Let the adventure begin…
This packet is even better than the movie!
Don’t believe me? Compare them yourself!
Water has a unique story because water is a unique compound.
Liquid water covers about 75% of the surface of the earth as oceans,
lakes, and rivers. Huge aquifers store water deep underground, and
solid water (ice!) dominates the vast polar regions of the globe. Water
appears as icebergs in the oceans, and it whitens the Polar Regions
with snow. Water vapor from the evaporation of surface water and
from steam spouted from geysers and volcanoes is ever-present in the
earth’s atmosphere. Water is the foundation of everything that lives:
without it, neither plant nor animal life as you know it would exist! It
is the most common substance on Earth, yet it is hard to find elsewhere
in the universe. Let’s learn about why water is so unique, so important,
and just so darn cool!
And so
darn cool!
Water is
important!
Planet Ocean?
THE WATER MOLECULE
Water is a simple three-atom molecule, with each oxygen-hydrogen bond being a highly
polar covalent bond. Because it is more electronegative, oxygen attracts hydrogen’s
electron, leaving hydrogen partially positive and making oxygen partially negative. Because of
oxygen’s 2 lone pairs of electrons, the water molecule is BENT in shape. And because of this,
the dipole moments of the polar bonds do not cancel out, so overall, water is a POLAR
MOLECULE. Remember all of this? 
As we have learned, there are intermolecular forces occurring between molecules
(wall-walking gecko!). The special dipole-dipole force that we witness with water is the
HYDROGEN BOND: a partially positive hydrogen is attracted to a more electronegative
element (in this case, the oxygen). Most of water’s unique properties are because of these
hydrogen bonds, such as high surface tension, low vapor pressure, high boiling point, high
heat of vaporization, high specific heat, solid form floating in its liquid form, and dissolving
so many substances. In addition, one of water’s oddest traits is that it is a liquid at room
temperature when other molecules it’s size, and even bigger, are gases! Observe:
*Look how water stands-out! Normal room temperature is 25 °C
Compound
Formula
Molar Mass
Boiling Point
Methane
CH4
16.0 g/mol
-33.4 °C
Ammonia
NH3
17.0 g/mol
-164 °C
Water
H2O
18.0 g/mol
100. °C
Carbon Dioxide
CO2
44.0 g/mol
-56.6 °C
SURFACE PROPERTIES of WATER
Have you ever gone off a diving board and “belly-smacked?” Ouch! Why did it hurt and
sting your skin? Also, have you ever seen a water-walking bug? How can they do that? Why
does water “bead” on a well-waxed car or on plant leaves? Well, the molecules on the surface
of water act like a thin skin, creating tension. Water molecules within a sample are completely
surrounded by, and hydrogen bonded to, adjacent water molecules. The water molecules at the
surface, however, experience an uneven attraction as they are only hydrogen bonded to the
water molecules below them, and not the air above them. As a result, these surface water
molecules are drawn into the rest of the liquid, minimizing the surface area. This inward force,
or pull, on these surface molecules creates a surface
tension. All liquids have a surface tension, but water’s
is much higher than most others because of the strong
hydrogen bonds that are pulling the surface molecules
inward. This high surface tension is at least partially
responsible for two of water’s other notable properties:
low vapor pressure and high boiling point.
VAPOR PRESSURE and BOILING POINT
With the water molecules at the surface being pulled inward, not many of them are
easily able to escape and become a vapor. While water does evaporate (becomes a gas below
its boiling point), it does not evaporate as easily as other liquids – acetone, gasoline, paint
thinner, finger nail polish, etc… When a liquid does evaporate easily, which typically means
they have a low boiling point as well, we call it volatile. Volatile liquids also have a high
vapor pressure – pressure caused by molecules escaping from the surface of a liquid and
colliding with air molecules above them. So, since water has such strong hydrogen bonds
holding the surface molecules, not many vaporize, so there are fewer collisions with air
molecules and, of course, less vapor pressure created! So, water is definitely a nonvolatile
liquid with a low vapor pressure.
The boiling point of a liquid is defined as the temperature at which its vapor pressure
equals the atmospheric pressure, aka there is a “free pass” for liquid molecules to escape into
the gas phase. The heat of vaporization is the amount of energy needed to convert a liquid
to a gas at its boiling point, and is usually defined for a specific amount, like 1 mole. So for
water, its heat of vaporization is 40.68 kilojoules per mole (kJ/mol), compared to acetone
(formerly found in fingernail polish remover) which is 31.27 kJ/mol, or acetonitrile (currently
in most fingernail polish removers) which is 33.23 kJ/mol. If you’ve ever had fingernail
polish remover on your hands, it feels “cold” because the molecules are escaping the liquid
phase (evaporating) fairly quickly, removing heat from your skin as they go. Overall, water
has an unusually high boiling point (100. °C) for a liquid, especially one with its light molar
mass (see table on the previous page), and high heat of vaporization. Why? Since the
hydrogen bonds are so strong between water molecules, much more energy is required to
separate the liquid water molecules so they can escape that phase and become free-moving gas
molecules, increasing the vapor pressure enough so bubbles can form!
The reverse of boiling, of course, is condensing. When steam
condenses back to liquid water, it releases all of the energy that it had
absorbed to vaporize. Essentially, it releases that heat of vaporization
(40.68 kJ/mole). This is why steam burns can be quite bad! For
example, if something is heating-up in the microwave, and it is not
allowed to vent, the steam builds-up in the container, and when you open it,
hot steam can land on your skin, releasing a ton of energy as it condenses
back to liquid water at 100 ºC, and you’re 37 °C! Ouch!!! Grab the aloe!!!
Don’t mess with the
Microwave kids.
SPECIFIC HEAT
We looked at and calculated this a little while back, but again specific heat is the
energy required to raise the temperature of exactly 1 gram of a substance 1 °C, and is
represented by the symbol c. This number is a constant, and each substance has its very own
specific heat. Water, again being unique, has a very high specific heat compared to other
substances. What does that mean? Well, water has the ability to absorb a relatively large
amount of energy before its temperature changes, and conversely, release a relatively large
amount of heat as its temperature drops. This is one reason why a large percentage of Earth’s
population lives near water, and why waterfront property is so valued – on a warm day, the
water absorbs heat so it is cooler near the water, and on a cool night, the water releases heat so
it stays warmer near the water (doesn’t frost as soon as inland areas do, for example). Let’s
look at this topic a little more, and do some of those calculations again! 
Water’s specific heat (c) is 4.184 J/g•˚C (Joules per gram-degree Celsius), compared to
other substances like iron at 0.444, aluminum at 0.897, copper at 0.385, vegetable oil at 2.0,
sand 0.703, and concrete at 0.881 – all of those being labeled J/g•˚C of course! So we can see
why an aluminum or copper pot on the stove gets really hot while the water inside of it takes a
lot longer to heat-up and get to its boiling point. Also, you can see why on a really hot summer
day, walking barefoot on the concrete or on the sand at the beach can be quite painful to the
feet, but walking through a puddle on that concrete or walking in the water near the sand is
much more pleasant, even though everything is being exposed to the same amount of heat
energy from the sun!
The awesome calculations we can do with specific heat involve calculating
the amount of heat absorbed or released by a substance as its temperature changes,
or solving for the specific heat of a substance (and if it’s an element, looking up what
element it is on the back of your Periodic Table!). Observe the following examples:
q = c • m • ∆T
c = q ÷ (m • ∆T)
q = heat in Joules, c = Specific Heat, m = mass
ΔT = the change in the objects temperature
A piece of metal weighing 341 g is heated up. Its
initial temperature is 25°C, and 15 minutes later it
is 150°C. If the metal absorbed 38,250 Joules of
energy, what is the specific heat of the metal, and
what metal is it?  Periodic Table! 
ΔT = 150-25 = 125
c = 38,250 ÷ (341 • 125)
c = 38,250 ÷ (42,625)
c = 0.897 J/g•°C
Aluminum!
How many Joules of energy are released from
222g of water as it cools from 64°C to 22°C?
q = (4.184) • (222) • (64-22)
q = (4.184) • (222) • (42)
q = 39,012 J
WHEN WATER FREEZES
Normally, when a liquid sample freezes and becomes a solid, the
particles get closer together, and the sample size contracts, reducing its
volume. Since the mass stays the same, the solid is then denser than
the liquid form, and because of this, most all solids sink in their own
liquid. Well, you guessed it, water is unique and this is not the case!
As water cools, the hydrogen bonds hold the molecules in a rigid,
honeycomb-like structure that actually causes the liquid to expand
from 4°C down to 0°C. Consequently, solid water, ice, has a larger
volume which changes its density to 0.917 g/mL compared to water at
1.00 g/mL. So, ice floats in liquid water! Ice is, indeed, one of the very,
very few solids that can float in its own liquid form.
The fact that ice floats has important consequences for
living organisms. A layer of ice on top of a pond or lake
acts as an insulator for the water beneath, preventing it
from freezing under normal wintry conditions. Thus, life
can still survive under that ice, and we can go ice fishing.
Hip-Hip Hooray!
SOLVATION and AQUEOUS SOLUTIONS
Water has the ability to dissolve so many substances. Because of this, it is often known
as the “Universal Solvent,” and it is extremely rare to find “pure” water in nature. Even our tap
water contains several dissolved gases and minerals and purifying chemicals. Water samples
containing dissolved substances are known as aqueous solutions (fancy way of saying
“dissolved in water,” remember?). In a solution, any dissolved particles are known as the
solute, while water (or any liquid) that is doing the dissolving is known as the solvent. Like
in our limiting reactant lab, we used a copper(II)sulfate solution that we made. The
copper(II)sulfate was the solute and water was the solvent.
While water can dissolve a lot of substances, it cannot dissolve every substance, for
example, oil! What determines whether or not something will dissolve? Well, we have this
little phrase, “like dissolves like.” Water, being a polar molecule, has partially positive and
partially negative parts to is. If a solute is ionic or also polar, it can dissolve in water because
the partially charged parts of the water molecules will be attracted to the oppositely charged
parts of the solute, surround them, and basically “rip” the solute apart! This process of
dissolving a solute is known as solvation, although when water is the solvent we can also call
it hydration. Nonpolar solutes, like oil, have no charged parts, so there is no interaction
between them and water. However, “like dissolves like,” so nonpolar solutes do dissolve in
nonpolar solvents, as the repulsive forces between the molecules break the solute apart.
*Rememeber, we can figure out what kind of bond we have from the electronegativity differences of the
elements in them, and then molecular shape plays a role in whether or not a molecule is polar or nonpolar.
For simplicity sake here, we will deal with simple diatomic molecules to see if they are polar or nonpolar.
The sodium chloride is “attacked” by the water, with the partially positive (δ+) hydrogen parts surrounding
and pulling away the negative chloride ions (Cl-), and the partially negative (δ-) oxygen parts surrounding
and pulling away the positive sodium ions (Na +), creating a sodium chloride solution.
Here we see “hydrated”
sodium and chloride
ions, completely
surrounded by the
oppositely charged
polar ends of water
molecules.
ELECTROLYES and NONELECTROLYTES
Refreshing our memories, compounds that conduct an electric current either when they
are a liquid or in an aqueous solution are called electrolytes. All ionic compounds, and polar
covalent compounds that break apart (like strong acids and bases) are elctrolytes.
Compounds that DO NOT conduct an electric current when they are a liquid or in an aqueous
solution are called nonelectrolytes, and these would be nonpolar molecular compounds or
polar covalent compounds that do not break apart (like sugars and alcohols).
Not all electrolytes conduct electricity to the same extent, so there are both strong and
weak electrolytes, depending upon how well the solute is dissolved and how many charged
particles are dissociated in the water. Some strong electrolytes include strong acids like
HCl, HNO3, and H2SO4, many ionic halides (NaCl, KI, NaBr, KF, etc.), and some strong bases,
which have hydroxides in them (NaOH, Ca(OH2), Mg(OH)2, KOH). Some weak electrolytes
include weak acids like acetic, citric, carbonic, and phosphoric, the weak base ammonia (NH3),
as well as some chlorides, such as HgCl2 and PbCl2. Some nonelectrolytes include a lot of
organic compounds like methane, propane, and butane, fats and oils, alcohols like ethanol,
methanol, and isopropyl, sugars (sucrose, glucose, fructose, dextrose), glycerol, and nonpolar
carbon compounds (CCl4, CH2Br2, etc.).
 On a final note, we can write an equation to show how solid ionic compounds dissolve
and become aqueous solutions by breaking apart into their component cations and
anions. Here are a couple of examples:
 The dissolving of potassium bromide:
KBr(s)
K+1(aq) + Br -1(aq)
 The dissolving of magnesium phosphate:
Mg3(PO4)2 (s)
3 Mg +2(aq) + 2 PO4-3(aq)
All finished!