Earth 540: Essentials of Oceanography for Educators
Published on Earth 540: Essentials of Oceanography for Educators (https://www.e-education.psu.edu/earth540)

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Lesson 3

About Lesson 3

In this lesson, we will return to the amazing properties of water and how those properties influence Earth surface processes.  We will explore the structure of water molecules and of water itself (yes, water has "structure"), and relate these concepts to important processes, such as the buffering of Earth surface temperature variations and weathering (breakdown) of rocks on land. In addition, we will investigate the composition of sea salt and its origin, as well as some of the exchanges of salt and water between various reservoirs (storage bins) at the Earth's surface. Although based on chemistry (scary or "boring" to some students), this topic has some fascinating elements that can really engage students, including thinking about the importance of salt in human history, debates about the age of the Earth from an ocean perspective, and the potential for extracting metallic riches from the ocean.  The topical coverage for Lesson 3 is as follows:

•    Ice, water and vapor
•    Latent heat,  heat capacity, and other unusual properties •    Water as the universal solvent
•    Why the sea is bitter
•    It’s all about cycles: from vapor to rain to snow to rivers and the ocean (the "Hydrologic Cycle").
•    Geochemical residence time
•    Salt as a commodity and the age of the oceans

What will we learn in Lesson 3?

By the end of Lesson 3, you should be able to: 

•    List the important characteristics of ice, water and vapor, and cite why they are important
•    Describe latent heat and heat capacity and apply concepts to understanding climate processes
•    Explain why water is the universal solvent
•    Speak knowledgeably about the origin of salts in seawater
•    List the factors that determine seawater salinity and the elemental composition of seawater salt
•    Explain the hydrologic cycle and its linkages to water properties
•    Explain the concept geochemical residence time in the oceans and what insights it provides.

What is due for Lesson 3?

The table below provides an overview of the requirements for Lesson 3. For  details, please see the Course Schedule

REQUIREMENT

LOCATION

SUBMITTED FOR GRADING?

Activity 1: Science Fiction Blog: If Water Behaved Differently

 

page 8 Yes -  Your discussion participation counts toward your overall class participation grade. This discussion will take place here, see Activity 1, rather than on Angel.

Activity 2: The Residence Time of Salt in the Ocean

 

page 9 Yes - Your discussion  participation counts toward your overall class participation grade. This discussion will take place here, see Activity 2, rather than on Angel.

Activity 3: End of Unit Quiz (ANGEL)

 

 

ANGEL

Yes. We will activate a set of questions to "test" your understanding of the material a few days before the end of this lesson.  The exam should take  no more than one hour.

Questions?

If you have any questions, please post them to our Questions? discussion forum (not e-mail), located under the Communicate tab in ANGEL. I will check that discussion forum daily to respond. While you are there, feel free to post your own responses if you, too, are able to help out a classmate.

Ice, Water, & Vapor

Introduction

We will first examine water as a molecule, and then explore the implications of water's molecular structure for its physical behavior and its importance as a "universal solvent."  This first section examines the phases of water.

Reading assignment

Read Chapter 6 (pp. 151-182) in Life's Matrix: A Biography of Water by Philip Ball, then read the material in this section.  Ball, an editor for Nature, has an elegant way of framing the "weirdness" of the water molecule that highlights water's unusual properties.  This chapter is a nice treatment to accompany our "drier" outline below. Click here for a PDF file of the reading [1]

The Water Molecule

A molecule of water is composed of two atoms of hydrogen and one atom of oxygen.  Now, the one and only electron ring around each hydrogen atom has only one electron.  The negative charge of the electron is balanced by the positive charge of the one proton in the nucleus.  Protons have mass, electrons do not.  One electron and one proton: hydrogen has an atomic number of one.  In the hydrogen nucleus is also one neutron; no charge but the weight of one proton.  One proton, one neutron: hydrogen has an atomic weight of two.  The electron ring of hydrogen would like to possess two electrons to create a stable configuration.  Oxygen, on the other hand, has an inner electron ring with two electrons, which is cool because that is a stable configuration.  The outer ring, on the other hand, has six electrons but it would like to have two more because in the second electron ring, eight electrons make the stable configuration.  To balance the negative charge of eight (2+6) electrons, the oxygen nucleus has eight protons.  Eight protons and eight electrons:  oxygen has an atomic number of eight.  The eight protons in the nucleus are matched by eight neutrons.  Eight protons, eight neutrons: oxygen has an atomic weight of 16.  Hydrogen and oxygen would like to have stable electron configurations but do not as individual atoms.  They can get out of this predicament if they agree to share electrons (a sort of an energy "treaty"?).  So, oxygen shares one of its outer electrons with each of two hydrogen atoms, and each hydrogen atom shares its one and only electron with oxygen.  This is called a covalent bond.  Each hydrogen atom thinks it has two electrons, and the oxygen atom thinks that it has eight outer electrons.  Everybody's happy, no?

Figure. 1:
                        O
                       /    \
                     H     H

However, the two hydrogen atoms are both on the same side of the oxygen atom with an angle of about 105 degrees between them (see Fig. 1) so that the positively charged nuclei of the hydrogen atoms are left slightly exposed, so to speak, leaving that end of the water molecule with a weak positive charge. Meanwhile on the other side of the molecule, the electrons of the oxygen atom give that end of the molecule a weak negative change. For this reason, a water molecule is called a "dipolar" molecule.  Water is an example of a polar solvent, capable of dissolving most other compounds. In solution, the weak positively charged side of one water molecule will be attracted to the weak negatively charged side of another water molecule and the two molecules will be held together by a weak "hydrogen bond," and so on.  At the temperature range of seawater, the weak hydrogen bonds are constantly being broken and re-formed. This gives water some structure, but allows the molecules to slide over each other easily, making it a liquid.

Water Properties

A calorie is the amount of heat it takes to raise the temperature of 1 g of pure water 1 degree C at sea level.  Therefore, it would take 100 calories to heat water from 0˚, the freezing point of water, to 100˚ C, the boiling point.  However, it would take 540 calories to convert that 1 g of water at 100˚ C to 1 g of water vapor still at 100˚ C.  This is called the heat of vaporization.  You would have to remove 80 calories from 1 g of water at the freezing point, 0˚ C, to convert it to 1 g of ice at 0˚ C.  This is called the heat of fusion.

Water does not give up or take up heat very easily. Therefore, it is said to have a high heat capacity. In Pennsylvania, it is common to have a difference of 20˚ C between day and night temperatures. During the same time frame, the temperature of lake water would hardly change at all.  

Water flows easily.  It is said to have a low "viscosity." Compare this with motor oil or honey that each have relatively high viscosities (or to the upper mantle that has an even higher viscosity as discussed in class during our "plate tectonics" lectures).  If you can't get the honey to flow out of the jar and onto your toast in the morning, you put it in the microwave and "nuke" it, then it flows easily, i.e. increasing the temperature lowers the viscosity. Similarly, warm water is less viscous than cold water.  

Pure water has a density of 1.0 g/cm3 at 4˚ C.  As you increase or decrease the temperature from 4˚ C, the density decreases.  In fact, if you measure the temperature of the deep water in large, temperate-latitude lakes that freeze over in the winter, you will find that the temperature is 4˚ C; that is because fresh water is at its maximum density at that temperature, and as surface waters cool off in the fall and early winter, the lakes overturn and fill up with 4˚ C water.  As you add dissolved solids to pure water to increase the salinity, the density increases.  The density of average seawater with a salinity of 35 o/oo (35 g/kg) and at 4˚ C is 1.028 g/cm3   As you add salts to seawater, you also change some other properties.  Increasing salinity increases the boiling point and decreases the freezing point.  Normal seawater freezes at -2˚ C, 2˚ C colder than pure water. Increasing salinity also lowers the temperature of maximum density.  

When Water Freezes, It Expands

Fig. 2 A rough sketch of water molecules in ice crystal (hexagonal form)
                     H2O
                   /          \
             H2O           H2O
              |                    |
            H2O            H2O
                   \            /
                      H2O

When water is a liquid, the water molecules are packed relatively close together, but can slide past each other and move around freely (that makes it a liquid).  When water freezes, however, bonds are formed that lock the molecules in place in a regular (hexagonal) pattern (Fig. 2).  For nearly every known chemical compound, the molecules are held closer together (bonded) in the solid state (e.g., in mineral form or ice) than in the liquid state.  Water, however, is unique in that it bonds in such a way that the molecules are held farther apart in the solid form (ice) than in the liquid form.  Water expands when it freezes making it less dense than the water from which it freezes.  In fact, its volume is a little over 9% greater (or density ca. 9% lower) than in the liquid state.  For this reason, ice floats on the water (like an ice cube in a glass of water). This latter property is very important for organisms in the oceans and fresh-water lakes.  For example, fish in a pond survive the winter because ice forms on top of the pond (it floats) and effectively insulates (does not conduct heat from the pond to the atmosphere as efficiently) the rest of the pond below, preventing it from freezing from top to bottom (or bottom to top).  If water did not expand when freezing, then it would be denser than liquid water when it froze; therefore, it would sink and fill lakes or the ocean from bottom to top.  Once the oceans filled with ice, life there would be impossible.  We are all aware that the expansion of liquid water to ice exerts a tremendous force.  If you have ever put a full container of water with a tight-fitting lid, or a can of soda in the freezer, you may have experienced this.   Ten cups of water will to turn into 11 cups of ice when it freezes.  The force of the crystallization of ice is capable of bursting water pipes and causing cracks in rocks to expand, thus accelerating the erosion of mountains!

The diagram below (Fig. 3) illustrates some changes in the freezing point and the density of pure water, and changes imposed by salt addition.  Note that the density of pure water is at a maximum at 4°C, whereas density continues to increase as temperature decreases when the salt content is 35 o/oo (parts per thousand)--near the average salinity of seawater.  Also, the freezing point of water is depressed when salt is added (that's why we put salt on icy roads and sidewalks!). We will explore the implications of this later in the course when we consider the circulation of the oceans and the production of "deep" water. Water is only slightly compressible as demonstrated by the small, but measurable, change in density as pressure increases.  Pressure is measured here in bars (not the kind you might be savoring about now), where 1 atmosphere (the pressure at sea level) equals = 1.01325 bar = 101.3 kPa.  Thus, the maximum pressure of 4000 bars on the graph is equivalent to the pressure at about 4000 meters depth in the ocean (the average depth of the ocean is about 3800 meters).  Pressure increases at the rate of one atmosphere every 10 meters (those of you who dive know this).  Air-filled volumes, such as human body cavities, cannot withstand the water pressures at depth (this includes most things except for the specially strengthened hulls of submersibles and submarines), but if we could fill our lungs with water and somehow still respire, like fish, we could do some interesting diving.  Even at 4000 meters, water has compressed only about 2.3% with respect to the volume it had at the surface (an increase in density of 2.3%, do the calculation given the graphical data below). What was the name of the movie that used that principle? OK, on to the next section...

 

Contact the instructor if you have difficulty viewing this image
Figure 3: The change in seawater density at 35 o/oo with pressure (bars) (top), and density and freezing points of pure water and saltwater (bottom).

 

Salt as a Commodity

We won't dwell too much on salt as a commodity here, but it is an interesting topic for students because salt has played such an important role in human history.  There is an excellent book on the subject by Mark Kurlansky titled Salt: A World History.  It has extensive information on the uses of salt through the ages, what makes good salt, trade routes, the historical monetary value of salt (Roman soldiers were paid in salt!), the role salt has played in wars, and so on. 

Although salt is mined from strata on land (this salt represents the remnants of ancient oceans that have evaporated and left their salts behind), much is now produced from the evaporation of seawater in carefully monitored salt pans. In the United States, salt mines can be found in New York state (around the Finger Lakes) and New Mexico (near Albuquerque). There is a saltworks (evaporative) in San Francisco Bay (see below). These are the Cargill saltworks near Hayward, California where salt production goes hand-in-hand with wetlands preservation. The reddish ponds are the most highly evaporated--red color is a "bloom" of a salt-tolerant species of photosynthetic dinoflagellates. Most of the active saltworks are in lower latitudes such as Baja, California. The oldest known salt mines are found near Krakow, Poland.

It is intriguing that we can't drink seawater as a means to get our salt (see later section on salt in seawater), but we are totally dependent on this source for the salt our bodies do (and don't) need, whether from ancient oceanic deposits or modern evaporative ponds.  There is a good source of information online about salt [2] (follow this link).

 

Source: USGS

 

Latent Heat & Heat Capacity

These are the relevant physical properties of water and their significance--all of which are shaped by the hydrogen bonding between polar water molecules:

  • Heat capacity (highest of all solids and liquids except NH3)
  • Latent heat of fusion (highest except NH3)
  • Latent heat of evaporation (highest of all substances)
  • Thermal expansion (in the first section we showed that the temperature of maximum density decreases with increasing salinity)     
  • Conduction of heat (highest of all liquids)   
  • Surface tension (highest of all liquids)
  • Dissolving power (the "universal solvent" dissolves more substances in greater quantities than any other liquid)
  • Transparency (large absorption of radiant energy)

Heat capacity and latent heat are key properties that allow water (the oceans in particular) to play a major role in "regulating" Earth's climate. Water absorbs solar energy and releases it slowly; thus, larger bodies of water do not change temperature rapidly. Likewise, the high latent heat of vaporization (see below), indicates that when water vapor (derived from evaporation of water at the ocean's surface driven by solar energy receipt at low latitudes) condenses into liquid droplets at high elevations or high latitude, the latent heat is released into the environment. In Lesson 4, we will examine this role in more detail, and we have already alluded to the fact that large lakes can help buffer temperature changes.

Heat Capacity
A direct result of the hydrogen bond in water is the high heat capacity of water. As noted, a calorie is the amount of heat required to raise the temperature of 1 g of water 1 °C. The heat capacity of water compared to that of most other substances is great.

Latent Heats of Melting and Vaporization (refer to figure below)
Closely related to water's unusually high heat capacity are its high latent heat of melting and latent heat of vaporization. A solid converts to a liquid at a temperature called its freezing point and a liquid is changed to a gas at a temperature defined as its boiling point.
When changing the state of any substance, there may be no increase of temperature at that point where a change of state occurs even though heat is continuously being added. All the heat energy is being used to break all of the bonds (e.g. between polar water molecules) required to complete the change of state. The heat that is added to 1 g of a substance at the melting point to break the required bonds to complete the change of state from solid to liquid is the latent heat of melting. The heat applied to effect a change of state at the boiling point is the latent heat of vaporization. The amount of heat required to convert 1 g of ice to 1 g of water, 80 Cal, is termed the latent heat of melting, and it is higher for water than for any other commonly occurring substance. The amount of heat required to convert water to vapor, 540 Cal, is termed the latent heat of vaporization. The figure illustrates energy input to a given mass of water that begins as very cold ice and the temperature path that that mass of water takes with continued heat input.  The path from condensation to cooling to ice formation returns energy to the environment.

Surface Tension
Next to mercury, water has the highest surface tension of all commonly occurring liquids. Surface tension is a manifestation of the presence of the hydrogen bond. Those molecules of water that are at the surface are strongly attracted to the molecules of water below them by their hydrogen bonds. If the diameter of the container is decreased, the combination of cohesion, which holds the water molecules together, and the adhesive attraction between the water molecules and the glass container will pull the column of water to great heights. This phenomenon is known as capillarity. This is, in part, what allows plants to stand up--when too much water is lost by evapo-transpiration, they wilt.

Water as the Universal Solvent

As indicated in previous sections, the polar water molecule allows water molecules to form bonds with one another. These bonds are referred to as hydrogen bonds. If we consider sodium chloride (salt), a compound containing ionic bonds, we could demonstrate that simply by placing table salt in water, for example, we can reduce the electrostatic attraction between the sodium and chloride ions by 80 times. As more and more ions of sodium and chlorine are freed by the weakening of the electrostatic attraction that is holding them together, they become surrounded by the polar molecules of water--what is termed "hydration."

 

Figure 1: Salt crystal being dissolved by water, individual ions hydrated.
Source: universe-review.ca)

Water dissolves more substances than any other common liquid by breaking "salts" into component "ions" (e.g.  NaCl  into  Na+ and Cl-) and hydrating those ions to keep them from interacting. Thus, polar water molecules have an attraction for ions (atoms or groups of atoms with a charge), where "cations" are ions with positive charge and "anions" have negative charge. Most elements have high solubilities in water, which means that large concentrations of those elements can build up before the capacity for water molecules to isolate the ions is exceeded.  The point at which Na and Cl, for example, would begin to precipitate a salt in seawater is termed "saturation."  For NaCl (the mineral "halite") this only occurs from present-day seawater when evaporation occurs and the volume of seawater is reduced to about 10% of its original volume.

Seawater is essentially an NaCl solution which averages a concentration of 35 g NaCl/kg water (or 3.5% salt). Na and Cl compose over 85% of the total dissolved solids (salt), but there are other important ions present. The relative abundance of ions in seawater ranks in order:  Cl, Na, SO4, Mg, Ca, K. Together, these ions make up >99% of the dissolved solids in seawater. With only four other elements--HCO3 (bicarbonate), Br, Sr, B, F--we have 99.99% of all dissolved solids. Charges must balance, so the positive charge associated with Na+, Mg+2, Ca+2, K+ equals the negative charge associated with Cl-, SO4-2 (and HCO3-).  We don't think you would want it any other way.  Think about what  the flow of the current would be from the sea to you, sitting on the beach,  if the charges were not balanced--shocking! 

Salinity varies over a range of about 32 to 37 o/oo in the open ocean as Figure 2 (below) illustrates.  Note that areas of highest salinity occur in regions of highest net evaporation, as one might expect.

Figure 2: Mean salinity at ocean surface 1990-94 (POP model).

All other dissolved substances in seawater are at very low concentrations (part per million or billion) (ppm or ppb; 10-6 to 10-9). This Includes important nutrients such as phosphate and nitrate that are cycled by organisms (elements called "bio-limiting") and essential for life.  Many metals have trace concentrations (wanna' get rich? There are about 9 million tons of gold dissolved in seawater, which is about equal to all the gold mined on earth throughout history).

As previously indicated, evaporation of seawater produces a predictable sequence of mineral salts (minerals become saturated at a certain point). After evaporation of a few % of water mass CaCO3 (calcite) precipitates; after evaporation of 81%, CaSO4 (gypsum) is fully precipitated; after evaporation of about 90.5%, NaCl (halite) is fully precipitated; at 96% evaporation, the K and Mg salts (w/ SO4 and Cl) drop out. There is enough salt in the ocean to cover land with a layer 170 m thick. Natural deposits from ancient oceans like this are called "evaporites."

Why the Sea is Bitter

                                              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 forebear 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 our “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. 

It's All About Cycles

Yes, water cycles (the hydrologic cycle) and geochemical cycles (tracing the paths of various elements to and from seawater).  Figure 1 is a conceptualization of the hydrologic cycle (source, USGS). As you are undoubtedly aware, solar energy drives the cycle of evaporation of water from the ocean surface (leaving salt behind), raining out on the continents, and returning to the ocean in rivers (surface runoff).  This water does do work on the land surface.  Eroding solids and dissolving minerals.  Eventually, much of the dissolved material becomes seawater salt.  In a geochemical cycle, plate tectonics causes uplift and exposure of "fresh" rocks, which can be weathered by water (and carbon dioxide). Carbon dioxide is driven out of the Earth's interior during volcanism. This is part of the cycle. So, if rock weathering is such an important process, does ocean chemistry simply reflect the chemistry of rivers, only more concentrated? Can the chemistry of the oceans be related to the inputs of rivers alone? We've already examined why water is a powerful solvent, now let's look at the whole picture.
The ocean is not simply concentrated river water.

Figure 1: Water Cycle diagram

 

Have you heard Bill Nye's rap version of the Hydrologic cycle? "Water Cycle Jump"www.youtube.com/watch [3]

In fact, the ocean has a much different chemical composition from average river water because, like water, salts are cycled as well.  Some salts build up in seawater over time, while other elements are rapidly used, stripped from seawater into organic matter and skeletons of marine organisms or extraction through alteration of near-seafloor basalt in the midocean ridge hydrothermal system.

Rivers supply a large proportion of dissolved solids to the oceans, but river chemistry is very different from seawater.
In rivers the abundance of elements is HCO3, Ca, SO4, SiO2 (1st 4= 80%), then Cl, Na, Mg, K. Rivers are essentially >35% dissolved inorganic carbon (HCO3 or CO3). Compare this with ocean chemistry in the previous section. The difference in chemical compositions between rivers and ocean reflects sedimentation (precipitation) processes and other inputs/exchanges, such as basalt-seawater reactions at midocean ridges. Activity 2 will help you develop an appreciation of geochemical cycling.

 

We can examine the "reactivity" of an element in the oceans by looking a the "residence time" of that element--on to the last section...

Geochemical Residence Time

 

 

Need a break? Listen to this from Tom Lehrer--it'll brighten your day... [4]

 

Every one of these elements is present at some concentration in seawater. As you have seen, some elements have high concentrations (e.g. Na, Cl) whereas others (e.g. Au or Fe, etc.) have very low concentrations.  Very few elements are near saturation (the maximum amount that could be held in seawater of a certain salinity, temperature, and pressure). The chemistry and behavior of elements differ among the various groups (for example redox-sensitive metals vs. alkaline earths).

Residence time is the average time that a substance remains in solution in seawater.  It can be calculated for any element by a standard equation. Note that this is cast in terms of the riverine input only (Activity 2 will ask you why this could be incorrect):

                   Residence Time (yrs.) = Total amount ion in sw (kg) / Input rate (kg/yr)
                   where Input rate = Avg. ion conc. in rivers (kg/km3) x  River discharge (km3/yr)

Let's consider an example:  Here's one for the residence time of water in the ocean. Click here for the ppt file [5] and here for the pdf file [6]

 

 

What is the residence time of all of the salt in seawater? This is an interesting consideration because, in the past, this question was used to argue something about the age of the Earth (How long would it take rivers to deliver all the salt in seawater today?). There are about 5 x 1022 g of dissolved solids in oceans, and rivers bring in about 2.5 x 1015 g of dissolved solids per year. Think about it.  It should only take about 2 x 107 years (20 million years) to bring the oceans to their present salinity, but we know that the oceans are 3.8 billion years old, and if rivers have been providing approximately the same input through time, and if the oceans have maintained approximately the same composition through time, there has to be an output of material that balances the inputs; otherwise, we are wrong about the age of the Earth and its oceans, and that, for various reasons, seems unlikely. This question is still worth exploring with your students because it gets them to think about the dynamic Earth. Interestingly, scientist John Joly (Irish), first tried this calculation around 1901 and obtained an age for the Earth of 90-100 million years. This was too long to suit Irish Archbishop Usher's (1654) supporters who, based on biblical genealogy, believed that the Earth was created in 4004 BC.

You will calculate the residence time for several elements to gain insights into their rate of cycling through the ocean system. Think about what it means to have a long residence time vs. a short residence time.  For example, we like to think of Penn State as a system.  Students come in; students go out. If we simply assume that all students graduate and that the total number of students allowed at the University Park campus does not change, we can calculate the average residence time of a student at the main campus. There are about 42 thousand undergraduate students, with just over 8 thousand students admitted per year. Residence time? Just over 5 years (ouch!). Of course, we have glossed over the details, right?  How many students simply left without their diplomas?  You get it--it's the same for geochemical cycle considerations of residence times. We tend to simplify, thereby missing some of the important stories. 

Figure 2: A full geochemical cycle, conceptualization highlighting the role of tectonic and water cycles.

Teaching and Learning About Seawater Properties I; Activity 1

Activity 1: Science Fiction Blog: if the properties of water were different

Directions

Write a 500-word blog (just one, could be longer if you think you need the space, but...) to add to your previous work.  This one will be a "science fiction" piece, envisaging how the oceans might be different if some fundamental property of water were very different from reality (e.g., what if viscosity were a lot higher or heat capacity a lot lower, or...?). Get creative.  Submit your blog to the group for comment by using the comment/reply box below. Reply briefly to everyone's post.

We can see this being a useful way to engage your students in thinking "outside the box" about the importance of water properties to the nature of the ocean and the continuity of life, and to give them an opportunity for creativity--not just reiterating what they read in textbooks.  But we will leave that evaluation up to you after you have tried this exercise.  We wrote an example below--a bit different vein, but a similar principle.  This is called "Earth without Oceans."  Earlier you read some speculation about what should happen if water ice (solid) were denser than liquid water, like most substances.  Wow, what a difference!

  1. Sample Blog from Mike and Chris:

    Ask yourself this question before reading the speculation below.  See if you can
    come up with other differences between the Earth with and without an ocean.  The
    bottom line is that the Earth would be very much different than it is today. 

    Critical Assumptions:  the answer would differ depending on: 1) whether oceans were
    created and then later destroyed (after evolution of life) or, 2) never existed on Earth at all. 
    A secondary consideration would be the nature and composition of the atmosphere
    with or without an ocean.

        Water is probably Earth's most important resource. Without water, perhaps
    only the simplest viruses could exist. Yes, the Earth would be there still, but life on it
    would have to adapt to very different conditions. In fact, life as we know it may not
    have evolved on earth without the oceans. 
        We must assume that the earth retained some atmosphere (mostly inert gases
    and carbon dioxide).  Without oceans, the earth's surface would be a dusty desert,
    typified by wild daily and yearly swings in temperature. The hydrologic cycle could
    not really function without some  perennial bodies of water. Having no water cycle
    would mean no rain.  Without rain (and plants--see below), the Earth's surface
    would weather much more slowly and surface topography and rock type would be
    much different.
         Earth surface temperatures would vary dramatically from equator to pole
    because energy from the sun would be distributed unequally on the earth. Solar
    energy would normally evaporate water from the oceans in the tropical regions and
    transport this vapor and its contained energy from low to high latitudes in the
    atmosphere; the ocean would normally transport heat energy to the polar regions as
    well.  Without the oceans, therefore, there would be no efficient mechanism for
    redistributing heat. 
        Life, as we understand it, would have a hard time evolving or prospering
    without oceans. All extant plants and animals would die if the oceans were
    suddenly removed from the Earth, although certain bacteria and very simple
    organisms might survive.  With no oceans, the oceanic phytoplankton and the land
    plants that provide our oxygen through photosynthesis would be eliminated.
    Normally respiring organisms would suffocate as the result of the lack of oxygen,
    the amount of carbon dioxide would increase in the atmosphere, and the Earth's
    temperatures in the lower latitudes might increase dramatically, probably enough to
    drive any other surviving animals or plants to extinction.

                Hey, maybe this is what happened on Mars!?

Submitting your work

  • You can begin the discussion activity by posting to the "Comments" space below. To respond to another student's posting, use the "reply" link that follows their posting.

    Don't see the Post new comment area? You need to be logged in to this site first! Do so by using the link at the top of the left-hand menu bar. Once you have logged in, you may need to refresh the page in order to see comment area.

Grading criteria

See the grading rubric [7] for specifics on how this assignment will be graded.

Teaching and Learning About Seawater Properties II; Activity 2

Activity 2: Is the Chemical Composition of Seawater Constant? Do All Elements Behave Alike?

For this activity we will investigate the behavior of some of the elements in seawater in order to understand their sources, sinks, and cycles through the ocean, and to help us decide whether ocean chemistry has been constant through time. For this purpose we will have you research two different chemical elements (partly of your choosing) to obtain the necessary information.

Directions

  1. Choose one element from each of two groups (please chat with one another so that you don't all pick the same elements):
    1. Group 1: Na, Cl, K, Mg, S
    2. Group 2: P, Ca, Si, N
  2. To start, go to the Web site at MBARI [8] for information about chemical elements in seawater. You will find a periodic table, "residence times," and other information on most of the elements.
  3. Find and list the following information for each:
    1. concentration in seawater (use g/kg)
    2. total mass in seawater (in metric tonnes or kg)
    3. amount that could be held in seawater at saturation (and percent saturation)
    4. known sources of the element to seawater (e.g. rivers, hydrothermal, meteoritic or cosmic, etc.)
    5. known sinks (extraction from seawater through hydrothermal alteration of seafloor, sedimentation, etc.)
    6. cycling (is the element bio-limiting, bio-intermediate, or bio-inert?)
  4. Do the following calculations and answer the following questions:
    1. Calculate the oceanic residence time for each of the two elements you have chosen to investigate.
    2. How does the Group 1 element you chose differ from the Group 2 element? Why are they different?
    3. Why might the inferences you draw from the "residence time" calculations be in error?
  5. Post your answers to 3 and 4 in the comment area below.
  6. Read the postings made by other EARTH 540 students.
  7. Respond to at least one other posting by asking for clarification, asking a follow-up question, expanding on what has already been said, etc.

Submitting your work

  • Begin by posting to the comments space below. To respond to another student's posting, use the "reply" link that follows their posting.

    Don't see the Post new comment area? You need to be logged in to this site first! Do so by using the link at the top of the left-hand menu bar. Once you have logged in, you may need to refresh the page in order to see the comment area.

Grading criteria

See the grading rubric [7] for specifics on how this assignment will be graded.

 

Activity 3: End of Unit Quiz and Review Questions

Directions

This will be the final activity for the first Unit: Creating the Seas: Ocean Basins and Water.  We will activate a set of questions to "test" your understanding of the material a few days before the end of this lesson.  The exam should take no more than one hour. You will log on to ANGEL to take the test. Stay tuned for additional details.

Additional Resources


Various Web sites with links to resources aimed at teachers and students:

  • NASA SeaWifs Ocean Color Page--Satellite-generated Maps of Ocean Productivity [9]
  • eWOCE Ocean Sections for various parameters (mostly chemical) [10]
  • Global Maps of Climatological Data (e.g. Sea-surface Temperature) [11]
  • Chemical Elements in Seawater--MBARI [8]
    [11]

A great (huge) book written about the importance of Salt in human history:

  • Kurlansky, Mark (2003)  Salt: A World History. New York: Penguin Books.

Tell us about it!

Have another Web site or printed piece on this topic that you have found useful? Share it in our Comment space below!

 

Summary & Final Tasks


Reminder - Complete all of the lesson tasks!

You have finished Lesson 3. Double-check the list of requirements on the first page of this lesson (Lesson 3 in the menu bar) to make sure you have completed all of the activities listed there before beginning the next lesson.

Tell us about it!

If you have anything you'd like to comment on, or add to, the lesson materials, feel free to post your thoughts below.  For example, what did you have the most trouble with in this lesson?  Was there anything useful here that you'd like to try in your own classroom?

Don't see the Post new comment area below? You need to be logged in to this site first! Do so by using the link at the top of the left-hand menu bar. Once you have logged in, you may need to refresh the page in order to see the comment area.

Authors: Michael Arthur and Chris Marone, Department of Geosciences, College of Earth and Mineral Sciences, The Pennsylvania State University.

Creative Commons License

© 2014 The Pennsylvania State University

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Source URL: https://www.e-education.psu.edu/earth540/content/c3.html

Links:
[1] https://courseware.e-education.psu.edu/courses/earth540/priv/Ball_Chapter.pdf
[2] http://www.saltinfo.com/
[3] https://www.youtube.com/watch?v=BayExatv8lE
[4] http://www.privatehand.com/flash/elements.html
[5] https://courseware.e-education.psu.edu/courses/earth540/WaterResTime.ppt
[6] https://courseware.e-education.psu.edu/courses/earth540/WaterResTime.pdf
[7] https://www.e-education.psu.edu/earth540/grading_rubric_problemsets
[8] http://www.mbari.org/chemsensor/pteo.htm
[9] http://oceancolor.gsfc.nasa.gov/SeaWiFS/
[10] http://www.ewoce.org/gallery
[11] http://iridl.ldeo.columbia.edu/maproom/.Global/