Water
 Water is a polar molecule composed of two polar covalent
O-H bonds in a bent or angular molecular geometry with
two pairs of nonbonding electrons.
 water has 4 pairs of electrons arranged tetrahedral around the central
oxygen atom
 75% of the earth’s surface is covered with water;
 about 97% of the total water available on Earth is salt water, about
2% is frozen as the polar ice caps, and the rest (1%) is fresh water.
 The amount of water on this planet is fairly constant and
cleans and replenishes itself via “The Hydrologic cycle”.
 Water vapor in the atmosphere (clouds & such) returns to the earth
via precipitation (rain, snow, etc.) where it flows into land pockets
(oceans, lakes, or rivers), or it is absorbed into the ground, or it
evaporates back up into the atmosphere (completing the cycle).
Water
 Rainwater collects dust particles and gases as it travels
from the atmosphere to the ground.
 Gases like O2, N2, and CO2 all dissolve to some degree in rainwater.
An equilibrium is established between dissolved CO2 in water with
carbonic acid making rainwater (about pH 5) more acidic than pure
water (pH 7). CO2 + H2O  H2CO3
 As water flows beneath or atop the surface of the planet, it
readily dissolves many substances from the soil and
rocks.
 Some common dissolved substances are Na+, K+, Ca2+, Mg2+, Fe2+,
Cl-, SO42-, and HCO3-. Ca2+, Mg2+, and Fe2+ salts are responsible for
“Hard water” (these positive ions react with the negative ions in
soap to form insoluble scum). Soft water contains soluble ions like
sodium and potassium.
Water
 There are different types of “Polluted” water:
 Pathogenic (disease-causing) microorganisms like cholera, typhoid, hepatitis, and
dysentery still effect over 70% of the world’s population.
 Aerobic biodegradation (aerobic oxidation) happens when microorganisms break
down organic material in the present of dissolved oxygen to produce, for example,
CO2, PO43-, NO3- SO42-, and HCO3-. The measure of oxygen needed to degrade
organic material is referred to as the BOD (biochemical oxygen demand).
 Anaerobic decay happens when the oxygen is depleted. The microorganisms reduce
organic material (instead of oxidizing it) to produce nasty smelling substances like
CH4, NH3, H2S, and amines. No life (excepts anaerobic microorganisms) can exist in
such water.
 Certain bacteria in water breaks down organic matter and, in the process, depletes the
dissolved oxygen (which marine life is dependent on) while enriching the amount of plant
nutrients (PO43-, NO3-) present. These nutrients promote algae growth. If the
concentration of plant nutrients (from natural and human contributions) is left
unchecked, it can lead to an excess of algae which, as it dies, increases the BOD eventually
leading to anaerobic biodegradation. This process called eutrophication.
 Industrial waste like VOC’s (volatile organic compounds like trichloroethylene),
heavy metal ions/compounds (like Hg, Pb, & Cd), and a number of organic and
inorganic materials from LUST (leaking underground storage tanks).
 Acid rain produced from dissolved SOx and NOx compounds from air pollution and
acid mine drainage from mining operations.
Some Physical Properties of Water
 Water is colorless, odorless, and tasteless.
 The normal boiling point is 100oC and the normal
melting point is 0oC.
 The heat of vaporization (DHvap) is 2259 J/g or 540
cal/g and the heat of fusion (DHfus) is 335 J/g or
80 cal/g.
 The vapor pressure of water at 20oC is 17.5 torr;
this is relatively low when compared to volatile ethyl
alcohol (43.9 torr) and very volatile ethyl ether (442.2 torr)
 The density of water at 4.0oC is 1.0 g/mL; the
density of ice at 0oC is 0.917 g/mL.
 The specific heat of water is 1.0 cal/g oC or 4.184
J/g oC.
The Unusual Properties of Water
 Water co-exists in all three states of matter naturally
on earth.
 The only common substance is a liquid at STP.
 As a solid, it is less dense than its liquid form, that is
“Ice floats”. Most substances contract upon
solidifying.
 It has a very high Heat Capacity. It stores a large
amount of energy with very little atomic or molecular
motion.
 It requires a lot of heat energy (enthalpy) to change
states.
 It has a high boiling point for such a low molecular
weight compound.
 It is a universal solvent, as a good dissolving medium
a large number of substances are soluble in water.
Water as a universal solvent
• Water is called the universal solvent because of its ability to
dissolve many substances. The general solubility rule is “like
dissolves like”. Since water is a polar molecule it will dissolve other
polar substances as well as ionic compounds. Water will not dissolve
or mix with nonpolar substances therefore water is immiscible in
nonpolar substances.
• Description of how water dissolves an ionic salt (like
NaCl) on the molecular level?
Although the attractive force from the partial charge of a single
polar molecule is not as strong as the charge from an ion, it is
plausible that a multitude of polar molecules could react on a single
ion effectively. The positive end (H+) of several water molecules are
attracted to the negative end of the salt crystal (Cl-) while the
negative end of several water molecules (O2-) are attracted to the
positive end of the crystal (Na+). The ionic bonds of the crystal
are weakened by the solvating effect of the water molecules and
the ions break away from the bulk crystal. The large number of
water molecules in the container prevent the salt ions from recombining.
Why is Water so unusual?
The fundamental explanation for water’s unusual properties relates to the
polarity of its bonds. Polarity describes the partial charge associated with
a bond or molecule. A polar bond or molecule has a charge distribution
present (one end positively charged and the other end negatively charged)
while a nonpolar bond or molecule has no distinct charge distribution
(neutral).
Water is composed of two polar covalent O-H bonds (the difference in
electronegativity is 1.4) arranged in a “bent” molecular geometry. Each
bond has a dipole moment pointing in an overall similar direction leading to
the existence of an overall dipole moment. The oxygen atom pulls the pair
of electrons closer towards itself (making it partially negative) and further
from the hydrogen atoms (making them partially positive).
-
+
This charge distribution allows the partially positive hydrogen atoms from one
molecule to be attracted to the partially negative oxygen atom of another
molecule. This strong interlocking network between neighboring molecules is
called HYDROGEN BONDING. The ability to form strong hydrogen bonds is
the main reason for water’s unusual properties.
PROPERTIES ASSOCIATED WITH WATER
HYDRATES: Solids that contain water molecules as part of
their crystalline structure. The water in the hydrate is
known as the water of hydration or the water of
crystallization.
HYGROSCOPIC: A substance is hygroscopic if it readily
absorbs water from the atmosphere and forms a hydrate.
DELIQUESCENT: A substance is deliquescent if it absorbs
water from the air until it forms a solution.
DESICCANTS: Compounds that absorb water and are used
as drying agents.
EFFLORESCENCE: The process by which crystalline
materials spontaneously lose water when exposed to air.
Water and the Changes of State
The energy required to heat (or cool) a solid (or heat/cool a liquid
or a gas) can be calculated using q = msDT. It requires
additional energy to change states. The energy required to
convert a specific amount of the solid to a liquid is known as the
heat of fusion (q = DHfus) and the energy required to convert a
specific amount of a liquid to a gas is the heat of vaporization (q
= DHvap).
Temperature oC
The total amount of energy can be calculated from qT = q1 + q2 +
q3...
Heating curve for water
Water and the Changes of State
Q. How many kilojoules of energy are needed to change 15.0 g of ice at
-5.00oC to steam at 125.0 oC?
The first step is to design a pathway:
q1 = msDT for ice from -5.0 to 0.0 oC, the specific heat of ice is 4.213 J/g oC
q2 = DHfus for ice to liquid at 0.0oC
q3 = msDT for liquid 0.0oC to 100.0 oC
q4 = DHvap for liquid to steam at 100.0oC
q5 = msDT for steam 100.0 to 125.0 oC; the specific heat of steam is 1.900 J/g oC
so qT = q1 + q2 + q3 + q4 + q5
The next step is to calculate each q:
q1= (15.0 g) (4.213 J/g oC) (0.0 - (-5.0) oC) = 316 J
q2 = (335 J / g) (15.0 g) = 5025 J
q3= (15.0 g) (4.184 J/g oC) (100.0 - (0.0) oC) = 6276 J
q4 = (2260 J / g) (15.0 g) = 33900 J
q5= (15.0 g) (1.900 J/g oC) (110 - 100 oC) = 285 J
qT = 316 J + 5025 J + 6276 J + 33900 J + 285 J = 45.8 kJ
1.
PRACTICE
PROBLEMS
#21a
Which contains less heat, ice at 0 C or water at 0 C? Explain your
o
o
answer.
Ice at 0oC contains less heat than liquid water at the same temperature.
Heat must be added to convert ice to water, so the water will contain that
much more additional heat energy. Also the liquid state is in motion much
more than the solid state. An increase in motion can only be accomplished
by an increase in energy.
2. On the basis of KMT, explain why vapor pressure increases with
temperature.
According to the kinetic molecular theory, the vapor pressure of a liquid should
increase with temperature because of the increase in collisions and kinetic
energy that always accompanies an increase in heat energy (temperature).
KEm = 3/2 RT. The increase in energy thus motion allows the liquid molecules to
escape (overcome the surface tension and other cohesive forces maintaining
the liquid state) from the surface of the liquid into the gas phase.
3. Write equations to show how the following metals react with water.
a) aluminum
a)
b)
c)
d)
b) calcium
Al (s) + 3H2O (g)  3H2 (g) + Al2O3
Ca (s) + 2H2O  H2 (g) + Ca(OH)2
2K (s) + 2H2O  H2 (g) + 2KOH + heat
3Fe (s) + 4H2O (g)  4H2 (g) + Fe3O4
c) potassium
d) iron
*requires steam
*slowly at ambient temperature
* vigorous at ambient temperature
*requires steam
PRACTICE PROBLEMS #21b
1. Explain the physical process of boiling.
See next slide for essay/answer
2. Why does ice float in water?
3. Why does water have a relatively high boiling point?
4. Explain if ice will float in ethyl alcohol (d = 0.789 g/L)?
5. How much energy is needed to change 62.74 g of water at
15.00oC to steam at 103.0 oC?
1.645 x 105 J or 3.931 x 104 cal
6. Magnesium carbonate, MgCO3, forms a hydrate containing
39.1 % water of hydration. Calculate the formula of this
MgCO3 . 3 H2O
hydrate.
1. Explain the physical process of boiling.
At room temperature the water molecules have
enough energy to allow the particles to move past each
other but not enough to escape the surface tension.
As the temperature of water increases, the heat energy
(from the burner) is transferred to kinetic energy (for
the molecules) leading to an increase in the molecular
motion of the molecules. This action results in an
increase in the vapor pressure above the surface of the
liquid. When the vapor pressure of the water equals
the external pressure, boiling begins. Now a sufficient
amount of the molecules have enough energy to resist
the attractive forces. Bubbles of vapor are formed
throughout the liquid and these bubbles rise to the
surface to escape.
2. Why does ice float in water?
Ice floats in its own liquid due to the intermolecular force,
hydrogen bonding. As water freezes, the molecular motion
of the molecules slow down and the partial positive end
(hydrogen) of one water molecule is attracted to the partial
negative end (oxygen) of another water molecule.
Combine this event with the bent shape of water and the
molecules become arranged in a 3-D hexagonal array.
This array creates pockets of vacuum (empty space) in the
lattice structure as well as a decrease in the number of
molecules per unit volume. The mass is directly related to
the number of molecules therefore, in the solid state, since
there are less particles then there must be less mass per
unit volume therefore the solid is less dense than the
liquid.
3. Why does water have a relatively high boiling point?
Water has a relatively high boiling point because of the
amount of intermolecular forces present. Water
experiences LDF (London Dispersion Forces) and d-d
(dipole-dipole) forces, along with the additional attractive
force, Hydrogen bonding. A large amount of heat energy is
required to break all of these forces in order for a phase
transition to occur, thus the high boiling point.
4. Explain if ice will float in ethyl alcohol (d = 0.789 g/L)?
Ice would not float in pure ethyl alcohol because the
density of water is 1.000 g/mL which is greater than 0.789
g/mL for ethyl alcohol. Yet since ethyl alcohol also
undergoes a small degree of hydrogen bonding, the
sinking effect is not as dramatic as it would be with a
nonpolar substance.
GROUP
STUDY
PROBLEMS
#21
Short Essay
1. Can ice be colder than 0.0oC? Justify your answer.
2. Why does a boiling liquid maintain a constant temperature when
heat is continually being added?
3. Why does a lake freeze from the top down?
Math
1. Suppose 50.0 g of ice at 0.0oC are added to 285g of water at 22.0oC.
Is there sufficient ice to lower the temperature of the system to 0.0oc
and still have ice remaining? Show all work.
2. A mixture of 70.0 mL of hydrogen and 50.0 mL of oxygen is ignited to
form water. Does any gas remain unreacted?
3. A 25.0 g sample of a hydrate of FePO4 was heated until all the water
was driven off. The mass of anhydrous sample is 16.9 g. What is
the formula of the hydrate?