Although alkalinity is primarily a term used by limnologists and oceanographers, 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. There can be long-term changes in the alkalinity of streams and rivers in response to human disturbances such as acid rain generated by SOx and NOx emissions.
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.
Alkalinity or measures the ability of a solution to neutralize acids to the equivalence point of carbonate or bicarbonate, defined as pH 4.5 for oceanographic/limnological studies. The alkalinity is equal to the stoichiometric sum of the bases in solution. In most Earth surface waters 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 (e.g., acetate), and sulfate. Solutions produced in a laboratory may contain a virtually limitless number of bases that contribute to alkalinity. Alkalinity is usually given as molar equivalents per liter or kilogram of solution. 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, thus reducing basicity while alkalinity remains unchanged (see example below).
A variety of titrants, endpoints, and indicators are specified for various alkalinity measurement methods. Hydrochloric and sulfuric acids are common acid titrants, while phenolpthalein, methyl red, and bromocresol green are common indicators.
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 divided by 2).
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
Physical processes can also affect alkalinity. Freshwater inputs like melting polar ice caps dilute seawater and 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 decrease AT. However, the actual effect would be much more complicated than this.[vague]
Biological processes have a much greater impact on oceanic alkalinity on short (decades to centuries) timescales. Aerobic respiration of organic matter can decrease alkalinity by releasing protons during the oxidation of organic nitrogen. Denitrification and sulfate reduction occur in oxygen-limited environments. Both of these processes consume hydrogen ions (thus increasing alkalinity) and release gases (N2 or H2S), which eventually escape into the atmosphere. Nitrification and sulfide oxidation both decrease alkalinity by releasing protons as a byproduct of oxidation reactions.
Throughout recent history, there have been many attempts to measure, record, and study oceanic alkalinity, together with many of the other characteristics of seawater, like temperature and salinity. These include: 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).
The ocean's alkalinity varies over time, most significantly over geologic timescales (millennia). Changes in the balance between terrestrial weathering and sedimentation of carbonate minerals (for example, as a function of ocean acidification) are the primary long-term drivers of alkalinity in the ocean. Over human timescales, mean ocean alkalinity is relatively stable. Seasonal and annual variability of mean ocean alkalinity is very low.
Alkalinity varies by location depending on evaporation/precipitation, advection of water, biological processes, and geochemical processes. 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.
River dominated mixing also occurs close to the shore; it is strongest close to the mouth of a large river. 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.
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.
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.
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
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.
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.
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
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.