they are often confused as to their relationship and interpretation. defines the form of alkalinity as being bicarbonate, carbonate or hydrate alkalinity. In boiler. typically deal with are hydroxide (OH-), bicarbonate (HCO3. -) and carbonate. ( CO3. =) ions. Take a look at the alkalinity and pH diagram presented here. Alkalinity in natural water is due to: * Salts of weak acids. * Carbonate, bicarbonate. * Borate, silicate, phosphate. * A few organic acids resistant to biological.
Example problems[ edit ] Sum of contributing species[ edit ] The following equations demonstrate the relative contributions of each component to the alkalinity of a typical seawater sample. Adding CO2 to the solution lowers its pH, but does not affect alkalinity. At all pH values: Only at high basic pH values: The dissolution or precipitation of carbonate rock has a strong influence on the alkalinity.
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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. 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.
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[ edit ] Processes that increase alkalinity[ edit ] There are many methods of alkalinity generation in the ocean.
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. 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 H2Swhich eventually escape into the atmosphere. Conversely, aerobic degradation can decrease AT. This process occurs in portions of the ocean where oxygen is present surface waters.
A third feature of this relationship involves the pH of seawater as the ambient CO2 level rises. If CO2 is allowed to double Figure 2the pH drops by 0. Consequently, in the future, the pH of seawater may actually drop into the upper 7's from the 8. The theoretical relationship between carbonate alkalinity and pH for seawater in equilibrium for preindustrial air green; ppm carbon dioxidecurrent air blue; ppm carbon dioxide and possible future air red; ppm carbon dioxide using equations 2 and 3.
Alkalinity - Wikipedia
An important point to keep in mind is that the relationship will be altered slightly if the tank is not in equilibrium with the air. Specifically, reef tanks are often not in equilibrium with the air, making the internal pCO2 for the tank something different than the surrounding air.
For example, tanks using limewater can have a pH value of 8. Looking at Figure 2, this puts them off of the theoretical relationship for seawater in ambient air. The fundamental explanation is that the tank is deficient in CO2. In effect, the tank has an internal pCO2 that is more like that for the preindustrial air with ppm CO2 Figure 2. In this case, driving more CO2 from "normal air" into the water would lower the pH to about 8. Again, that set of values falls off of the theoretical curve shown in Figure 2.
In this case, the tank has an artificially high internal pCO2 of more than twice "normal air".
Driving more CO2 from the tank into "normal air" would raise the pH to about 8. A third way that reef tanks can present unusual combinations of pH and alkalinity is if the tank is in an environment where the ambient CO2 is far from normal. Rarely would such a situation involve reduced CO2, but homes and businesses are frequently elevated with respect to CO2.
Such levels as those represented by the ppm line in Figure 2 are frequently encountered by aquarists, especially those living in newer, "tighter" homes and some have proven this fact to themselves with carbon dioxide detectors.
Aquarists that experience chronic low pH despite adequate alkalinity and aeration may do so because their homes have such elevated levels of carbon dioxide. Many of these aquarists have found that the pH of their tanks rises substantially by simply leaving a window near the tank open to permit better exchange with exterior, "normal" air. Finally, pCO2 fluctuates within a reef tank every day because of the activities of the organisms present.
Some are producing CO2 as a waste product of metabolism, including all organisms in the dark. Those that photosynthesize consume CO2 during the day. As a consequence, the pCO2 rises during the night and declines during the day. This change in pCO2 is largely responsible for the pH fluctuation over the course of a day.
For all of these reasons, a tank may move between the red and green lines of Figure 2 or further in extreme cases without the alkalinity changing at all.
Typical diurnal pH fluctuations in a reef tank and in some natural lagoons, for that matter are about 0. For tanks with a larger fluctuation than about 0. This minimization is best accomplished by maximizing the gas exchange between the tank and "normal" air through better circulation, better aeration through devices such as skimmers, having part of the tank system, such as a refugium, on a reverse photocycle so some organisms are always photosynthesizing, or by more rapidly exchanging the room air with exterior air.
One can also impact the diurnal pH fluctuation by adding high pH additives like limewater or other high pH alkalinity additives during the nightly pH minimum, and by adding low pH additives like sodium bicarbonate during the daily pH maximum. The magnitude of the alkalinity itself, of course, can influence pH stability, and that is the focus of the next section.
What is "Buffering" Buffer and buffering are terms that are thrown around indiscriminately in the world of reefkeeping, and the actual meaning of these terms is often lost. Many aquarists refer to any alkalinity supplement as a buffer, but this isn't the case.
For example, neither sodium bicarbonate nor sodium carbonate, taken alone, is a true buffer. A buffer is something that helps minimize pH changes in the presence of added acid or base. No buffer can completely stop the pH from changing when acid or base is added. The change in pH, however, is made smaller when an appropriate buffer is used.
A buffer is almost always comprised of two different chemical entities. Bicarbonate and carbonate together, for example, form a buffer in the pH range from about 8 to 11 in seawater, though the buffering is best between about 8. Here's what is happening on a chemical level. When a base such as OH- is added to the system in an effort to raise pHsome of the bicarbonate is converted to carbonate.
Chemistry And The Aquarium: The Relationship Between Alkalinity And pH
This process effectively "uses up" some of the OH- that was added, and the pH does not rise as much as it would without the "buffer". At about pH 8. At lower pH, there is less CO, and at pH 8. Consequently, seawater is not especially well buffered against substantial pH drops when the pH is already less than 8.
It is, however, well buffered against substantial pH rises. Here's an actual experiment. The results of an immediate pH measurement before atmospheric carbon dioxide has a chance to equilibrate are: This result shows that the water is better buffered against a pH rise than a pH drop, and the reason for this difference is simply that there is more bicarbonate than carbonate at pH 8. The only reason that the drop stops at pH 6. Nevertheless, there is much more to fully understanding how a buffer works.
Chemists have chosen the term "buffer intensity" symbolized by b to reflect the buffering capacity of a solution at any given pH. While it has an exact mathematical definition, it is beyond the scope of this article to describe b in detail.
There are, however, a few details worth mentioning in the context of reefkeeping with additional details are provided in "Aquatic Chemistry Concepts" by James Pankow.
M Alkalinity and P Alkalinity
The most important fact to reefkeepers is that the buffering due to bicarbonate and carbonate, at a given pH, is directly related to the carbonate alkalinity. If you double the alkalinity, you double b, and hence have twice as much buffering due to the carbonate and bicarbonate system.
Consequently, marine aquaria with higher alkalinity tend to have greater buffering against pH swings. Other, more esoteric tidbits arise from this system as well. For example, while the buffering against substantial changes can be different in the two different directions as shown experimentally abovethe buffering against very small changes is necessarily exactly the same.
That's what b represents: