7.7: Seawater Constituents

For this section, you may want to download a black-and-white periodic table or refer to the periodic table above. It helps to mark—and even color code—the following major categories of chemical constituents in seawater:

• Major constituents, the elements whose concentrations exceed 1 part per million in seawater
• The trace elements, or minor constituents, elements whose concentrations in seawater are less than 1 part per million, including elements that are important in biological reactions, the so-called biologically important nutrients
• Dissolved organic matter (DOM), a heterogeneous group of very dilute, dissolved carbon–containing substances, and particulate organic matter (POM), living and dead particles of organic matter
• Dissolved gases, gas-phase elements in seawater

The Major Constituents

The major constituents comprise a group of 11 elements that make up 99.9 percent (by weight) of all the elements dissolved in seawater. Put another way, if you evaporate all the water from seawater so all that remains is salt, 99.9 percent of that salt will be composed of these 11 elements.

In order of decreasing concentration, we find chlorine, sodium, magnesium, sulfur, calcium, potassium, carbon, bromine, boron, strontium, and fluorine. Of course, seawater is mostly water, so don’t be misled. The dissolved salts make up only a small fraction of the total weight and volume of seawater, about 3.5 percent.

Perhaps most remarkably, these major constituents (except carbon) occur in identical proportions throughout the ocean, even though salinity may vary. Establishment of the Principle of Constant Proportions, as the relationship between major constituents is known, marked a key advance in chemical oceanography.

While a number of chemists worked on aspects of seawater chemistry from the late 1600s to the mid-1800s, three men in particular set in motion discoveries that would establish the foundations of chemical oceanography: Alexander Marcet (1770–1822), a Swiss-born chemist and physician; Georg Forchhammer (1794–1865), a chemist and geologist from Denmark; and William Dittmar (1833–1894), a German-born chemist who worked in Scotland. These men contributed techniques, sampling approaches, principles, and terminology that remain in use today. (See Wallace 1974—followed here—for extensive details of their contributions.)

Marcet (in 1819) was the first to suggest that seawater from different parts of the Atlantic Ocean and the Mediterranean Sea “contained the same ingredients” and in “very nearly the same proportions to each other,” regardless of their salinity. In 1843 Forchhammer received several samples of Mediterranean seawater from a friend and set out to refine the methods of analysis. His painstaking work was summarized in a single publication, On the Composition of Sea-water in the different parts of the Ocean (Forchhammer 1865). Dittmar (1884) built on Forchhammer’s findings in his work on samples collected during the Challenger Expedition. Dittmar’s findings mirrored Forchhammer’s, with several important improvements. He adopted Forchhammer’s term “salinity” to describe total dissolved salts but modified it to the “parts of total salt in 1000 parts of sea water.” Here he established a commonly used unit for salinity—parts per thousand. He confirmed that six elements comprised the greatest percentage of salts, and that those salts occurred in constant proportions in all depths of all oceans. Dittmar’s result firmly established the Principle of Constant Proportions as a useful assumption when analyzing seawater.

The benefit of this principle comes when applied to the chemical analysis of seawater. If you know the concentration of just one of the major constituents, you know the concentrations of all of them (except carbon). Consider, for example, the color distribution of M&Ms (a very popular topic among math folks on the internet). According to StatsMedic (2023), M&Ms are produced in the following proportions: Brown (12.4%), Red (13.1%), Yellow (13.5%), Green (19.8%), Orange (20.5%), and Blue (20.7%). So if you had a bowl of M&Ms and wanted to know how many blue M&Ms were in the bowl, what might you do? Well, according to the Principle of Constant Proportions of M&Ms, you could count the number of brown M&Ms (the smallest number) to calculate the total number of M&Ms because

(Total brown M&Ms) ⁄ 0.124 = Total number of M&Ms

(Eq. 10.1)

Once you know the total number of M&Ms, you only need multiply by the percentages to calculate the number of each color.

It’s much simpler and much faster, which is why the Principle of Constant Proportions made such a splash with marine chemists. Determination of the salinity of a sample could be based on analysis of a single constituent. And it could be accomplished in hours instead of days.

Because nine other major constituents (except carbon) occur in a constant ratio—regardless of the salinity—major constituents are considered conservative elements; their ratios are conserved at all salinities. Most of the next group of elements, the minor constituents, obey no such rules.

The Trace Elements

Pilson (2013) confesses that

a great deal of time and material effort is required to achieve reliable results for some of the important elements . . . that are present at very low concentrations in the central water masses of the oceans. The real concentrations are so vanishingly small that the quantities present are easily swamped by contamination.

Thus, we remain uncertain about the source–sink pathways of many of the minor constituents. Nevertheless, we can make some useful generalizations.

The minor constituents—elements whose concentrations are less than 1 part per million in the ocean—include elements too dilute to be considered major. Pilson (2013) lists some 63 minor constituents and their concentrations, excluding gases and radioactive compounds. An element’s designation as “minor” should not mislead you. A good number of these elements prove vital to living things, and many trace the motions of water masses and other useful oceanographic processes. What unites them primarily are their low to extremely low concentrations in seawater.

Like major constituents, minor constituents originate from riverine, atmospheric, and subseafloor processes (Chapter 11). And like major constituents, their concentration and distribution in the ocean depend on the nature of their sources and sinks. For the most part, however, chemical oceanographers categorize minor constituents according to processes that remove them. On that basis the minor constituents can be divided into five types based on their vertical distribution in the water column (following Bruland and Lohan 2006).

Conservative-Type

Just like major constituents, some minor constituents exhibit conservative-type distributions; that is, their concentration in seawater varies proportionally with salinity. Metals such as cesium (Ce⁺) and rubidium (Rb⁺) exist as unreactive cations that spend a great deal of time in the ocean, 300,000 and 3 million years, respectively. Others, such as molybdenum and tungsten, play key roles in biological processes, but their high concentrations relative to their use by organisms make their concentrations largely invariable. Less than a dozen minor constituents are included in this category.

Scavenged-Type

As noted by Jeandel et al. (2015), “particles act as an essential regulator of ocean chemistry because they determine the residence time of many dissolved elements in seawater.” Particles prove especially important for the removal of trace elements from the ocean. Jeandel and colleagues broadly distinguish two types of particles in the ocean: (1) external particles—those delivered via rivers, the atmosphere, the seafloor (shelf currents and hydrothermal vents), and even outer space (e.g., micrometeorites); and (2) internal particles, those formed within the water column via biological activities, aggregation of organic matter, or chemical precipitation. Dissolved trace elements interact with particles of both types, but here we focus on the removal of dissolved trace elements by sinking particles, a process called scavenging.

Scavenged elements easily attach to sinking particles, a process called adsorption. The attraction of trace elements to particles involves physical forces, or physisorption—the same weak forces that allow a gecko to climb a wall. It also involves chemical forces, or chemisorption, which pertains to chemical bonding between atoms and the surface of particles. Chemisorption occurs, for example, in the removal of harmful substances in water (or air) by activated carbon filters. Note that adsorption—the sticking to or adherence of dissolved elements to particles—differs from absorption, in which the dissolved elements become part of the particle. It’s kind of like the difference between hanging on to the outside of a cable car in San Francisco (adsorption) or getting inside it (absorption). It’s not a perfect simile, but you get the idea.

Particles stick together, too, a process called aggregation. Thus, scavenging of trace elements by particles in the ocean is a three-part process: (1) adsorption of an element onto a particle; (2) aggregation; and (3) sinking of the particle to the seafloor.

Bruland and Lohan (2006) note that scavenged-type trace metals tend to occur in their greatest concentrations near their sources: the atmosphere, rivers, vents, or sediments. Aluminum, for example, shows its highest concentrations in the Atlantic Ocean and Mediterranean Sea due to a high input of atmospheric dust in these regions. Scavenged-type elements show their greatest depletion in older waters—those with less frequent atmospheric interaction. In the Pacific Ocean, for example, deep water remains out of contact with the atmosphere for hundreds of years, so scavenging can take place over a longer period of time. Pilson (2013) lists three dozen scavenged and “surface-depleted” minor constituents. However, he cautions that any of a number of geological or biological causes may generate a vertical distribution for a given constituent. As he puts it, “There is certainly much more to learn, and one wonders if there are more real surprises yet waiting.”

Hybrid-Type

Hybrid-type trace elements occur where a minor constituent participates in more than one pathway in its removal from seawater. Iron, for example, may exhibit surface minima—where it is strongly utilized by phytoplankton (a nutrient-type distribution)—or surface maxima—where low phytoplankton activity and significant atmospheric deposition elevate its concentrations. In deep waters, iron concentrations tend to remain constant, unlike other nutrients. Variations in dissolved iron in the ocean appear to represent a balance of recycling of organic matter, which releases iron, and scavenging of iron by particles, which removes it.

Mixed-Type

Some elements exist in multiple forms in the ocean, a chemical process known as speciation. Germanium (the element, not the flower, which is spelled geranium) takes a form very much like dissolved silica; it forms a compound called germanic acid (H₄GeO₄). But it also binds with methyl groups (CH₃) to form what chemical oceanographers call the “Teflon of the sea,” mono- and dimethylgermanium—CH₃Ge(OH)₃ and (CH₃)₂Ge(OH)₂, respectively. As germanic acid, germanium acts like a nutrient. It may be absorbed by phytoplankton and other plankton that form “shells.” As methylated germanium, however, it’s very stable, like a conservative minor constituent. Thus, the same element exhibits a mixed-type distribution, one that depends on the form of the element.

Biologically Important Nutrients

Biologically important nutrients, or simply nutrients, represent a group of dissolved substances necessary for the growth of photoautotrophs—photosynthetic organisms, such as phytoplankton, seaweeds, marine plants, and terrestrial plants. When oceanographers talk about nutrients, they aren’t talking about the stuff that comes in a box of Wheaties or even the kind referred to by nutritionists. An oceanographer’s nutrients are the dissolved stuff phytoplankton (and other photoautotrophs) absorb across their cell walls. Let’s get this straight: they do not “eat” nutrients like Pac-People eat dots; they absorb them.

As many as 28 different nutrients may be required by phytoplankton (e.g., Moore et al. 2013; Supplemental). Elements required in greater quantities are generally referred to as macronutrients, while those needed in smaller amounts are called micronutrients. Reynolds (2006) defines macronutrients as those that exist at an abundance of greater than or equal to 1 percent of the ash-free dry mass of a cell. By this definition macronutrients include carbon (C), oxygen (O), hydrogen (H), nitrogen (N), phosphorus (P), and sulfur (S). Other authors, however, restrict the definition of macronutrients to those abundant elements that may limit the biomass of phytoplankton in the ocean, namely, phosphorus and nitrogen (e.g., Jones 2011). To this list some add silicon (e.g., Falkowski et al. 1998). Silica (SiO₂) is required for the growth of diatoms—who use it to manufacture their cell walls. Surprisingly cyanobacteria also take up silica, as recently discovered by Baines et al. (2012). The accumulation of silica by cyanobacteria has major implications for the global silicon cycle (Ikeda 2021).

Micronutrients include several trace metals, such as iron (Fe), manganese (Mn), cobalt (Co), nickel (Ni), copper (Cu), zinc (Zn), cadmium (Cd), and boron (B), among others (e.g., Carrano et al. 2009; Twining and Baines 2013). Iron has received a lot of attention in recent decades for its role in limiting the biomass of phytoplankton in regions of the ocean that otherwise should be very productive. Micronutrients also include organic nutrients, such as thiamine, biotin, and vitamin B₁₂.

Given their use by phytoplankton, biologically important nutrients tend to be depleted in the upper ocean, where sunlight drives photosynthesis. Conversely, they tend to be highest in the deepest and oldest waters that remain out of sunlight, where photosynthesis doesn’t occur.

DOM and POM

Greek philosopher-scientist Aristotle (385–323 BCE) was the first to report the presence of fatty substances on the sea surface, evidence for the presence of organic matter—biologically produced, carbon-containing molecules. Organic matter in the ocean broadly divides into two types: dissolved organic matter (DOM) and particulate organic matter (POM).

Operationally, DOM includes anything that passes through a 0.2 micrometer filter (e.g., Libes 2009). Technically, that includes small microbes and viruses—which are not dissolved—but that’s why it’s called an operational definition. POM includes anything alive or once living. As expressed somewhat tongue-in-cheek by Pilson (2013), POM can be “collected with harpoons . . . hook and line . . . dragging a net . . . or filtering,” referring to the different sizes of living things—from blue whales to microbes—which make up POM.

DOM and POM mostly consist of carbon atoms. So oceanographers focus on measurements of dissolved organic carbon (DOC) or particulate organic carbon (POC) to estimate DOM or POM, respectively. By far, the greatest portion of organic matter in the ocean belongs to the DOC pool. Some 662 gigatons are reported (e.g., Hansell et al. 2009), slightly more than exists in terrestrial plants (e.g., Schlesinger and Barnhardt 2013). Libes (2009) reports an estimated 55 gigatons of POC in the ocean, 50 gigatons of which is detritus—particles of dead organic matter (think finely ground leaf litter). Note that this amount includes everything you think of when you think of marine organisms.

Most of the biomass of the ocean can be found in the small things. Biomass-wise, the largest organisms, such as whales, represent only a tiny fraction of the ocean’s organic matter, around a tenth of a percent of the total POC (Smith 2007). Because DOM and POM play a significant role in primary productivity and marine food webs, we’ll defer our discussion of them until later chapters.

Dissolved Gases

Have you ever shaken a can of soda and asked an unsuspecting person to open it for you? To their surprise and your amusement, the can erupts like a geyser, spraying soda everywhere.

Why do soft drinks fizz and froth when you open them after shaking? It’s because they contain a gas—carbon dioxide—that has been dissolved in the liquid (consisting of about 90 percent water and a combination of sugar, flavors, sodium, phosphorus, and other ingredients). Soft drink manufacturers pump carbon dioxide gas into the liquid at cold temperatures and high pressures, effectively trapping the gas in the container and creating a supersaturated solution. When that pressure is released—when the tab or cap is opened—the reduced pressure allows the gas to come out of solution.

Shaking the bottle or can ahead of time accelerates the process because shaking disrupts the arrangement of the water molecules holding the gas in place. As Live Science puts it, shaking releases dissolved gases into the space above the liquid, increasing its pressure (Choi 2022). Carbon dioxide also reacts chemically with water, a process that affects everything from lakes to coral reefs. In fact, increases in the atmospheric concentration of carbon dioxide are changing the chemistry of the ocean, as we shall see. But first, let’s take a look at some fundamental properties of gases in seawater. Feel free to pop your own favorite beverage while reading this section.

Factors Affecting Dissolved Gases

Three main physical or chemical factors affect the dissolution of gases in seawater: temperature, salinity, and pressure:

• The colder the water temperature, the greater the solubility of the gas.
• The lower the salinity (the fresher the water), the greater the solubility of the gas.
• The higher the pressure, the greater the solubility of the gas.

As the temperature of water decreases, the ability of the water to dissolve gases increases. Put another way, the solubility of gases increases with decreasing temperature. Conversely, as water temperature increases, less gas can be dissolved in a liquid. The solubility of gases decreases with increasing temperature. This relationship may seem at odds with what you learned previously about the solubility of solids in water. But the dissolution of gases in water works a bit differently from the dissolution of solids. Most gases (except for CO₂) do not separate into their individual elements, like NaCl does. Instead, the water molecules surround the gas molecules, essentially trapping them in a net. At lower temperatures the water molecules are closer together and form a tighter net, and gas molecules cannot escape as easily. At higher temperatures water molecules are spaced farther apart, so the “mesh” of the water net expands, allowing more gas molecules to escape.

Higher salinity reduces the solubility of gases. That means that freshwater holds a greater concentration of gases than saltwater. In effect the water net is more crowded in saltwater; Na and Cl take up space and reduce the capacity of water to surround gas molecules. In oceanic regions with higher salinity—regions with high rates of evaporation relative to precipitation or where sea ice forms, causing rejection of the seawater’s salts—the saturation concentration of dissolved gases may be lower. In regions with lower salinity—regions with high rates of precipitation relative to evaporation, in estuaries, or near the mouths of rivers—the saturation concentration of dissolved gases may be higher. However, temperature may counteract salinity effects. Colder temperatures in polar oceans reduce the effects of higher salinity in these waters, and warmer temperatures in estuaries may act against any increase in solubility as a result of fresher water.

Pressure affects the solubility of gases in water, as we might expect from our soft drink example above: The higher the pressure, the greater the solubility of the gas. So, the deeper you descend in the ocean, the greater the solubility of gases in seawater. In fact, every 33 feet (10 m) increases solubility by about 10 percent. Thus, seawater at the surface (100% saturation) may become undersaturated with oxygen and other gases at depth. So at 33 feet (10 m), the saturation concentration of a gas drops to 90 percent. At the same time, pressure acts to capture bubbles of air that are swirled downward by breaking waves. As a result, surface waters of the ocean are often supersaturated with gases, just like a soft drink.

Molecular interactions help explain why oxygen dissolves in water. As polar water molecules approach, the oxygen molecule’s normally symmetrical electron cloud rearranges into an asymmetrical, weakly polar configuration—what is known as a dipole. As it does so, oxygen takes on the characteristics of a polar molecule. The positive and negative ends of the water molecule are now attracted to the negative and positive ends, respectively, of the oxygen dipole, and the oxygen molecule dissolves. Other nonpolar gases, like nitrogen, may also be dissolved in this way. Think of it as being like when you tell your parents your sibling made you do it. One molecule influences another molecule to do something it normally wouldn’t.

The Air–Sea Interface

The boundary between the ocean and the atmosphere—where the sky and sea meet—is called the atmosphere–ocean interface, or simply the air–sea interface. It’s a part of the ocean that few have heard of, yet it just might be the most important feature of our planet. Engel et al. (2017) call it “the ocean’s vital skin,” the sea surface microlayer, the thin film of water no deeper than 1 millimeter (the width of a pinhead). This interface acts as the gatekeeper for what enters the ocean from the atmosphere, and vice versa. A brief look at the role of this layer in gas exchange across the air–sea interface provides insights into the dynamic set of physical processes that control the concentration of dissolved substances in the ocean.

Gas Exchange

The simplest model of gas exchange across the air–sea interface (e.g., film theory; see Sarmiento and Gruber 2006; Liao and Wang 2013) assumes that transfer occurs as a result of simple diffusion—the movement of molecules across a boundary due to a concentration gradient (i.e., where the concentration of molecules across the boundary differs). For example, a drop of food coloring in a glass of water diffuses throughout the glass until the same concentration of dye exists everywhere in the water. If diffusion was the only process controlling the exchange of gases across the air–sea interface, then the concentration of gases in the ocean would be directly proportional to their concentration in the atmosphere (according to a relationship known as Henry’s law). However, dissolved gases in the ocean, even near the surface, exhibit undersaturation or oversaturation relative to their atmospheric concentrations. In other words, most gases dissolved in the ocean are not in equilibrium with the atmosphere. That means other processes contribute to air–sea exchanges.

Anyone who has watched breaking waves at the shore or experienced them from beneath the surface as a swimmer or diver has likely noticed the effect of the waves on the water: They create foam, bubbles, and sea spray. Indeed, breaking waves in the open ocean also accelerate the exchange of gases across the air–sea interface. Waves break through the sea surface microlayer and allow gases to be transferred faster, both in and out of the atmosphere. Bubbles also penetrate the surface waters, where higher pressures compress and capture the gas molecules. The penetration and dissolution of bubbles of air caused by breaking waves is called bubble injection, a process that may lead to supersaturation.

Waves also contribute to turbulence in the upper ocean, the chaotic and irregular changes in the speed and direction of flowing water or air. Perhaps you’ve experienced turbulence in an aircraft or watched the turbulent chaotic spume of an incoming wave. Turbulence carries gases deeper and in more directions than simple diffusion. Bubble injection is an example of a turbulent process.

One way to think about turbulence is to consider the difference between pouring a dash of milk into a cup of coffee and letting it sit until the milk is distributed throughout the coffee or stirring the coffee to homogenize the mixture. Turbulence acts like a stir stick that transfers heat, particles, gases, and a whole host of other things throughout the ocean.

Dissolved Gases in Ocean Research

Research on seawater chemistry and dissolved gases in the ocean takes center stage for understanding a number of oceanographic processes, including the movements of water (e.g., using gases as tracers), the chemistry and biology of the seafloor (e.g., methane seeps), human impacts on marine life (e.g., dead zones), and ocean chemistry (e.g., ocean acidification, the increase of acidity in the ocean). Here we take a look at two gases: chlorofluorocarbons, used as tracers of water mass movements, and methane, a potential energy source and potent greenhouse gas. In later chapters we look more closely at research on carbon dioxide and oxygen, as a full understanding of these gases requires a bit more knowledge of the physics and biology of the ocean.

Chlorofluorocarbons (pronounced klor-oh-FLOOR-oh-car-buns), also known as CFCs, represent a class of gas molecules containing chlorine and fluorine (i.e., halogens) used widely as refrigerants and propellants from the 1930s through the 1980s. When carried high into the atmosphere, these gases break apart and destroy ozone (O₃), most of which is found in the ozone layer—a part of the stratosphere, which is the atmospheric layer above the one in which we live. Near the ground ozone can be harmful (“bad” ozone). In the stratosphere ozone absorbs DNA-damaging ultraviolet radiation (just like a good pair of sunglasses). So the ozone layer (“good” ozone) acts as a shield protecting life—including ocean life—from the harmful effects of ultraviolet radiation.

The reduction of the ozone layer occurs mostly in polar regions (most notably Antarctica), where cold temperatures and other atmospheric conditions favor the ozone-destroying reactions. The net result has been an ozone hole, a region where stratospheric ozone concentrations are greatly reduced. Because CFCs destroy the ozone layer, they were banned in 1987 under an international treaty called the Montreal Protocol on Substances That Deplete the Ozone Layer. Former UN Secretary-General Kofi Annan (1938–2018) called this global effort “the single most successful international environmental agreement to date” (Annan 2000, p. 56). Unfortunately, CFCs and their byproducts (especially chlorine) can remain in the atmosphere for hundreds of years. Thus, while most CFCs have been completely phased out and are no longer released into the atmosphere, their long lifetimes mean that the ozone hole will remain with us for several more decades. In 2022, the ozone hole reached its smallest maximum extent in decades (Gillespie 2022). A report released by the World Meteorological Organization in 2022 highlighted that “actions taken under the Montreal Protocol continued to decrease atmospheric abundances of controlled ozone-depleting substances” (WMO 2022).

Despite their negative impacts, CFCs offer an opportunity to study and better understand deep ocean circulation. CFCs act as a kind of time stamp for the origin and movements of water in the ocean’s interior, a group of substances known as ocean tracers. When surface waters sink, the gases they contain are transported with them. Because these waters are in contact with the atmosphere, the soon-to-be-deep water takes on the properties of the atmosphere. As a result, deep water masses record the composition of the atmosphere at the time of their sinking. Dissolved gases, other dissolved substances, particulate matter, and even temperature and salinity can act as tracers. These tracers allow oceanographers to figure out when a particular mass of water formed and how it mixed with other water masses as it moved. As the Water Encyclopedia (2023) puts it, “Tracers serve as a ‘dye’ with which to follow the circulation of ocean waters.”

The discovery of hydrothermal vents reinvigorated seafloor exploration and paved the way for the discovery of other chemosynthetic systems, such as cold seeps—places in the seafloor where methane and other energy-rich fluids leak through the sediments. Cold seeps provide an energy source for chemosynthetic ecosystems that is every bit as bizarre and wonderful as hydrothermal vents. In fact, some of these ecosystems may survive for hundreds—maybe even thousands—of years. There has been considerable speculation about the mining of the reserves of the methane associated with cold seeps. But the environmental consequences and wisdom of unleashing yet another carbon-emitting energy source remain to be determined.

Methane (CH₄), a little-heralded molecule composed of one carbon and four hydrogen atoms, exerts a tremendous influence on Earth’s atmosphere and plays an important role in food webs and carbon cycling in lakes and the ocean. In the atmosphere methane acts as a greenhouse gas, the second-most potent after carbon dioxide.

Methane caught scientists by surprise when it was found in frozen form at the bottom of the ocean as methane hydrates, molecules of methane gas trapped in crystals of water ice that form at reduced temperatures and elevated pressures. Nearly all the methane and methane hydrates found in ocean sediments originate from the activities of methanogens, a type of microbe that converts organic matter into methane. Methanogens can be found in the water column and in sediments at the seabed surface (e.g., Max and Johnson 2016; Niu et al. 2018). Reeburgh (2007) estimates that methanogens living in seafloor sediments produce a billion tons of methane annually.

The primary interest in methane hydrates—though undoubtedly important in deep-sea ecology—lies with their vast potential as an energy source. Natural gas hydrate, or NGH, as it is known in the energy industry, represents potentially the largest reservoir of unconventional gas on our planet (Max 2018). Even if only 15 percent of the world’s reserves can be economically extracted, there is enough NGH on the seafloor to meet global energy needs for the next 200 years or longer (Max and Johnson 2016).

As of 2019, Japan and China have successfully extracted deep-sea methane in experimental test wells, but no technology currently exists for commercial extraction of NGH from the seabed. Nevertheless, given the vast reserves, the energy industry views NGH as a “bridge” energy source, one that will enable countries to make the transition from fossil fuels to renewable sources. As Max and Johnson (2016) put it:

A completely renewable energy future may be in our future, but its timing is very uncertain. . . . Natural gas is the clean hydrocarbon fuel that will reach into the renewable energy future. Its continued availability at affordable prices becomes increasingly important as coal and oil power plants are retired and energy demand becomes increasingly filled by development of renewable or intermittent power sources.

Clearly, oceanographers have a lot more to discover and learn about the ocean and its resources. It is a frontier ripe for exploration and adventure—and a lot closer to home than the moon or Mars.