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Water in the Universe | Part 3

Water in the Universe | Contents | Intro | Part 1 | Part 2 | Part 3

Earth's Oceans

A Brief Early History – Formation to 3 Billion Years Ago

At some point in the very early history of the Earth, perhaps as early as 4.4 billion years ago and certainly by 4.2 billion years ago, the surface of the Earth became cool enough that water could condense upon it and accumulate into what became Earth’s first global ocean.  Earth would have been a true water world back then with perhaps a few hot spot islands but no continental land masses.  This first global ocean may have formed at a temperature of as much as 400 °F which was possible because the very early Earth had a CO2 atmosphere at least 100 times, maybe 250 times, thicker than now which would also have made the ocean acidic.  

Whatever the source of that water, internal or external, the hot mantle of that time could only have held about 0.7 times the mass of today’s oceans in comparison to today’s mantle capacity of about 2.3 oceans equivalent.  Estimates are that the amount of surface water was much greater in the time of the early Earth with something on the order of four times the ocean volume then than now.  That means that a lot of surface water from the time of the early Earth must have moved into the mantle, especially considerably later when plate tectonics and subduction made possible large transfers.  Lacking a protective ozone layer, photolysis of water vapor may also have been significant.  

By 3.7 billion years ago there were significant land masses, probably volcanic islands and the first continental micro-continents (cratons).  The early cratons were not stable; they could founder down into the mantle.  Between 3.5 and 3.2 billion years ago, it is estimated that the amount of land surface was from 2 to 12% of the Earth’s surface compared to the 29% of today.  By 3 billion years ago, there is definite evidence of subduction; early plate tectonics had begun.  

The moon formed about two Earth diameters away, probably within 50 million years of Earth’s formation because of a collision between Earth and a Mars-size body (Theia).  Since then the moon has been moving away from the Earth (the moon is now about thirty Earth diameters away), slowing the Earth’s rotation (conservation of angular momentum) from about 4 hours in a day to 24.  Some 3 billion years ago the Earth had a day of 15 hours and some significant land areas.  This ancient land would have had low relief (limited plate tectonics) and  the moon-generated tides would have been much higher and more frequent than they are today.  They may even have swept across most of the land areas of the time, greatly enhancing their erosion.  

There is evidence of microbes (microscopic organisms) in marine sediments 3.7 billion years ago.  Cyanobacteria began to release O2 into the atmosphere but it would take hundreds of millions of years to become significant in the atmosphere and ocean; the ocean remained anoxic (no oxygen) for most of its history.  But before continuing this story, let us consider some ocean chemistry.  

Some Chemistry of Earth’s Current Oceans

Since water is so good at dissolving so many things, it is not surprising that oceans everywhere, be they surface oceans like on Earth or subterranean oceans such as in Europa (one of the big moons of Jupiter), would be loaded with various dissolved salts.  At this time, we only have access to one world’s oceans, those of the Earth.  It may be that other oceans on other worlds might have a drastically different chemical profile than those of the Earth but, surprises aside, there are probably broad similarities.  Even in Earth’s oceans there are significant variations in the amounts and presence of some dissolved substances in different locations and at different times in Earth’s history.  As a start, however, we will consider what the water chemistry of typical (if there is such a thing) well-oxygenated salty ocean water looks like today, comparing it to the chemical profile of some continental ‘fresh’ water but first, a little background.  

Dissolved substances in water are mostly ions, that is, they started out as elements/compounds which have lost or acquired one or more electrons.  They are not neutral; they have an electric charge (gas can dissolve in water too without ionizing).   Those ions which have a positive charge are known as cations; those with a negative charge are anions.  Most water samples can be characterized by a small group of major ions which are those ions that make up most of the ions in the water.  The major cations are usually sodium (Na+), calcium (Ca++), potassium (K+), and magnesium (Mg++).  The major anions are usually chloride (Cl), bicarbonate (HCO3), and sulfate (SO4=).  In some water, iron (Fe++) and nitrate (NO3) and even hydrogen phosphate (HPO4=) can become important although usually at much lower levels.  

The table below lists the major ions and their amounts found in ‘typical’ ocean water, in human blood plasma, and in a ‘typical’ water sample of groundwater in the Catskill Formation (mostly sandstone) in northeastern Pennsylvania, along with drinking water standards.  The units for water have been all converted to mg/L for comparison; if, for example, you were to look up blood concentrations, they would likely be given as mEq/L (milli-equivalents per liter) or mg/dL (mg per 100 ml).  The drinking water standards are a mix of primary and secondary standards with that of Na+ being a mere recommended limit for those with high blood pressure.  K+ has no standard as, even though it is usually a major cation, it is not normally present in a high enough amount in drinking water to be a potential problem for human health.  

For comparison, TDS (total dissolved solids) and pH are also given.  The TDS minus the sum of the major ions gives an indication of the amount of non-major ions.  The pH is important as it can influence what form of major ion is present.  Except at very high pH, the bicarbonate (HCO3) anion, for example, will be much more abundant than the carbonate (CO3=) anion.  

Also shown are the amounts of the elements, the counterparts of the major ions, present in crustal and mantle rock which are the sources of the ions found in the water.  Their amounts are given in mg/kg.  

Depending on the source of information, there can be a wide variation in the amounts reported.  In some cases, the amount listed is probably not much better than an order of magnitude.  And, of course, the values for the Pennsylvania water sample are certainly not representative of many other groundwater sources.  It happens, for example, that the area represented by that water sample has little or no limestone.  Groundwater samples from limestone areas would have much higher values of Ca++.  Nevertheless, in spite of all of these caveats, some interesting observations can be made.  

Ion Seawater Blood
Water Std
Catskill Fm Element Crust Mantle
Cl 19,000 3550 <250 5 Cl 0.3 0.03
Na+ 11,000 3265 < 20 8 Na 25 3
SO4= 2,600 48 < 250 10 S 0.4 0.2
Mg++ 1,300 37 < 10 5 Mg 25 230
Ca++ 400 100 < 180 23 Ca 40 25
K+ 380 196 1 K 22 0.3
HCO3 140 1650 30-400 90 C 0.7 0.1
Fe+++ 0.002 3 0.3 0.08 Fe 50 60
TDS 35,000 7000 < 500 130 O 470 450
pH 8.1 7.4 6.5-8.5 7.2 Si 280 215
          Al 82 22

Considering the major cations (positively-charged ions), it is clear that Na+ is by far the most abundant cation in sea water and Fe+++ the least.  And yet, in the Earth’s crust, iron (Fe) is the most abundant of the potential cations, closely followed by calcium.    The crustal abundances are explained by the crustal mineral abundances.  How much of these elemental or compound ions are in sea water would depend on their solubility.  

The crust, both types, is essentially some kind of aluminum silicate which explains the very high crustal abundance of oxygen, silicon, and aluminum and yet, neither silicon nor aluminum show up as major cations (Si++++ and Al+++) in sea water.  Those two cations, with so much positive charge, are tightly linked to oxygen, forming compounds which are almost insoluble in water.  SiO2 (silica) can be present in water on the order of ppm (parts per million) which is measurable but certainly doesn’t make it a major cation.  It is, however, even at those very low concentrations in sea water, significant in forming the tests (tiny shells) of marine plankton such as diatoms and radiolarians.  

Al2O3 is completely insoluble in water although other aluminum compounds, such as the alum [in the form of Al2(SO4)3] used in tertiary sewage treatment to remove phosphate from the treated sewage, can make aluminum more soluble as aluminum-water complexes such as [Al(H2O)6]+++.  Incidentally, dissolved aluminum in water is not a good thing for most living organisms.  

Iron, in the form of the ferric cation (Fe+++), has a very low solubility of about 0.3 mg/L (about the same in ppm at that low concentration).  The ferrous cation (Fe++) is more soluble (pH matters) and can be present in significant amounts in water with low oxygen.  Mine water and well water often have little oxygen and can be saturated with ferrous iron.  When that iron-saturated water reaches the surface and is exposed to the oxygen in the air, it will, within seconds, oxidize to the much less soluble ferric iron and promptly precipitate as colorful ferric hydroxide on stream beds, plumbing fixtures, and laundry which is why there is a secondary drinking water standard of 0.3 mg/L for iron in drinking water.  The level of Fe+++ in sea water suggests that the solubility of the ferric iron is on the order of 0.002 mg/L, much less than that of Fe++.  

The Ca++, Na+, K+ cations ultimately come from the chemical weathering of the feldspars in igneous rocks like granite (continental crust) and andesite (intermediate between continental and oceanic crust); Mg++ is derived from other minerals in basalt (oceanic crust) and peridotite (upper mantle).  Within the crust, these elements are abundant on the same order of magnitude.  In sea water, however, Na+ is a magnitude more abundant than Mg++ which is some three times more abundant than Ca++ or K+.  Sedimentary rocks whose constituents are ultimately derived from igneous rocks rich in feldspars, can be an important source of these cations too.  Limestone, for example, is a very important source of Ca++ and dolomite and gypsum for magnesium.  

It is tempting to blame that discrepancy on differences in solubility but it is more complicated than that.  One approach might be to compare the solubility of these elements as components of a compound.  Since Cl is the most abundant anion (negatively-charged ion) in sea water, consider the chloride compounds of those elements (chloride compounds tend to be quite soluble in water):  

Solubility of Four Chloride Compounds in g/L

Ca Cl2 MgCl2 NaCl KCl
745 529 360 278

The solubility of these compounds in water depends on temperature, among other factors, but it is clear that they are all soluble in the same order of magnitude and, surprisingly, CaCl2 is the most soluble of the lot.  However, this is misleading.  In sea water, while the Mg++, Na+, and K+ will precipitate out as chlorides, the Ca++ will precipitate out, first as the much less soluble CaCO3 (as low as 0.013 g/L in pure water) and then as CaSO4•2H2O (gypsum) leaving nothing to precipitate out as CaCl2.  

If you were to start out with a sample of sea water and allow it to evaporate, the first major precipitate would be CaCO3 (limestone), followed by CaSO4•2H2O (gypsum), and then, after some 90% of the water had evaporated, by NaCl followed by KCl and other K compounds. After all of the water had evaporated and the gypsum dehydrated to CaSO4 (anhydrite), the composition of the precipitate would be something like:

% Composition of Precipitate from the Evaporation of Seawater

NaCl MgCl MgSO4 CaSO4 CaSO4 KCl CaCO3
77.8 10.9 4.7 3.6 2.5 2.5 0.3

Of course, not all of the cations weathered from crustal rocks and minerals end up dissolved in sea water.  Many of them become part of clays and the shales that form from the clays but that is a future story.  

The sea water abundance of anions is a great surprise: Cl is a magnitude more abundant than the other anions but is only a very tiny percentage of the composition of crustal rock.  Why that is is that the chloride anion is extremely soluble in water; it was mentioned above how soluble chloride compounds are.  The crust may have only a tiny amount of chloride in its minerals but once released by weathering, almost all of the Cl ends up in the oceans.  If the sea water should become sufficiently concentrated (brine), NaCl (halite, the mineral) may crystallize out to form rock salt deposits which are stable only if buried or in a very arid environment.  

The other major anions are sulfate (SO4=) and bicarbonate (HCO3).  There are nitrates (NO3) and biphosphates (HPO4=) but they are normally present at much lower levels although they can pose problems in drinking water and as nutrients promoting too much algae growth.  Remember that at the beginning of this section, it was mentioned that this discussion was about well-oxygenated ocean water (also at a slightly alkaline pH).  If the oxygen levels in the water are very low or nonexistent, the sulfur would be in the form of sulfide (S=, or hydrogen sulfide gas, H2S), not sulfate.  pH is also important as it determines the relative abundance of the bicarbonate and carbonate (CO3=) forms.  Calcium combines with the carbonate, producing CaCO3 as calcite or aragonite (the minerals) and as limestone (the rock).  In a similar manner, biphosphate would become phosphate (PO4≡) and precipitate as some phosphate compound. 

It has been suggested that the abundances of dissolved ions in blood plasma mimic those of Earth’s early oceans in a time before the oceans were as salty as they are today after billions of years of accumulating salt.  If that is true, then the proportion of ions in seawater should be similar to that of the ions in the blood.  This would assume that the ions were accumulating at the same rate in the oceans for billions of years which is probably not realistic but let’s make the comparison anyway:

Proportion of Major Ions:  Sea Water to Blood Plasma

Cl Na+ SO4= Mg++ Ca++= K+ HCO3
5.4 3.4 54 35 4 2 0.08

The proportions of chloride, sodium, calcium, and potassium are similar, at least of the same order of magnitude, but sulfate, magnesium, and bicarbonate are quite different; blood plasma has, proportionately, much less sulfate and magnesium and much more bicarbonate than modern sea water.  The big difference in bicarbonate is understandable as metabolism produces carbon dioxide gas which would dissolve in the blood plasma, some of which would react with the water to become the bicarbonate ion.  As for sulfate and magnesium, such proportionately high amounts in the blood would be detrimental and there probably is some mechanism to ensure that the blood keeps their levels much lower.  The drinking water standard for sulfate and magnesium certainly suggests that is true.  So, do the ionic amounts of the major ions in blood plasma reflect the ionic amounts of an earlier, less salty ocean?  Maybe.  

The concentration of ions in drinking water sources varies wildly depending on the source so the values given for the Catskill Formation water sample are absolutely not typical except in that its values all meet drinking water standards (of course, some sources don’t meet the standards).  More importantly, none of the values are zero.  Some minimum level of some ions in drinking water is very desirable, both for health and taste.  Distilled water is great for steam irons but not for humans, or dogs, or cats…  Sometimes, impurity is a good thing.  

TDS (Total Dissolved Solids) is, purportedly, the sum of the masses of all of the ions in the water, including minor ions.  Theoretically, if you were to sum up all of the individually determined amounts of all of the ions in the water (the testing of which could be rather expensive), it should match the TDS.  In actuality, that is sometimes not the case.  Should there be a big discrepancy between TDS and the sum of the ions, that would suggest one of several possible reasons:  All of the tests were accurate but a significant ion (possibly an ion that is usually a minor ion but not in this particular sample) was left out.  One or more tests were inaccurate.  All the tests were accurate but there was a mistake in transcribing the results.  One surprisingly common transcription mistake is a misplaced decimal.  Whatever the reason, any big discrepancy is a flag.  

The proportion of the TDS values for sea water and blood plasma is 5, in line with that of sodium and chloride, suggesting that all significant ions have been accounted for.  The drinking water standard for TDS is < 500 mg/L which is an order of magnitude less than that of blood plasma (7000 mg/L) and two orders of magnitude less than sea water (35,000 mg/L).  Needless to say, it is not a good idea to drink either sea water or blood.  

The pH of modern oceans is 8.1 compared to the 7.4 of blood plasma.  Both values are slightly alkaline (7 is neutral; anything less is acidic) and since the pH scale is logarithmic, there is not a huge difference between a pH of 7.4 and 8.1.  However, if today’s sea water pH dropped to the 7.4 of blood plasma, it would be catastrophic for some modern marine life.  Sometimes even small changes can be very significant; we are currently worried about a much smaller drop in oceanic pH.  However, it should be mentioned again that the pH of Earth’s early oceans was considerably acidic.  

The drinking water standard for pH has a range of 6.5 to 8.5.  Humans have a much greater tolerance for drinking acidic beverages than for alkaline beverages.  The drink with the highest pH you are likely to encounter is green tea with a pH on the order of 9.  With a pH of 2.9 to 3.3, grapefruit juice is notoriously acidic but not many people drink it.  Coke, on the other hand, has a pH of 2.6 to 2.7, not because of its carbonation but because of its phosphoric acid (it can etch teeth).  Coke is great for cleaning car battery terminals and toilets but would never pass the drinking water standard.  Lemon juice is even worse at 2.25. 

You might wonder why, if people can enjoy drinking such acidic things as sodas and citric juices, the drinking water standard for pH is as high as 6.5. The reason is that water with a lower pH can be corrosive, both corroding the pipes through which it travels and picking up in solution undesirable ions that are more soluble at the lower pH. Water at a pH of 6 or lower, for example, will corrode cement pipes over time.

Land and Water

3.8 billion years ago continental crust made up about 3% of the Earth’s surface.  The first small continent (Vaalbara) formed about 3.7 billion years ago.  By 3.0 billion years ago, large-scale plate tectonics had begun and the first supercontinent, Kenorland, appeared although, it must be admitted, it wasn’t all that big.  At the end of the Archaen Eon, 2.5 billion years ago, some two thirds of the Earth’s continental crust had formed.  It wasn’t until 1.5 billion years ago that the continental crust reached today’s 29% of the Earth’s surface.  Although the oceans still make up the majority of the Earth’s surface, Earth was no longer a complete water world with important consequences to global climate.   

The face of Jupiter is marked by a pronounced banding which is produced by alternating east-west zones of atmospheric circulation.  Although Jupiter has neither land nor water, such a circulation pattern may have been similar to the Earth’s early ocean currents when Earth had little land to block such zonal global flows.

One consequence of a zonal ocean current flow can be seen in today’s oceans: the Antarctic Circumpolar Current (ACC) which is the strongest ocean current the Earth has.  This current today flows unimpeded clockwise around Antarctica, insulating it from warmer water currents to the north.  But until about 35 million years ago, the southern tip of South America remained attached to Antarctica, blocking an unimpeded current flow around Antarctica.  The counterclockwise current flow circulation (a gyre) in the South Atlantic between South America and Africa was able to reach all the way down to Antarctica, bringing warmer water to that polar continent, keeping it largely free of glacial ice.  Once South America broke away from Antarctica, however, the ACC could establish itself and isolate Antarctica from the warmer waters of the South Atlantic gyre.  Thus began the enormous accumulation of glacial ice on Antarctica which persists today.  

Antarctic Circumpolar Current

North-south ocean currents are very important in transporting heat from the warmer regions of the Earth to the cooler regions as northwestern Europe can appreciate.  The Gulf Stream, part of the clockwise North Atlantic gyre, transports enormous amounts of heat to northwestern Europe, keeping it considerably warmer than it would otherwise be.  Consider Stockholm, Sweden and Churchill, Canada, both at a latitude of 59° N.  Churchill, located on the western shore of Hudson Bay and famous for its polar bears, is far from the influence of the North Atlantic gyre and has an average annual temperature of 19 °F.  In contrast, Stockholm enjoys a relative toasty average annual temperature of 45 °F.  Were there no continents to force north-south ocean current flows, all of the higher latitudes would be considerably colder, more like Churchill.  Most of the north-south heat transfer would be left to the atmosphere which would probably mean stronger winds.  

Land and Water Temperature Changes

Water has an enormous heat capacity of 4184 J/kg•°C as compared to rock (granite) with a heat capacity of 790 J/kg•°C.  This means that it would take about five times as much heat to raise the temperature of water the same amount as it would for land.  Consequently, the land will heat up much faster than the water will.  The reverse also applies.  The air temperature above the ocean will change much less from night to day (diurnal) as it will over land.  Humid land areas will see less diurnal change than will arid land.  

During the day the land heats up the air above it, causing the heated air to rise.  If the air rises high enough and cools enough, clouds will form.  The ocean water, which remains cooler because of its greater heat capacity, is less likely to form clouds.  An early ocean explorer seeing a few clouds in the distance in an otherwise cloudless sky, might conclude that there was land under the clouds.  The difference in temperature between land and sea also creates characteristic winds.  

At dawn along a shore line (such as the New Jersey beaches) the land temperature and ocean temperature are about the same and there is no wind (ignoring things like frontal systems).  As the land heats up faster during the day, the land-sea temperature differential creates a wind.  The heated air above the land rises and cooler air from the ocean moves inland to replace it, producing a cool sea breeze.  As the sun sets, the land cools down, nearing the temperature of the ocean water, and the sea breeze stops.  Going into the night, the land cools faster than the ocean water.  Air still rises from the now warmer ocean water and cool air from the land blows out to sea to replace it, creating a land breeze.  As the sun rises, the land temperature increases to that of the ocean water and the land breeze dies; the cycle begins anew.  

On a much larger scale a similar process produces monsoons.  Instead of day and night temperature differences think of seasonal temperature differences and for the greatest effect, consider a subtropical land area adjacent to an equatorial ocean, the best example of which is subtropical India with the equatorial Indian Ocean south of it.  The Spring in India is equivalent to dawn and early morning on the New Jersey beach.   The land and water temperatures are similar so there is not much movement of air between land and sea.  Moving into late Spring and early Summer, the Indian subcontinent warms up much faster than the water.  Heated air begins to rise over India and cooler air loaded with moisture moves from the ocean inland to replace it only to be heated up in turn and caught up in the rising convection, producing ever-thickening clouds and lots of rain – the monsoon.  The Himalayas contribute to this effect as they block cooler continental air from coming in from the north.  

In most places around the world, the hottest part of the year is mid to late Summer but not in India.  Its hottest time of the year is in late Spring before the monsoon clouds block the sun.  As India moves into Fall, the land cools, approaching the temperature of the ocean water.  The monsoon winds and rain die down, corresponding to sunset on the New Jersey beach.  But the cycle does not end here.  Moving into winter, the Indian land area cools below the ocean water temperature.  Air is still rising above the now warmer Indian Ocean and cool, dry air begins to move from India south into the Indian Ocean, producing what is known as the dry monsoon, the equivalent of the New Jersey beach land breeze.  As Winter is succeeded by early Spring, the land and sea temperatures become similar and the dry monsoon ends; the cycle then begins again.  

Monsoons can also be seen in northeastern Africa, bolstered by the Ethiopian highlands which force the monsoon wind to rise (orographic effect), and generating the annual flooding of the Nile River.  Weaker monsoons occur in northern Australia and even the southeastern United States.  

The presence of land in oceans is an important generator of clouds and precipitation, at least near their boundaries.  An early Earth with little land above sea level probably had far fewer clouds which meant that less sunlight would be reflected back out into space, making the early Earth warmer than it would otherwise be.  It is interesting to note that over its billions of years of history, the Earth’s oceans have never completely frozen solid nor become horribly hot even though the sun has steadily become brighter.  The early sun was as much as 40% dimmer than it now is and yet the oceans remained liquid, probably mostly because of changes in the composition of the Earth’s atmosphere (the early greenhouse effect was much stronger then than it is now) but at least in part because of its sparse cloud cover.  As the sun became brighter, because of the appearance of land and the clouds it generates, the Earth became more reflective.  

The Oceans from 3 Billion Years Ago

Earth was no longer a purely water world 3 billion years ago but had substantial land areas.  However, its chemistry was considerably different than it now is, the prime difference being that the atmosphere and, hence, the oceans, had very little oxygen.  Much like mine pool water and some groundwater, the oceans were saturated with dissolved iron (Fe++) and were slightly acidic with a pH of 6.5 to 7 as compared to today’s 8.1.  The gradual appearance of oxygen was to result in enormous changes in the oceans.  

To paraphrase Samuel Taylor Coleridge (The Rime of the Ancient Mariner, 1798), Oxygen, oxygen everywhere, nor any whiff to breathe.  It seems a bit odd to say that there is little or no oxygen in the water when water, H2O, is mostly oxygen.  Even the silicate rock and sediment is mostly oxygen.  Of course, what is meant is that there is no free oxygen, O2, oxygen gas not combined and locked up with anything else.

A Tangent On Some Oxygen Compounds

We are now in a time when there is talk of establishing bases on the moon and colonies on Mars.  Two extremely important resources necessary on such worlds are water and oxygen.  If you have water, you can extract oxygen from it by electrolysis.  Lacking water, you could generate free oxygen by breaking the oxygen bonds in silicate rock, probably by using solar power.  Mars has substantial quantities of water in the form of ice, much of it below ground.  The moon has a surprising amount of water ice in its polar areas but little or none in other lunar areas.  There is plenty of oxygen to be had, one way or another, anywhere on the moon, if only from the rocks, but in order to make water, you would need hydrogen and that is in very short supply on the moon.  For that reason, early moon colonies will probably be restricted to the polar areas.  

Breaking oxygen bonds can be fairly easy or nearly impossible depending on the oxygen compound and the method.  Heat alone can, at relatively low temperatures, break the oxygen bonds in many copper compounds which is why the stone age was followed by the copper age.  It was then discovered that a tin oxide, SnO2, (cassiterite) could also be smelted into tin metal which, combined (alloyed) with copper (about 12% tin, 88% copper) would produce bronze leading to the bronze age.  Smelting iron from an iron oxide such as Fe2O3 (hematite) required a much hotter temperature, a technology which took a while to develop, finally resulting in the iron age succeeding the bronze age.  And then there is aluminum.  

Aluminum metal is ubiquitous in civilization today.  Many people think that aluminum doesn’t rust which is why there are things such as aluminum siding.  On the contrary, aluminum rusts and very quickly.  If you were to take a piece of aluminum and scratch it with some steel wool, look quickly.  Within seconds, the shiny scratches will dull slightly as the air produces a thin coating of aluminum oxide, Al2O3.  Aluminum actually oxidizes much faster than iron does, the difference is that the thin aluminum oxide coating prevents the oxygen from penetrating deeper into the metal, greatly slowing its oxidation, unlike iron whose iron oxide coating is much easier to penetrate.  

Consider that just as there are no native iron (unoxidized) deposits on Earth, neither are there native (natural) aluminum metal deposits.  Tin cans (actually now steel cans with an interior coating of plastic instead of expensive tin) will rust away in years to decades.  Aluminum cans may take centuries to rust away completely but rust they will.  So where does the aluminum metal come from?  

Under the right conditions, primarily tropical with seasonal rains, intense and prolonged weathering can convert an aluminum silicate rock to a mix of iron and aluminum oxides called bauxite.  The bauxite can be processed to separate the iron and aluminum oxides to produce alumina (just the aluminum oxide) which can be further processed, using great amounts of electricity, to produce aluminum metal.  The cost in electrical energy to process the alumina is why it makes good sense to recycle aluminum but why is so much electricity required?  

The oxygen bond with aluminum is much stronger than its bond with iron.  You could vaporize aluminum oxide (~ 3000 °C) without breaking the oxygen bond; heat alone is not enough.  The core of the Earth is mostly iron, not iron oxide, because the core is much too hot.  In contrast, the crust is made up of aluminum silicates, not aluminum metal.  The oxygen bond with silicon is another very strong bond that cannot be broken by heat alone, hence, the silicates of the mantle.  As with aluminum, there is no native silicon metal.  Breaking such strong oxygen bonds requires electricity which is why the aluminum age  and the later production of silicon for electronics came so late.  

Now consider water which is a hydrogen oxide.  How strong is that oxygen bond?  Like aluminum and silicon, you could heat it so much that it actually vaporizes (100 °C) without breaking the oxygen bond.  OK, so it doesn’t take much heat to vaporize water but the principle is the same.  If you want to break that oxygen bond in water, you need electricity (electrolysis).  Incidentally, although you can generate oxygen and hydrogen gas by the electrolysis of water, don’t try it with seawater.  Because of the salt in the water, the electrolysis of seawater would produce chlorine gas.  

Even at the very high temperatures and pressures deep within the ice giants, Uranus and Neptune, water remains intact as a hydrogen oxide.  It can even remain a solid at ridiculous temperatures within those planets ––> hot ice (ice XVIII). 

A Digression on Dissolved Gases in Water

If there is gas in the air, then some of that gas will enter (dissolve into) water.  One important example of that is dissolved oxygen (DO) although in seawater, much of the oxygen is being generated in the water by photosynthesis; the oxygen then moves from the water to the air.  Most gases, including oxygen, do not ionize in the water (CO2 is a big, important exception), simply remaining a mix of gas and water.  How much gas can dissolve in water depends upon what the gas is, whether the gas reacts with the water, whether the gas is generated or consumed by living things, what the gas (partial) pressure is in the atmosphere, how much other things are dissolved in the water (salinity), and what the temperature of the water is.  In general, the higher the temperature and the greater the salinity of the water, the less gas will dissolve in the water.  For dissolved gas in seawater, temperature is usually more important in explaining variations in dissolved gas levels than is salinity.  

Consider three gases: nitrogen, oxygen, and carbon dioxide.  Nitrogen gas (N2), which makes up some 78% of today’s atmosphere, is dissolved in seawater at a level of about 10 mg/L which is all the water can hold (the water is saturated with the gas).  There is some variation, mostly because of temperature.  Cold arctic water can hold up to 14.5 mg/L nitrogen gas compared to hot equatorial water which may be saturated at only 8.4 mg/L.  Since nitrogen gas does not react with water and it is not generated or consumed in globally significant amounts by living organisms in the water, dissolved nitrogen in seawater is pretty much the same everywhere with small differences because of temperature differences.  

All else being equal, dissolved oxygen gas is more than twice as soluble as nitrogen gas in water.  Of course, all else is not equal.  The atmosphere today has almost four times as much nitrogen as oxygen, living things produce oxygen in water and living and dead things consume oxygen in water.  The result is, unlike nitrogen, oxygen levels in seawater vary from zero to saturation (20+ mg/L depending on water temperature).  

Dissolved CO2 has even more variation because in addition to biologic additions and subtractions, it is much more sensitive to temperature and pressure differences and it reacts with the water. Cold arctic seawater can hold as much as fifty times as much (~ 1500 mg/L) dissolved CO2 as hot tropical water (~ 30 mg/L). More importantly, it reacts with water:

CO2 (in air) <==> CO2 (dissolved in water) + H2O <==> H2CO3 (carbonic acid)
==> H+ + HCO3– (bicarbonate ion) <==> H+ + CO3= (carbonate ion)

All of these different forms of CO2 are in equilibrium and easily and reversibly interconvert.  The amount of each form depends on the environmental conditions.  At a pH of near neutrality (~ 6.5 to 8.5) characteristic of seawater and most groundwater, the most abundant form will be the bicarbonate ion.  Being in equilibrium means that if you were to increase or decrease any one of the forms, the amounts of all of the other forms would change so as to maintain an equilibrium.  

If, for whatever reason, more CO2 is introduced into the atmosphere, more CO2 would dissolve into the seawater, producing more carbonic acid, bicarbonate, and carbonate.  The bicarbonate ion would still dominate but notice that more H+ would also be produced; the seawater would become more acidic; the pH would become lower.  

Under the right conditions and with an abundance of Ca++, limestone (CaCO3) can be precipitated out of the seawater.  This would reduce the concentration of the carbonate which would cause the equilibrium to shift so as to balance the loss of carbonate ion.  More CO2 in air would dissolve in seawater, producing more carbonic acid etc.  The formation of limestone becomes an important sink for atmospheric CO2.  

Limestone formation is much more common in the warmer waters of the world because CO2 is less soluble in warm water.  This can be a problem for people who have hard water (water loaded with Ca++).  Take that water and heat it up in a hot water heater and the Ca++ may react with the bicarbonate in the water to produce scale (CaCO3) deposits, your very own local example of limestone formation in your hot water heater and hot water pipes.  

Pressure plays a part too in how much dissolved CO2 et al there can be.  Carbonated drinks are created by forcing more CO2 into the soda under pressure (~ 2.5 atmospheres).  Once the pressure is released by opening the bottled carbonated soda, most of the CO2 will slowly come out of the soda, making it flat.  Shaking the bottle can hasten the process, sometimes with spectacular results.  The loss of the CO2 would also mean that the soda should become less acidic although most of the acidity of a carbonated drink like coke is actually from another additive, phosphoric acid, so there is actually little reduction in acidity (pH ~ 2.4).  

In Saratoga Springs, New York, there are some deep artesian wells that tap into some Ordovician limestone.  That deep, cold water becomes saturated with dissolved Ca++ and bicarbonate but when the groundwater reaches the surface, the drop in pressure reduces the ability of the water to hold bicarbonate in solution which promptly begins to combine with the Ca++ and precipitate, forming a milky cloud of CaCO3.  Notice that there are two ways to get excess bicarbonate out of the now supersaturated water.  CO2 gas could leave the water, like coke going flat, or CaCO3 could precipitate out.  But it takes a while for coke to go flat.  The CaCO3 precipitates out much more quickly.  

You might expect pure water to have a neutral pH of 7 but the moment you expose it to the air, CO2 in the air will dissolve in the water, creating carbonic acid, lowering the pH to 5.5.  Fortunately, water is not normally pure and usually contains other dissolved substances (buffers) that keep the water more pH neutral.  

Microbial Energy and Organic Matter in a Changing Ocean


One of the most familiar chemical reactions related to life is photosynthesis whose net two-way reaction is often represented as:

6 CO2 + 6 H2O <==> C6H12O6 + 6 O2

Carbon dioxide, using solar energy (certain wavelengths of light), is combined with water to produce organic matter with the release of oxygen gas.  It is a reversible reaction (it goes both ways) such that the oxidation of organic matter releases energy.  It is a net reaction in that it is the sum of many individual chemical reaction steps.  

The C6H12O6 is a general representation of organic matter in the form of a common carbohydrate. A carbohydrate is a combination of carbon, hydrogen, and oxygen, in this case, in the ratio of 1:2:1 which is why the carbohydrate is sometimes represented as (CH2O) which would simplify the equation to:

CO2 + H2O <==> (CH2O) + O2

although it should be understood that CH2O does not exist as a simple compound.  There is another class of organic compounds called hydrocarbons whose similar name can be confusing.  ‘Hydro’ can sometimes mean ‘water’ but can also mean ‘hydrogen.’  A hydrocarbon is a compound of carbon and hydrogen, no oxygen; in this case, the ‘hydro’ stands for ‘hydrogen,’ not water.  The simplest hydrocarbon is methane, CH4.  

An organic compound is a carbon compound, usually created by some life activity, with three exceptions: carbon dioxide, carbon monoxide (CO), and methane.  Those three carbon compounds are considered to be inorganic because they are so simple and so often created without the involvement of life.  Other carbon compounds usually (on Earth) are created, mediated, promoted, controlled by some life activity although some surprisingly complex ‘organic’ compounds can be found in distant gas/dust clouds in space.  

The specific organic compound given in the formula for photosynthesis, C6H12O6, is the chemical formula for glucose sugar, starch, and cellulose; the difference between the three is in how individual glucose units are linked to one another.  Most organisms cannot break the glucose bonds in cellulose so cellulose becomes a very important structural element in land plants.  Nevertheless, in the reverse reaction (the reaction goes to the left), either the controlled breakdown of glucose or starch (respiration) or the less controlled oxidation of cellulose (slow rot or rapid burning of wood), release energy.  Some of that energy can be used in other reactions along with additional elements to generate things like other carbohydrates, proteins, fats, and DNA.  

Photosynthesis (the reaction goes to the right) is inherently an activity of some living organism; you would not expect it to occur without life.  The many coordinated and carefully controlled individual steps that sum up to the net reaction just would not happen without some living thing.  

But now consider the long-term global consequences of photosynthesis.  On the left side of the equation, both carbon dioxide and water are being removed.  The Earth has oceans of water and so the loss of water to photosynthesis is negligible.  However, there is much less carbon dioxide and so billions of years of photosynthesis could and did significantly lower atmospheric levels of that gas with a great impact on global climate.  The carbon originally present as CO2 in the atmosphere ended up as organic matter in sediments and sedimentary rocks, later as limestones.  Concentrated carbon deposits include coal, petroleum, and natural gas (mostly methane) which, as impressive as they are, are greatly surpassed by the less concentrated carbon present in the much more numerous ordinary sediments and sedimentary rocks.  

While the very early Earth may have had an atmosphere as much as 250 times as thick as it is today, most of it carbon dioxide, similar to the atmospheric compositions of Venus and Mars, that collision with a Mars-size body early in Earth’s history may have blown off most of that original atmosphere.  Subsequent volcanism replenished some of that atmosphere but one dominated by nitrogen gas although carbon dioxide was still as much as several percent of a now much thinner atmosphere.  Volcanism also brought huge amounts of water vapor to the surface, aided by some water-rich asteroids, which, as soon as the surface had cooled enough, formed the Earth’s oceans.  Billions of years of photosynthesis then reduced the atmospheric carbon dioxide levels from a few percent to a tiny fraction of a percent (~ 0.04% today).  

On the right side of the equation, organic matter and oxygen are being produced.  As mentioned above, much of the organic matter ended up in the sediments of the time, removing carbon from what would otherwise have been carbon dioxide.  For billions of years, the oxygen, as fast as it was produced, ended up oxidizing many other things, especially iron and so did not build up to any appreciable amount in the atmosphere which raises an interesting question about Venus and Mars.  

Mars, notoriously red, is red because of iron oxide (hematite).  Were you able to see the surface of Venus, it would look reddish too.  Where did the oxygen come from that oxidized all that iron?  An early Mars definitely had seas if not oceans of water but did it have photosynthesis?  If it did, there should be some evidence of it in the form of organic matter for which the Mars rovers are looking.  However, it is probably more likely that for both Mars and Venus, the oxygen that oxidized the iron in the surface rocks and sediments came from the photolysis of water vapor in their early atmospheres although that would not preclude life in the early seas of Mars.  Time will tell.  

One last caveat about the net reaction for photosynthesis is that, while generally representative of photosynthesis on land, the equation is not so accurate for photosynthesis in the oceans.  In water, CO2 is most abundant in the ionic form of bicarbonate:  CO2 + H2O  <==>  HCO3 + H+.  If this is substituted into the net equation for photosynthesis, the result is:  

6 HCO3 + 6 H+ <==> C6H12O6 + 6 O2

Now it can be seen that photosynthesis also neutralizes the acidity generated by the dissolving of carbon dioxide into the water.


Photosynthesis is only possible because the Earth has surface oceans into which light can penetrate. There is an analogue to photosynthesis, chemosynthesis, that does not require light but which, instead, uses the thermal and chemical energy of hydrothermal solutions. If written in the same form for the general net equation for photosynthesis, one version of chemosynthesis would look like this:

18 H2S + 6 CO2 + 3 O2 ––> C6H12O6 + 12 H2O + 18 S

Hydrogen sulfide gas (H2S) and carbon dioxide are common components of volcanic gas and hydrothermal solutions although the free oxygen seems a little problematic. Nevertheless, Earth’s deep oceans have many hydrothermal vents, mostly associated with plate boundaries, that host communities of giant tube worms which use chemosynthesis to generate organic matter, forming the base of a food chain independent of solar energy.

Tube Worms

Although the Earth is the only world known to have surface water oceans, many of the bodies in the outer Solar System are known or suspected to have subsurface water oceans. Without access to sunlight to drive photosynthesis, if there is life in such subsurface oceans, it would require some form of chemosynthesis to generate organic matter. It seems unlikely that such worlds would have a version of plate tectonics but there could still be hydrothermal vents that could conceivably host some form of life analogous to the tube worm communities on Earth. Once again, time will tell.

Photoferrotrophs and Iron Bacteria

Another important version of photosynthesis, especially in the Earth’s early oceans, was used by a group of bacteria known as photoferrotrophs who generated organic matter using sunlight to oxidize dissolved iron:

24 Fe++ + 6 CO2 +66 H2O ––> 24 Fe(OH)3 + C6H12O6 + 54 H+

The Oxygenation of the Oceans

The early atmosphere and ocean had no oxygen which meant no colorful hematite (red) or limonite (yellowish-brown) on land (rocks of this age are mostly shades of gray) and no limestone in the ocean.  Starting about 3.8 billion years ago, peaking at ~ 2.5 billion years ago, and continuing to 1.7 billion years, banded iron formations (BIFs) appeared which were marine sediments consisting of alternating layers of sediment from a few microns to a few meters thick.  A layer loaded with hematite (Fe2O4) and a little magnetite (Fe3O4) with an iron content of up to 55% would alternate with a less iron-rich, less colorful layer.  (That formula for magnetite can be a little misleading; it is actually a mix of ferrous and ferric oxide:  FeO•Fe2O3 = Fe3O4)  

What was happening was that the photoferrotrophs were oxidizing the ferrous iron in the seawater and precipitating it as ferric hydroxide, incorporating it in the marine sediments of the time; the ferric hydroxide later becoming hematite.  The reason for the banding was that the thinner layers were caused by seasonal changes in ocean temperatures:

summer (~ 25 °C)
          photoferrotrophs very active ––> produce oxidized iron
          higher temps inhibit precipitation of silica (jasper and chert)
winter (~ 5 °C)
          photoferrotrophs much less active ––> much less oxidized iron production
          lower temps enhance precipitation of grayish silica, less red hematite

The thicker layers are attributed to longer-term climate fluctuations.  The oxygen in the BIFs represents some twenty times as much oxygen as in the atmosphere today and it was only when most of the iron in the oceans had been precipitated out that a trace of free oxygen could appear in the atmosphere ~ 3 billion years ago.   

The hematite in the BIFs is the source of all major iron ore deposits in the world.  Originally marine sediments, these deposits were incorporated into continents as a result of plate tectonic movements, sometimes becoming mountain ranges of high grade iron ore (hematite) such as in the Lake Superior Ranges of Minnesota, Wisconsin, and Michigan.  

Around 3.1 billion years ago the first microbial aerobes appeared (they used the free oxygen as what was to become a very important energy source).  Small but significant amounts of free oxygen began to form regionally and sporadically, leading to the Great Oxidation Event of 2.7 to 2.6 billion years ago.  Although anoxic (no oxygen) parts of the ocean persisted, mostly at deeper depths, by 2.316 billion years ago atmospheric O2 had reached ~ 2 ppm which was enough to begin to form a weak ozone layer in the stratosphere which was to prove crucial for the further development of life both in the oceans and, eventually, on land.  

Absent an ozone layer, ultraviolet levels at the Earth’s surface would be lethal to most forms of life.  Fortunately for marine life, a few feet of water blocks the UV but with even a weakly protective ozone layer, marine photosynthesizers could rise closer to the surface where there was more sunlight, leading to a dramatic rise in free oxygen production.  Cyanobacteria became more important than the photoferrotrophs and within ~ 10 million years the atmospheric O2 had risen to ~ 0.2% (20,000 ppm), further bolstering the stratospheric ozone layer.  This is a great example of a positive feedback.  

Of course there were tradeoffs.  Oxygen was being introduced into a world that had had none before.  That oxygen, a byproduct of photosynthesis, was poisonous to most of the microbes of the time, including the photoferrotrophs.  Cyanobacteria, on the other hand, were able to tolerate and flourish in the oxygenating water which is why they supplanted the photoferrotrophs who were hit by the double whammy of poisonous oxygen and plummeting dissolved iron.  

The oxygen was also changing the atmosphere and even the global climate.  An important component of the atmosphere up until the Great Oxidation Event was methane (CH4), a greenhouse gas considerably more potent than carbon dioxide.  Even with a dimmer sun 3 billion years ago, methane kept the Earth considerably warmer than it otherwise would have been.  Without oxygen, methane is stable in the atmosphere.  With oxygen, methane will oxidize into the much less potent carbon dioxide.  Adding to the loss of atmospheric methane was the precipitation of dissolved nickel in the oceans.  As the dissolved ferrous iron oxidized to less soluble ferric iron and precipitated, it took dissolved nickel with it.  Marine methanogens (methane-producing microbes) need nickel.  As the dissolved nickel levels dropped, the methanogens produced less methane.  

In today’s well-oxygenated atmosphere, methane has a half-life of about nine years and is present at a level of about 1.8 ppm (compared to ~ 410 ppm for CO2).  Although methane is about 25 times better at trapping infrared radiation, 1.8 ppm methane would be equivalent to only 45 ppm CO2 which is why, in today’s atmosphere, CO2 is an order of magnitude more important greenhouse gas than methane.   That was not true of Earth’s atmosphere 3 billion years ago which had much greater quantities of both, especially methane.  

An interesting question is just how much oxygen can be dissolved in water.  Many factors play a role in answering that question including: the temperature of the water, how many ions are already dissolved in the water, and the (partial) pressure of the oxygen in the air above the water.  To get an idea of how much that is, consider current conditions.  The dissolved oxygen (DO) of seawater today is about 7-8 mg/L.  Cold well-oxygenated river water suitable for trout can contain up to 10 mg/L.  The much lower atmospheric oxygen levels of billions of years ago suggest that the DO of seawater of that time was correspondingly much lower but still sufficient to oxidize and precipitate iron and encourage the development of marine aerobic microbes.  

Another interesting question is whether water or air contains more free oxygen.  Using today’s values, the oxygen content of the atmosphere is about 21% which seems to be a lot more than the paltry < 10 mg/L of water.  However, the density of air (~ 1250 mg/L) is much less than that of water (~ 1,000,000 mg/L).  21% of 1250 mg/L ≈ 260 mg/L oxygen in air which does mean that the air has substantially more oxygen in it per volume than the water can hold.  

Several billion years ago and for a considerable time afterward, atmospheric oxygen levels fluctuated between 0.25% to 2.5%, typically ~ 1%.  Assuming an air density of 1250 mg/L, this would be something on the order of 12 mg/L oxygen in the atmosphere, probably not all that much more than the DO content of the seawater.  The Great Oxidation Event boosted atmospheric oxygen from essentially nothing to only some 1% of the atmosphere but that was enough to trigger consequential changes in both the oceans and the atmosphere.  

These variations in atmospheric oxygen were correlated with massive swings in the rate of organic matter deposition in the oceans.  When the rate of organic matter deposition in marine sediments exceeded the oxidative weathering of the land, there would be a net release of O2 to the atmosphere.  When continents collided, producing mountains with exposed organics, the oxidative weathering of the land would increase, decreasing the atmospheric level of O2.  Atmospheric levels of CO2 were inversely proportional to atmospheric levels of O2.  The combination promoted the fluctuations of both the gasses and organic deposition.  

However, oxygen could and did later accumulate in the atmosphere at levels far beyond that of the water, making the much later transition of water-breathing marine animals to air-breathing land animals easier.  Microbial aerobes reverse photosynthesis in respiration to generate much more energy than would be possible without oxygen in other energy-producing reactions, eventually leading to multicellular nucleated forms of life (eukaryotes).  Multi-celled animals require > 4% atmospheric oxygen which level was reached about 0.8 billion years ago.  

A Global Ice Age

Lowering levels of both atmospheric methane and carbon dioxide resulted in global cooling, culminating in Snowball Earth ~ 635 million years ago.  The name can be a bit misleading as the Earth’s oceans certainly did not freeze completely solid although surface sea ice probably reached as far towards the equator as the tropics.  An alternative name has been suggested for this time, Slushball Earth.  Atmospheric CO2 levels probably dropped to around 150 ppm during this greatest of all ice ages.  For comparison, the CO2 level of the recent ice age was about 180 ppm.  

Another supercontinent, Rodinia, a predecessor to the last supercontinent Pangaea, formed 1 billion years ago in the southern hemisphere.  By this time, the amount of continental crust on the Earth had pretty much reached how much continental crust the Earth has now.  

More continental crust, especially newly-formed mountainous supercontinents, would lead to more CO2 removal from the atmosphere by weathering which would promote global cooling.  Atmospheric CO2 eventually dropped to as low as ~150 ppm, < half the current level of ~ 420 ppm.  More continental crust would lead to more clouds due to convection from land heating and would raise Earth’s albedo (reflectivity), also promoting global cooling.  Finally, at the time of Rodinia, the sun was 6% less bright than it is today.  

The subsequent breakup of Rodinia some 650 million years ago is associated with tropical flood basalts which have significant trace amounts of phosphorus.  Rapid tropical weathering of the flood basalts released phosphate into the oceans which promoted rapid, extensive algae growth leading to drawdown of atmospheric CO2.  Dead algae sequestered carbon on the sea floor, promoting anoxic sea floor conditions ––> global cooling.  The oceanic ice cover prevented air exchange between the atmosphere and oceans resulting in widespread anoxia (no oxygen) in the oceans.  

As a consequence of all this, major global cooling ensued.  Global glaciation extended even to the tropics although ocean currents kept some tropical oceans partly open even while tropical continents were ice-covered.  Thus began the Big Chill (Varangian glaciation) ~ 635 million years ago.  

However, the situation eventually turned around. Marine organisms, many of which were photosynthesizers, took a big hit which greatly reduced the removal of CO2 from the atmosphere (more feedbacks). The flood basalts associated with the breakup of Rodinia also released copious amounts of CO2, so much that the Earth swung from a global ice age to a global hothouse for a while. The weathering of tropical flood basalts released nutrients which ended up in the oceans, providing a big boost to O2-producing photosynthesizers after the end of the Big Chill, reversing their decline which meant that oxygen kept accumulating in the oceans and then into the atmosphere.

End of the Precambrian and the Beginning of Limestone

Centuries ago when geologists were trying to make a geologic time scale, paleontologists noticed that rocks of a certain age and older seemed to be devoid of fossils after which came rocks full of fossils, especially fossils of shelled organisms.  Such a change seemed to mark a great divide in the geologic/biologic history of the Earth which was then divided into two Eons, the Phanerozoic Eon (Greek for ‘visible’ life) and the Cryptozoic Eon (Greek for ‘hidden’ life).  The Cryptozoic Eon, now replaced by the Hadean, Archaeozoic, and Proterozoic Eons, was eventually determined to end about 570 million years ago.  

Eons are subdivided into Eras which are subdivided into Periods.  The first Era of the Phanerozoic Eon is the Paleozoic and its first Period is the Cambrian.  What was once the Cryptozoic Eon is informally referred to as the Precambrian which is a testament to how important this great divide, between visible fossils and hidden (microscopic) fossils is.  It is not an equal divide as the Precambrian covers the first 4 billion years of Earth’s geologic history and the Phanerozoic Eon barely a half billion years.  So what happened to create this great fossil divide?  

As described earlier, the Earth’s oceans have undergone big changes in their chemistry such as their oxygenation.  One consequence of that oxygenation is that the oceans were slowly becoming less acidic.  By 570 million years ago, the ocean acidity had reduced to the point that the seawater became slightly alkaline and that made possible the beginning of the precipitation of CaCO3 as limestone (the rock) or calcite/aragonite (the minerals).  That is when certain forms of life were able to precipitate protective shells of calcite/aragonite and many marine sedimentary rocks of that time and since became full of very visible fossils.  

Precipitating CaCO3 from seawater removes dissolved CO2 from the water which is replaced by CO2 from the air, reducing the CO2 content of the atmosphere.  The formation of limestone became and remains an important sink for atmospheric CO2.  Limestone in contact with seawater acts to counter any acidification of the seawater.  Should the seawater, for whatever reason, become more acidic, some limestone will dissolve and neutralize (buffer) it, keeping the pH nearly constant, slightly alkaline.  

There are limits to this buffering action, for one, the limestone has to be in contact with the water.  If it is buried, it can’t react with (buffer) the water.  The oceans can absorb a lot of atmospheric CO2 but so much has gotten into the air today that we now have to worry about acidification of the oceans.  

Although limestone can precipitate out of seawater under the right conditions without the intermediation of some form of life, most of the world’s limestone deposits are of biogenic (some life form did it) origin.  More generally, you might have noticed that most of the big chemical changes in the ocean relied on some form of life. 

The Last Half Billion Years

A strengthening stratospheric ozone layer reduced surface UV levels enough that photosynthesizers could thrive all the way to the surface of the oceans.  Fueled by the energy of photosynthesis and the oxygen it generated, ever more sophisticated marine life appeared, especially in the Ediacaran Period (635-539 million years ago) which was the last period of the Precambrian.  The Cambrian Period (539-485 million years ago) saw an explosion of diversity of life and the appearance of all the phyla of life.  

Although bacteria and fungi had reached the land considerably earlier, by 470 million years ago the first land plants appeared, adding land photosynthesis to marine photosynthesis.   At first, land photosynthesis was not very important as life on land was essentially two dimensional unlike the three dimensional oceans.  The first land animals were ancient millipedes beginning ~ 420 million years ago.  400 million years ago the first vascular plants developed (Lycophytes).  The significance of vascular plants is that they use lignin to give their cells enough structural strength to allow them to grow tall – land photosynthesis became three dimensional.  

Only 20 million years later (late Devonian Period) the first forests appeared.  Now land photosynthesis became a potent oxygen-generating partner to marine photosynthesis.  Atmospheric oxygen levels soared by more than an order of magnitude, peaking at as much as 30% of the atmosphere.  All of that atmospheric oxygen tended to keep the oceans more deeply and stably oxygenated although anoxic areas can still exist even

Atmospheric oxygen levels fluctuated considerably during the Phanerozoic Eon but never dipped below double digits.  Carbon dioxide levels tended to vary inversely with the changes in oxygen levels with consequent variations in global climate.  Such changes sometimes greatly affected both marine and land life but that fascinating history is beyond the scope of this series.  

Nevertheless, it can be seen that without life, Earth’s oceans would be very different than they are now and that extends to the atmosphere too.  It is amazing the extent to which life has shaped the near-surface of the Earth.  We now worry about the impact that the human use of fossil fuels will have on global climate.  That is a valid concern but life has been affecting the Earth in many global ways ever since it first appeared, sometimes to the detriment of some organisms which could now include us.  Life and plate tectonics are responsible for the ocean’s, continents, and atmosphere as they are today.  Earth is the only body we know of that has a surface-water ocean, life, and plate tectonics.  Are we truly unique?  

Water in the Universe | Contents | Intro | Part 1 | Part 2 | Part 3

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