Alkalinity

Alkalinity (from Arabic "al-qalī"[1]) is the capacity of water to resist changes in pH that would make the water more acidic.[2] (It should not be confused with basicity which is an absolute measurement on the pH scale.) Alkalinity is the strength of a buffer solution composed of weak acids and their conjugate bases. It is measured by titrating the solution with a monoprotic acid such as HCl until its pH changes abruptly, or it reaches a known endpoint where that happens. Alkalinity is expressed in units of meq/L (milliequivalents per liter), which corresponds to the amount of monoprotic acid added as a titrant in millimoles per liter.

Although alkalinity is primarily a term invented by oceanographers, [3] it is also used by hydrologists to describe temporary hardness. Moreover, measuring alkalinity is important in determining a stream's ability to neutralize acidic pollution from rainfall or wastewater. It is one of the best measures of the sensitivity of the stream to acid inputs.[4] There can be long-term changes in the alkalinity of streams and rivers in response to human disturbances.[5]

WOA05 GLODAP pd ALK AYool
Sea surface alkalinity (from the GLODAP climatology).

History

In 1884, Professor Wilhelm (William) Dittmar of Anderson College, now the University of Strathclyde, analysed 77 pristine seawater samples from around the world brought back by the Challenger expedition. He found that in seawater the major ions were in a fixed ratio, confirming the hypothesis of Johan Georg Forchhammer, that is now known as the Principle of Constant Proportions. However, there was one exception. Dittmar found that the concentration of calcium was slightly greater in the deep ocean, and named this increase alkalinity.

1884 was also the year when Svante Arrhenius submitted his PhD theses in which he advocated the existence of ions in solution, and defined acids as hydronium ion donors and bases as hydroxide ions donors. For that work, he received the Nobel Prize in Chemistry in 1903.

Thus Dittmar's alkalinity is the hydronium cations which exist to balance electrically the increase in calcium anions in deep ocean water, although now the meaning alkalinity has expanded.

Simplified summary

Alkalinity roughly refers to the amount of bases in a solution that can be converted to uncharged species by a strong acid.[6] The cited author, James Drever, provides an equation expressed in terms of molar equivalents, which means the number of moles of each ion type multiplied by (the absolute value of) the charge of the ion. For example, 1 mole of HCO31− in solution represents 1 molar equivalent, while 1 mole of CO32− is 2 molar equivalents because twice as many H+ ions would be necessary to balance the charge. The total charge of a solution always equals zero.

Quoting from page 52, "Ions such as Na+, K+, Ca2+, Mg2+, Cl, SO42−, and NO3 can be regarded as "conservative" in the sense that their concentrations are unaffected by changes in the pH, pressure, or temperature (within the ranges normally encountered near the earth's surface and assuming no precipitation or dissolution of solid phases, or biological transformations)."

On the left-hand side of the equation is the sum of conservative cations minus the sum of conservative anions. Balancing this on the right side is the sum of the anions that could be neutralized by added H+ ions (non-conservative anions) minus H+ ions already present, as indicated by the pH. All numbers are molar equivalents.

This right side term is called total alkalinity. It is, quoting Drever, "formally defined as the equivalent sum of the bases that are titratable with strong acid (Stumm and Morgan, 1981)".[7] The listing of ions shown on the right in Drever was "mHCO3 + 2mCO32− + mB(OH)4 + mH3(SiO)4 + mHS + morganic anions + mOH - mH+". Total alkalinity is measured by adding a strong acid until all the anions listed above are converted to uncharged species. The total alkalinity is not (much) affected by temperature, pressure, or pH, though the values of individual constituents are, mostly being conversions between HCO3 and CO32−.

Drever further notes that in most natural waters, all anions except HCO3 and CO32− have low concentrations. Thus carbonate alkalinity, which is equal to mHCO3 + 2mCO32− is also approximately equal to the total alkalinity.

Detailed description

Alkalinity or AT measures the ability of a solution to neutralize acids to the equivalence point of carbonate or bicarbonate. The alkalinity is equal to the stoichiometric sum of the bases in solution. In the natural environment carbonate alkalinity tends to make up most of the total alkalinity due to the common occurrence and dissolution of carbonate rocks and presence of carbon dioxide in the atmosphere. Other common natural components that can contribute to alkalinity include borate, hydroxide, phosphate, silicate, dissolved ammonia, the conjugate bases of some organic acids, and sulfate. Solutions produced in a laboratory may contain a virtually limitless number of bases that contribute to alkalinity. Alkalinity is usually given in the unit mEq/L (milliequivalent per liter). Commercially, as in the swimming pool industry, alkalinity might also be given in parts per million of equivalent calcium carbonate (ppm CaCO3).

Alkalinity is sometimes incorrectly used interchangeably with basicity. For example, the addition of CO2 lowers the pH of a solution. This increase reduces the basicity; however, the alkalinity remains unchanged (see example below).For total alkalinity testing, N/10 H2SO4 is used by hydrologists along with phenolphthalein indicator.

Theoretical treatment

In typical groundwater or seawater, the measured alkalinity is set equal to:

AT = [HCO3]T + 2[CO32−]T + [B(OH)4]T + [OH]T + 2[PO43−]T + [HPO42−]T + [SiO(OH)3]T − [H+]sws − [HSO4]

(Subscript T indicates the total concentration of the species in the solution as measured. This is opposed to the free concentration, which takes into account the significant amount of ion pair interactions that occur in seawater.)

Alkalinity can be measured by titrating a sample with a strong acid until all the buffering capacity of the aforementioned ions above the pH of bicarbonate or carbonate is consumed. This point is functionally set to pH 4.5. At this point, all the bases of interest have been protonated to the zero level species, hence they no longer cause alkalinity. In the carbonate system the bicarbonate ions [HCO3] and the carbonate ions [CO32−] have become converted to carbonic acid [H2CO3] at this pH. This pH is also called the CO2 equivalence point where the major component in water is dissolved CO2 which is converted to H2CO3 in an aqueous solution. There are no strong acids or bases at this point. Therefore, the alkalinity is modeled and quantified with respect to the CO2 equivalence point. Because the alkalinity is measured with respect to the CO2 equivalence point, the dissolution of CO2, although it adds acid and dissolved inorganic carbon, does not change the alkalinity. In natural conditions, the dissolution of basic rocks and addition of ammonia [NH3] or organic amines leads to the addition of base to natural waters at the CO2 equivalence point. The dissolved base in water increases the pH and titrates an equivalent amount of CO2 to bicarbonate ion and carbonate ion. At equilibrium, the water contains a certain amount of alkalinity contributed by the concentration of weak acid anions. Conversely, the addition of acid converts weak acid anions to CO2 and continuous addition of strong acids can cause the alkalinity to become less than zero.[8] For example, the following reactions take place during the addition of acid to a typical seawater solution:

B(OH)4 + H+ → B(OH)3 + H2O
OH + H+ → H2O
PO4−3 + 2H+ → H2PO4
HPO4−2 + H+ → H2PO4
[SiO(OH)3] + H+ → [Si(OH)40]

It can be seen from the above protonation reactions that most bases consume one proton (H+) to become a neutral species, thus increasing alkalinity by one per equivalent. CO3−2 however, will consume two protons before becoming a zero level species (CO2), thus it increases alkalinity by two per mole of CO3−2. [H+] and [HSO4] decrease alkalinity, as they act as sources of protons. They are often represented collectively as [H+]T.

Alkalinity is typically reported as mg/L as CaCO3. (The conjunction "as" is appropriate in this case because the alkalinity results from a mixture of ions but is reported "as if" all of this is due to CaCO3.) This can be converted into milliEquivalents per Liter (mEq/L) by dividing by 50 (the approximate MW of CaCO3/2).

Example problems

Sum of contributing species

The following equations demonstrate the relative contributions of each component to the alkalinity of a typical seawater sample. Contributions are in μmol.kg−soln−1 and are obtained from A Handbook of Methods for the analysis of carbon dioxide parameters in seawater "[1],"(Salinity = 35 g/kg, pH = 8.1, Temperature = 25 °C).

AT = [HCO3]T + 2[CO32−]T + [B(OH)4]T + [OH]T + 2[PO43−]T + [HPO42−]T + [SiO(OH)3]T − [H+] − [HSO4] − [HF]

Phosphates and silicate, being nutrients, are typically negligible. At pH = 8.1 [HSO4] and [HF] are also negligible. So,

AT = [HCO3]T + 2[CO32−]T + [B(OH)4]T + [OH]T − [H+]
= 1830 + 2 × 270 + 100 + 10 − 0.01
= 2480 μmol.kg−soln−1

Addition of CO2

Addition (or removal) of CO2 to a solution does not change its alkalinity, since the net reaction produces the same number of equivalents of positively contributing species (H+) as negative contributing species (HCO3 and/or CO32−). Adding CO2 to the solution lowers its pH, but does not affect alkalinity.

At all pH values:

CO2 + H2O ⇌ HCO3 + H+

Only at high (basic) pH values:

HCO3 + H+ ⇌ CO32− + 2H+

Dissolution of carbonate rock

Addition of CO2 to a solution in contact with a solid can (over time) affect the alkalinity, especially for carbonate minerals in contact with groundwater or seawater . The dissolution (or precipitation) of carbonate rock has a strong influence on the alkalinity. This is because carbonate rock is composed of CaCO3 and its dissociation will add Ca+2 and CO3−2 into solution. Ca+2 will not influence alkalinity, but CO3−2 will increase alkalinity by 2 units. Increased dissolution of carbonate rock by acidification from acid rain and mining has contributed to increased alkalinity concentrations in some major rivers throughout the Eastern U.S.[5] The following reaction shows how acid rain, containing sulfuric acid, can have the effect of increasing river alkalinity by increasing the amount of bicarbonate ion:

2CaCO3 + H2SO4 → 2Ca+2 + 2HCO3 + SO4−2

Another way of writing this is:

CaCO3 + H+ ⇌ Ca+2 + HCO3

The lower the pH, the higher the concentration of bicarbonate will be. This shows how a lower pH can lead to higher alkalinity if the amount of bicarbonate produced is greater than the amount of H+ remaining after the reaction. This is the case since the amount of acid in the rainwater is low. If this alkaline groundwater later comes into contact with the atmosphere, it can lose CO2, precipitate carbonate, and thereby become less alkaline again. When carbonate minerals, water, and the atmosphere are all in equilibrium, the reversible reaction

CaCO3 + 2H+ ⇌ Ca+2 + CO2 + H2O

shows that pH will be related to calcium ion concentration, with lower pH going with higher calcium ion concentration. In this case, the higher the pH, the more bicarbonate and carbonate ion there will be, in contrast to the paradoxical situation described above, where one does not have equilibrium with the atmosphere.

Oceanic alkalinity

Processes that increase alkalinity

There are many methods of alkalinity generation in the ocean. Perhaps the most well known is the dissolution of CaCO3 (calcium carbonate, which is a component of coral reefs) to form Ca2+ and CO32− (carbonate). The carbonate ion has the potential to absorb two hydrogen ions. Therefore, it causes a net increase in ocean alkalinity. Calcium carbonate dissolution is an indirect result of ocean pH lowering. It can cause great damage to coral reef ecosystems, but has a relatively low effect on the total alkalinity (AT) in the ocean.[9] Lowering of pH due to absorption of CO2 actually raises the alkalinity by causing dissolution of carbonates.

Anaerobic degradation processes, such as denitrification and sulfate reduction, have a much greater impact on oceanic alkalinity. Denitrification and sulfate reduction occur in the deep ocean, where there is an absence of oxygen. Both of these processes consume hydrogen ions and releases quasi-inert gases (N2 or H2S), which eventually escape into the atmosphere. This consumption of H+ increases the alkalinity. It has been estimated that anaerobic degradation could be as much as 60% of the total oceanic alkalinity.[9]

Processes that decrease alkalinity

Anaerobic processes generally increase alkalinity. Conversely, aerobic degradation can decrease AT. This process occurs in portions of the ocean where oxygen is present (surface waters). It results in dissolved organic matter and the production of hydrogen ions.[9] An increase in H+ clearly decreases alkalinity. However, the dissolved organic matter may have base functional groups that can consume these hydrogen ions and negate their effect on alkalinity. Therefore, aerobic degradation has a relatively low impact on the overall oceanic alkalinity.[10]

All of these aforementioned methods are chemical processes. However, physical processes can also serve to affect AT. The melting of polar ice caps is a growing concern that can serve to decrease oceanic alkalinity. If the ice were to melt, then the overall volume of the ocean would increase. Because alkalinity is a concentration value (mol/L), increasing the volume would theoretically serve to decrease AT. However, the actual effect would be much more complicated than this.[11]

Global temporal variability

Researchers have shown oceanic alkalinity to vary over time. Because AT is calculated from the ions in the ocean, a change in the chemical composition would alter alkalinity. One way this can occur is through ocean acidification. However, oceanic alkalinity is relatively stable, so significant changes can only occur over long time scales (i.e. hundreds to thousands of years).[12] As a result, seasonal and annual variability is generally very low.[9]

Spatial variability

Researchers have also shown alkalinity to vary depending on location. Local AT can be affected by two main mixing patterns: current and river. Current dominated mixing occurs close to the shore in areas with strong water flow. In these areas, alkalinity trends follow current and have a segmented relationship with salinity.[13]

River dominated mixing also occurs close to the shore; it is strongest close to the mouth of a large river (i.e. the Mississippi or Amazon). Here, the rivers can act as either a source or a sink of alkalinity. AT follows the outflow of the river and has a linear relationship with salinity. This mixing pattern is most important in late winter and spring, because snowmelt increases the river’s outflow. As the season progresses into summer, river processes are less significant, and current mixing can become the dominant process.[9]

Oceanic alkalinity also follows general trends based on latitude and depth. It has been shown that AT is often inversely proportional to sea surface temperature (SST). Therefore, it generally increases with high latitudes and depths. As a result, upwelling areas (where water from the deep ocean is pushed to the surface) also have higher alkalinity values.[14]

Measurement data sets

Throughout recent history, there have been many attempts to measure, record, and study oceanic alkalinity. Some of the larger data sets are listed below.

  • GEOSECS (Geochemical Ocean Sections Study)
  • TTO/NAS (Transient Tracers in the Ocean/North Atlantic Study)
  • JGOFS (Joint Global Ocean Flux Study)
  • WOCE (World Ocean Circulation Experiment)
  • CARINA (Carbon dioxide in the Atlantic Ocean)

See also

References

  1. ^ "the definition of alkali". www.dictionary.com. Retrieved 2018-09-30.
  2. ^ "What is Alkalinity?". Water Research Center. 2014. Retrieved 5 February 2018.
  3. ^ Dickson, Andrew G. (1992). "The development of the alkalinity concept in marine chemistry". Marine Chemistry, 40, 1: 49–63. doi:10.1016/0304-4203(92)90047-E.
  4. ^ "Total Alkalinity". United States Environment Protection Agency. Retrieved 6 March 2013.
  5. ^ a b Kaushal, S. S.; Likens, G. E.; Utz, R. M.; Pace, M. L.; Grese, M.; Yepsen, M. (2013). "Increased river alkalinization in the Eastern U.S". Environmental Science & Technology: 130724203606002. doi:10.1021/es401046s.
  6. ^ Drever, James I. (1988). The Geochemistry of Natural Waters, Second Edition. Englewood Cliffs, NJ: Prentice Hall. pp. 51–58 [52]. ISBN 0-13-351396-3.
  7. ^ Stumm, W. & J.J Morgan (1981). Aquatic Chemistry, 2n Ed. New York: Wiley-Interscience. p. 780.
  8. ^ Benjamin. Mark M. 2015. Water Chemistry. 2nd Ed. Long Grove, Illinois: Waveland Press, Inc.
  9. ^ a b c d e Thomas, H.; Schiettecatte, L.-S.; et al. Enhanced Ocean Carbon Storage from Anaerobic Alkalinity Generation in Coastal Sediments. Biogeosciences Discussions. 2008, 5, 3575-3591
  10. ^ Kim, H.-C., and K. Lee (2009), Significant contribution of dissolved organic matter to seawater alkalinity, Geophys. Res. Lett., 36, L20603, doi:10.1029/2009GL040271
  11. ^ Chen, B.; Cai, W. Using Alkalinity to Separate the Inputs of Ice-Melting and River in the Western Arctic Ocean. Proceedings from the 2010 AGU Ocean Sciences Meeting, 2010, 22-26.
  12. ^ Doney, S. C.; Fabry, V. J.; et al. Ocean Acidification: The Other CO2 Problem. Annu. Rev. Mar. Sci., 2009, 69-92. doi:10.1146/annurev.marine.010908.163834
  13. ^ Cai, W.-J.; Hu, X. et al. Alkalinity Distribution in the Western North Atlantic Ocean Margins. Journal of Geophysical Research. 2010, 115, 1-15. doi:10.1029/2009JC005482
  14. ^ Millero, F. J.; Lee, K.; Roche, M. Distribution of alkalinity in the surface waters of the major oceans. Marine Chemistry. 1998, 60, 111-130.

External links

  • Holmes-Farley, Randy. "Chemistry and the Aquarium: What is Alkalinity?," Advanced Aquarist's Online Magazine. Alkalinity as it pertains to salt-water aquariums.
  • DOE (1994) "[2],"Handbook of methods for the analysis of the various parameters of the carbon dioxide system in sea water. Version 2, A. G. Dickson & C. Goyet, eds. ORNL/CDIAC-74.
  • GEOSECS data set [3]
  • JGOFS data set [4]
  • WOCE data set [5]
  • CARINA data set [6]

Carbonate system calculators

The following packages calculate the state of the carbonate system in seawater (including pH):

Alkali soil

Alkali, or Alkaline, soils are clay soils with high pH (> 8.5), a poor soil structure and a low infiltration capacity. Often they have a hard calcareous layer at 0.5 to 1 metre depth. Alkali soils owe their unfavorable physico-chemical properties mainly to the dominating presence of sodium carbonate, which causes the soil to swell and difficult to clarify/settle. They derive their name from the alkali metal group of elements, to which sodium belongs, and which can induce basicity. Sometimes these soils are also referred to as alkaline sodic soils.

Alkaline soils are basic, but not all basic soils are alkaline.

Base (chemistry)

In chemistry, bases are substances that, in aqueous solution, release hydroxide (OH−) ions, are slippery to the touch, can taste bitter if an alkali, change the color of indicators (e.g., turn red litmus paper blue), react with acids to form salts, promote certain chemical reactions (base catalysis), accept protons from any proton donor or contain completely or partially displaceable OH− ions. Examples of bases are the hydroxides of the alkali metals and the alkaline earth metals (NaOH, Ca(OH)2, etc.—see alkali hydroxide and alkaline earth hydroxide).

In water, by altering the autoionization equilibrium, bases yield solutions in which the hydrogen ion activity is lower than it is in pure water, i.e., the water has a pH higher than 7.0 at standard conditions. A soluble base is called an alkali if it contains and releases OH− ions quantitatively. However, it is important to realize that basicity is not the same as alkalinity. Metal oxides, hydroxides, and especially alkoxides are basic, and conjugate bases of weak acids are weak bases.

Bases can be thought of as the chemical opposite of acids. However, some strong acids are able to act as bases. Bases and acids are seen as opposites because the effect of an acid is to increase the hydronium (H3O+) concentration in water, whereas bases reduce this concentration. A reaction between an acid and a base is called neutralization. In a neutralization reaction, an aqueous solution of a base reacts with an aqueous solution of an acid to produce a solution of water and salt in which the salt separates into its component ions. If the aqueous solution is saturated with a given salt solute, any additional such salt precipitates out of the solution.

For a substance to be classified as an Arrhenius base, it must produce hydroxide ions in an aqueous solution. Arrhenius believed that in order to do so, the base must contain hydroxide in the formula. This makes the Arrhenius model limited, as it cannot explain the basic properties of aqueous solutions of ammonia (NH3) or its organic derivatives (amines). There are also bases that do not contain a hydroxide ion but nevertheless react with water, resulting in an increase in the concentration of the hydroxide ion. An example of this is the reaction between ammonia and water to produce ammonium and hydroxide. In this reaction ammonia is the base because it accepts a proton from the water molecule. Ammonia and other bases similar to it usually have the ability to form a bond with a proton due to the unshared pair of electrons that they possess. In the more general Brønsted–Lowry acid–base theory, a base is a substance that can accept hydrogen cations (H+)—otherwise known as protons. In the Lewis model, a base is an electron pair donor.

Calcium reactor

In marine and reef aquariums, a calcium reactor creates a balance of alkalinity. An acidic solution is produced by injecting carbon dioxide into a chamber with salt water and calcium rich media. The carbon dioxide lowers the pH by producing a solution high in carbonic acid, and dissolves calcium. The effluent is returned to the reef aquarium where the calcium is consumed by organisms, primarily corals when building skeletons. A calcium reactor is an efficient method to supply calcium to a reef aquarium. Reactors may be used in elaborate freshwater and brackish aquariums where freshwater clams and other invertebrates need a constant supply of calcium.

The reactor dissolves the calcium-laden media to provide bicarbonates HCO3− (alkalinity) and calcium (Ca++) ions at the sames rate as consumed during calcification. Effectively dissolving the media requires an acidic pH. Saltwater may have a pH of 7.8 or higher, so to reduce the pH carbon dioxide (CO2) is used. The reaction formula is:

CaCO3 + H2O + CO2 ⟷ Ca2+ + 2 HCO3−Inside the reaction chamber, a calcium rich media (aragonite), mainly CaCO3, is forced into contact with water injected with carbon dioxide (CO2) in order to create carbonic acid (H2CO3). This increases the solubility of the calcium carbonate. The reaction frees the calcium and carbonate, supplying the aquarium with water rich in Ca2+ and CO32−, important for maintaining alkalinity and calcium levels.

The bubble counter measures carbon dioxide. The flow rate of carbon dioxide is monitored so that the dissolved gas goes into the solution, with a minimum unconsumed. A needle valve or solenoid valve regulates the CO2 bubble rate. Valves with precise adjustment abilities improve bubble control.

The feed pump controls the volume of water exchange. This is important because a high rate of water flow into the reactor reduces its efficiency, thus resulting in underproduction and a waste of CO2.

Some reactors siphon water into the input of the reactor's re-circulation pump. A potential complication is the medium in the reactor becoming compacted, increasing back pressure onto the pump and reducing water into the reactor. Placing a gate or needle valve on the reactor's outlet side will improve flow characteristics compared to control from the inlet side.

Peristaltic pumps are effective operating against pressure, capable of supplying an adjustable and continuous flow over flow rates with minimal maintenance.

The pH control is connected to a probe in the reactor and adjusts the rate at which the calcium media dissolves. This probe monitors the pH level in the calcium reactor. The pH range for the typical calcium reactor is 6.5–6.8. When the pH rises above a certain level, a valve opens, allowing carbon dioxide to enter the reactor. The control closes the valve as the pH falls below this level.

Some pH controllers have an interface for an air pump. This air pump is connected to an airstone in the sump or main tank. If the probe detects a low pH level, the pump activates. The bubbles raise the pH by dissipating the CO2 gas.

Carbonate hardness

Carbonate hardness, is a measure of the water hardness caused by the presence of carbonate (CO2−
3
) and bicarbonate (HCO
3
) anions. Carbonate hardness is usually expressed either in degrees KH (dKH) (from the German "Karbonathärte"), or in parts per million calcium carbonate ( ppm CaCO
3
or grams CaCO
3
per litre|mg/l). One dKH is equal to 17.848 mg/l (ppm) CaCO
3
, e.g. one dKH corresponds to the carbonate and bicarbonate ions found in a solution of approximately 17.848 milligrams of calcium carbonate(CaCO
3
) per litre of water (17.848 ppm). Both measurements (mg/l or KH) are usually expressed as mg/l CaCO
3
– meaning the concentration of carbonate expressed as if calcium carbonate were the sole source of carbonate ions.

An aqueous solution containing 120 mg NaHCO3 (baking soda) per litre of water will contain 1.4285 mmol/l of bicarbonate, since the molar mass of baking soda is 84.007 g/mol. This is equivalent in carbonate hardness to a solution containing 0.71423 mmol/L of (calcium) carbonate, or 71.485 mg/l of calcium carbonate (molar mass 100.09 g/mol). Since one degree KH = 17.848 mg/L CaCO3, this solution has a KH of 4.0052 degrees.

Carbonate hardness should not be confused with a similar measure Carbonate Alkalinity which is expressed in either [milli[equivalent]s] per litre (meq/L) or ppm. Carbonate hardness expressed in ppm does not necessarily equal carbonate alkalinity expressed in ppm.

whereas

However, for water with a pH below 8.5, the CO32− will be less than 1% of the HCO3− so carbonate alkalinity will equal carbonate hardness to within an error of less than 1%.

In a solution where only CO2 affects the pH, carbonate hardness can be used to calculate the concentration of dissolved CO2 in the solution with the formula CO2 = 3 × KH × 10(7-pH), where KH is degrees of carbonate hardness and CO2 is given in ppm by weight.

The term carbonate hardness is also sometimes used as a synonym for temporary hardness, in which case it refers to that portion of hard water that can be removed by processes such as boiling or lime softening, and then separation of water from the resulting precipitate.

Citrus reshni

Citrus reshni also known as Cleopatra mandarin is a citrus tree that is commonly used in agriculture as a rootstock of different cultivated species of citrus, mostly orange, grapefruit, tangerine and lemon. It originated in India and later was introduced to Florida from Jamaica in the mid-nineteenth century.The Cleopatra mandarin fruit belong to the "acidic" group of mandarins, which are too sour to be edible. When they are grown it is for the rootstock or for juice production. The rootstock can handle multiple soil conditions including tolerance to the presence of limestone, salinity and soil alkalinity along with being suitable for shallow soils. It is resistant to citrus tristeza virus and exocortis but is sensitive to root asphyxia and Phytophthora. One of the down sides to using the rootstock is it grows slow in the early years. In the right conditions it can induce high productivity and excellent fruit quality, although these are usually somewhat smaller than with others.

Cliff Pond

Cliff Pond is a 204-acre (830,000 m2) kettle pond in Brewster, Massachusetts. It is the largest pond in Nickerson State Park and is quite popular with swimmers and fishermen in summer months.

Cliff Pond was totally reclaimed in 1960 and, like many kettle ponds has been treated for alkalinity over the years. The Massachusetts Division of Fisheries and Wildlife stocks the pond in spring and fall with various trout species. It also has smallmouth bass and various other species. In 1992, a world record American eel (8 pounds, 9 ounces, 46 inches

long, 10.5 inches in girth) was caught.The pond has had repeated problems with bluegreen algae blooms causing closures over the years and was treated in spring 2016 with aluminum sulfate. Water transparency has improved and the treatment is expected to last twenty years.

Dealkalization of water

The dealkalization of water refers to the removal of alkalinity ions from water. Chloride cycle anion ion exchange dealkalizers remove alkalinity from water.

Chloride cycle dealkalizers operate similar to sodium cycle cation water softeners. Like water softeners, dealkalizers contain ion exchange resins that are regenerated with a concentrated salt (brine) solution - NaCl. In the case of a water softener, the cation exchange resin is exchanging sodium (the Na+ ion of NaCl) for hardness minerals such as calcium and magnesium.

A dealkalizer contains strong base anion exchange resin that exchanges chloride (the Cl– ion of the NaCl) for carbonate (CO−3), bicarbonate (HCO−3) and sulfate (SO2−4). As water passes through the anion resin the carbonate, bicarbonate and sulfate ions are exchanged for chloride ions.

"Higher capacities can be realized by use of type II rather than type I strong base anion resins. Although bicarbonates are not held as tightly as chlorides on the SBA (strong base anion) resins in the hydroxide form, when the resin is predominantly in the chloride form the pH has been raised by a small addition of caustic to the brine regenerant, there will be a favorable exchange of bicarbonate for the chloride. This exchange works well only with high alkalinity waters (40% to 80%), with capacities of 4 to 10 Kg/CF being obtained. The advantages of SBA resin dealkalization is that low-cost salt is used in place of the acid necessary for the SAC (strong acid cation)and un-lined steel tanks can be used."

Dingle Reservoir

Dingle Reservoir is an artificial, low alkalinity, shallow reservoir near to the town of Egerton, Greater Manchester. The reservoir itself is a little under 400 metres from the border between Lancashire and Greater Manchester, bring found on the Lancashire side.

Geochemical Ocean Sections Study

The Geochemical Ocean Sections Study (GEOSECS) was a global survey of the three-dimensional distributions of chemical, isotopic, and radiochemical tracers in the ocean. A key objective was to investigate the deep thermohaline circulation of the ocean, using chemical tracers, including radiotracers, to establish the pathways taken by this.Expeditions undertaken during GEOSECS took place in the Atlantic Ocean from July 1972 to May 1973, in the Pacific Ocean from August 1973 to June 1974, and in the Indian Ocean from December 1977 to March 1978.Measurements included those of physical oceanographic quantities such as temperature, salinity, pressure and density, chemical / biological quantities such as total inorganic carbon, alkalinity, nitrate, phosphate, silicic acid, oxygen and apparent oxygen utilisation (AOU), and radiochemical / isotopic quantities such as carbon-13, carbon-14 and tritium.

Kapurthala district

Kapurthala district is a district of Punjab state in northern India. The city of Kapurthala is the district headquarters.

Kapurthala District is one of the smallest districts of Punjab in terms of both area and population, with 754,521 people by the 2001 census. The district is divided into two noncontiguous parts, the main Kapurthala-Sultanpur Lodhi portion and the Phagwara tehsil or block.

The Kapurthala-Sultanpur Lodhi part lies between north latitude 31° 07' and 31° 22' and east longitude 75° 36'. In the north it is bound by Hoshiarpur, Gurdaspur, and Amritsar districts, in the west by the Beas River and Amritsar district, and in south by the Sutlej River, Jalandhar district, and Hoshiarpur district.

Phagwara tehsil lies between north latitude 31° 22' and east longitude 75° 40' and 75° 55'. Phagwara lies on the National Highway No 1, and the tehsil is much more industrially developed than the remainder of Kapurthala District. Phagwara is situated at a distance of 19 kilometers (12 mi) southwest of Jalandhar, and the tehsil is bounded on all sides by Jalandhar District except in the northeast, where it is bounded by Hoshiarpur district.

The district has three subdivisions/tehsils: Kapurthala, Phagwara, and Sultanpur Lodhi. The total area of the district is 1633 km² (630 mi²) of which 909.09 km² (350.91 mi²) is in Kapurthala tehsil, 304.05 km² (117.36 mi²) in Phagwara tehsil and 451.0 km² (174.1 mi²) in Sultanpur Lodhi tehsil. The economy of the district is still predominantly agricultural. The major crops are wheat, rice, sugarcane, potato and maize. The major portion of Kapurthala district lies between the Beas River and the Kali-Bein River and is called the ‘BET’ area. This area is prone to floods. Water logging and alkalinity in the soil is the major problem of the area. A flood protection bundh called ‘Dhussi Bundh’ has been constructed along the left bank of the Beas River, and it has saved the area from the ravages of flood. The entire district is an alluvial plain. To the south of the river Kali-Bein lies the tract known as ‘Dona’ meaning the soil formed of two constituents i.e. the sand and clay.

The climate is typical of the Punjab plains i.e. hot in summers and cold in winters. It has sub-tropical continental monsoon type climate. Intensive cultivation in the district leaves no scope for forest cover and wildlife is practically nonexistent.

Maud Menten

Maud Leonora Menten (March 20, 1879 – July 17, 1960) was a Canadian bio-medical and medical researcher who made significant contributions to enzyme kinetics and histochemistry. Her name is associated with the famous Michaelis–Menten equation in biochemistry.

Maud Menten was born in Port Lambton, Ontario and studied medicine at the University of Toronto (B.A. 1904, M.B. 1907, M.D. 1911, Ph.D., 1916). She was among the first women in Canada to earn a medical doctorate. She completed her thesis work at University of Chicago. At that time women were not allowed to do research in Canada, so she decided to do research in other countries such as the United States and Germany.

In 1912 she moved to Berlin where she worked with Leonor Michaelis and co-authored their paper in Biochemische Zeitschrift which showed that the rate of an enzyme-catalyzed reaction is proportional to the amount of the enzyme-substrate complex. This relationship between reaction rate and enzyme–substrate concentration is known as the Michaelis–Menten equation.

After studying with Michaelis in Germany she entered graduate school at the University of Chicago where she obtained her PhD in 1916. Her dissertation was titled "The Alkalinity of the Blood in Malignancy and Other Pathological Conditions; Together with Observations on the Relation of the Alkalinity of the Blood to Barometric Pressure". Menten worked at the University of Pittsburgh (1923–1950), becoming Assistant Professor and then Associate Professor in the School of Medicine and head of pathology at the Children's Hospital of Pittsburgh. Her final promotion to full Professor, in 1948, was at the age of 69 in the last year of her career. Her final academic post was as a research fellow at the British Columbia Medical Research Institute.

Organic base

An organic base is an organic compound which acts as a base. Organic bases are usually, but not always, proton acceptors. They usually contain nitrogen atoms, which can easily be protonated, for example amines have a lone pair of electrons on the nitrogen atom and can thus act as proton acceptors (bases).. Amines and nitrogen-containing heterocyclic compounds are organic bases. Examples include:

pyridine

alkanamines, such as methylamine

imidazole

benzimidazole

histidine

guanidine

phosphazene bases

hydroxides of quaternary ammonium cations or some other organic cations

PH indicator

A pH indicator is a halochromic chemical compound added in small amounts to a solution so the pH (acidity or basicity) of the solution can be determined visually. Hence, a pH indicator is a chemical detector for hydronium ions (H3O+) or hydrogen ions (H+) in the Arrhenius model. Normally, the indicator causes the color of the solution to change depending on the pH. Indicators can also show change in other physical properties; for example, olfactory indicators show change in their odor. The pH value of a neutral solution is 7.0 at 25°C (standard laboratory conditions). Solutions with a pH value below 7.0 are considered acidic and solutions with pH value above 7.0 are basic (alkaline). As most naturally occurring organic compounds are weak protolytes, carboxylic acids and amines, pH indicators find many applications in biology and analytical chemistry. Moreover, pH indicators form one of the three main types of indicator compounds used in chemical analysis. For the quantitative analysis of metal cations, the use of complexometric indicators is preferred, whereas the third compound class, the redox indicators, are used in titrations involving a redox reaction as the basis of the analysis.

Paleoceanography

Paleoceanography is the study of the history of the oceans in the geologic past with regard to circulation, chemistry, biology, geology and patterns of sedimentation and biological productivity. Paleoceanographic studies using environment models and different proxies enable the scientific community to assess the role of the oceanic processes in the global climate by the re-construction of past climate at various intervals. Paleoceanographic research is also intimately tied to paleoclimatology.

Soil pH

Soil pH is a measure of the acidity or basicity (alkalinity) of a soil. pH is defined as the negative logarithm (base 10) of the activity of hydronium ions (H+ or, more precisely, H3O+aq) in a solution. In soils, it is measured in a slurry of soil mixed with water (or a salt solution, such as 0.01 M CaCl2), and normally falls between 3 and 10, with 7 being neutral. Acid soils have a pH below 7 and alkaline soils have a pH above 7. Ultra-acidic soils (pH < 3.5) and very strongly alkaline soils (pH > 9) are rare.Soil pH is considered a master variable in soils as it affects many chemical processes. It specifically affects plant nutrient availability by controlling the chemical forms of the different nutrients and influencing the chemical reactions they undergo. The optimum pH range for most plants is between 5.5 and 7.5; however, many plants have adapted to thrive at pH values outside this range.

Stop bath

Stop bath is a chemical bath usually used in processing traditional black-and-white photographic films, plates, and paper used after the material has finished developing. The purpose of the stop bath is to halt the development of the film, plate, or paper by either washing off the developing chemical or neutralizing it. With the former, a simple water rinse can be used between developer and fixer, but the development process continues (though possibly at a very low level) for an indefinite and uncontrolled period of time during the rinsing.

Where an immediate stop of development is desired, a stop bath will usually consist of some concentration of acetic acid, commonly around 1 to 2%. Since organic developers only work in alkaline solutions, stop bath halts the development process almost instantly and thus provides more precise control of the development time. It also cuts overall processing time, because the required immersion time in the stop bath—typically fifteen to thirty seconds—is much shorter than the time required for an adequate plain-water rinse. As well, by neutralizing the alkalinity of basic developers, it can help to preserve the strength of the fixer, making it last longer.

Stop bath accounts for the characteristic vinegar-like odor of the traditional darkroom. In its concentrated form it can cause chemical burns, but is harmless when diluted to a working solution. Stop bath becomes exhausted when bases carried over from the developer cause the solution to become alkaline. For indicator stop bath—a stop bath that changes colours to indicate when the stop bath is exhausted and no longer effective—a pH indicator like bromocresol purple is used to determine when the solution has become too alkaline to use. Low-odor stop baths use citric acid or sodium bisulfite in place of acetic acid.

Universal indicator

A universal indicator is a pH indicator made of a solution of several compounds that exhibits several smooth colour changes over a wide range pH values to indicate the acidity or alkalinity of solutions. Although there are several commercially available universal pH indicators, most are a variation of a formula patented by Yamada in 1933. Details of this patent can be found in Chemical Abstracts. Experiments with Yamada's universal indicator are also described in the Journal of Chemical Education.A universal indicator is typically composed of water, propan-1-ol, phenolphthalein sodium salt, sodium hydroxide, methyl red, bromothymol blue monosodium salt, and thymol blue monosodium salt. The colours that indicate the pH of a solution, after adding a universal indicator, are:

The colours from yellow to red indicate an acidic solution, colours blue to violet indicate alkali and green colour indicates that a solution is neutral.

Wide-range pH test papers with distinct colours for each pH from 1 to 14 are also available. Colour matching charts are supplied with the specific test strips purchased.

Wallace Run (Bald Eagle Creek tributary)

Wallace Run is a tributary of Bald Eagle Creek in Centre County, Pennsylvania, in the United States. It is 12.1 miles (19.5 km) long and is a low-alkalinity stream. The stream flows through Union Township and Boggs Township in Centre County. Most of the watershed is in Boggs Township. North Branch Wallace Run is one tributary of the stream. The watershed has an area of 24 square miles. Oaks, maples, ash trees, birches, hemlocks, and rhododendrons all exist in the upper reaches of the stream, which is mostly forested. The lower reaches of the stream are mostly developed.

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