SMS-491: Physical solutions of everyday problems in
aquatic sciences.
Lecture 12: Properties of water
The water molecule
A single water molecule has an oxygen atom bonded to two hydrogen atom (Fig. 1). The
nature of the bonds is such that the hydrogen atoms are arranged on one side of the
oxygen atom resulting in a molecule that has electric polarity; the oxygen side is slightly
negative and the hydrogen side is slightly positive. This polarity means that adjacent
water molecules (or ions in water) attract/repulse each other electrostatically depending
on the arrangement of the molecules. The negatively charged oxygen of one molecule
attracts the positively charged hydrogen molecule of another molecule forming a
hydrogen bond. This bond is not as strong as the intermolecular (covalent) bond between
oxygen and hydrogen. At room temperature the energy associated with this hydrogen
bond (the work needed to break it) is about five times the average kinetic energy of the
molecules. Thus, at this temperature, the water molecules are attached to their neighbors
rigidly but not firmly, consistent with it being liquid. When the temperature is raised
above 100°C the average kinetic energy is higher than the energy of the hydrogen bonds
and molecules move about without being tied to their neighbors.
Figure 1: water molecule and charge distribution around it. From:
http://www.sbu.ac.uk/water/index.html
The attraction between the molecules in liquid water results in the charge of a water
molecule being canceled by its neighbors inside the water. However, the charge of a
water molecule located at the edges of a water volume is not canceled and it experiences
an imbalance of forces on it. The molecule is preferentially attracted into the fluid,
resulting in the formation of concave drops when water is spilled, to minimize the size of
the interface (spheres have the minimum surface to volume ratio, but remember that
gravity also affect the spilled liquid). Surface tension expresses the amount of energy a
molecule at the interface between a fluid and its surrounding has (or the work needed to
place a molecule there) relative to the extra surface area created at the interface. It is
denoted by g and has units of J m-2 (=N m-1). Water has the highest surface tension when
in contact with air than all other liquids in room temperature. It varies little with
temperature (10% for 40°C) or salinity (1% for 35psu). Surface tension can vary a lot if
surfactants (organic films which are found at the water-air interface) are present.
Figure 2: water molecule within the fluid and at the interface. From:
http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch14/property.html#tension
There is a force of attraction between molecules in liquids, and liquids can flow until they
take on the shape that maximizes this force of attraction. Below the surface of the liquid,
the force of cohesion (literally, "sticking together") between molecules is the same in all
directions, as shown in the Fig. 2. Molecules on the surface of the liquid, however, feel a
net force of attraction that pulls them back into the body of the liquid. As a result, the
liquid tries to take on the shape that has the smallest possible surface area
the shape of
a sphere. The magnitude of the force that controls the shape of the liquid is called the
surface tension. The stronger the bonds between the molecules in the liquid, the larger the
surface tension.
There is also a force of adhesion (literally, "sticking") between a liquid and the walls of
the container. When the force of adhesion is more than half as large as the force of
cohesion between the liquid molecules, the liquid is said to "wet" the solid. A good
example of this phenomenon is the wetting of paper by water. The force of adhesion
between paper and water combined with the force of cohesion between water molecules
explains why sheets of wet paper stick together.
Water wets glass because of the force of adhesion that results from interactions between
the positive ends of the polar water molecules and the negatively charged oxygen atoms
in glass. As a result, water forms a meniscus that curves upward in a small-diameter glass
tube, as shown in the figure below. (The term meniscus comes from the Greek word for
"moon" and is used to describe anything that has a crescent shape.) The meniscus that
water forms in a buret results from a balance between the force of adhesion pulling up on
the column of water to wet the walls of the glass tube and the force of gravity pulling
down on the liquid.
Figure 3. Water climbs the walls of a small-diameter tube to form a meniscus that curves
upward, whereas mercury forms a meniscus that curves downward.
The force of adhesion between water and wax is very small compared to the force of
cohesion between water molecules. As a result, rain doesn't adhere to wax. It tends to
form beads, or drops, with the smallest possible surface area, thereby maximizing the
force of cohesion between the water molecules. The same thing happens when mercury is
spilled on glass or poured into a narrow glass tube. The force of cohesion between
mercury atoms is so much larger than the force of adhesion between mercury and glass
that the area of contact between mercury and glass is kept to a minimum, with the net
result being the meniscus shown in the above figure.
Capillarity
Immersing an open tube into water we observe water climbing into the tube (Fig. 4). The
vertical force pulling the water up equals F=2πrγcosθ, where 2πr is the circumference of
the tube, γ the surface tension and θ the angle of the meniscus. The pressure exerted due
to this force (which balances the additional gravitational pressure) is given by the force
divided by the surface area within the tube (πr2) Æ p=2γcosθ/r. The gravitational
pressure is the hydrostatic pressure, p=hgρ, where h is the height of the column, ggravity and ρ the density. Equating the two (each pulling in a different direction) and
solving for h: h=2γcosθ/(rρg). The smaller the radius of the tube, the higher up the fluid
climbs. Similarly, the higher surface tension the higher the fluid climbs (for a similar
meniscus angle).
Figure 4. Water (or other fluid) climbing through a tube.
http://student.dcu.ie/~copains/fr489_02/experiences/LaChromatographie.html
Hydrogen Bonding and the Anomalous Properties of Water
We are so familiar with the properties of water that it is difficult to appreciate the extent
to which its behavior is unusual.
•
•
•
•
•
•
•
•
Most solids expand when they melt. Water expands when it freezes.
Most solids are more dense than the corresponding liquids. Ice (0.917 g/cm3) is
not as dense as water.
Water has a melting point at least 100oC higher than expected on the basis of the
melting points of H2S, H2Se, and H2Te.
Water has a boiling point almost 200oC higher than expected from the boiling
points of H2S, H2Se, and H2Te.
Water has the largest surface tension of any common liquid except liquid
mercury.
Water has an unusually large viscosity.
Water is an excellent solvent. It can dissolve compounds, such as NaCl, that are
insoluble or only slightly soluble in other liquids.
Water has an unusually high heat capacity. It takes more heat to raise the
temperature of 1 gram of water by 1oC than any other liquid.
These anomalous properties all result from the strong intermolecular bonds in water
discussed above. The hydrogen bonds in water are particularly important because of the
dominant role that water plays in the chemistry of living systems. Hydrogen bonds are
not limited to water, however.
Hydrogen-bond donors include substances that contain relatively polar H-X bonds, such
as NH3, H2O, and HF. Hydrogen-bond acceptors include substances that have
nonbonding pairs of valence electrons. The H-X bond must be polar to create the partial
positive charge on the hydrogen atom that allows dipole-dipole interactions to exist. As
the X atom in the H-X bond becomes less electronegative, hydrogen bonding between
molecules becomes less important. Hydrogen bonding in HF, for example, is much
stronger than in either H2O or HCl.
The hydrogen bonds between water molecules in ice produce the open structure shown in
Fig. 5. When ice melts, some of these bonds are broken, and this structure collapses to
form a liquid that is about 10% denser. This unusual property of water has several
important consequences. The expansion of water when it freezes is responsible for the
cracking of concrete, which forms potholes in streets and highways. But it also means
that ice floats on top of rivers and streams. The ice that forms each winter therefore has a
chance to melt during the summer.
Figure 5. The structure of ice. Note that the hydrogen atoms are closer to one of the
oxygen atoms than the other in each of the hydrogen bonds. Figure from
http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch14/critical#critical.
Figure 6 below shows another consequence of the strength of the hydrogen bonds in
water. There is a steady increase in boiling point in the series CH4, GeH4, SiH4, and
SnH4. The boiling points of H2O and HF, however, are anomalously large because of the
strong hydrogen bonds between molecules in these liquids. If this doesn't seem important,
try to imagine what life would be like if water boiled at -80oC.
The surface tension and viscosity of water are also related to the strength of the hydrogen
bonds between water molecules. The surface tension of water is responsible for the
capillary action that brings water up through the root systems of plants. It is also
responsible for the efficiency with which the wax that coats the surface of leaves can
protect plants from excessive loss of water through evaporation.
Figure 6. Melting point and boiling point of different material in comparison to water.
The unusually large heat capacity of water (see last lecture) is also related to the strength
of the hydrogen bonds between water molecules. Anything that increases the motion of
water molecules, and therefore the temperature of water, must interfere with the hydrogen
bonds between these molecules. The fact that it takes so much energy to disrupt these
bonds means that water can store enormous amounts of thermal energy. Although the
water in lakes and rivers gets warmer in the summer and cooler in the winter, the large
heat capacity of water limits the range of temperatures that would otherwise threaten the
life that flourishes in this environment. The heat capacity of water is also responsible for
the ocean's ability to act as a thermal reservoir that moderates the swings in temperature
that occur from winter to summer.
Hydrophobic and hydrophilic substances
Substances that dissolve readily in water are termed hydrophilic. They are composed of
ions or polar molecules that attract water molecules through electrical charge effects.
Water molecules surround each ion or polar molecule on the surface of a solid substance
and carry it into solution.
Ionic substances such as sodium chloride dissolve because water molecules are attracted
to the positive (Na+) or negative (Cl-) charge of each ion. Polar substances such as urea
dissolve because their molecules form hydrogen bonds with the surrounding water
molecules.
Molecules that contain a preponderance of nonpolar bonds are usually insoluble in water
and are termed hydrophobic. This is true, especially, of hydrocarbons, which contain
many C–H bonds. Water molecules are not attracted to such molecules and so have little
tendency to surround them and carry them into solution.
Many substances, such as household sugar, dissolve in water. That is, their molecules
separate from each other, each becoming surrounded by water molecules. When a
substance dissolves in a liquid, the mixture is termed a solution. The dissolved substance
(in this case sugar) is the solute, and the liquid that does the dissolving (in this case
water) is the solvent. Water is an excellent solvent for many substances because of its
polar bonds.
Capillary waves
Figure 7. Capillary waves on a beach. Figure from
http://www.eng.vt.edu/fluids/msc/gallery/gall.htm
Capillary waves (or ripples) are surface waves where the primary restoring force is
surface tension. The force per unit area (dynamic pressure, in addition to the hydrostatic
pressure due to gravity) is proportional to the surface tension [γ, N m-1] and the curvature
of the wave, the second derivative of the wave height with respect to distance [m-1].
Assuming a sinusoidal surface wave, η=Asin{2p(x/λ-w/T)}, where T is the period and λ
the wavelength, the pressure at the surface is:
p={ρg+γ(2π/λ)2}η
The only difference with surface gravity wave is in the added surface tension term. The
phase speed of these waves (assuming wavelength<<depth of water) is:
c=[(2π/λ) (g+γ(2π/λ)2/ρ]1/2
The group speed (speed of energy propagation) is faster than the phase speed (3c/2). Thus
capillary-gravity waves are faster than gravity waves. As the waves become large
(λ>>2π(γ/ρg)1/2 (about 2cm for pure water), surface tension is negligible and the wave is
a gravity wave.
Capillary waves are strongly attenuated and thus do not propagate very far. In the oceans
they are generated by the wind and cause the surface of the oceans to be much steeper
than it would be were they not present. This, in turn, has strong effect on light penetration
in the oceans, as the steep surface refract light resulting in the strong focusing of light
observed in bottom of swimming pools and small ponds. It also affects the properties of
RADAR waves bouncing off the ocean’s surface. When hydrocarbon leak from the
seabed they appear as oil slicks on the sea surface. Gas bubbles surrounded by oil rise
through the water, burst on the sea surface and create a small, thin oil slick. Such natural
oil seepage is known to occur from time to time in potential oil reservoirs, but the
frequency and the amount of seepage may vary. When oil is forming a thin layer on the
sea surface it will dampen the capillary waves. Due to the difference in back scatter from
areas with waves and areas with no or dampened waves (from it the expression ‘oil on
troubled waters’), radar satellites may detect natural oil seepage at the sea surface.
Another sources of material for slick may be derivatives of algal blooms.
References and additional reading
http://faculty.uca.edu/~johnc/wat1440.htm
http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch14/property.html#tension
http://www.sbu.ac.uk/water/index.html
http://www.accessexcellence.org/AB/GG/garland_PDFs/Panel_2.02a.pdf
http://www.accessexcellence.org/AB/GG/garland_PDFs/Panel_2.02b.pdf
Denny, M. W., 1993, Air and Water, Princeton U. Press, Chapter 12.
Lighthill, J., Waves in fluid, Cambridge University Press, Chapter 3.
Boss and Jumars, 2003
This page was last edited on 4/25/2003