MAR 110: Introductory
Oceanography
Properties of ocean water
Salt and history, part 1
• The ancient city of Jericho – one of the oldest
surviving settlements in human history – is located on
the Dead Sea.
– The city was apparently founded as a salt trading center,
and the Dead Sea has long been a source of salt.
– The Dead sea is in a depression along a transform fault
complex originating in the Gulf of Aqaba.
– The salt of the Dead Sea originated in an arm of the
Mediterranean Sea that was cut off by tectonic movements
about 6 million years ago.
– In time, most of the water evaporated and left the salt
behind.
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Salt and history, part 2
• Salt has long been important to humans as a
preservative and spice; at times in human history, salt
was as valuable as gold, if not more so.
• Besides its use as a preservative and seasoning, salt is
used to clear roads and walkways, in manufacturing,
and as water softeners to remove other dissolved
substances during water treatment.
• Salt is typically obtained either by evaporation
processes or by solution mining of rock salt.
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Water and life
• Water is essential for life as we know it.
• Astronomers search for water on distant planets in the
effort to find life elsewhere in the universe.
• Life on Earth began in water and evolved there for 3
billion years before spreading onto land.
• Terrestrial organisms are tied to water, too.
• Most cells are surrounded by water and cells are
about 70-95 percent water.
• Water exists in three states: ice, liquid, and vapor.
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Water is a polar molecule
• The chemical formula of water is H2O.
• Water is a polar molecule because oxygen has a much
higher affinity for electrons (is more electronegative)
than hydrogen.
– Because oxygen is more electronegative, the electrons
spend more time around the oxygen, giving the oxygen end
of the molecule a partial negative charge.
– The hydrogen end of the molecule likewise has a partial
positive charge.
– Thus, opposite ends of the water molecule have opposite
charges; such molecules are called polar molecules.
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Hydrogen bonds, part 1
• The slightly negative regions of one molecule are
attracted to the slightly positive regions of nearby
molecules, forming a hydrogen bond.
– Each water molecule can form hydrogen bonds with up to
four neighbors.
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Hydrogen bonds, part 2
• The hydrogen bonds joining water molecules are
weak, about one-twentieth as strong as covalent
bonds (which involve electron sharing between
atoms).
– They form, break, and reform with great frequency.
– At any instant, a substantial percentage of all water
molecules are bonded to their neighbors, creating a high
level of structure.
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Unusual properties of water
• Water’s high polarity and ability to form hydrogen
bonds gives the molecule peculiar properties:
– Cohesion (and adhesion);
– Unique thermal properties that make water resistant to
temperature change (and change of state);
– Highest density in liquid state; and
– A versatile solvent.
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Cohesion
• Hydrogen bonds hold water molecules together; this
ability to stick together is called cohesion.
• Cohesion makes it possible to transport water against
force of gravity.
– A related property, adhesion – the ability to stick to another
polar or charged substance like glass or xylem vessels –
likewise makes it possible to transport water against force
of gravity.
• Cohesion between water molecules creates high
surface tension makes it difficult to break surface.
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Kinetic energy
• Kinetic energy is the energy of motion.
• Heat is the total quantity of kinetic energy in a
substance.
• Temperature is the average amount of kinetic energy
in a substance.
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Measuring heat
• Kinetic energy often measured in calories (cal), the
amount of heat energy required to raise the
temperature of 1 g of water 1 °C
• The joule (J) is another way to measure kinetic
energy.
– 1 cal = 4.187 J; 1 J = 0.239 cal.
• The British Thermal Unit (BTU) is the amount of
energy it takes to raise the temperature of 1 pound of
water 1 °F (from 62 °F to 63 °F).
– 1 BTU = 252 cal = 1055 J
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Measuring temperature, part 1
• There are a number of instruments for measuring
temperature. All work on the principle that most
substances expand when heated, calibrating this
change in volume to measure temperature.
• There are three temperature scales used in the United
States: the Fahrenheit Scale, the Celsius Scale, and
the Kelvin Scale.
– The Fahrenheit scale is used by public weather reports from
the National Weather Service and the news media; few
other countries than United States use it.
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Measuring temperature, part 2
• Three temperature scales (continued):
– The Celsius scale is used either exclusively or
predominately in most countries other than United States,
which uses it for scientific work. It is slowly being
established to supersede the Fahrenheit scale.
• 0 °C = 32 °F
• 100 °C = 212 °F
– The Kelvin scale is used in scientific research, but not by
climatologists and meteorologists. It is similar to the
Celsius scale, but the zero point is set to absolute zero, the
temperature at which all molecular motion ceases.
• 0 K = -273.15 °C = -459.67 °F
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Unique thermal properties, part 1
• Based on its molecular weight and the properties of
similar substances, pure water should freeze at -90 °C
and boil at about -70 °C.
– Instead, pure water freezes at 0 °C and boils at 100 °C.
– The unique thermal properties are a result of water’s
structure, and its tendency to form hydrogen bonds.
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Latent heat, part 1
• Latent heat is energy used to change the state of water
rather than change the temperature, thus it is latent, or
hidden.
– When water is at a temperature at which it should change
state, latent heat is added or removed until the change of
state is complete; the temperature, however, does not
change.
• Latent heat of fusion is the hidden heat given off or
absorbed as water freezes or melts.
– The latent heat of fusion of water is 80 cal/g at 0 °C.
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Latent heat, part 2
• Latent heat of vaporization is the amount of energy it
takes for 1 g of a substance evaporates.
– Latent heat of condensation is the amount of energy given
off as 1 g of a substance condenses.
– The latent heat of vaporization/condensation of water
varies with temperature, ranging from 597 cal/g at 0 °C to
540 cal/g at 100 °C.
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Latent heat, part 3
• Latent heat of sublimation is the amount of energy it
takes for 1 g of a substance changes directly from the
solid state to the gaseous state.
– Latent heat of deposition is the amount of energy given off
as 1 g of a substance changes directly from the gaseous
state to the solid state.
– The latent heat of sublimation/deposition of water is
relatively constant, ranging from 677 cal/g at 0 °C to 678
cal/g at -30 °C.
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Specific heat
• Specific heat is the amount of energy it takes to raise
or lower the temperature of 1 g of a substance 1
degree C.
– The specific heat of water is 1 cal/g/°C.
– The specific heat of ice is 0.5 cal/g/°C.
• Water changes its temperature less than most other
substances when it absorbs or radiates a given amount
of energy.
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Climate effects
• Water’s thermal properties explains its thermal
inertia, or its resistance to temperature change when it
absorbs or radiates a given amount of energy.
– Water stabilizes air temperatures by absorbing heat from
warmer air and releasing heat to cooler air.
– Water can absorb or release relatively large amounts of heat
with only a slight change in its own temperature.
• In turn, water’s thermodynamic properties explain
why maritime climates have more moderate
temperature ranges than arid climates, and why
sweating is so important to cooling the body.
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Evaporative cooling
• The evaporation of water removes a lot of kinetic
energy from the body, thus bringing the temperature
down.
• When a person becomes dehydrated, this cooling
mechanism breaks down.
• This is why it is important to drink plenty of liquids
during hot weather, intense exercise or while sick.
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Solutions
• A solution is a liquid that is a homogenous mixture of
two or more substances.
– The solvent is the dissolving agent.
– The solute is the substance being dissolved.
– An aqueous solutions is one in which water is the solvent.
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Versatile solvent
• Pure water rarely occurs naturally, as water typically
contains a variety of dissolved and suspended
substances.
– Water’s polar properties make it a good solvent for polar
and ionic compounds.
• The negatively charged oxygens attract positively charged ions or
poles, while the positively charged hydrogens attract negatively
charged ions or poles.
• Polar molecules are also soluble in water because they can also
form hydrogen bonds with water.
• Even large molecules, like proteins, can dissolve in water if they
have ionic and polar regions.
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Example of a solution
• When a crystal of salt (NaCl) is placed in water, the
Na+ cations form hydrogen bonds with partial
negative oxygen regions of water molecules.
• The Cl- anions form hydrogen bonds with the partial
positive hydrogen regions of water molecules.
• Each dissolved ion is surrounded by a sphere of water
molecules, a hydration shell.
• Eventually, water dissolves all the ions, resulting in a
solution with two solutes, sodium and chloride.
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Chemistry of seawater, part 1
• The constituents of seawater comes from several
sources:
– Weathering and erosion of rock on land and transport by
rivers to the oceans;
– Volcanic activity, such as at hydrothermal vents;
– Removal via sedimentation, subduction and other
geological processes;
– Solution of gases from the atmosphere.
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Chemistry of seawater, part 2
• Seawater is a salt solution of nearly uniform
composition.
– Salinity is a measure of the amount of salt dissolved in
seawater.
– The salinity of seawater averages about 3.5 percent.
– If all the water in the oceans evaporated, sea salts would
cover the Earth to a depth of 45.5 m.
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Chemistry of seawater, part 3
• William Dittmar, analyzing water samples from the
Challenger expedition, confirmed earlier observations
that the ratio of the concentrations of the major
constituents of seawater was similar.
– Dittmar thus discovered the principle of constant
proportions, which allows one to estimate total salinity of
seawater based on a measurement of the concentration of
one ion, such as chloride (Cl-).
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Chemistry of seawater, part 4
• The concentration of seawater constituents such as
chloride and sodium (Na+) occur in constant
proportions and change slowly by mixing; thus they
are conservative properties of seawater.
• The concentrations of other constituents are more
variable; thus they are non-conservative properties of
seawater.
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Chemistry of seawater, part 5
• The salinity of seawater was originally measured by a
titration method, but after World War II scientists
observed that measurments of the electrical
conductivity of seawater can provide much more
accurate measurements.
– Salinity is now determined from the ratio of the
conductivity of a water sample to that of standard seawater;
it is given in practical salinity units (psu).
• Standard seawater is prepared by Ocean Scientific International.
• Evaporation of seawater increases salinity; freshwater
runoff decreases salinity.
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Chemistry of seawater, part 6
• While more than 70 minerals are dissolved in
seawater, six make up more than 99 percent of all sea
salts: chloride, sodium, sulfate (SO42-), magnesium
(Mg2+), calcium (Ca2+), and potassium (K+).
– Table salt (NaCl) makes up more than 86 percent of sea
salts.
• Important trace elements include aluminum (Al),
chromium (Cr), gold (Au), lead (Pb), nickel (Ni), and
zinc (Zn).
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Chemistry of seawater, part 7
• Gases, such a carbon dioxide (CO2), nitrogen (N2),
and oxygen (O2), also dissolve in water.
• The solubility of gases in water is typically inversely
correlated with temperature.
– In seawater, salinity and pressure also affect the solubility
of gases, but temperature is the major driver.
• Waves on the surface encourage solution of gases by
increasing the roughness and surface area of the
water, thus increasing the interface between
atmosphere and ocean.
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Chemistry of seawater, part 8
• Biochemical processes also affect the concentrations
dissolved gases in seawater.
– In the photic zone (defined by the depth to which light
penetrates), photosynthesis removes carbon dioxide and
adds oxygen.
– In the aphotic zone (below the depth to which light
penetrates), no photosynthesis is possible; decomposers
predominate; the process of cellular respiration removes
oxygen and adds carbon dioxide.
• Carbon dioxide is more abundant in the ocean than
the atmosphere.
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Chemistry of seawater, part 9
• Gases not involved in biochemical processes remain
at a relatively constant concentration; any changes in
concentration typically result from diffusion of the
gases.
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Dissociation, part 1
• Occasionally, a hydrogen atom shared by two water
molecules shifts from one molecule to the other.
• The hydrogen atom leaves its electron behind and is
transferred as a single proton - a hydrogen ion (H+).
• The water molecule that lost a proton is now a
hydroxide ion (OH-).
• The water molecule with the extra proton is a
hydronium ion (H3O+).
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Dissociation, part 2
• A simpler way to view this process is that a water
molecule dissociates into a hydrogen ion and a
hydroxide ion:
– H2O <=> H+ + OH-
• This reaction is reversible.
• In pure water at equilibrium, only one water molecule
in every 554 million is dissociated.
• At equilibrium the concentration of H+ or OH- is
10-7M (25 degrees C).
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Dissociation, part 3
• Because hydrogen and hydroxide ions are very
reactive, changes in their concentrations can
drastically affect the proteins and other molecules of
a cell.
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Acids and bases, part 1
• Adding certain solutes, called acids and bases,
disrupts the equilibrium and modifies the
concentrations of hydrogen and hydroxide ions.
• An acid is a substance that increases the hydrogen ion
concentration in a solution.
• Any substance that reduces the hydrogen ion
concentration in a solution is a base.
• Some bases reduce H+ directly by accepting hydrogen
ions.
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Acids and bases, part 2
• Other bases reduce H+ indirectly by dissociating to
OH- that combines with H+ to form water.
• Strong acids or bases dissociate completely in water.
• In weak acids or bases, the dissociation is reversible;
at equilibrium there will be a fixed ratio of products
and reactants.
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Example reactions
•
•
•
•
HCl => H+ + ClNaOH => Na+ + OHNH3 + H+ <=> NH4+
H2CO3 <=> HCO3- + H+
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The pH scale
• In any solution the product of their H+ and OHconcentrations is constant:
– [H+] [OH-] = 10-14.
• The pH is the negative logarithm of the hydrogen ion
concentration, and can range from 1 (very strong
acid) to 14 (very strong base)
• Neutral pH is 7.
• Most biological fluids range from pH 6 to pH 8.
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Buffers
• Buffers minimize changes in pH by accepting
hydrogen ions from the solution when they are in
excess and donating hydrogen ions when they have
been depleted.
• Buffers typically consist of a weak acid and its
corresponding base.
• Carbonic acid (H2CO3) serves as an important buffer
in biological systems.
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Acid precipitation
• Acid precipitation harms water quality, thus the
environment for all life where it occurs.
– Uncontaminated rain has a slightly acidic pH of 5.6.
– The acid is a product of the formation of carbonic acid
from carbon dioxide and water.
• Acid precipitation occurs when rain, snow, or fog has
a pH that is more acidic than 5.6.
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Primary causes
• Acid precipitation is caused primarily by sulfur
oxides and nitrogen oxides in the atmosphere.
• These molecules react with water to form strong
acids.
• These fall to the surface with rain or snow.
• The major source of these oxides is the burning of
fossil fuels (coal, oil, and gas) in factories and
automobiles.
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Widespread effects
• Tall smokestacks allow pollution to spread from its
site of origin to contaminate relatively pristine areas.
• For example, rain in the Adirondack Mountains of
upstate New York averages a pH 4.2.
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Effects on aquatic systems, part 1
• The effects of acids in lakes and streams is more
pronounced in the spring during snowmelt.
• As the surface snows melt and drain down through
the snow field, the meltwater accumulates acid and
brings it into lakes and streams all at once.
• The pH of early meltwater may be as low as 3.
• Acid precipitation has a great impact on the eggs and
the early developmental stages of aquatic organisms
that are abundant in the spring.
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Effects on aquatic systems, part 2
• Thus, strong acidity can alter the structure of
molecules and impact ecological communities.
• Freshwater ecosystems are more vulnerable to acid
precipitation than marine ecosystems because they
lack the buffering capacity of seawater.
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Physical properties, part 1
• Several properties change with salinity.
• Dissolved substances lower the freezing point of
water.
– The freezing point of average seawater is -1.9 °C.
– As seawater freezes, salt crystals are excluded from ice
crystals that form; as a result the temperature of seawater
changes as it freezes.
• Some brine may remain trapped in the ice, however, so the
unfrozen water becomes saltier; this lowers the freezing point even
more.
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Physical properties, part 2
• Density is mass per unit volume.
– The density of fresh water is 1 g/cm3.
• Water is only slightly compressible.
• Objects less dense than surrounding water will rise or float at the
surface (positive buoyancy); objects more dense than surrounding
water will sink (negative buoyancy); objects of the same density as
surrounding water will remain at the same level (neutral buoyancy).
– Generally, as an object cools, it contracts and becomes
more dense; as an object warms, it expands and becomes
less dense.
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Physical properties, part 3
• Fresh water is unusual because it is less dense as a
solid than as a liquid.
– Most materials contract as they solidify, but water expands.
– At temperatures above 4 ° C, water behaves like other
liquids, expanding when it warms and contracting when it
cools.
• Water begins to freeze when its molecules are no
longer moving vigorously enough to break their
hydrogen bonds.
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Physical properties, part 4
• When water reaches 0 °C, water becomes locked into
a crystalline lattice with each molecule bonded to the
maximum of four partners.
• As ice starts to melt, some of the hydrogen bonds
break and some water molecules can slip closer
together than they can while in the ice state.
• Ice is about 10 percent less dense than water at 4 °C.
• Therefore, ice floats on the water below.
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Physical properties, part 5
• When it freezes, seawater does not behave as
freshwater does.
– Density increases with increasing salinity (and vice versa).
– The temperature of maximum density increases with
increasing salinity.
• For salinities below 24.7, the temperature of maximum density is
above the freezing point.
• At a salinity of 24.7, the temperature of maximum density equals
the freezing point.
• At a salinity below 24.7, the temperature of maximum density is
below the freezing point.
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Physical properties, part 6
• Freezing seawater (continued):
– Density increases as temperature decreases, such that
colder waters sink until they reach bottom, or they reach a
level in which their density equals that of surrounding
water.
– Consequences for life:
• If ice sank, eventually all ponds, lakes, and even the ocean would
freeze solid; During the summer, only the upper few inches
of the ocean would thaw.
• Instead, the surface layer of ice insulates liquid water below,
preventing it from freezing and allowing life to exist under the
frozen surface.
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Physical properties, part 7
• Freezing seawater (continued):
– Density increases as temperature decreases, such that
colder waters sink until they reach bottom, or they reach a
level in which their density equals that of surrounding
water.
• Air pressure at sea level is 1013.25 mbar (millibar).
– Water is much denser, so that it exerts much more pressure
than air.
• In general, with each 10 m of depth change, you increase or
decrease pressure by one atmosphere (one bar).
• In general one decibar (0.1 bar) is equivalent to 1 m in depth, so
depth is essentially equivalent to pressure.
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Physical properties, part 8
• Pressure (continued):
– Water is relatively incompressible, so that density does not
vary much with pressure.
• Seawater freezes at high latitudes in the respective
hemisphere’s winter to form sea ice.
– Microscopic crystals change into hexagonal needles about
1 to 2 cm long.
– Ice crystals form a surface like a blanket of wet snow.
– Ice crystals then grow downward to form a plastic-like ice
layer containing chambers filled with increasingly saltier
brine.
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Physical properties, part 9
• Sea ice (continued):
– The salt content of newly formed sea ice depends on
temperature.
• The colder the temperature, the more quickly the ice freezes, thus
trapping more brine in the ice.
– An ice layer 1 to 3 m thick can form in one winter. This is
called first-year ice.
– Where sea ice rarely melts, multi-year ice dominates.
• Sea ice can grow to 3.5 m thick, but currents and winds can cause
ice to pile up as pressure ridges.
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Physical properties, part 10
• Sound travels more quickly through water than
through air.
– The speed of sound through water is a bout 1500 m/sec,
about four times that of the speed of sound in air.
– Sound, as sonar, is used to remotely sense the water
beneath the surface.
• Depth gauges, fish finders, etc.
• SONAR (SOund Navigation And Rangin) differs from echo
sounders in that SONAR operators can aim the beam.
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Physical properties, part 11
• Sound (continued):
– Refraction of sound waves through water gives rise to the
SOFAR channel.
• The SOFAR (SOund Fixing And Ranging) channel is about 1000 m
deep; it can propagate sound worldwide as sound waves are
refracted off the top and bottom of the layer.
• The SOFAR channel can be used to monitor the temperature of
seawater through a process called acoustic tomography.
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