Skip to main content
Geosciences LibreTexts

3.6: Why the Sea is Bitter

  • Page ID
  • \( \newcommand{\vecs}[1]{\overset { \scriptstyle \rightharpoonup} {\mathbf{#1}} } \)

    \( \newcommand{\vecd}[1]{\overset{-\!-\!\rightharpoonup}{\vphantom{a}\smash {#1}}} \)

    \( \newcommand{\id}{\mathrm{id}}\) \( \newcommand{\Span}{\mathrm{span}}\)

    ( \newcommand{\kernel}{\mathrm{null}\,}\) \( \newcommand{\range}{\mathrm{range}\,}\)

    \( \newcommand{\RealPart}{\mathrm{Re}}\) \( \newcommand{\ImaginaryPart}{\mathrm{Im}}\)

    \( \newcommand{\Argument}{\mathrm{Arg}}\) \( \newcommand{\norm}[1]{\| #1 \|}\)

    \( \newcommand{\inner}[2]{\langle #1, #2 \rangle}\)

    \( \newcommand{\Span}{\mathrm{span}}\)

    \( \newcommand{\id}{\mathrm{id}}\)

    \( \newcommand{\Span}{\mathrm{span}}\)

    \( \newcommand{\kernel}{\mathrm{null}\,}\)

    \( \newcommand{\range}{\mathrm{range}\,}\)

    \( \newcommand{\RealPart}{\mathrm{Re}}\)

    \( \newcommand{\ImaginaryPart}{\mathrm{Im}}\)

    \( \newcommand{\Argument}{\mathrm{Arg}}\)

    \( \newcommand{\norm}[1]{\| #1 \|}\)

    \( \newcommand{\inner}[2]{\langle #1, #2 \rangle}\)

    \( \newcommand{\Span}{\mathrm{span}}\) \( \newcommand{\AA}{\unicode[.8,0]{x212B}}\)

    \( \newcommand{\vectorA}[1]{\vec{#1}}      % arrow\)

    \( \newcommand{\vectorAt}[1]{\vec{\text{#1}}}      % arrow\)

    \( \newcommand{\vectorB}[1]{\overset { \scriptstyle \rightharpoonup} {\mathbf{#1}} } \)

    \( \newcommand{\vectorC}[1]{\textbf{#1}} \)

    \( \newcommand{\vectorD}[1]{\overrightarrow{#1}} \)

    \( \newcommand{\vectorDt}[1]{\overrightarrow{\text{#1}}} \)

    \( \newcommand{\vectE}[1]{\overset{-\!-\!\rightharpoonup}{\vphantom{a}\smash{\mathbf {#1}}}} \)

    \( \newcommand{\vecs}[1]{\overset { \scriptstyle \rightharpoonup} {\mathbf{#1}} } \)

    \( \newcommand{\vecd}[1]{\overset{-\!-\!\rightharpoonup}{\vphantom{a}\smash {#1}}} \)

    The primeval ocean... must have been only faintly salt. But the falling rains
    were the symbol of the dissolution of the continents. From the moment the
    rains began to fall, the land began to be worn away and carried to the sea. It is
    an endless, inexorable process that has never stopped--the dissolving of the
    rocks, the leaching out of their contained minerals, the carrying of the rock
    fragments and dissolved minerals to the ocean. And over the eons of time,
    the sea has grown ever more bitter with the salt of the continents.
    --Rachel Carson, The Sea Around Us

    Rachel Carson provided this poetic statement about the evolution of seawater chemistry over time in her book, first written in 1950. It is an interesting statement about the prevailing thought of the time--that ocean salinity evolved slowly and progressively and that rivers were the only source of salt. Both of these ideas are incorrect in light of more recent scientific investigations. We will highlight these issues in this section of Lesson 3 as they lead us to some interesting concepts and calculations. In defense of Rachel Carson, a native Pennsylvanian and the forebearer to the modern environmental movement, her failure to correctly describe the system is a function of the huge scientific advances that have been made in the geosciences, beginning in the early 1960s. The concept of plate tectonics was in its nascency in the 1950s and was not widely accepted by the geoscience community until definitive evidence in support of it in the 1970s. Rachel Carson had no idea that the mid-ocean ridge hydrothermal system existed because no one observed a submarine hot spring until 1977 (Galapagos at 2500 meters depth). We can forgive her her ignorance, right?

    As the Earth cooled over 4 billion years ago and water began to condense in the oceans (it probably originally condensed and fell as rain), that first water probably did not have a very high salt content. This water was outgassed along with other volatiles from the Earth's interior (mantle) and possibly also accumulated from cometary impacts. Some geologic evidence suggests that the bulk of the oceans were already formed by about 3.8 billion years ago(Ga). But very quickly various chemical ions must have dissolved in water as it bathed or passed over freshly formed igneous rocks (probably mostly basaltic in composition initially), and began to be washed into the pools that eventually grew into the oceans. Water is a remarkable substance (see write up on the "Physical Properties of Water”). The "polar" water molecule allows it to interact with and isolate charged chemical ions (elements with unfilled electron shells that are dissolved in the water), such as Na+ and Cl- in solution. These chemical ions, when dissolved in water, are commonly called "salts." It perhaps took hundreds of millions of years for the ocean to accumulate significant amounts of these salts as the result of the operation of the global hydrologic cycle. In a nutshell: ever since atmospheric water vapor could condense into rain, water has fallen onto the land surface and drained eventually, through rivers and groundwater, into the oceans. The water that falls on the land dissolves minute amounts of salt (called "rock weathering") during its passage over the land. It carries that salt to the ocean. In the meantime, heat from the sun provides the energy to cause more evaporation. The evaporated water then condenses and falls again as rain on land (essentially replacing water that flowed into the sea), and thus continues the cycle. Seawater salts essentially cannot evaporate and, therefore, when the ocean water evaporates, salt remains behind.

    The ocean, of course, is constantly losing pure fresh water through evaporation and receiving small amounts of dissolved salt from the river and groundwater coming in. While it would seem that the oceans should be getting saltier over time, the record of sedimentary deposits, called "evaporites” (see the experiment below, also discussed in class), from ancient oceans and the continuity of life as evidenced in the fossil record, indicate that this does not occur. Interestingly, the salinity of seawater appears to have remained relatively constant (but we will see about this!) at about 3.5 % (35 ppt by weight or 35 grams of salt dissolved in 965 grams of fresh water), at least over the past 500 million years or so, but possibly even since sometime earlier (e.g., probably since about 2 billion years ago or more), after formation of the oceans. Thus various chemical, biological, and tectonic processes must act to remove salts from seawater in the amounts necessary to keep the ocean salt content from varying much.

    What Determines the Composition of Seawater Salt?

    Although much of the ocean's salt has ultimately come from the weathering of continental rocks, there are other important sources and chemical exchanges between seawater and the Earth. The chemical composition of river water and salty inland lakes is, surprisingly, not very similar to that of the oceans. Average river water contains mostly calcium and bicarbonate ions, while seawater consists largely of sodium and chloride; in fact, only five chemical elements make up more than 99% of salt dissolved in seawater. Why does the chemistry of seawater differ from that of the runoff from the continents? This difference must reflect the other sources of "salt" to the oceans, as well as the dominant processes that remove certain salts by "precipitation." (We will explore this for various elements in Lesson 3, Activity 2).

    For example, the upper-mantle layer of the Earth contains huge reserves of the elements found in seawater. Deep sea vents, rift vents, and volcanoes, which expel heat and fluids from the Earth's interior, supply large amounts of certain salts through outgassing. In the case of Na (sodium) and Cl (chloride), rock-weathering supplies most of the sodium ions, whereas outgassing of volatiles supplies chlorine. Na and Cl are so strongly enriched in seawater though because they are not used by organisms and do not precipitate out very easily except under highly evaporative conditions in salt ponds or isolated basins where they precipitate as evaporite minerals. These kinds of salts are said to have long "residence times" in seawater compared to other elements (e.g. nutrients such as nitrogen and phosphorus, silica, bicarbonate, and certain others are cycled very rapidly). Interaction (chemical exchange) of seawater and hot basalts at mid-ocean ridges (remember the "hydrothermal circulation" discussed in Lesson 2?) supplies a significant amount of Ca (calcium) to the ocean, while leaving behind an equivalent amount of seawater Mg (magnesium) in the resulting altered basalts. This process constantly modifies the amounts of Ca and Mg in seawater. In addition, seawater contains a lower relative proportion of dissolved silica (SiO2), Ca, and bicarbonate (HCO3-) than river water does. This is because certain groups of marine plants and animals remove these components very rapidly to form their hard parts (skeletal material such as shells or "tests").

    Keep in mind that during evaporation or dilution by fresh water, the salt content (salinity) increases or decreases respectively. However, the ratio of each salt component in seawater to another (e.g Na/Cl or Ca/Mg) remains constant as long as the salinity does not increase to the extent that mineral precipitation begins. This is called the “Principle of Constant Proportion” and is useful for understanding external inputs or outputs of various elements that might change the ratio of one element to another.

    An Experiment:

    Here is a simple experiment that illustrates the process of evaporation and precipitation of salt from seawater that might reinforce this concept for students (we commonly see college students who don’t think about evaporation leaving the salt behind as a mechanism for increasing saltiness). This experiment will only work in a reasonable time during a warm, dry period (in your “not so fair” state of Pennsylvania, these are few and far between). You should use any sort of clear glass jar and fill it to some line that you have marked on the side of the vessel. First, fill the jar with pure (distilled, not tap) water to the line. When the water evaporates completely (look for any residue), there should not be any. Now mix a mild salt-water solution (use common table salt to 3.5 g in about 100 ml of pure water) or use seawater if available, and again fill the jar to the line with the solution. When that evaporates, again look for residue (if you could scrape it all out and weigh it, the weight of salts left behind as precipitates should be 3.5 g). Of course, the water evaporates into the air and the salt remains behind. If seawater (even artificial aquarium sea salt) is used, one might even observe salts of different minerals precipitating out as the water level in the glass drops. Minerals of different salt components have different saturation points (lower solubility), such that calcite (calcium carbonate, CaCO3) precipitates first, followed by gypsum (CaSO4), halite (NaCl), sylvite (KCl), and finally some small amounts of various magnesium sulfate salts, etc.

    This page titled 3.6: Why the Sea is Bitter is shared under a CC BY-NC-SA 4.0 license and was authored, remixed, and/or curated by Eliza Richardson (John A. Dutton: e-Education Institute) via source content that was edited to the style and standards of the LibreTexts platform; a detailed edit history is available upon request.

    • Was this article helpful?