Water

1
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 earths 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).

2
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.

3
Water

• There are different types of Polluted water
• Pathogenic (disease-causing) microorganisms like
  cholera, typhoid, hepatitis, and dysentery still
  effect over 70 of the worlds 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 VOCs (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.

4
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.

5
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.

6
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
  re-combining.

7
Why is Water so unusual?

• The fundamental explanation for waters 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 waters
  unusual properties.

8
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.
9
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). The total amount of energy can be
calculated from qT q1 q2 q3...
Heating curve for water
Temperature oC
10
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
11
PRACTICE PROBLEMS 21a

• 1. Which contains less heat, ice at 0oC or
  water at 0oC? Explain your answer.
• 2. On the basis of KMT, explain why vapor
  pressure increases with temperature.
• 3. Write equations to show how the following
  metals react with water.
• a) aluminum b) calcium c)
  potassium d) iron

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.
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.
a) Al (s) 3H2O (g) ? 3H2 (g) Al2O3
requires steam b) Ca (s) 2H2O ? H2
(g) Ca(OH)2 slowly at ambient
temperature c) 2K (s) 2H2O ? H2 (g)
2KOH heat vigorous at ambient
temperature d) 3Fe (s) 4H2O (g) ? 4H2 (g)
Fe3O4 requires steam
12
PRACTICE PROBLEMS 21b

• 1. Explain the physical process of boiling.
• 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?
• 6. Magnesium carbonate, MgCO3, forms a hydrate
  containing 39.1 water of hydration. Calculate
  the formula of this hydrate.

See next slide for essay/answer
1.645 x 105 J or 3.931 x 104 cal
MgCO3 . 3 H2O
13
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.

14
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.

15
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.
16
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?