Monitoring for coastal and estuarine acidification is difficult due to the living and non-living resources that influence these waters. It is important that monitoring takes place at diverse locations and that data is collected on a variety of factors that contribute to changing water chemistry across ecosystems.
To build a robust monitoring network in the Mid-Atlantic, MACAN is working to develop a monitoring plan that tracks acidification in estuarine, coastal, and ocean waters of the region. In order to create such a plan, MACAN first developed an understanding of existing acidification monitoring. This information is represented on a series of acidification monitoring maps found under the Oceanography theme on the Mid-Atlantic Ocean Data Portal.
A Mid-Atlantic Ocean Data Portal map depicting acidification monitoring efforts in the region.
As MACAN develops a monitoring plan for a physical monitoring network in the Mid-Atlantic, these maps will serve as a jumping off point. The monitoring plan will seek to:
For more information on the development of the monitoring plan, please contact MACAN at info@MidACAN.org.
There are four parameters most relevant to monitoring for acidification, Partial Pressure Carbon Dioxide (pCO2), Dissolved Inorganic Carbon (DIC), pH, and Total Alkalinity (TA). To better understand the carbonate chemistry of a given location, at least two of the four parameters must be measured. The remaining two parameters can be calculated. Additionally, other components of the carbonate system can be calculated including saturation state (Ω), which is a measure of whether calcium carbonate will dissolve or form (At equilibrium, Ω = 1; Saturated, Ω > 1, calcium carbonate will form; Undersaturated, Ω < 1, calcium carbonate will dissolve).
To understand more about each of the components of carbonate chemistry and how they influence seawater keep reading for some specifics. The key take away is that each piece of the carbonate chemistry interacts with the other pieces through a variety of naturally occurring processes. As the amount of carbon dioxide in the atmosphere increases, the ocean also uptakes more carbon dioxide. This uptake can result in a shift in the extent to which various processes occur and a shift in the ratios of the components of seawater towards more acidified conditions. Understanding the dynamics of these components now and into the future can help researchers, natural resource managers, impacted industries, and the general public understand how the estuarine, coastal, and ocean waters of the Mid-Atlantic may be changing overtime and what that might mean for those ecosystems.
Partial Pressure Carbon Dioxide (pCO2) refers to the idea that in a mixture of gases, the total pressure exerted by that mixture is equal to the sum of the individual pressures, or partial pressures, of the gases. According to Henry’s law, the concentration of a gas dissolved in a liquid is directly proportional to the partial pressure of the gas above the liquid. Meaning, that when CO2 fills the atmosphere, more CO2 will get dissolved in water. Often, however, estuaries and coastal areas have very high concentrations of dissolved CO2 compared to the overlaying atmosphere. This is due to organisms respiring during the night when photosynthesis (and the uptake of CO2) is not occurring. When this happens, CO2 leaves the water and enters the atmosphere. Think about a can of soda. Before it is opened there is a headspace at the top of the can that is pressurized with CO2. That high CO2 partial pressure of the headspace ensures that the CO2 in the soda remains dissolved. When the can is opened, the CO2 in the headspace encounters the outside atmosphere (which has a relatively low concentration of CO2). The result is that the soda fizzes and bubbles float to the surface, releasing previously dissolved CO2 into the atmosphere.
Thus, coastal systems can either act as a source of CO2 or a place where CO2 gets absorbed, depending on their unique partial pressures. pCO2 can be thought of as a measurement of the amount of CO2 in a liquid (i.e. seawater). A higher level of pCO2 will increase acidification.
When carbon is dissolved into water it reacts to form four forms called species: aqueous carbon dioxide (CO2aq), carbonic acid (H2CO3), bicarbonate (HCO3− ) and carbonate (CO32−). Collectively these species are known as dissolved inorganic carbon (DIC). The four substances undergo a series of reactions simultaneously until an equilibrium is reached. When CO2 is dissolved in water, either from the atmosphere or from respiration from coastal organisms, it reacts with H2O to produce H2CO3:
CO2aq + H2O <-> H2CO3 (Eq. 1)
H2CO3 further breaks down to release a hydrogen ion and form HCO3− :
H2CO3 <-> HCO3− + H+ (Eq. 2)
The amount of CO2 entering the system can vary (i.e. the pCO2 varies) and the equilibrium can shift. If pCO2 continues increasing, HCO3− reacts further by losing another H+ :
HCO3− <-> CO32− + H+ (Eq. 3)
Again, these reactions occur simultaneously until an equilibrium, or balance, is reached. If more CO2 is added (a rise in pCO2) H2CO3 breaks down into HCO3− and CO32− and more H+ ions get released, increasing acidification.
The continued forcing of CO2 into the water can alter the relative abundance of carbonate forms. For instance, under high pCO2 conditions the concentration of CO32− (the molecule used by most calcifying organisms to build their shells) actually gets reduced, in favor of HCO3 and CO2aq / H2CO3. See the plot below to see how changes in CO2 effect other forms of carbon. When CO2 concentration is high, CO32− concentrations are low and the pH decreases, thus making it more difficult for calcifying organisms to form.
The plot above depicts the change in the carbonate system of seawater from ocean acidification. This plot shows that as CO2 concentrations increase in water, the amount of CO32- (the molecule used by most calcifying organisms to build their shells) and pH get reduced. By BeAr [Public domain], via Wikimedia Commons
pH or “potential of hydrogen” is a measure of the acidity or alkalinity of a solution. It is a measurement of the amount of hydrogen ions (H+) in a solution. An acid is characterized by its propensity to donate hydrogen ions. More H+ ions means the solution is more acidic. An alkaline substance, or base, on the other hand, accepts hydrogen ions. Fewer H+ ions mean the solution is more basic or alkaline.
Acidity and alkalinity are measured on a logarithmic pH scale from 0 (acid) to 14 (base). Strongly acidic solutions can have millions or trillions of times more H+ ions than strongly basic solutions, so each unit change on the pH scale corresponds to a 10-fold change in H+ concentration.
Alkalinity is the measurement of the ability of a solution (i.e. seawater) to accept positively charged ions (H+) and bind them to a negatively charged base ion or molecule, thus neutralizing the acid. Alkalinity is important as it is a measurement of the ocean’s ability to resist acidification. Total Alkalinity (TA) refers to the difference between the negatively (Cl−, SO42−, Br−, and F−) and positively (Na+, Mg2+, Ca2+, K+ and Sr2+) charged ions in water. Together these ions have a net positive charge because overall there are more positively charged ions in water. The overall charge of the water is neutralized by the HCO3− and CO32− ions. HCO3− and CO32− are defined as carbonate alkalinity (AC). The ions that influence TA are conservative, meaning that their concentrations are consistent in water. AC is the main factor in determining alkalinity. If the AC is high enough, the HCO3− and CO32− ions can act to buffer the water from dramatic changes in pH accepting hydrogen ions. The increased forcing of CO2 (high pCO2) into the water, however, is resulting in lower pH and could alter carbon speciation (see Dissolved Inorganic Carbon section above) which can reduce the water’s ability to buffer changes in pH. In addition, estuaries and other coastal waters typically have lower alkalinity due to dilution from freshwater sources, compared to the open ocean.
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