8.4 Water

1
8.4 Water

• Focus 1 Water is distributed on Earth as a
  solid, liquid and gas

2
The significance of water on Earth

• Necessary for living things
• A transport medium for nutrients in cells
• A raw material for plants in photosynthesis
• A solvent for nutrients and O2 in blood
• A solvent for wastes (e.g. CO2, sweat)

• Alters landforms
• Moving water in rivers forms canyons
• Water can erode rocks by dissolving minerals
• Freezing water expands in small cracks creating
  fragments
• Glaciers slowly change the landscapes from
  mountains to the sea

• A natural resource
• Drinking and food preparation
• Washing of clothes, dishes, etc.
• Irrigation of crops
• Recreation
• Transportation (ferries)
• Hydroelectricity

• A habitat for life
• A place where aquatic flora and fauna live (e.g.
  fish, coral, algae, phytoplankton)
• There is much less fluctuation in temperature
  within an aquatic environment than a terrestrial
  habitat. This is due to the high heat capacity of
  water

3
Distribution on Earth

• Biosphere
• Liquid (as a solvent for nutrients, etc.)
• 60-90 in most living things (50-75 in humans)
• Lithosphere
• Solid, liquid or chemically bound as waters of
  hydration (e.g. CaSO4.2H2O)
• Variable percentages in groundwater, aquifers and
  rocks
• Hydrosphere
• Liquid and solid
• Approx. 95 in the oceans and greater in lakes,
  rivers and the polar icecaps
• Atmosphere
• Gas and liquid droplets
• 0-5 in the air, variable depending upon
  environment

4
Solutions

• Solutions are homogeneous mixtures that contain a
  solvent and solute.
• Solvent a substance that dissolves a solute in a
  solution.
• Solute a substance that is dissolved by a
  solvent in a solution
• Examples
• Sea water (solvent-water solute-salt)
• Blood (solvent-water solutes-oxygen, nutrients)

5
Density

• The density of any substance is defined as
• Most substances contract as they cool due to a
  decrease in kinetic energy of the particles. This
  generally leads to an increase in density.
• Water behaves differently. As the temperature of
  water decreases to around 50C, the H-bonds
  arrange themselves so that there is more space
  between water molecules. This means that the
  density drops until solid ice is formed.
• The lower density of ice means aquatic organisms
  can survive under floating sheets of ice.

There is a sharp decrease in the density of water
due to the rigid hexagonal shape that is formed
when water freezes, leaving more space between
molecules.
6
8.4 Water

• Focus 2 The wide distribution and importance of
  water on Earth is a consequence of its molecular
  structure and hydrogen bonding

7
Property Comparisons
The structure of water can be understood by
comparing it with some isoelectronic molecules
(i.e. the same number of electrons)
Water Ammonia Hydrogen Sulfide

Construct Lewis dot structures of water, ammonia,
and hydrogen sulfide to identify the distribution
of electrons
8
Property Comparisons
Water Ammonia Hydrogen Sulfide
Shape of molecule bent pyramidal bent
m.p./b.p.(0C) 0/100 -78/-33 -86/-60
Construct Lewis dot structures of water, ammonia,
and hydrogen sulfide to identify the distribution
of electrons
9
Hydrogen Bonding
ammonia

• Hydrogen bonding occurs between hydrogen atoms
  and unshared pairs of electrons on N, O or F
  atoms.
• This results in an unequal sharing of electrons
  leading to a partial positive charge on the H
  atom.
• These bonds are stronger than dipole-dipole
  forces and dispersion forces.

water
10
Dipole-dipole

• Previously, we defined electronegativity as an
  atoms tendency to attract electrons to itself.
  The top right of the periodic table has the
  highest values. Fluorine is the most
  electronegative atom.
• Polar covalent bonding
• When two atoms that have a difference in
  electronegativity bond with each other, the one
  with the higher value tends to have a partially
  negative charge due to the stronger attraction to
  the electrons. For the same reason, the other
  atom will tend to have a partially positive
  charge. This bond is said to be polar covalent.
• For example, Cl is more electronegative than H.
  This means that Cl will have a slight negative
  charge given the symbol d (Greek letter delta)
  meaning small or slight.

d H --- Cl d-
Chlorine is partially negative due to a higher
attraction to e-
Hydrogen is partially positive due to a lower
attraction to e-
11
Dipole-dipole

• Net dipole
• An HCl molecule has an obvious negative end and
  positive end. However, not all molecules with
  polar covalent bonds are polar overall, or have a
  net dipole. Some of these molecules contain polar
  bonds that cancel out the effect of any single
  polar bond. For example, BeF2

This molecule contains two polar covalent bonds,
but no overall net dipole.

• Questions
• Draw out the structures for water, ammonia and
  beryllium trifluoride.
• Indicate any polar bonds with the arrow head
  pointing towards the negative end.
• Which of these molecules has a net overall dipole?

12
Dipole-dipole
Answers
No net dipole
Net dipole towards O
Net dipole towards N
13
Dipole-dipole

• Dipole-dipole interactions
• When molecules have an overall net dipole, the
  interactions between them is largely
  electrostatic. As one might expect, the positive
  end of one of these molecules is attracted to
  negative end of another molecule.
• If we imagine a polar molecule as a positive and
  negative end, the interaction between molecules
  becomes simple to visualise.

Note dipole-dipole interactions are generally
not as strong as hydrogen bonds
14
Explaining the properties of water

• Surface Tension
• At the surface, water appears to have a skin
  that resists deformation. This can be seen when
  water striders sit on top of the water (see
  left). This is due to a property known as surface
  tension.
• Water molecules are attracted to one another in
  all directions. This is known as cohesive forces.
  At the surface, the molecules can only be
  attracted to either side and down. This results
  in an overall force in towards the liquid. This
  unbalanced force results in surface tension. This
  also explains why water forms droplets as all of
  the force is directed inwards.

Water molecules attract each other in all
directions except at the surface leading to
surface tension.
Question What are adhesive forces and how do
they help to explain a meniscus in a glass
cylinder?
15
Explaining the properties of water

• Viscosity
• The resistance of a liquid to flow is known as
  its viscosity.
• Viscosity can be high if the molecules of the
  liquid are quite large, for instance in motor
  oil.
• Viscosity is also directly related to the
  strength of the intermolecular forces.
• Water has a relatively high viscosity due to
  strong hydrogen bonding.

• Melting/Boiling point
• Generally, the larger the molecule, the higher
  the melting and boiling point.
• Water has unusually high melting and boiling
  points relative to its low molecular weight.
• As stated previously, hydrogen bonds between
  water molecules are relatively strong.
• Since melting and boiling involve breaking
  intermolecular forces, the strength of hydrogen
  bonds in water leads to higher relative melting
  and boiling points.
• This explains why hydrogen sulphide and other
  similar sized molecules have much lower mp/bp.
  There are no hydrogen bonds. (See slide 8)

16
8.4 Water

• Focus 3 Water is an important solvent

17
Soluble ionic compounds

• Solubility of table salt (NaCl)
• Common table salt (NaCl) is made up of ve Na
  ions and ve Cl ions. Their attraction to one
  another is what holds them together in a lattice.
• Water is polar and has positive and negative ends
  as well. Therefore, water is attracted to the
  charged ions in the salt lattice. The water
  molecules surround the ions one at a time,
  overcoming the forces between the solute
  particles until the salt is dissolved. This
  process is known as dissociation where ions in a
  lattice are dispersed in a solvent.
• It is important to note that not all ionic
  compounds are soluble in water. More on this
  later.

Salt lattice
18
Soluble molecular compounds

• Solubility of Sucrose
• Sucrose is common table sugar and readily
  dissolves in water. As can be seen below, the
  sucrose molecule contains many polar OH groups
  that attract water molecules. This breaks the
  crystal structure of the sugar and distributes
  the individual sugar molecules through the water.
  Normally, the only soluble molecular compounds
  are highly polar or can form hydrogen bonds with
  water.

Sucrose
19
Soluble or partially soluble molecular
elements/compounds

• Covalent molecular elements and compounds are
  held together by either dipole-dipole
  interactions or dispersion forces. These are
  relatively weak compared to the hydrogen bonds in
  water. These species are therefore generally
  insoluble or only slightly soluble.
• Solubility of iodine
• Iodine (I2) is a non-polar substance that are
  held together by weak dispersion forces compared
  to waters strong hydrogen bonds. For this
  reason, iodine is not very soluble in water.
  Rather, it is more soluble in non-polar solvents
  such as liquid bromine (Br2).
• Solubility of oxygen gas
• Many gases such as O2, N2 and CO2 are non-polar
  and, therefore, are not very soluble in water. In
  fact, oxygen has a solubility of 0.004 g in a 100
  g of water.
• Solubility of hydrogen chloride
• Hydrogen chloride is a covalent molecular gas
  that is highly soluble in water. Hydrogen
  chloride molecules undergo a process known as
  ionisation because the HCl reacts with water to
  produce H and Cl ions in solution. Acids produce
  H ions and bases produce OH- ions in water.
• HCl(g) ? H(aq) Cl-(aq)
• Whether a covalent substance is soluble in water
  is largely dependant upon whether it reacts with
  water.

20
Solubility of covalent network

• Solubility of silicon dioxide
• Covalent network compounds such as silicon
  dioxide (sand) have very strong covalent bonds
  that hold the atoms in a rigid lattice.
• The strong covalent bonds in silicon dioxide
  cannot be broken by attraction to water
  molecules.
• Therefore, covalent network compounds such as
  silicon dioxide are not soluble in water.

Sand (mostly silicon dioxide) is insoluble in
water. Source www.thesand.net/ csalva/Index.htm
21
Solubility of large molecules

• Solubility of polyethene
• Polyethene is a polymer, meaning that it has
  repeating units called monomers.
• It is used to make plastic drink bottles and
  other containers.
• The diagram below shows the repeating unit for
  this molecule.

• Solubility of cellulose
• Cellulose is a very large molecule that contains
  repeating units of C6H10O5.
• There are polar OH groups on this molecule, but
  the long chains lie beside each other and the OH
  groups are involved in hydrogen bonding.
  Therefore, the OH groups are not accessible to
  the water molecules.

Polyethene is non-polar and is therefore
insoluble in water.
Cellulose is insoluble in water.
In general, the larger the molecule, the less
likely it will be soluble in water unless there
are a large number of accessible polar groups
(i.e. OH, NH2, COOH)
22
Water as a solvent summary

• Water is a polar molecule, with a slight negative
  charge on the oxygen end and a slight positive
  charge on the hydrogen end.
• Like dissolves like polar solvents (e.g. water)
  tend to dissolve polar solutes and non-polar
  solvents (e.g. hexane) tend to dissolve non-polar
  solutes.
• Dissolution is a process where ions in a solid
  are dispersed in a solvent such as water e.g.
  NaCl(s) ? Na(aq) Cl-(aq)
• Ionisation is a process where a covalent molecule
  reacts with water to form ions in solution e.g.
  HCl(g) ? H(aq) Cl-(aq)
• Covalent network molecules are highly insoluble
  in water due to strong covalent lattice
  structures.
• Large molecules tend to be insoluble in water
  unless they have a large number of accessible
  polar sites.
• Most non-polar molecules are insoluble or are
  only slightly soluble in water e.g. O2 and I2.

23
8.4 Water

• Focus 4 The concentration of salts will vary
  according to their solubility, and precipitation
  can occur when the ions of an insoluble salt are
  in solution together

24
Solubility of ionic compounds

• Not all ionic compounds are soluble in water. If
  the ionic bonds are stronger than the interaction
  with water, the compound will be insoluble
• If two solutions of different ionic compounds are
  mixed, an insoluble compound may form. This
  insoluble compound falls to the bottom of the
  container and is called a precipitate.
• For example
• barium nitrate(aq) sodium sulphate(aq) ? barium
  sulphate(s) sodium nitrate(aq)

(solution) (solution)
(precipitate)
(solution)
Activity Write a net ionic equation for this
reaction.
Ba2(aq) SO42-(aq) ? BaSO4(s)
25
Solubility Rules for common saltssummary

• Soluble Ionic compounds Important
  exceptions
• Compounds containing NO3- None
• C2H3O2- (acetate) None
• Cl- Cmpds of Ag, Hg22, Pb2
• Br- Cmpds of Ag, Hg22, Pb2
• I- Cmpds of Ag, Hg22, Pb 2
• SO42- Cmpds of Sr2, Ba2, Hg22, Pb2
• Insoluble Ionic compounds Important exceptions
• Compounds containing S2- Cmpds of NH4, IA
  cations, Ca2, Ba2
• CO32- Cmpds of NH4, IA cations
• PO43- Cmpds of NH4, IA cations
• OH- Cmpds of IA cations, Ca2, Ba2

26
Some precipitation examples
Aluminium ion reacts with aqueous ammonia to
produce a white gelatinous precipitate of
Al(OH)3 Al 3(aq) 3NH3 (aq) 3H2O (aq) ??
Al(OH)3(s) 3NH4 (aq)
Aluminium hydroxide
Copper hydroxide
Lead iodide
Activity write out possible reactions that could
have formed the two precipitates above.
27
Equilibrium

• We normally think of chemical reactions as
  reactants forming products (i.e. from left to
  right). However, sometimes the products can also
  react to form the reactants (i.e. right to left).
  These are known as reversible reactions. We use a
  double arrow, such as the one below, to indicate
  these reactions.
• In reversible reactions, both are taking place
  simultaneously. Once the forward and reverse
  reactions are happening at the same rate, we have
  what is known as dynamic equilibrium. At this
  stage there are no observable changes in the
  system.
• Essentially all reactions are somewhat
  reversible, but some favour one direction so
  strongly that they seem to go in one direction
  only.
• Equilibrium is also possible where a vapour is in
  equilibrium with a liquid or when a precipitate
  is in equilibrium with a saturated solution.

28
Saturated solution equilibrium

• Recall that a saturated solution is when the
  maximum amount of solute is dissolved in a given
  amount of solvent.
• Solution equilibria
• In a saturated solution, the dissolved solute is
  in dynamic equilibrium with any undissolved
  solid. Any excess solid that is added to the
  system will not affect the dynamic equilibrium.
  The rate of crystallisation and dissolution will
  be the same.

In a saturated solution, dissolved and
undissolved solute particles are in dynamic
equilibrium.
29
Concentration of solutions

• For quantitative analysis it is important to know
  exactly how much solute is dissolved in a given
  amount of solution. This is known as the
  concentration of the solution. There are many
  ways of expressing concentration depending upon
  the nature of the solution and the information
  that is available.
• Some ways of expressing concentration are
• g/L mass per unit volume of solution (normally
  g/L)
• (w/v) meaning percent weight per volume of
  solution (g/100mL)
• (v/v) meaning percent volume per volume of
  solution (mL/100mL)
• (w/w) meaning weight percent (g/100g)
• ppm or parts per million (g/million g solution)
• Molarity (M) or moles per litre of solution
• Questions
• Which of these concentrations would be most
  useful when the solute is a liquid?
• Which units would be most useful if the
  concentration was very low, such as heavy metal
  concentrations in drinking water?
• What if we wanted to calculate quantities used in
  chemical reactions?

30
Molarity

• Recall that molarity (M) (read as molar) is
  defined as
• M moles of solute / L of solution
• or
• M n/V
• Remembering that the number of moles is
• Moles mass / molar mass
• We use molarity in the laboratory because the
  basic unit of measure in chemical reactions is
  the mole. We give concentrations for solutions in
  mol/L. Lets consider some examples.

31
Molarity examples
Example 1.  What is the molarity of a 5.00 liter
solution that was made with 10.0 moles of  KBr ?

• Solution  Note that in this particular example,
  where the number of moles of solute is given, the
  identity of the solute (KBr) has nothing to do
  with solving the problem.
•                     of moles of soluteMolarity
  ----------------------                  Liters
  of solution
• Given  of moles of solute 10.0
  moles           Liters of solution 5.00 liters
•                   10.0 moles of KBrMolarity
  --------------------------  2.00
  M                 5.00 Liters of solution
• Answer 2.00 M

32
Molarity examples
Example 2.  What is the volume of 3.0 M solution
of NaCl made with 526g of solute?

• Solution
• First find the molar mass of NaCl.
• Na 23.0 g x 1 ion per formula unit 23.0 g
• Cl 35.5 g x 1 ion per formula unit  35.5
  g                                                
  ----------                                      
              58.5 g
• Now find out how many moles of NaCl you have
•                      mass of sample of moles
  -----------------                      Molar
  mass
• Given  mass of sample 526 g           Molar
  mass 58.5 g
•                                    526 g
  of moles of NaCl ------------                  
                 58.5 g
• Answer  of moles of NaCl 8.99 moles
•  

• Finally, go back to your molarity formula to
  solve the problem
•                               of moles of
  soluteLiters of solution    --------------------
                                       Molarity
• Given  of moles of solute 8.99 moles
  Molarity of the solution 3.0 M (moles/L)
•                                       8.99
  moles of Liters of solution
  -------------                                  
  3.0 moles/L
• Final Answer 3.0 L

33
Dilutions

• It is often important in the laboratory to dilute
  solutions to obtain a desired concentration. This
  is known as a dilution and can be accomplished
  using a straightforward calculation.
• M1V1 M2V2
• Where M1 is the initial concentration in mol/L
  and V1 is the initial volume. M2 and V2 are the
  final concentration and volume respectively.
• Example
• What volume of 25M NH3 is needed to prepare 500mL
  of 1M solution?
• M1V1 M2V2
• 25mol/L X V1 1mol/L X 0.5L
• V1 0.02L
• V1 20 mL

34
8.4 Water

• Focus 5 Water has a higher heat capacity than
  many other liquids

35
Specific Heat Capacity

• Some materials resist changes in temperature more
  than others. Water is a substance that has a high
  resistance to changes in temperature.
• This can be seen on a hot summer day. The air
  temperature may increase to over 40OC, while the
  temperature of a large pool of water will rise
  only slightly.
• This resistance to changes in temperature is due
  to water having a large specific heat capacity
  (C). In other words, water can absorb a large
  amount of heat energy before changing temperature.

On a hot summer day, water stays cool due to a
high specific heat capacity.
Specific heat capacity (C), is the amount of heat
energy in Joules (J), required to raise the
temperature of 1g of a substance by 1 kelvin (K).
36
Specific Heat Capacity
Substance Specific Heat Capacity J K-1g-1 Substance Specific Heat Capacity J K-1g-1
Water 4.18 Ethanol 2.41
Ethylene glycol 2.39 Hexane 2.26
50/50 water/E. glycol 2.86 Chloroform 0.96
Acetone 2.17 Aluminium 0.90

• Compared to other solvents, the relatively high
  specific heat of water makes it very useful in
  the removal of heat energy.
• Radiator coolant
• Water can absorb 10 times as much heat as cast
  iron and five times as much heat as aluminium.
• Ethylene glycol has a lower heat capacity than
  water, but it lowers the freezing point and
  raises the boiling point of the coolant making it
  a useful additive to car cooling systems.

Aquatic ecosystems Bodies of water maintain
stable temperatures allowing aquatic organisms to
thrive. These organisms can only survive and
reproduce within very narrow temperature ranges.
Living systems/Climate The water in living cells
is used to regulate temperature in organisms due
to waters high heat capacity and high thermal
conductivity (ability to remove heat). Climates
are moderated by large bodies of water acting as
thermal reservoirs.
37
Energy Changes

• Recall that changes in temperature are related to
  changes in the amount of heat energy that is
  being absorbed or released. This change in heat
  energy, or enthalpy (H), can be calculated using
  the following equation

? H m C ?T
Change in Temperature (final initial)
Heat energy released or absorbed in Joules (J)
Mass (g)
Specific heat Capacity
Example What quantity of energy is required to
raise the temperature of 0.5L of water from 200C
to 1000C? ? H m C ?T 500g X 4.18 J
g-1K-1 X 800C or K 167,200 J
167 kJ NB 0C K - 273
38
Energy changes in chemical reactions

• In addition to heat changes in a particular
  substance, it is also possible to measure the
  changes in heat content in chemical reactions. In
  a chemical reaction, the change in enthalpy is
  known as the heat of reaction

? Hrxn H (products) H (reactants)

• From this equation, we can see that there are two
  possibilities for the heat of reaction values.
  These are
• The enthalpy of the reactants is higher than the
  enthalpy of the products.
• The enthalpy of the products is higher than the
  enthalpy of the reactants.
• Considering the above information,
• Which will have a negative value and which will
  have a positive value.
• Which will result in a rise in temperature and
  which a fall in temperature.

39
Energy changes in chemical reactions

• Exothermic reactions
• The enthalpy of the reactants is higher than the
  products.
• Reactants
• Products

• Endothermic reactions
• The enthalpy of the reactants is lower than the
  products.
• Products
• Reactants

? H is -ve
? H is ve
Heat absorbed
Heat released
Enthalpy (H)
40
Calorimetry

• What is calorimetry
• Water can be used to measure the change in heat
  energy in a chemical reaction due to its ability
  to absorb heat. This is known as calorimetry.
• A calorimeter is a device that is used to measure
  the enthalpy change that occurs during a chemical
  reaction.

• Measuring heat in the laboratory
• If a known quantity of water is placed in the
  calorimeter and a reaction is carried out, the
  change in temperature due to the reaction is
  transferred to the water and is measured using a
  thermometer placed in a hole in the lid.
• The heat of reaction is
• ? Hrxn -m C ?T
• The amount of heat lost or gained in the reaction
  is equal in size but opposite in sign to the
  amount of heat lost or gained by the water, thus
  the negative sign in the above equation.

Coffee cup calorimeter The simplest calorimeter
makes use of two polystyrene (Styrofoam) cups,
one inside the other with a lid on top. This
minimises the amount of heat lost to the
surroundings. This is the most significant source
of error.
41
Heat of solution

• When an ionic substance dissolves, it may be
  endothermic (absorbs heat/gets cold) or it may be
  exothermic (releases heat/warms up). The change
  in heat (enthalpy) when one mole of a solute
  dissolves in a solvent is called the molar heat
  of solution, ? Hsoln. The values for these
  dissolutions are given in kJ/mol.
• Energy changes involved in dissolution
• Breaking the bonds that hold ions together
  requires energy (endothermic)
• Forming bonds with water (hydration) releases
  energy (exothermic)
• So,
• ? Hsoln heat absorbed when breaking bonds
  heat released when forming bonds
• Exothermic examples (when the solvent-solute
  interactions are stronger)
• NaOH(s) CaCl2(s) H2SO4(l) HCl(g)
• Endothermic examples (when the solvent-solute
  interactions are weaker)
• NH4Cl(s) KCl(s) NH4NO3(s)

42
Calorimetryexperiment
Procedure

1. Prepare the coffee cup calorimeter by placing one
   polystyrene cup inside another. Trim off the top
   of a third cup to make a lid. (See diagram)
2. Make two holes in the lid for the thermometer and
   stirrer just big enough to fit (no gaps).
3. Accurately weigh 100mL of water and pour into the
   calorimeter. Record initial temperature.
4. Accurately weigh approximately 5g of reagent for
   dissolution.
5. Add the reagent to the calorimeter and
   immediately replace the lid and start stirring.
6. Record the temperature every 10 sec until no
   change is observed.

• Questions
• Calculate the heats of solution (? Hsoln -m C
  ?T) for each reagent in kJ/mol.
• Which dissolution was endothermic/exothermic?
• Explain why the sign is reversed in the above
  calculation.
• What are the sources for error in this procedure?
  How could you reduce them? Calculate the error.
• Find the accepted values for molar heats of
  solution and compare with the experimental values.

43
Calorimetryexperiment (Solutions)

1. ? Hsoln -m C ?T
2. M mass of water
3. C specific heat of water (4.18 J/g.K)
4. ?T final temp initial temp
5. To calculate heat of solution, multiply by the
   reagents molar mass
6. Depends upon reagent (see table)
7. Because we are measuring the quantity of heat
   change of the water and we want to know the heat
   of solution for the reagent. The quantity is the
   same, but the sign is reversed because as one
   gives up heat the other absorbs the same
   quantity.
8. The sources for error include any potential for
   loss of heat to the surroundings (e.g. gaps,
   inadequate insulation). These can be reduced by
   reducing the gaps, using better insulating
   material.
9. Accepted Heats of Solution (kJ/mol) for some
   substances in table at right.

Compound Heat of soln (kJ/mol)
calcium chloride, anyhydrous - 81.3
sodium hydroxide -44.0
sodium chloride 3.9
ammonium chloride 14.8
ammonium nitrate 25.7
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Thermal pollution

• When human activities significantly change the
  temperature of waterways, it is known as thermal
  pollution. A significant change can be as little
  as 20C.
• The most common example of thermal pollution is
  the use of natural bodies of water for industrial
  cooling such as in power plants.
• Thermal pollution can have detrimental effects on
  aquatic life.
• Solubility of oxygen decreases as temperature
  increases. Less than 5 ppm oxygen can threaten
  the survival of fish.
• Sharp increases in temperature may directly kill
  fish and other aquatic organisms.
• High temperatures can prevent the
  development/hatching of fish eggs.

Source http//www.isgs.uiuc.edu/servs/pubs/geobit
s-pub/geobit12/gb12a.htm
45
Water

• Compiled by Robert Slider (2006)
• Please share this resource with others