Chemistry and the Aquarium: What is Alkalinity? — Advanced Aquarist | Aquarist Magazine and Blog H3PO4 H2PO4- HPO4-- PO4--- Alkalinity using Test Kits Of course, most reefkeepers measure alkalinity with a test kit, not with a pH titration. How does that work? Well, in effect test kits do a pH endpoint titration. They all include pH indicating dyes (providing a color change) and an acid (frequently dilute sulfuric acid) to lower the pH. You typically add acid until the dyes turn color. Since these dyes are selected to have a color change in the pH = 4 to 5 range, what you get is a measurement of how much acid it takes to lower the pH to that range. This color change is used to approximate the endpoint of the titration. Interestingly, many test kits use more than one pH indicating dye. Using more than one dye at the same time permits the endpoint to be sharper.

For example, bromcresol green has a broad color transition between pH 3.8 (yellow) and 5.4 (blue-green) and methyl red has a broad transition between pH 4.4 (red) and 6.2 (yellow). A mixture of the two (used in the Hach alkalinity kit) has a sharp transition (orange to blue-green) around pH 5.1 in fresh water (which may be slightly different in salt water). Five point 1 you say? Based on the discussion above, is that low enough? Well, the Hach kit was designed for use in fresh water where the pKa of the bicarbonate is much higher than in seawater, and in that situation, it is appropriate. In seawater, however, it is marginal. My tank water took 3.4 meq/L to get down to pH = 5.03, and then an additional 0.4 meq/L to get down to pH 4.00. Consequently, this kit (and others with a similar dye mix) may be missing out on 10% of the alkalinity simply because it isn't titrating low enough.

This difference obviously isn't significant to most reef keepers, but is something to keep in mind when doing such things as comparing test kits to standards (in seawater) or to each other. Some test kits also provide a different dye for a different measure of alkalinity. Frequently, this other dye is phenolphthalein. This dye has a color change between pH 8.2 and pH 9.8. In fresh water, carbonate is almost completely converted into bicarbonate at pH 8.3, and that is the purpose of phenolphthalein titrations : to determine alkalinity in freshwater due to carbonate only (discussed in detail below). This test serves no purpose in a reef tank or seawater for two reasons: 1) the water is probably already more acidic than the endpoint of this dye, and 2) the carbonate in seawater is not completely converted into bicarbonate at this pH anyway. That is, even if the pH were higher than 8.3 (say, 8.6), titrating down to the phenolphthalein endpoint will not effectively "count" all of the carbonate because in saltwater there will still be substantial carbonate present at the phenolphthalein endpoint.

Why is Alkalinity Important? Now that we know what alkalinity is, we can understand why it is an important measure for reef tanks. Corals and other organisms deposit calcium carbonate in their skeletons and other body parts. In order to do this they must generate calcium and carbonate at the surface of the growing calcium carbonate crystal. While it is far beyond the scope of this paper to describe this process, it is readily apparent that if corals deposit these chemicals, they are using them up from the water that they inhabit. So, if that's the case, why not just measure carbonate as we do calcium? Well, there are two answers. The first is that there is no simply way to measure carbonate with a kit without doing a pH titration as an alkalinity test kit does. Second, corals may actually use bicarbonate instead of carbonate as their ultimate source of carbonate (which they split into H+ and CO3--). If we could easily measure bicarbonate, we'd probably be doing just that. Unfortunately, we can't do either of those things easily. So what we are doing is using a very simple alkalinity test as a surrogate measure for bicarbonate and carbonate.

Since these two substances comprise the great majority of alkalinity in seawater, it is safe for most people to equate alkalinity with "availability of bicarbonate and carbonate for my corals". There are, however, some important caveats to that equation. Some of these were described above, such as salt mixes that have excessive borate. Such complications make it difficult to know how much of the measured alkalinity is bicarbonate and carbonate, and thus it is difficult to know if you are satisfying the needs of the corals [Hence the unusually high alkalinity recommendations by Seachem]. Unusual Contributors Reef tanks can also have contributors to the total alkalinity that are simply not present in seawater at any appreciable concentration. This result comes from the fact that we have a closed system in which organics (e.g., acetate, polygluconate, EDTA; citric acid) and other ions may be unusually high. As an example, consider those people who are dosing limewater with organic acids such as vinegar . Acetic acid is a complication to an alkalinity test that may or may not be significant to people using it, but the more vinegar that is used, the more confounding it may become.

Ultimately, the acetate that is added in this fashion will be oxidized into CO2 and OH- (equation 4), with the OH- providing alkalinity in the same fashion that the original limewater would have. (4) 2 O2 + CH3COO- ==> H2O + 2 CO2 + OH- The issue at hand is how fast this conversion takes place, or alternatively, how much acetate is present in such a system when one measures the alkalinity. Since I've seen no studies of acetate levels in reef tanks, the question remains unanswered (at least to me). Figure 2. A pH titration of pure water, water from the author's reef tank, and water from the author's reef tank with 30 mM of sodium acetate added (titrated with 0.1 N HCl). The potential for a problem comes about because acetate is partially "counted" in a total alkalinity titration of tank water. The extent to which it is counted will depend upon what pH is being used as the titration endpoint. Figures 2 and 3 show the pH titration of tank water with a huge excess of acetate added (30 mM).

This excessively large amount was added not because a reef tank would contain such a large amount (after all, the measured total alkalinity is about 20 meq/L), but because it makes the acetate titration clearly visible in the presence of carbonate and bicarbonate. If the endpoint of the alkalinity titration is at pH 5, then about 25% of the acetate is counted. With the endpoint at pH 4, about 80% is counted. Figure 3. An extended pH titration of pure water, water from the author's reef tank, and water from the author's reef tank with 30 mM of sodium acetate added (titrated with 0.1 N HCl). Consequently, if a tank has marginal alkalinity and some substantial portion of this alkalinity is acetate (or some other organic), then the availability of bicarbonate and carbonate may be less than optimal for corals and other calcifying organisms. Note that the acetate does not impact the titration of carbonate between the native pH and about 7.3.

If one is using large amounts of vinegar, it might be worthwhile to titrate the carbonate down to 7.3 to verify that the total alkalinity is not being dominated by acetate (by observing at least 0.2-0.4 meq/L alkalinity down to pH = 7.3. My tank water without acetate had 0.6 meq/L for this titration (Figure 1) and the same when a large amount of acetate was added (Figures 2 and 3). Alkalinity Facts There are several facts about total alkalinity that follow directly from the definition. Unfortunately, some of these have been misunderstood by some hobby authors. One of these facts is termed The Principle of Conservation of Alkalinity by Pankow ("Aquatic Chemistry Concepts", 1991). He shows mathematically that the total alkalinity of a sample CANNOT be changed by adding or subtracting CO2. Unfortunately, there is an article available on line, which claims otherwise, and encourages people to "lower alkalinity" by adding CO2 in the form of seltzer water. This is simply incorrect.

Forgetting the math for the moment, it is easy to see how this must be the case. If carbonic acid is added to any aqueous sample with a measurable alkalinity, what can happen? Well, the carbonic acid can release protons by reversing equations 1 and 2: (5) H2CO3 ==> H+ + HCO3- (6) HCO3- ==> H+ + CO3-- These protons can go on to reduce alkalinity by combining with something that is in the sample that provides alkalinity (carbonate, bicarbonate, borate, phosphate, etc). However, for every proton that leaves the carbonic acid and reduces alkalinity, a new bicarbonate or carbonate ion is formed that adds to alkalinity, and the net change in total alkalinity is exactly zero. The pH will change, and the speciation of the things contributing to alkalinity will change, but not the total alkalinity. This is not true for strong acids, however. If you add hydrochloric, sulfuric or phosphoric acids (or any acid with a pKa lower than the carbonic acid endpoint), there will be a reduction in the alkalinity.

Another interesting result of the Principle of Conservation of Alkalinity is the equation for determining the total alkalinity when two different aqueous solutions are mixed together. If you mix (a) parts of a solution with total alkalinity A with (b) parts of a solution of total alkalinity B, the resulting alkalinity is just the weighted average of the two samples: TAmix = [a(A) + b(B)]/[a + b] Equation 7 can be used to calculate changes in TA for water changes in a tank, for additions of limewater, for dilution of tank water with pure water, and a host of other situations where you might want to know what the final alkalinity will be. It can also be used for calculating reductions in alkalinity caused by strong acids, where the alkalinity of the acid is just the normal strength of the acid as a negative number. Other Definitions of Alkalinity Any definition of alkalinity other than the total alkalinity seems to lead to confusion. For example, Millero defines the carbonate alkalinity (AC) as the alkalinity coming from just bicarbonate and carbonate (equation 8). Some test kits use this definition as well.

(8) AC = [HCO3-] + 2[CO3--] Unfortunately, another leading author, Pankow, defines carbonate alkalinity (CO3-- - Alk) as the total alkalinity down to the pH where all carbonate is converted into bicarbonate (the bicarbonate equivalence point or endpoint; about pH 8.3 in fresh water; about pH 7.3 in seawater). Consequently, it doesn't count bicarbonate at all, and does count borate and other ions that take up acid above the carbonate endpoint. For freshwater, this type of alkalinity is represented by the phenolphthalein endpoint used in the Hach and other kits. Others define carbonate alkalinity as just that portion of total alkalinity down to the carbonic acid endpoint that comes from carbonate ions, exclusive of bicarbonate, hydroxide, borate, etc. And there are still other definitions of alkalinity. The hydroxide alkalinity (OH- - Alk), sometimes called the caustic alkalinity, is defined by some as the total alkalinity down to the carbonate equivalence point (about pH 10.7 in fresh water).

One test kit ( Seachem ) provides a test for borate and hydroxide alkalinity. I have not tested this kit to know whether it is effective or not. Because of these potential points of confusion, in any discussion of alkalinity other than the total alkalinity, one needs to be very clear about the definitions being used. Units of Alkalinity The various units used for alkalinity are themselves cause for confusion. The clearest unit, and that used by most scientists is milliequivalents per L (meq/L). For a 1 millimolar solution of bicarbonate, the alkalinity is 1 meq/L. Since carbonate takes up two protons for each molecule of carbonate, it "counts" twice, and a 1 millimolar solution of carbonate has an alkalinity of 2 meq/L. A unit that is used by many kits and some industries involves representing alkalinity in terms of the amount of calcium carbonate that would need to be dissolved in fresh water to give the same alkalinity. Typically, it is reported as ppm calcium carbonate. Of course, it has nothing to do with calcium, and there may be no carbonate in the water at all. Nevertheless, it is frequently used.

Since calcium carbonate weighs 100 grams/mole (100 mg/mmole), then a solution that has an alkalinity of 100 ppm calcium carbonate equivalent contains 100 mg/L calcium carbonate divided by 100 mg/mmole calcium carbonate = 1 mmol/L calcium carbonate equivalent. Since carbonate has 2 equivalents per mole, this 100 ppm of alkalinity is equivalent to 2 meq/L. So to convert an alkalinity expressed as ppm CaCO3 to meq/L, divide by 50. Figure 4. Two Trachyphyllia geoffroyi in the refugium of the author's tank. The pH of the refugium is often higher than the tank due to the photosynthetic processes reducing the CO2, but the alkalinity is largely unchanged as described by the Principle of Conservation of Alkalinity. Finally there is the German term dKH (degrees of carbonate hardness), or just KH (carbonate hardness).Strictly speaking, it is the same as the carbonate alkalinity (AC in equation 8). Unfortunately, it is a very confusing term, as it has nothing to do with hardness.

Further, it has been corrupted by the marine aquarium hobby to mean the same as total alkalinity, and every test kit that tests for dKH with a single titration is giving total alkalinity. The only kit that I am aware of that even makes a distinction between carbonate alkalinity and total alkalinity is one of the Seachem kits (Reef Status: Magnesium, Carbonate, & Borate) and it thankfully doesn't use the term dKH at all. Consequently, most hobbyists should think of dKH as simply another measure of total alkalinity. The results obtained with such a kit (dKH) can be divided by 2.8 to yield the alkalinity in meq/L. For those who are mathematically challenged, here is an alkalinity conversion table for all three units. Conclusion I hope this article provides a detailed understanding of alkalinity, from what it is and how it is measured, to why it is important in coral reef tanks. I also hope that it serves to clear up some of the confusion about alkalinity and how it is impacted by carbon dioxide and other acids. Document Actions pH Measurements - What Is pH? pH Measurements pH Measurements What Is pH and What Does It Measure?

Litmus paper is a type of pH paper that is used to test the acidity of water-based liquids. David Gould, Getty Images pH is a logarithmic measure of the hydrogen ion concentration of an aqueous solution: pH = -log[H+] where log is the base 10 logarithm and [H+] is the hydrogen ion concentration in moles per liter pH describes how acidic or basic an aqueous solution is, where a pH below 7 is acidic and a pH greater than 7 is basic. pH of 7 is considered neutral (e.g., pure water). Typically, values of pH range from 0 to 14, although very strong acids may have a negative pH , while very strong bases may have a pH exceeding 14. The term "pH" was first described by Danish biochemist Søren Peter Lauritz Sørensen in 1909. pH is an abbreviation for "power of hydrogen" where "p" is short for the German word for power, potenz and H is the element symbol for hydrogen . Why pH Measurements Are Important Chemicals reactions in water are affected by the acidity or alkalinity of the solution.

This is important not only in the chemistry lab, but in industry, cooking, and medicine. continue reading below our video 10 Facts About the Titanic That You Don't Know pH is carefully regulated in human cells and blood. The normal pH range for blood is between 7.35 and 7.45. Variation by even a tenth of a pH unit may be fatal. Soil pH is important for crop germination and growth. Acid rain caused by natural and man-made pollutants changes the acidity of soil and water, greatly affecting living organisms and other processes. In cooking, pH changes are used in baking and brewing. Since many reactions in everyday life are affected by pH, it's useful to know how to calculate and measure it. How pH Is Measured There are multiple methods of measuring pH. The most common method is a pH meter, which involves a pH-sensitive electrode (usually made of glass) and a reference electrode. Acid-base indicators change color in response to different pH values. Litmus paper and pH paper are used for quick, relatively imprecise measurements. These are strips of paper that have been treated with an indicator. A colorimeter may be used to measure the pH of a sample.

A vial is filled with a specimen and a reagent is added to produce a pH-dependent color change. The color is compared against a chart or standard to determine the pH value. Problems Measuring Extreme pH Extremely acidic and basic solutions may be encountered in laboratory situations. Mining is another example of a situation that may produce unusually acidic aqueous solutions. Special techniques must be used to measure extreme pH values below 2.5 and above around 10.5 because the Nernst law isn't accurate under these conditions when glass electrodes are used. Ionic strength variation affects electrode potentials . Special electrodes may be used, otherwise it's important to remember pH measurements won't be as accurate as those taken in ordinary solutions. Information on swimming pool water balance, testing and chemical management Pool Library/Links Pool Info: Pool Water Balancing Pool Water balancing is not such a complicated exercise. It is simply the relationship between different chemical measurements in your pool water. Your pool water is constantly changing, year round. Everything from weather to oils, dirt, and cosmetics affect your pool water balance - in short, anything that comes in contact with your pool water. You will probably not change the water in your pool for many years.

Continuous filtration and disinfection removes contaminants which keep the water enjoyable but this is does not balance your water. A pool that is "balanced" has proper levels of pH , Total Alkalinity , and Calcium Hardness . It may also be defined as water that is neither corrosive or scaling . This concept of water balance is derived from the fact that water will dissolve and hold minerals until it becomes saturated and cannot hold any more water in solution. It reminds me of that school science experiment, where you dissolve spoonful after spoonful of salt (or sugar) in a beaker, until at last - the water will dissolve no more salt into solution, and the grains just sit on the bottom of the beaker. When water is considerably less than saturated it is said to be in a corrosive or aggressive condition. When water is over saturated and can no longer hold the minerals in solution it is in a scaling condition. So then, balanced water is that which is neither over or under-saturated. The cliché that "water seeks its own level" certainly applies here. Water which is under-saturated will attempt to saturate itself by dissolving everything in contact with it in order to build up its mineral content.

Water which is over-saturated will attempt to throw off some of its content by precipitating minerals out of solution in the form of scale. How do we know when our water is over or under saturated? Use a good test kit (with fresh testing reagents ) to measure the chemical levels of pH, alkalinity, and calcium hardness. Any good discussion on pool water balance would certainly cover some definitions on pH, Total Alkalinity and Calcium Hardness levels in pools. pH pH is a measure of how acidic or basic the water is. pH is a logarithmic scale from 0-14, with 7 being neutral. Below 7.0 and a substance is defined as being acidic, while levels above 7 are said to be basic or alkaline. Everything that enters your pool has a pH value. Ever heard of acid rain? This is rainfall with a very low pH. The human eye at a pH value of 7.35, is just slightly basic. This is, coincidentally, in range with proper pH levels for your pool.

To have pH in balance we adjust the water with additions of pH increasers (bases) or pH decreasers (acids) to achieve the range of 7.2 - 7.6, or more precisely, 7.4. If your testing (recommended daily) of the water shows a pH value below 7.0; the water is in a corrosive (acidic) condition and you will need to add a base to bring the pH into a more basic range to prevent corrosion. It doesn't take too long for a low pH condition to weaken vinyl, strip heat exchangers and erode plaster. Conversely, if the pH is above 7.8, we are in a scaling (basic) condition and must add an acid to bring down the pH to prevent the formation of scale or calcium deposits on our tile, in the filter, or even precipitation out of solution into a cloudy water condition. Another effect of high pH is that chlorine becomes less active, almost sluggish. At a pH of 8.0, only about 20% of your chlorine is available to kill pathogens and contaminants.

Total Alkalinity A close cousin of pH, the level of alkalinity in the water is a measurement of all carbonates , bicarbonates , hydroxides , and other alkaline substances found in the pool water. pH is alkaline dependent; that is, alkalinity is defined as the ability of the water to resist changes in pH. Also known as the buffering capacity of the water, alkalinity keeps the pH from "bouncing" all over the place. Low alkalinity is raised by the addition of a base (similar to pH); sodium bicarbonate is commonly used. High levels of alkalinity are lowered by the addition of an acid (similar to pH). We recommend "pooling" the acid in a small area of low current for a greater effect on alkalinity. That is, adding an acid underwater, in a deep end corner will create a greater hydrocarbon exchange and have a greater effect on lowering TA in your pool. Acids will lower both pH and alkalinity, but “walking the acid” around the pool in a highly distributed manner is said to have a greater effect on lowering the pH than the alkalinity. Pooling the acid has the opposite effect.

Some manufacturers notably have created acids specifically designed to lower Alkalinity strongly, while having less effect on pH level. A very important component of water balance, alkalinity should be maintained in the 80-120ppm range for "gunite" and concrete pools and 125-170ppm for painted, vinyl, and fiberglass pools. Levels should be tested weekly. Follow package label for treatment guidelines. Calcium Hardness When we speak of scale, we are talking about calcium carbonate which has come out of solution and deposited itself on surfaces. It is a combination of carbonate ions , a part of total alkalinity and calcium, and a part of the Calcium Hardness level. The test for Calcium Hardness is a measure of how "hard" or "soft" the water is. Hard water can have high levels of calcium and magnesium . If these levels are too high, the water becomes saturated and will throw off excess particles out of solution which then seek to deposit themselves on almost any surface inside the pool. They can be attracted to ladders, lights and in extreme cases deposit themselves as very small crystalline clumps - all over the pool surfaces.

Calcium Carbonate scale; a "white-ish," crystallized rough nodule. If the Calcium Hardness levels are too low, the water is under-saturated. If under-saturated, the water will become aggressive as it attempts to obtain the calcium it needs. Such "soft-water" will actually corrode surfaces inside the pool which contain calcium (like pool plaster) and other minerals to maintain its hardness demand. If your Calcium Hardness levels are too high you can use a product called CalTreat to correct, with varying levels of success. In most cases you need not worry if your calcium levels are below 500, but much higher than that and it can cause pool problems. Hardness can also be reduced by dilution (adding water to the pool which has a lower calcium hardness content). Levels which are too low require the addition of calcium chloride . Recommended range for calcium hardness is 200-400ppm. Calcium Hardness levels should be tested weekly with fresh reagents or strips. The Saturation Index Also called the Langelier Index, this chemical equation or formula is used to diagnose the water balance in aquatic environments (pools).

The formula is "SI = pH + TF + CF + AF - 12.1". To calculate the Saturation Index, test the water for pH, temperature, calcium hardness, and total alkalinity. Refer to a chart for assigned values for your temperature, hardness, and alkalinity readings then add these to your pH value. Subtract 12.1, which is the constant value assigned to Total Dissolved Solids and a resultant number will be produced. A result between -0.3 and +0.5 is said to indicate balanced water. Results outside of these parameters require adjustment to one or more chemical components to achieve balance. This formula is not guaranteed; however, some readings for pH, calcium, and alkalinity which, if taken individually would be considered to be well beyond recommendations, can combine within the formula to produce "balanced water". The SI can be used to pinpoint potential water balance problems, and the water’s propensity for precipitation or the chance of corrosion. What about Total Chlorine and Cyanuric Acid? These are important measurements of your water cleanliness, but have nothing to do with water balance, really.

Chlorine tests should be done daily, or as needed, to maintain a 1-3 ppm level at all times, depending on your level of cyanuric acid. Higher cyanuric acid levels will require higher chlorine levels, because cyanuric acid has a suppressive effect on the activity of free chlorine. Cyanuric acid is much more stable, and can be tested monthly, to maintain a 20-40 ppm range. Cyanuric acid levels above 50 ppm should be reduced by dilution; that is, draining a portion of pool water and refilling with unstabilized water. What Is the pH Scale? Acidity, Alkalinity & Buffers Contact What Is an Acid? An acid is any ionic compound that releases hydrogen ions (H+) in solution. Ionic compounds are molecules held together by ionic chemical bonds that result from the attraction of oppositely charged atoms, called ions. Article Summary: The pH scale is a measurement of the acidity or alkalinity of a solution. To understand the scale, it is important to first understand what acids and bases are. Acids, Bases, Buffers, and the pH Scale Need Help with Chemistry?

SPO offers free PowerPoint lectures, sample test questions, review questions and assignments on many different science topics. Below are links to those that relate to chemistry: Interactive pH Scale Simulation UNIVERSITY OF COLORADO Test the pH of things like coffee, spit, and soap to determine whether each is acidic, basic, or neutral. Visualize the relative number of hydroxide ions and hydronium ions in solution. Switch between logarithmic and linear scales. Must have updated version of Java to run simulation. Weak acids release fewer hydrogen ions and have a sour taste. Strong acids release more hydrogen ions and are highly corrosive. The more hydrogen ions in the solution, the more acidic it is. What Is a Base? Bases, also called alkaline substances, are compounds that release hydroxyl ions (OH-) in solution. Weak bases have a bitter taste (opposed to sour taste of acids and sweetness of aldehydes and ketones) and are slimy to the touch. The more hydroxyl ions in a solution, the more alkaline it is. Classroom You have free access to a large collection of materials used in a college-level introductory Cell Biology Course.

The Virtual Cell Biology Classroom provides a wide range of free educational resources including Power Point Lectures, Study Guides, Review Questions and Practice Test Questions. Page last updated: 8/2015 Analysis Method for pH and Alkalinity | UMass Amherst Water Resources Research Center Analysis Method for pH and Alkalinity ** Background Information pH is a measure of the hydrogen ion concentration of the water as ranked on a scale of 1.0 to 14.0. The lower the pH of water, the more acidic it is. The higher the pH of water, the more basic, or alkaline, it is. pH affects many chemical and biological processes in the water and different organisms have different ranges of pH within which they flourish. The largest variety of aquatic animals prefer a pH range of 6.5 - 8.0. pH outside of this range reduces the diversity in the stream because it stresses the physiological systems of most organisms and can reduce reproduction. Low pH can also allow toxic elements and compounds such as heavy metals to become mobile and "available" for uptake by aquatic plants and animals. Again, this can produce conditions that are toxic to aquatic life, particularly to sensitive species like trout.

Changes in acidity can be caused by atmospheric deposition (acid rain or acid shock from snowmelt), surrounding rock, and wastewater discharges. Technically, the pH scale measures the logarithmic concentration of hydrogen (H+) and hydroxide (OH-) ions, which make up water (H+ + OH- = H20). When both types of ions are in equal concentration, the pH is 7.0 or neutral. Below 7.0, the water is acidic (there are more hydrogen ions than hydroxide ions). When the pH is above 7.0, the water is alkaline, or basic (there are more hydroxide ions than hydrogen ions). Since the scale is logarithmic, a drop in the pH by 1.0 unit is a 10-fold increase in acidity. So, a water sample with a pH of 5.0 is ten times as acidic as one with a pH of 6.0. pH 4.0 is 100 times as acidic as pH 6.0. Alkalinity is a measure of a river's "buffering capacity," or its ability to neutralize acids.

Alkaline compounds in the water such as bicarbonates (baking soda is one type), carbonates, and hydroxides remove H+ ions and lower the acidity of the water (which means increased pH). They do this usually by combining with the H+ ions to make new compounds. Without this acid neutralizing capacity, any acid added to a river would cause an immediate change in the pH. Measuring alkalinity is important to determining a river's ability to neutralize acidic pollution (as measured by pH) from rainfall or snowmelt. It's one of the best measures of the sensitivity of the river to acid inputs. Alkalinity comes from rocks and soils, salts, certain plant activities, and certain industrial wastewater discharges. Total alkalinity is measured by collecting a water sample, and measuring the amount of acid needed to bring the sample to a pH of 4.2. At this pH all the alkaline compounds in the sample are "used up." The result is reported as milligrams per liter (mg/l) of calcium carbonate.

The Massachusetts Acid Rain Monitoring Project ranks waters according to their alkalinity as follows: <0* mg/l: Acidified Cooler and Koolits or ice Zip-loc bag (1 gallon size) Method Rinse a 500 ml bottle (including cap) three times with surface water at the sampling site. Be sure to empty your rinse water away from your sampling location. Lower open sample bottle upside down to arm's length. Underwater, turn bottle upside right and let it fill up. When no more bubbles come up, cap the bottle underwater and bring back up to the surface. Put sample in cooler for delivery to lab , record sample ID and write 'pH/ANC' in the Chemistry column Holding time is 8 hours for pH and 14 days for alkalinity. Note: Put enough ice or equivalent in cooler to reduce sample temperature to 4° C as soon as possible. Sampling Method for Rivers Cooler and Koolits or ice Zip-loc bag (1 gallon size) Sampling Protocol Samples should be take from representative, flowing water. The water must be deeper than the sample bottles and free of surface scum and debris.

If the water is not deep enough at your regular sampling site, look for another location nearby which is equally representative of the site but deeper. If there is none, do not collect a sample and indicate on your field sheet that water level is too low. Note that sampling from the streambank is discouraged, as it can result in non-representative samples. Carefully wade into the stream, walking upstream and avoiding to stir up bottom sediment. Wait for pre-disturbance (from wading in) conditions to return before taking sample. If you are in a canoe, have your partner steady it. Take sample in mid-stream, if possible. If not, get as far out from shore as is safe. Walk upstream and collect sample so that you are not standing or floating upstream of the bottle. Uncap sample bottle and rinse three times with river water: fill bottle partially, cap, shake, and empty downstream. To take sample, dip bottle completely under water, filling to overflowing. Cap bottle while it is still underwater, in order to eliminate any air from the sample bottle. Return to shore and place sample in cooler with ice.

Stir Bar 100ml Graduated Cylinder (optional for pH measurement) After calibrating your meter with the buffers, rinse the electrode(s) and glassware with distilled or deionized water. Carefully measure 100 ml of your sample and place in a 150 ml beaker for the pH and alkalinity part. Place the rinsed electrode in the test sample. We strongly encourage letting all samples come to room temperature in the tightly capped bottle before analyzing. If you are conducting other analyses with the sample water, keep in mind that pH should be analyzed within 5 minutes of uncapping the sample bottle. The sample should be stirred very gently, preferably with a magnetic stirrer. It may take up to 3 minutes for the reading to become stable. When stable, but not in excess of 5 minutes, record the sample pH to the nearest 0.01 pH unit Alkalinity Measurement Digital Titrator 0.16N Sulfuric Acid Cartridge After placing the sulfuric acid cartridge in position in the Hach Digital Titrator, be sure to advance the plunger manually until titrant is forced out of the delivery tip.

Do this as you would a hypodermic syringe, with the delivery tip up to remove bubbles. Get all the bubbles out! Then advance the plunger using the delivery knob on the end of the titrator until you are sure that the delivery tip is filled with solution. Check for leaks where the tip connects to the cartridge. Rinse the tip WELL with distilled water or sample; this is important because the titrant is concentrated and a little bit goes a long way. Reset the counter to zero and you are ready to titrate. After completing a titration and recording the digits of titrant used, rinse the delivery tip with distilled water or the next sample, reset the counter (THIS IS EASILY FORGOTTEN WHEN BUSY), and you are immediately ready for the next sample. Titrations go better if the delivery tip is positioned under the surface of the solution being titrated. For one or two samples, the titrator can be held in the hand, however, it is easier to mount the titrator on a ring stand using a clamp.

Try to keep the titrator vertical through all titrations; putting the titrator horizontally on the bench between titrations may introduce bubbles in the tip. The acid cartridges provided are 0.16N sulfuric acid. Our waters are typically quite low in alkalinity, so we use a special double end-point alkalinity procedure to accurately measure alkalinity below 20 mg L-1. After reading and recording the pH as described above, titrate with the digital titrator and sulfuric acid cartridge to pH 4.5; record titrant used to this point as A. Continue the titration to pH 4.2. Record the titrant used to this point as B. If the initial pH is less than 4.5, record the initial pH value. Titrate until the pH is 0.3 units below the starting point. Enter the digits of titrant used as B; A = 0. Write down the pH reading where you stopped (as an accuracy check). We will use computers to calculate the alkalinity, but you may do your own calculations using the formulas below. The examples will help to clarify what can be somewhat confusing formulas.

A = digits used to pH 4.5 B = digits used to pH 4.2 or 0.3 pH units below initial value (total titrant including A) Double end-point alkalinity= (2A - B) x 0.1 EXAMPLE: A sample required 120 digits to reach pH 4.5. An additional 15 digits were required to reach pH 4.2, for a total of 135 digits. Therefore, A = 120 and B = 135. Double end-point alkalinity = (240 - 135) x 0.1 = 10.5 mg/l EXAMPLE: A sample had an initial pH of 4.3. The sample required 22 digits to lower the pH to 4.0. Therefore, A = 0; B = 22. Double end-point alkalinity = (0 - 22) x 0.1 = -2.2 mg/l Although the negative alkalinity value may seem not to make much sense, it is an extremely important measurement for assessment of acidification.

pH Electrode Care and Maintenance pH & alkalinity QA/QC General electrode care and handling procedures are very important in your lab because pH measurements will only be as good as the condition of your electrode(s). For greater accuracy in your measurements and longer electrode life, there are a few areas of electrode care with which you should be familiar. Storage Glass combination or separate pH and reference electrodes should be kept wet. The reference electrode requires a free-flowing junction, so be sure to maintain the reference filling solution at a level significantly above the storage or sample solution level at all times This will provide a positive head pressure, which forces the filling solution out through the junction rather than the storage solution into the probe. For dry storage, the sleeve or plug should cover the filling hole to reduce the flow of filling solution. During the measurement or storage in pH 4 buffer, however, this sleeve or plug must be slid away or removed to allow flow of the reference solution into the sample. To obtain a faster electrode response, the glass electrode should be stored in a slightly acidic solution. In the protective cap for the glass electrode, put a drop or two of pH 4 buffer and put the cap on the electrode, carefully.

Distilled water extracts ions from the bulb causing a slower response; pH 7 buffer over a long time period ages the electrode slightly. If using a separate reference electrode, the best solution would be to place the reference electrode in its own filling solution but this can be messy. Providing KCl to both sides of the junction keeps it flowing freer. To reduce the salt crust of saturated solution, an approximately 0.1 M KCl solution may be used, but for storage only. Experience indicates that simply covering the filling hole with the protective sleeve and storing dry suffices in most instances as long as the soaking procedure is followed. For combination electrodes, store the electrode in a combined solution of approximately 0.1 M KCl in pH 4 buffer. One day or more prior to analysis, soak both electrodes in pH 4 buffer and, during analysis, place the electrodes in the same buffer when not in use. Reference Electrode Filling Solution Read the instructions that came with your electrodes carefully. Saturated calomel reference electrodes such as those used by the Acid Rain Monitoring Project must not be filled with filling solutions containing silver chloride (AgCl). We use 4M KCl solutions only.

However, the most common filling solution for combination electrodes is 4 M KCl saturated with AgCl. Be sure to ascertain which filling solution is correct for your electrode(s) and double check that your filling solution matches these requirements. Permanently filled or Gel electrodes Due to their unique micropore junction, it is recommended that they be stored hanging dry. Preliminary Electrode Response Testing If your electrode exhibits slow response, poor span between two buffer values or undue sensitivity to movement of the electrode, rejuvenation may be necessary to improve performance. Response varies with the electrode and the solution it is in. Generally working electrodes reach 0.05 pH units of the final reading in buffer within 10 seconds. A stable reading (less than 0.01 pH units per minute change) should be reached in fresh water samples within a minute or two. If you have to wait too long (5 minutes or more) then the pH itself may change due to the contact of the water sample with air. Electrodes may also require adjusting the slope to values significantly different from 100% for two point calibration.

Perform the following test if in doubt: Set your meter to 100% slope and room temperature, then standardize as usual with pH 7 buffer. Without moving the slope dial, read a pH 4 buffer. It should read between 3.85 and 4.15; set the slope to read pH 4, the slope should be 95% to 105%. If your electrode exhibits either of the above problems or is sensitive to movement, rejuvenation is in order. Glass Electrode Rejuvenation To treat the bulb of the pH electrode: Use 1 bottle each of acid and base (0.1N). BE CAREFUL WHEN HANDLING THESE SOLUTIONS. IF YOU GET ANY ON YOU RINSE OFF WITH LOTS OF WATER. Dip the electrode bulb into the acid and wait until meter reads pH ~1 Rinse electrode, then dip into the base.

Wait until meter reads pH ~ 13 Repeat this rinsing and dipping cycle several times (at least 3 times, 6 times are better) For the last cycle, you can leave electrode in base 5 min, then in acid briefly until it reaches pH ~ 1 Then rinse the electrode under tap water and let sit in pH 4 buffer for 2 hours Rinse electrodes and re-calibrate meter as you normally do with pH 7 and pH 4 buffers. To treat the reference electrode: Replace the 4M KCl solution in the reference electrode and get rid of crystals that may have formed. If there are lots of crystals, then shake out the solution and put deionized pure water into the filling hole and soak the electrode tip in hot tap water for 15 minutes or so until the crystals have dissolved. Then shake all the liquid out of the filling hole in the reference electrode and refill with fresh 4 M KCl. Let the electrode sit at room temperature for ½ hour before use. Frequently add more 4M KCl solution to the reference electrode since it will continually leak out and evaporate. The solution in the electrode should be within ½ inch of the filling hole.

The hole should be open when reading pH but close it when you are through for the day or else the solution will evaporate and new crystals will form (but do not close the hole if you will be storing the electrode soaking in pH 4 solution). If you still have problems with slow response, try rubbing the tip on your blue jeans or on very fine (600 grit) sandpaper. For combination electrodes, do both treatments described above. Final Test For Linearity Standardize the meter as described below. Rinse the electrodes and your sample cup with pure deionized water. Then titrate 100.0 ml of deionized water with your 0.16N acid as follows: Make sure your digital titrator is working and reset to zero. Add 10 digits of acid, record digits and pH, increase acid to 20 digits, record pH; repeat until you have added 100 digits of acid and stop. Send the results to us and we will send you a report. If you want to see the results yourself, try plotting the hydrogen ion concentration (H = 10(-pH)) vs. digits and see if the line is straight.

Movement Sensitivity If your meter gives wild readings and is sensitive to your touch, it may not be properly grounded. Try using a three prong power plug or attach a wire from the meter to a cold water pipe. Sometimes a problem of fluctuating readings or consistency wrong readings can be solved by disconnecting and reconnecting the electrode connectors several times. Apparently an oxide layer can sometimes cause these symptoms. Calibration The pH meter should be standardized (calibrated) prior to sample analyses and after every 25 sample analyses. Buffers should be at room temperature (68°F). Remove the electrodes from the pH 4 buffer solution where they have been soaking for at least one day. Rinse with deionized water. Insert the electrodes in pH 7.00 buffer and adjust the calibration dial until exactly pH 7.00 shows on the meter. Remove the electrodes and rinse with deionized water. Place the electrodes in pH 4.01 buffer and adjust the slope until the meter shows pH 4.01. Rinse with deionized water. A note on buffers. The accuracy of your pH measurement is in direct relation to the accuracy of the standard buffer solution used to calibrate your pH meter.

In order to maintain a reasonable degree of accuracy when making a pH measurement, a number of precautions concerning the care and use of buffers should be observed. These include: 1. Do not use buffers after their expiration date. Mold growth, CO2 absorption and contamination cause changes in the buffer pH. 2. Do not use buffers which have mold growth floating in the buffer. 3. Always cap the buffer container when storing to prevent contamination and reduce CO2 pickup. 4. pH buffer values change with temperature. Be sure to measure the temperature of the buffer and look up its value at that temperature before standardizing the meter (see below). 5. Do not pour used buffer back into the bottle. Buffer Values at Various Temperatures Temperature 6.998 pH & Alkalinity QA/QC Quality control for pH and alkalinity consists of normal pH measurement and titration of a sample prepared by the WRRC and sent to you prior to field collection. There will be three of these samples. Several days prior to sampling, you will receive the first QA/QC sample from us, along with a postcard for reporting your results. This is a diagnostic sample. Follow the procedures described for pH and alkalinity measurement.

Analyze two separate aliquots of this sample and report your results to us on the postcard. You will be called if we find a significant discrepancy between what we expect and what you measured. We will work with you to troubleshoot the problem so that you are confident of quality analysis for the field samples. Two other QA/QC samples will arrive just before field sampling. Unlike the first QA/QC sample, these are used to document data quality by helping us to statistically define the accuracy and precision of your analyses. Analyze two separate aliquots of one of these immediately prior to measuring pH and alkalinity on field samples; analyze two separate aliquots of the second QA/QC sample immediately after analyzing the field samples. In other words, the first two samples analyzed should be from one of the QA/QC bottles, then analyze the field samples, and finally analyze two samples from the other QA/QC bottle. Results should be reported on the pH & alkalinity lab data sheet. The Straight Dope: In chemistry, what do the letters "PA" stand for? A Straight Dope Classic from Cecil's Storehouse of Human Knowledge In chemistry, what do the letters "PA" stand for?

June 15, 1990 Dear Cecil: What do the letters PA stand for? I know some, we'll say, amateur chemists who concoct all sorts of things and consistently have to check the PA to get it right. But none seems to know what PA stands for or why it's such an important factor. The makers of the "Strong Dope" have let me down, so I turn to you, the writer of the Straight Dope. PS: Could you mail me a reply? I'd send you a stamp but I'm in jail awaiting transfer to the Texas Department of Corrections. If you can't, put it in the next book and I'll read it depending on which of us (the book or me) is out first. — Duncan Staggs, San Antonio, Texas Cecil replies: Too bad about prison, Dunc, but I told you to get that census form in on time. I imagine what your friends are referring to is not PA but pH, a measure of the acidity or alkalinity of a solution. The pH scale ranges from 0 to 14, 0 being highly acidic, 14 highly alkaline, and 7 neutral.

Maintaining a proper level of acidity is important in many chemical processes, and while my experience in home-brew drug manufacture is limited, I imagine it's important there, too. The pH scale was invented in 1909 by one S.P.L. Sorenson and stands for the exponential power (originally German Potenz) of hydrogen ion concentration. Hydrogen ions are the positively charged particles that make acids acidic. They like to mix it up with other molecules, so the more of them you have, the more potent the acid. For the record — I know this is going to sound like high school chemistry, but I feel obliged to be thorough — a pH of 1 means you've got 10-1 moles of hydrogen ions per liter, or one tenth of a mole; a pH of 14 means you've got 10-14 of a mole, or one hundred-trillionth. What's a mole, you inquire? Science talk for "whole bunch of," of course. Don't ask silly questions. Now we know why they call it the Texas Department of Corrections Dear Cecil: In reference to the question you answered from Mr. Duncan Staggs concerning what PA means, you are wrong.

Mr. Staggs most likely is asking about PA as it relates to the manufacture of methamphetamine. PA stands for lead acetate. Trust me, I know. — Speedy from Irving, Texas Cecil replies: Far be it from me to tell large gentlemen with weapons they're mistaken about their terminology, but Dunc said his buds have to "consistently check the PA to get it right," which sounds like the litmus-paper pH test. Also, the chemical notation for lead acetate is Pb(C2H3O2)2.3 H2O, not PA. — Cecil Adams Ph | Define Ph at Dictionary.com ph Chemistry. 1. the symbol for the logarithm of the reciprocal of hydrogen ion concentration in gram atoms per liter, used to express the acidity or alkalinity of a solution on a scale of 0 to 14, where less than 7 represents acidity, 7 neutrality, and more than 7 alkalinity. Ph Examples from the Web for ph Expand Contemporary Examples One ph direction is no better or more natural than the other.

Alkalinizing Someone to Improve His or Her Health Is Simple-Minded, Fatuous, and Dangerous Kent Sepkowitz January 28, 2014 As we were taught in high school Chem, any ph below 7.4 is acidic. Alkalinizing Someone to Improve His or Her Health Is Simple-Minded, Fatuous, and Dangerous Kent Sepkowitz January 28, 2014 Things on the other side of the ledger, at a ph above 7, are alkaline (also called “basic” as in acid-base). Alkalinizing Someone to Improve His or Her Health Is Simple-Minded, Fatuous, and Dangerous Kent Sepkowitz January 28, 2014 Mammalian cellular metabolism, aka the basic rules of being alive on planet Earth, runs best at a ph between 7.38 and 7.42. British Dictionary definitions for ph Expand noun 1. potential of hydrogen; a measure of the acidity or alkalinity of a solution equal to the common logarithm of the reciprocal of the concentration of hydrogen ions in moles per cubic decimetre of solution.

Pure water has a pH of 7, acid solutions have a pH less than 7, and alkaline solutions a pH greater than 7 Ph Collins English Dictionary - Complete & Unabridged 2012 Digital Edition © William Collins Sons & Co. Ltd. 1979, 1986 © HarperCollins Publishers 1998, 2000, 2003, 2005, 2006, 2007, 2009, 2012 Word Origin and History for ph Expand pH 1909, from German PH, introduced by S.P.L. Sörensen, from P, for German Potenz "potency, power" + H, symbol for the hydrogen ion that determines acidity or alkalinity. now in English usually representing "f," originally it was the combination used by Romans to represent Greek letter phi (cognate with Sanskrit -bh-, Germanic -b-), which at first was an aspirated "p," later the same sound as German -pf-. But by 2c. B.C.E.

had become a simple sound made by blowing through the lips (bilabial spirant). Roman "f," like modern English "f," was dentilabial; by c.400, however, the sounds had become identical and in some Romanic languages (Italian, Spanish), -ph- regularly was replaced by -f-. This tendency took hold in Old French and Middle English, but with the revival of classical learning the words subsequently were altered back to -ph- (except fancy and fantastic), and due to zealousness in this some non-Greek words in -f- began to appear confusedly in -ph-, though these forms generally have not survived. Online Etymology Dictionary, © 2010 Douglas Harper pH (pē'āch') A numerical measure of the acidity or alkalinity of a solution, usually measured on a scale of 0 to 14. Neutral solutions (such as pure water) have a pH of 7, acidic solutions have a pH lower than 7, and alkaline solutions have a pH higher than 7. The pH of lemon juice is 2.4; that of household ammonia is 11.5. The normal pH for human blood is 7.4.

◇ The letters pH stand for potential of hydrogen, since pH is effectively a measure of the concentration of hydrogen ions (that is, protons) in a substance. The pH scale was devised in 1923 by Danish biochemist Søren Peter Lauritz Sørensen (1868-1969). The American Heritage® Science Dictionary Copyright © 2002. Published by Houghton Mifflin. All rights reserved. What does pH stand for? - Simply Hydroponics Simply Hydroponics Learn How - Hydroponics and Greenhouse Gardening Main navigation Home / Insider Gardening Secrets / What does pH stand for? What does pH stand for? August 6, 2007 By Charlotte 32 Comments by: Charlene Rennick Have you ever wondered why the “p” in “pH” is a lower-case letter while the “H” is capitalized? What does it mean? The “p” stands for potential and the “H” stands for Hydrogen. Okay, so that makes it as clear as mud. What is potential Hydrogen? A scientific explanation would state that pH refers to the plant’s ability to attract hydrogen ions.

A less scientific explanation says pH is the acid/alkaline balance. Translated into a language those of us without the Ph D can understand, pH level refers to the amount of acid and alkaline contained inside of both the water and the growing medium or soil. If the environment is too acidic, that means the plant will not attract enough hydrogen, while an environment that is too alkaline will attract too much hydrogen. An environment that continually fluctuates from one extreme on the pH scale to the other is unhealthy for the plant. Hydrogen is one of four elements any living plant needs to survive. Without hydrogen, the plant would wilt and not be able to take in nutrients. Because the plant contains hydrogen, it continues to absorb hydrogen through the water via a process called osmosis . This hydrogen-osmosis cycle is what keeps the nutrients traveling from the water into the plant. Once a plant has died due to lack of water, there is no amount of water that can be added to it that will cause the plant to be revived. The level of pH is measured on a scale of 0 to 14 with 0 representing the highest concentration of acid and 14 representative of the most alkaline.

Seven is the magic figure for pH because it means that there is a balance of acid and alkaline in the solution and is often referred to as pH neutral. It is usually sufficient to say that a pH neutral environment is perfect for most plants, but some vegetation requires water or a growing medium that is more acidic than alkaline in order to flourish or have the right colour of blossoms, while other plants prefer the opposite. Testing strips for pH can be purchased to determine exactly what the acid/alkaline balance is in your water, growing medium or soil. This makes it easy to adjust the level for home gardeners or for mixing nutrients for your own hydroponic garden .

Question: In chemistry, the measure of acidity and alkalinity is the pH. For what does the letter H stand?
Answer: (POTENTIAL) HYDROGEN