make an acetic acid/acetate buffer you can purchase acetic acid and the salt, sodium acetate, from a chemical supplier and make a solution containing the proportions which will give the desired pH, based on the formula given above. The buffer would be most efficient at a pH near the acid's pKa value of 4.7. In natural waters and in wastewater treatment plants, the water most often relies on the carbonic acid/bicarbonate system for buffering near neutral pH (pK = 6.3). The carbonic acid is formed when carbon dioxide dissolves in water. It is a product of aerobic or anaerobic respiration by microorganisms living in the water, and is also present in air; carbonates are present is some minerals, such as limestone, with which the water may come in contact. In laboratories, phosphate buffers are often used in chemistry or bacteriology to keep pH conditions constant. Phosphoric acid is a tribasic acid, with pK's of 2.1, 7.2, and about 12.0. You can see that the middle one, corresponding to a mixture of the ions H2PO4- and HPO4=, would be very useful for making neutral buffers. In wastewater analysis, phosphate buffers are used in the BOD test, the DPD method for total chlorine residual and the colorimetric test for cyanide, for diluting and rinsing in coliform bacterial testing, and for calibrating pH meters. As an example of the protective effect of buffer solutions, consider the 0.01 M (moles per liter) carbonic acid/bicarbonate/carbonate system shown in the last of the six titration curves. If we take the case of this system at a pH of 7.0, the graph shows that the value of f equals about 0.83. This means that of the total concentration of 0.01 moles per liter, 83% (0.0083 moles per liter) is in the form of bicarbonate (HCO3-) -- so that 17% (0.0017 moles per liter) is iin the form of carbonic acid (H2CO3). [ The ratio, 0.0083 / 0.0017, equals 4.88-- the logarithm of which is 0.69, or about 0.7. Add this to the pK1 of 6.3, according to the formula above, and you get a pH of 7.0] Now, let's say we add 0.001 moles of a strong acid to a liter of this solution. [Remember that adding this amount of strong acid to pure water will lower the pH from 7 down to a value of 3.] The reaction which would occur, HCO3- + H+ =====> H2CO3 would convert 0.001 moles of bicarbonate ion to carbonic acid, so the new concentrations would be 0.0073 moles per liter of bicarbonate and 0.0027 moles per liter of carbonic acid. The new ratio of concentrations would be (0.0073 / 0.0027), or 2.70, the logarithm of which equals 0.43, or about 0.4. So the new pH would be 6.3 plus 0.4, or 6.7-- a decrease of only 0.3 units. This would have much less of an effect on the chemistry and microbiology of the water than the drop of four pH units that would occur in an unbuffered solution. The protective species here is bicarbonate, and its "acid neutralizing capacity" would be referred to as the bicarbonate alklinity of the solution. [In terms of "mg/L as calcium carbonate", the amount of bicarbonate alkalinity in the original solution would have been equivalent to 0.0083 X 50,000, or 415 mg/L. After the addition of the strong acid, it would have been lowered to 365 mg/L.] Buffering is clearly an important feature of water quality control.
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Post by Robot on Jul 27, 2004 10:05:53 GMT -5
Acid Titration CurvesThe graphs on this page show what happens to the pH of an acid solution as it is neutralized with a strong base, such as sodium hydroxide. The letter, "f ", is used to represent the fraction of the acid which has been neutralized. Choose the acid in the left frame to display its titration curve in this frame. Scroll below the graph or graphs for descriptions. (The curves are simplified in that any increase in volume due to adding the base has been ignored.) These graphs were developed in Lotus 1-2-3 Release 5 for Windows© using equilibrium equations and constants which can be found in the textbook, "Aquatic Chemistry" by Werner Stumm and James J. Morgan, Wiley-Interscience, 1970. The graphs were copied to the clipboard and into a graphics program, then saved in gif format. These two graphs show the behavior of hydrochloric acid, a strong acid, at two concentrations: 1 molar and 0.005 molar. The pH's of the initial solutions are equal to the negative logarithms of the acid concentration-- 0 (zero) and 2.3, respectively. After 90% of the acid has been neutralized ( f = 0.9), the acid concentration is one tenth of the its original value, so the pH has gone up by one unit-- to 1.0 and 3.3, respectively. Similarly, after 99% has been neutralized, the pH is 2 units higher than it was at the start. After this, the pH rapidly approaches neutrality. As soon as more base has been added than is needed to neutralize all of the acid ( f > 1), the solution would rapidly become alkaline, in sort of a mirror image of the previous part of the curve. With a strong acid /strong base system like this the pH will be very low or very high most of the time. It is hard to control the pH at an intermediate value because a small addition of acid or base causes large changes in pH near the equivalence point-- that is, at f values near 1. These two curves show the behavior of a weak acid, acetic acid at the same two concentrations. Aside from the pH near the pure acid (f = 0) range, the curves are not very different from one another. The more concentrated acid solution has a lower pH. As strong base is added, the hydroxide ion reacts with acetic acid to form water and acetate ion, which is a weak base. Over most of the range, the pH is determined by the ratio of the concentration of undissociated acetic acid to the concentration of acetate ion. At the f = 0.5 point, where the acetic acid and acetate concentrations are equal, the pH is equal to the pK value of the acid, 4.7. On the other hand, the partially neutralized 1M acid will have a much greater " buffering" ability than the 0.005M acid. For example, with the 0.005M solution, if you have a liter of an acetic acid/acetate mixture with a pH of 5.0 (f = 0.7) and you add 0.002 moles of a strong acid, you will be moving the system to the point where f = 0.3)-- and the pH will be lowered to about 4.3 or 4.4. (See graph.) But with a 1M acetic acid/acetate buffer at pH7, adding the same amount of strong acid will only lower the f from 0.7 to 0.698, and the pH change will be almost undetectable. This curve illustrates hypchlorous acid, HOCl, the weak acid which forms -- along with hydrochloric acid-- when chlorine gas is dissolved in water. The curve is for a concentration equivalent to 1 mg/L chlorine. From the curve you can see that at a pH of 7, only 20% of the chlorine is in the basic form of hypochlorite ion, OCl- (f = 0.2), meaning that 80% is in the form of HOCl. At pH8, the situation is reversed. This has importance in water and wastewater treatment, as HOCl is a more potent disinfectant than OCl-. This last example is of a dibasic acid, carbonic acid, H2CO3. This is the acid which forms when carbon dioxide gas, CO2, dissolves in and reacts with water, H2O. As strong base is added, it reacts with the H2CO3 to form water and HCO3-, which is called bicarbonate ion. The pK of carbonic acid is 6.3 (making it a weaker acid than acetic acid). So a pH of 6.3 represents the middle of the first "buffer range" of this acid. If we continue to add strong base after all of the carbonic acid has been converted to bicarbonate ion, (at f = 1), the HCO3- reacts with the hydroxide to form water and carbonate ion, CO3=. So bicarbonate ion is both a weak base and a weak acid. As an acid, it has a pK of 10.3, so this is the middle of the second buffer range. The pH of a pure carbonate solution, such as could be made by dissolving powdered sodium carbonate in water, would depend on the concentration: the higher the concentration, the higher the pH (just as a higher concentration of pure carbonic acid would result in a lower pH.) A pure bicarbonate solution has a pH of about 8.3, and is less affected by the concentration.
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Post by Robot on Jul 27, 2004 10:07:25 GMT -5
Dissolved Oxygen (D.O.): Like solids and liquids, gases can dissolve in water. And, like solids and liquids, different gases vary greatly in their solubilities, i.e, how much can dissolve in water. A solution containing the maximum concentration that the water can hold is said to be saturated. Oxygen gas, the element which exists in the form of O2 molecules, is not very water soluble. A saturated solution at room temperature and normal pressure contains only about 9 parts per million of D.O. by weight ( 9 mg/L). Lower temperatures or higher pressures increase the solubility, and visa versa. Significance: Dissolved oxygen is essential for fish to breathe. Many microbial forms require it, as well. The oxygen bound in the water molecule (H2O) is not available for this purpose, and is in the wrong "oxidation state", anyway. The low solubility of oxygen in water means that it does not take much oxygen-consuming material to deplete the D.O. As mentioned before, the biodegradation products of bacteria which do not require oxygen are foul-smelling, toxic, and/or flammable. Sufficient D.O. is essential for the proper operation of many wastewater treatment processes. Activated sludge tanks often have their D.O. monitored continuously. Low D.O.'s may be set to trigger an alarm or activate a control loop which will increase the supply of air to the tank. Measurement: D.O. can be measured by a fairly tricky wet chemical procedure known as the Winkler titration. The D.O. is first trapped, or "fixed", as an orange-colored oxide of manganese. This is then dissolved with sulfuric acid in the presence of iodide ion, which is converted to iodine by the oxidized manganese. The iodine is titrated using standard sodium thiosulfate. The original dissolved oxygen concentration is calculated from the volume of thiosulfate solution needed. Measurements of D.O. can be made more conveniently with electrochemical instrumentation. "D.O. meters" are subject to fewer interferences than the Winkler titration. They are portable and can be calibrated directly by using the oxygen in the air.
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Post by Robot on Jul 27, 2004 10:10:04 GMT -5
Oxygen Demand: The biochemical oxygen demand, abbreviated as BOD, is a test for measuring the amount of biodegradable organic material present in a sample of water. The results are expressed in terms of mg/L of D.O. which microorganisms, principally bacteria, will consume while degrading these materials. As this method is a fairly long-term bioassay test (5 days), a more rapid (2 hour) test is often used to estimate the BOD; it is known as the COD, or chemical oxygen demand test. An even more rapid test, known as the TOC, or total organic carbon test takes only a few minutes, but requires expensive instrumentation. In the United States, most regulatory agencies specify the BOD test for permit reporting, especially for biological treatment plants. Significance: For reasons discussed earlier, the depletion of oxygen in receiving waters has historically been regarded as one of the most important negative effects of water pollution. Preventing these substances from being discharged into our waterways is a key purpose of wastewater treatment. Monitoring BOD removal through a treatment plant is necessary to verify proper operation. However, because the test takes too long to be useful for short-term control of the plant, the chemical or instrumental surrogate tests are often used as guides. Measurement: The BOD test is performed in a specially designed bottle with a flared cap which forms a water seal to keep out air. The bottle is filled completely with sample, which must be near neutral pH and free of toxic materials. After an initial measurement of the D.O., the bottle is sealed and stored in a dark incubator at 20C for five days. The D.O. is measured again after this incubation period. The difference is the BOD. (The bottles are kept in the dark because algae which may be present in the sample will produce oxygen when exposed to light.) Since most wastewaters have BOD's which are much higher than the limited solubility of oxygen in water, it is necessary to make a series of dilutions containing varying amounts of sample in a nutrient-containing, aerated "dilution water." The measured BOD's are then multiplied by the appropriate dilution factors. A variation of this test, called the carbonaceous BOD, adds an inhibitor which prevents the oxidation of ammonia, so that the test is a truer measure of the amount of biodegradable organic material present. Samples which do not contain enough bacteria to carry out the BOD test can be "seeded" by adding some from another source. Examples of samples which would need seeding are industrial wastewaters which may have been at high temperatures or high or low pH, or samples which have been disinfected. (If there is residual disinfectant present, it must be neutralized before testing.) BOD Calculator (300 mL bottle)www.geocities.com/RainForest/5161/bod1_js.htmThe calculations: Seed correction = Seed BOD x mL seed per bottle / 300 BOD = (initial D.O. - final D.O. -seed correction) x 300/mL sample in bottle If < 2 mg/L D.O. consumed, TOO WEAK. If < 1 mg/L D.O. remains, TOO STRONG. Otherwise, BOD is valid. The COD test is done by heating a portion of sample in an acidic chromate solution, which oxidizes organic matter chemically. The amount of chromate remaining (measured by a titration), or the amount of reduced chromium produced (measured spectrophotometrically), is translated into an oxygen demand value. Biodegradability, toxins, and bacteria are not important, and the test is complete in about two hours. The figure will be higher than the BOD. The TOC is done instrumentally. The organic carbon is oxidized to carbon dioxide by burning or by chemical oxidation in solution. The carbon dioxide gas is swept out and measured by infrared spectrometry or by redissolving it in water and measuring the pH change (the gas is acidic.) Both COD and TOC can often be correlated with BOD for a specific wastewater sample, but each wastewater is different. As a rough guide, the COD of a raw domestic wastewater is about 2.5 times the 5-day BOD.
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Post by Robot on Jul 27, 2004 10:14:31 GMT -5
Solids: Water, a liquid, can contain quite a bit of solid material, both in dissolved and suspended forms. The term "dissolved" implies that the individual molecules of a substance are mixed in among the water molecules. In practice, solids are classified as "dissolved" if they pass through a standard glass-fiber filter with about one micrometer pore size. Solids captured on the filter are, by definition, "suspended" solids.. Solids which settle out of a water sample on standing for a period of an hour are defined as "settlable." . Solids are also further classified as "fixed" or "volatile." Fixed solids are basically the ash left over after burning the dried solids; volatile solids are those that are lost in this procedure. The sum of the two is referred to as "total". (This can be confusing, as the word "total" is also used in describing the sum of suspended and dissolved solids.) Volatile solids are often used as an estimate of the organic matter present. Significance: Solids in wastewater contribute to sediment formation; volatile solids may be associated with oxygen demand. Measurement: Total solids (TS) are determined by drying a known amount of a sample at a temperature of 103 to 105 C in a tared (pre-weighed) vessel, such as a porcelain dish, cooling in a dry atmosphere (in a container known as a desiccator), weighing on an analytical balance, subtracting the tare weight, and dividing by the original amount of sample. Results can be expressed in mg/L if the sample was originally measured out by volume; or percent by weight, if the sample was originally weighed. If the sample is then burned in a furnace at about 500 C, cooled, and weighed, the fixed (FS) or volatile solids (VS) can be determined. . If the original sample is filtered through a tared glass-fiber filter, which is then dried, the weight of the material captured on the filter is used to figure the total suspended solids (TSS). Burning the filter in the furnace allows measurement of volatile suspended solids (VSS) or fixed suspended solids (FSS). % Solids and % Volatile Solids Calculatorwww.geocities.com/RainForest/5161/pctsv_js.htmThe calculations: Percent solids = [(Dish plus dry sample - tare) / (Dish plus wet sample - tare)] * 100% , which is the same as (dry sample weight / wet sample weight) * 100% Percent volatile solids (i.e., percent of dry solids which are volatile) = [(Dish plus dry sample - Dish plus ash) / (Dish plus dry sample - tare)] * 100%, which is the same as (weight lost on ignition) / (dry sample weight) * 100% Total and Volatile Suspended Solids Calculatorwww.geocities.com/RainForest/5161/ss_js.htmThe calculations: TSS, mg/L = [(Filter plus dry sample,grams) - (Filter tare weight,grams)] / (Volume filtered,mL) * 1,000,000 VSS, mg/L = [(Filter plus dry sample,grams) - (Filter plus ash,grams)] / (Volume filtered,mL) * 1,000,000 The dissolved solids (DS) can be estimated from the difference between the total solids and the total suspended solids, but the official method calls for drying the filtrate (the liquid which passes through the filter) in a dish at 180C. (And, of course, there are TDS, FDS and VDS). An estimate of total suspended solids can be obtained by an optical/instrumental measurement known as turbidity. The sample is placed in a glass tube; a beam of light is shined through it, and the light scattered at right angles to the beam is measured photometrically. In the same way that COD can be correlated with BOD, turbidity can be correlated with TSS; but the correlation will hold only for the particular sample from which it was derived. Similarly, an estimate of dissolved solids is often made by measuring the water's electrical conductivity. Pure water does not conduct electricity. If substances which dissociate into electrically charged ions are dissolved in the water, they will conduct a current, roughly proportional to the amount of dissolved substances. Conductivity can be used to track sewage pollution. Note, however, that many organic materials dissolve in water without producing ions. So, while a salt solution may have a high electrical conductivity, a concentrated solution of sugar would go undetected by this method.
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