The chemistry of marine aquaria is a complex subject and one that is not easily described in a short article. Previous articles on marine chemistry in Aquarium Frontiers authored by Craig Bingman have dealt with selected topics of interest to marine aquarists. In particular, these articles have focused on the biochemistry taking place in aquaria. In this article I will endeavor to provide an understanding of seawater itself, rather than how the components are used by the tank inhabitants.
What’s In Seawater?
Major Species
Seawater has been found to contain virtually every chemical element, although some of them are found in very small concentrations. Water is, of course, the most abundant molecule, comprising about 97 percent of seawater. Water itself is far more complicated than is generally recognized and has been an active area of chemical research for more than a hundred years.
One of the remarkable things about water is that it is liquid at room temperature. Based simply on its molecular weight, it ought to be a gas. Nitrogen (N2) and oxygen (02) are much heavier than water (H2O), and yet they are gasses and water is a liquid. Why?
The reason involves the hydrogen bonding that takes place in water. The hydrogen atom of one molecule of water interacts strongly with the oxygen atom of a nearby water molecule. This interaction is much weaker than the bond between atoms within a single water molecule, but it is strong enough to make the water molecules “prefer” to be surrounded by each other, rather than floating around individually, as they would in a gas. Hydrogen bonding is best viewed as a fleeting interaction between water molecules that lasts only a tiny fraction of a second before breaking. Once broken, however, they quickly reform, perhaps to a different water molecule. On balance, each water molecule is bonded to one or two other water molecules almost all of the time.
Major Ions
Most of the remaining constituents of seawater are inorganic ions. The major components of seawater — all ions present at greater than 1 part per million (ppm) or 1 milligram per liter (mg/L) — are shown in Figure 1 and Table I. A different definition of major ions based on the numbers of ions present, rather than the weight of those ions, has a slightly different list, with lithium being added. Together, these ions account for 99.9 percent of the dissolved solutes in seawater.
It is clear from Figure 1 that seawater contains mostly table salt (sodium and chloride). In fact, sodium and chloride comprise 86 percent of the ions present in seawater, by weight.
One important point about these concentrations: they are correct for typical seawater, which contains about 35 parts of salt by weight per thousand parts of seawater (35 ppt). This seawater has a specific gravity of around 1.027, so it may be higher than is maintained in many marine aquaria. As the salinity of seawater is varied, these concentrations move up and down together. Consequently, if an aquarium contains water with a specific gravity of 1.023, the salinity is about 30 ppt and all of the concentrations in Table I are reduced by about 14 percent.
A logical question to ask is why do we not hear much discussion about chloride, sulfate or sodium levels in marine aquaria, if they are among the most abundant ions? The answer is that while they are very important, their abundance makes it difficult for them to become significantly depleted or enriched without altering the salinity. Of course, one could start out with a salt mix that did not contain the correct proportions, but assuming one starts out correctly, there isn’t any normal activity in a marine aquarium that will significantly change the levels of these ions (without changing salinity).
All of these major ions are essentially unchanged in concentration at different locations in the ocean, except as salinity changes move them all up or down together. Ions that do not change concentration from place to place are referred to as “conservative type” ions, a description that also applies to some of the minor and trace elements that are discussed below.
I have. also included organics on this list, though they traditionally are not considered a major specie. As will be discussed below, organics are important in seawater, but are poorly understood.
Minor Ions
There are various definitions, of which ions in seawater constitute the “minor ions.” By some definitions, the list of constituents is rather long. Table II shows just a few of the constituents of seawater that are often labeled as minor ions. The more abundant of these are sometimes lumped with the major ions (such as lithium), while the least abundant (such as iron) are often lumped in with trace elements. Ions in this category often vary significantly with location in the ocean. That is primarily because many of them are tightly linked to biological activity. These ions can be locally depleted if biological activity is high enough. Ions that vary in this fashion are referred to as “nutrient type” ions, because they are consumed by one or more types of organism.
Trace Elements
There is much discussion about trace elements in marine aquaria and for good reason. Most chemicals dissolved in seawater are classified as trace elements simply because there are so many ions and molecules present at very low concentrations. In many cases, these ions are quite variable in concentration from place to place and also as a function of depth. Anyone wishing to view extensive lists of these ions is advised to check out one of the references given at the end of this article.
Many of these trace elements are metals. While people typically view dissolved heavy metals as toxic, a great many of them are essential for organisms. Their toxicity is primarily related to their concentration: a happy medium is essential, where enough of each of these metals is present for life to exist, but not so much is present as to be toxic.
A perfect example is copper. It is present in natural seawater at about 0.25 parts per billion (ppb), which is about a thousand times less than the toxic levels often used to kill microorganisms in the treatment of sick marine fish. It is, however, absolutely necessary for many animals to have copper available to them to survive.
Some of the most important trace elements to marine aquarists are those involved in the nitrogen cycle (ammonia/nitrite/nitrate). These are discussed in detail below.
Organics
The nature of organic molecules is certainly the most complicated aspect of seawater chemistry. Organics comprise about 2 ppm of seawater. Of this 2 ppm, the majority is in the form of dissolved organic carbon (DOC). DOC includes all fully dissolved organic compounds and any particulates that are small enough to pass through a 0.45-micron (?m) glass fiber filter. Strictly speaking then, it is not all fully dissolved. Any organic particles greater than 0.45 ?m are called particulate organic carbon (POC). The POC is about a factor of 10 lower in concentration than DOC and is composed of living and dead organisms, as well as assemblies of organic molecules.
DOC is an incredibly complicated mixture of molecules that represents billions of years of biological waste products from uncounted numbers of different organisms, combined with reactions catalyzed by light, heat, inorganic catalysts (metals), biological processes, and many other factors. It includes carbohydrates (20 to 35 percent of the total), humic substances (10 to 30 percent of the total), amino acids and proteins (2 to 3 percent), hydrocarbons (less than 1 percent), carboxylic acids (1 percent) and steroids (trace).
There is also a great deal of uncharacterized organic material. In fact, the study of seawater organics is an active area of research. Additionally, the summation of all dissolved organics in the ocean is a pool of carbon larger than carbon dioxide in the atmosphere, so it cannot be ignored by those looking at the planetary carbon cycle. In addition to carbon, these organics contain significant amounts of oxygen, nitrogen, phosphorus, and sulfur.
It is probably also safe to say that most, if not all, closed marine systems have higher organic levels than the ocean, although hard numbers are difficult to come by. The desire to reduce these organic levels is one of the reasons for the popularity of skimmers with marine aquaria.
What Forms Do Ions Take In Seawater?
In the previous sections I have described what ions are present in seawater, but I have not presented the forms they typically take. Contrary to popular belief, many of these ions are attached to each other in solution and do not act as completely individual species. This tendency to form ion pairs in solution is much more prevalent for some ions (e.g., Ca2+, Mg2+, CO32-, F-, OH-) than it is for some others (e.g., Na+, K+, Cl-, Br-). In general, the tendency to form ion pairs is higher for ions with a higher net charge. In the next few sections, I will present an overview of some of these interactions and why they are important.
Simple Ions
The simplest positively charged ions in solution are sodium (Na+) and potassium (K+). They are primarily free ions, with a shell of three to four tightly bound water molecules attached to them. This is known as the “primary hydration sphere.” These water molecules are fairly tightly bound, but are rapidly exchanged with other water molecules from the bulk solution (at a rate of about a billion exchanges per second for each ion!). Beyond this first shell are another 10 to 20 water molecules that are less tightly bound, but that are still strongly influenced by the metal ion. These types of hydrating water molecules are present for all ions in solution and won’t be mentioned further for each ion in turn.
A small proportion of both sodium and potassium (about 5 percent) exists as ion pairs with sulfate, forming NaSO4- and KSO4-. This type of ion pair is best viewed as a temporary association between the two ions and may only last for a very small fraction of a second before the ions move apart. Nevertheless, this type of association can have very important implications for the behavior of these ions, as will be shown below. Ions forming such pairs actually “touch” each other. That is, most or all of the hydrating water molecules that are in between them have been temporarily removed. This removal of the intervening water molecules is the primary distinction between ion pairs and ions that are simply near each other.
The simplest negatively charged ions, chloride (Cl-) and bromide (Br-), form few ion pairs in solution. They are primarily present in the form of hydrated free ions, with two and one tightly bound water molecules, respectively.
Carbonate
One of the more complex interactions, and one that is very important for marine reefkeepers, involves carbonate (CO32-). Carbonate is primarily ion paired in solution, with only about 15 percent of it actually present as free CO32- at any given point in time. This fact is very important to the maintenance of calcium and alkalinity levels in aquaria, because it is the free carbonate concentration that “wants” to precipitate with calcium as calcium carbonate (CaCO3). If the free carbonate levels rise too much, the calcium levels will drop due to CaCO3 precipitation.
So, what is carbonate ion paired with? Primarily magnesium, forming soluble MgCO3. This is the reason why magnesium levels are so important in marine aquaria for maintenance of simultaneously high levels of alkalinity and calcium. If magnesium is too low, more carbonate will be in the free form and will “want” to precipitate as calcium carbonate.
Carbonate is also ion paired to sodium and calcium, forming soluble NaCO3- and CaC03, respectively. The soluble calcium ion pair sounds odd, but it is essentially one individual molecule of CaCO3 that is soluble in water: it is not precipitated out of the solution. The fact that carbonate is also ion paired by sodium is one of the reasons that salinity has an impact on the amount of calcium and alkalinity that can be maintained in solution: lower salinity means lower sodium, which means more free carbonate and a greater likelihood of precipitation of CaCO3.
Ion pairing has another large effect on carbonate that is more subtle. In water, carbon dioxide hydrates to form H2CO3, which can then break up (ionize) into protons (H+), bicarbonate (HCO3-) and carbonate CO32-).
When CO2 is added to water, the system will come to equilibrium with specific concentrations of each of the species shown above. By LeChatelier’s principle, if one takes away something from one side of the equilibrium, the equilibrium will shift in that direction. For example, if carbonate is removed from the system, then each of the reactions shown will proceed to the right, effectively replacing some of the carbonate that was removed.
Importantly, that is exactly the effect that takes place in seawater when carbonate is “removed” by forming ion pairs. It is only the “free” concentration of these species that determines the position of the chemical equilibrium, so carbonate in the form of an ion pair does not “count,” and the equilibrium shifts strongly to the right. If one then counts carbonate in all forms (free and ion paired) it is found to be far higher in seawater than in freshwater at the same pH and ion pairing is the primary reason.
The exact same effect can be seen in the solubility of CaCO3.
In this case, if CaCO3 is added to water, it breaks apart into Ca2+ and CO22-. Eventually, an equilibrium is reached where no more CaCO3 will dissolve. However, if some of the carbonate is removed by ion pairing (and some of the Ca2+ as well), then additional CaCO3 can dissolve to replace those that were “lost.” This is the primary reason that CaCO3 is approximately 15 times more soluble in seawater than in freshwater.
Calcium, Magnesium and Strontium
Calcium, magnesium and strontium are primarily present in the free form, hydrated by six to eight tightly bound water molecules. A small percentage (about 15 percent) is present as an ion pair with sulfate. Much smaller percentages are present as ion pairs with carbonate and bicarbonate. Importantly, while these complexes involve only a small percentage of the total calcium and magnesium, they involve a large portion of the total carbonate (which is possible because there is so much calcium and magnesium compared to carbonate).
Sulfate
As mentioned above, sulfate forms ionic interactions with most positively charged species in seawater. In fact, more than half of it is in the form of an ion pair, with NaSO4- and MgSO4 dominating.
Phosphate
Phosphate in marine aquaria is of tremendous importance because it is often a limiting nutrient for algae growth. In seawater, the amount of phosphate present is typically quite low (usually less than 0.1 ppm) and often varies significantly from location to location. In many marine aquaria, however, the phosphate concentration can be significantly higher (up to several ppm).
The ability to export phosphate from marine aquaria has been the topic of lengthy discussion and is the object of numerous commercial products. The nature of the inorganic phosphate present in marine aquaria, however, is certainly more complicated than traditionally credited.
Inorganic phosphate can exist in a number of forms, in a manner analogous to carbonate.
Ignoring ion pairing and complex formation for the moment, phosphate is primarily found in the HPO42- and PO43- forms in seawater. This is quite different than in freshwater at the same pH, where the H2PO4- and HPO42- forms predominate. Table III shows the forms of phosphate present in seawater at a pH of 8.1.
To a large extent, the high proportion of phosphate present in the PO43- form in seawater is due to ion pairing, just as in the case of carbonate. These various phosphate species pair extensively with magnesium and calcium in seawater. PO43- is nearly completely ion paired (96 percent), while only 44 percent of HPO42- is paired. This is what causes the shift in the equilibrium to more of the PO43- form in seawater compared to freshwater (just as it does for carbonate).
Additionally, phosphate will interact with certain ions in a manner that is much stronger than simple ion pairs. Phosphate can, for example, complex with a number of positively charged species, including both metals (e.g., iron) and organics. These interactions further serve to reduce the concentration of free phosphate.
Phosphorus is also contained in dissolved organics. While natural seawater has more inorganic phosphate than organic forms, this may not be true in aquaria where much higher organic levels prevail.
Metals
The metals, in particular, are strongly ion paired in solution. Copper primarily forms soluble CuCO3, iron forms soluble Fe(OH)3 and silicon (not strictly a metal) forms (Si(OH)4. Some of the other metals that are biologically important (e.g., zinc, molybdenum, manganese, cobalt) form a wide variety of ion pairs with different ions in solution. In some cases, the number of different species that form is extensive. Table IV shows the speciation of copper in seawater at a pH of 8.1.
In recent years, however, it has become more and more apparent that certain metals are largely complexed to organic materials, even in natural saltwater where the level of organics is low. In a marine aquarium, the level of organics can be higher than in the ocean, so such complexes are even more likely to form.
In addition to complexation of metals to the widespread organics present in the oceans (e.g., humic acids), there is also the possibility of complexation to specific organics that were made exclusively for that purpose. For many microorganisms, metals such as iron are limiting nutrients for growth and these creatures have designed systems to bring iron to them.
Bacteria and fungi, for example, release organic compounds called siderophores into the environment. They are large organic molecules with a very high affinity for iron. The released siderophores eventually encounter an iron atom and bind very strongly to it. The organisms themselves have enzymes in their outer membranes that interact strongly with siderophores that contain iron, and transport them into the cell. Consequently, the siderophores can be viewed as collection devices for iron.
Of course, many of the siderophores released into the ocean are not quickly reabsorbed by the microorganisms and remain in solution. In a closed marine aquarium with a large population of microorganisms, one would expect that such molecules would be present in solution. Consequently, many metals in solution may be bound by such molecules.
Additionally, many aquarists intentionally add complexing agents in the various supplements they add to their aquaria. These include EDTA and citrate, which are two common forms for adding iron. These will equilibrate with other metals already in the tank and the tank will then contain a variety of metals complexed to these organics.
Nitrogen Compounds
The primary nitrogen compound in seawater is nitrogen gas (N2). It is present at about 11 ppm at 25 degrees Celsius (77 degrees Fahrenheit), although its solubility is a strong function of temperature, with nearly twice as much dissolving in near freezing seawater. Nitrogen gas is present at a higher concentration than any other dissolved gas, with oxygen (02) at 7 ppm, argon (Ar) at 0.4 ppm and all others at sub-ppb levels (not including carbon dioxide, which is primarily ionized in seawater).
There are certain organic and inorganic forms of nitrogen at concentrations lower than nitrogen gas. The organic forms are poorly defined, but include such molecules as proteins.
The inorganic forms are much more familiar to aquarists as components of the nitrogen cycle. The concentrations of these components in seawater are highly variable. In natural seawater, ammonia (NH3) ranges in concentration from 0.02 to 8 ppm (as ammonia), nitrite (NO2-) ranges from 0.005 to 0.2 ppm (as nitrite) and nitrate (NO3-) ranges from 0.06 to 30 ppm (as nitrate). These values vary by location, depth and time of year. Other inorganic forms present at much lower concentration include hydroxylamine (NH2OH), nitrous oxide (N2O), and hyponitrite (N2O22-).
Ammonia exists in two forms in seawater. The primary form is ammonium (NH4+), which accounts for about 95 percent of the total in seawater at a pH of 8.1. The secondary form is free ammonia (NH3), which accounts for the remaining 5 percent. These proportions vary strongly with pH and the free ammonia form rises as pH rises, to about 50 percent of the total at a pH of 9.5.
The toxicity of ammonia towards fish has been found to depend upon pH, with some researchers observing lower toxicity at lower pH. It has been suggested that this relationship between toxicity and pH is due to the proportion of ammonia in each form at a given pH. While these ideas seem to have been accepted by many in the aquarium hobby, the exact cause of this relationship is unclear and is beyond the scope of this article. This topic is discussed in more detail in Captive Seawater Fishes (Spotte 1992).
Nitrite and nitrate are both interesting molecules in that they exist in a number of resonance forms. If one draws a simple structure for these molecules it appears that the oxygen atoms are not all exactly the same, with one carrying a negative charge, while the others do not. Experimentally, however, this has not been found to be the case: all oxygen atoms are exactly equivalent.
How can this be? Resonance forms are a simple way of thinking about this, with the various forms interconverting extremely rapidly. The only thing required to convert one form to another is to move electrons around within the ion, so it can happen essentially instantly. In reality, the electrons are spread around these ions in such a way that each oxygen on average carries a partial negative charge (-? in the case of nitrite; -1/3 in the case of nitrate).
Iodine
Iodine seems to get an amazingly disproportionate amount of discussion with respect to marine aquaria and much of it is incorrect. The reasons for this are many, but are primarily related to its chemical and biochemical complexity. In fact, its chemical complexity is far greater than many aquarists are aware.
Iodine takes two primary forms in seawater: iodide (I-) and iodate (IO3-). The often quoted value for the total concentration of iodine in seawater (0.06 ppm) is reasonably accurate, although the value varies significantly. This value, however, is a combination of both iodide and iodate. It is not correct to state that seawater contains 0.06 ppm of iodide. The value for iodide is more typically around 0.01 ppm or less, although it is sometimes as high as 0.03 ppm and sometimes as low as 0.002 ppm. The remainder is iodate.
Additionally, the interconversion between iodide and iodate in seawater is very slow. This reaction is believed to be mediated in a number of ways, including catalysis by light and microorganisms. It is probably safe to say, however, that the two are not in equilibrium in marine aquaria. One effect of this lack of equilibrium is that dosing one type does not necessarily give you any of the other type.
It is not well known which forms are used by which organisms, so I will not comment on the necessity of maintaining specific levels of iodide or iodate. There is good evidence, however, that iodide is rapidly depleted in marine aquaria, although it is not well established where it goes. Conversion of iodide to iodate has been observed in aquaria, but this may not represent a significant sink. Iodate itself is much slower to become depleted from marine aquaria and can build up to toxic levels if it is being actively dosed.
An additional complication is that some aquarists dose a third form of iodine: I2. Lugol’s solution, for example, is a combination of iodide and iodine. When iodine (as I2) is added to seawater, it quickly reacts to form other iodine species that probably end up as both iodide and iodate in marine tanks.
Conclusion
There are, of course, many other details of seawater chemistry that may be of interest to marine aquarists. This article is only a first pass at understanding the chemistry behind what is happening in our tanks.
For those wanting a more in depth exposure to marine chemistry, I recommend two books: Captive Seawater Fishes. Science and Technology by Stephen Spotte (Wiley-Interscience, New York. Pp. 942.) and Chemical Oceanography, Second Edition by Frank J. Millero (CRC Press, Boca Raton, FL. Pp. 469.).
The Spotte book is excellent, with sections directed specifically toward aquarium chemistry. It covers chemistry from the standpoint of aquarium keeping, rather than understanding of the natural ocean. It is also practically oriented, rather than directed toward a deep chemical understanding of phenomena.
The Millero book will only be of interest to those who are undaunted by chemical reactions and jargon. It is, however, the best marine chemistry book I have encountered. It gives a tremendous amount of detail about natural marine systems, but has no discussion about aquaria. Most of the chemical data in this paper was pulled from this book.
Previous “Biochemistry of Reef Aquariums” columns in Aquarium Frontiers magazine have also dealt with selected topics of interest to marine aquarists, especially the column on “Ion Pairing, Buffer Perturbation and Phosphate Export in Marine Aquariums” (Bingman, C. 1996. Aquarium Frontiers 3[1]:10-17).