Chapter 5: Water for Life
“Water has never lost its mystery. After at least two and a half
millennia of philosophical and scientific inquiry, the most vital of
the world’s substances remains surrounded by deep uncertainties.
Without too much poetic license, we can reduce these questions to
a single bare essential: What exactly is water?”
Philip Ball, in Life’s Matrix: A Biography of Water,
University of California Press,
Berkeley, CA, 2001, p. 115
Do you know where your drinking water comes from?
Do you know if your drinking water is safe to drink?
How would you know?
Different Representations of Water
Lewis structures
Space-filling
5.1
Water is a very unique molecule.
It has a very high boiling point for such a
small molecule
It is an excellent solvent for many types of
compounds
The solid form is less dense than the liquid
(a very rare property)
Has a very high heat capacity

The properties of water are due to:

It's molecular geometry

It's small size

And the type of bonds it contains
Bond – an attractive force that holds two
atoms together.
Atoms bond to obtain a more stable
electronic configuration.
They do this by gaining, losing, or sharing
electrons with other atoms
Covalent (molecular) bonds
 Between non-metals and non-metals
 Atoms tied together by sharing electrons
 Forms molecules – fixed numbers of atoms
in a particular geometry
Ionic bonds
 Between metals and non-metals
 Attraction between oppositely charged
atoms/molecules (ions)
 Formula shows the ratio of ions


Metallic bonds
 Between two or more metals
 Outer electrons are not linked to a
particular atom, as in covalent and ionic
bonding
 These electrons are shared between all
atoms in a 'sea of electrons'
 This is why metals conduct electricity and
heat so well
For Covalent and Ionic Bonding:
The electronegativity determines what type
of bond will form between two atoms
Electronegativity (EN)– attraction for shared
electrons
Fluorine is the most EN element
Francium the least EN element
Larger EN – stronger attraction for
bonding electrons
Covalent Bonds
There are two types of covalent bonds:
Non-polar (covalent)
Bonding electrons are equally shared
When a non-metal bonds to itself, or to
another non-metal with a similar EN
Polar (covalent)
Bonding electrons are shared, but not
equally
One atom has a larger EN than the other
Covalent Bonds
Polar (covalent)
Usually between non-metals two or more
spaces apart on the periodic table
The atom with greater EN (closer to F)
pulls harder on the shared electrons
This pull creates a polarity, or dipole,
across the bond:
The atom with the higher EN has a
slight negative charge
The other has a slight positive charge
A difference in the electronegativities
of the atoms in a bond creates a
polar bond.
O
H
H
Partial charges result from
bond polarization.
A polar covalent bond is a
covalent bond in which the
electrons are not equally shared,
but rather displaced toward the
more electronegative atom.
5.1
Polarity of hydrogen covalent bonds is difficult
to tell from the PT:
H-C – non-polar (some books list as very
slightly polar)
H-O – very polar
H-N – very polar
H-F – very polar
Ionic Bonds
If the electronegativity difference is large
enough between two atoms, they will not
share the 'bonding' electrons:
The atom with the greater EN takes the
'bonding' electrons from the other atom
This atom then has a negative charge
The atom that lost the electron(s) has a
positive charge

Ions
Cations – lost one or more electrons
 positively charged

Anions – gained one or more electrons
 negatively charged

When forming ions, atoms usually want to get to
the same number of electrons as the nearest
Noble Gas

Group
Ion usually
formed
1 (1A)
+1
H can form -1 ion
2 (2A)
+2
13 (3A)
+3
14 (4A)
+/-4
15 (5A)
-3
16 (6A)
-2
17 (7A)
-1
18 (8A)
0
Naming Ions:
●Cation (metal) – name is the same as the
element, + 'ion'
●Fixed charge cations – metals that only form one
cation (such as Group 1 and 2 metals):
Li+1 → lithium ion, Ca+2 → calcium ion
●Variable charged cations – metals that may form
different cations (most transition metals)Use
Roman numerals to show the charge:
Fe+2 → iron (II) ion
Fe+3 → iron (III) ion
●Anion (non-metal) – use the root of the
element name, change the ending to 'ide', +
'ion':
S → S-2
sulfur → sulfide ion
N → N-3
nitrogen → nitride ion
O → O-2
oxygen → oxide ion
Naming Binary Ionic Compounds:
●List the cation first, then the anion
●Do not include 'ion' in the name
●Names must be distinctive, in order to
distinguish between similar compounds, such as
with variable-charged metals
NaCl – sodium chloride
CaF2 – calcium fluoride
FeI2 – iron (II) iodide
FeI3 – iron (III) iodide
Writing formulas for binary ionic compounds:
●The formula shows a ratio of one ion to the other.
●The ionic charges must cancel out so that the
overall charge is neutral
●Always list the metal first, then the non-metal
●Select subscripts to balance charges
●Reduce subscripts if needed to obtain the lowest
whole number ratio between ions
To determine the charge on a variable charge cation,
treat the formula as an algebraic expression:
To determine the iron charge in Fe2O3
●let Fe = x and O = y (x and y are ionic charges)
●the charges of the ions must add up to the overall
charge, which is 0 in this case, so
2x + 3y = 0
●we know that y = -2 (oxide ion)
2x + 3 (-2) = 0
x = +3
●so Fe2O3 is named iron (III) oxide
H
H
H2 has a nonpolar
covalent bond.
A water molecule is polar – due to
polar covalent bonds and the shape of
the molecule.
NaCl
NaCl has an ionic
bond – look at the
EN difference.
Na = 1.0
Cl = 2.9
EN = 1.9
5.1
Polyatomic Ions
These are covalently bonded atoms with an overall
charge (an ionic molecule):
NO3-1 – nitrate ion
ClO3-1 – chlorate ion
C2H3O2-1 – acetate ion
OH-1 – hydroxide ion
SO4-2 – sulfate ion
CO3-2 – carbonate ion
PO4-3 – phosphate ion
H3O+1 – hydronium ion
NH4+1 – ammonium ion (NH3 – ammonia)
Intermolecular Forces



Polar bonds can result in polar molecules.
For molecules like CO2, the polar bonds
cancel each other out.
For other molecules, like water, the polar
bonds cause slight positive and negative
ends on each molecule.
Intermolecular Forces




The dipoles on one molecule are attracted
to the dipoles on other molecules.
This is an example of intermolecular
attractive force.
Water molecules are extremely polar, and
so have strong intermolecular attraction.
This is why water has such a high boiling
point. N2 is a heavier molecule, but with
little intermolecular attraction, it's boiling
point is 300°C lower than that of water
Polarized bonds
allow hydrogen
bonding to occur.
A hydrogen bond is an electrostatic attraction between an atom
bearing a partial positive charge in one molecule and an atom
bearing a partial negative charge in a neighboring molecule. The
H atom must be bonded to an O, N, or F atom.
Hydrogen bonds typically are only about one-fifteenth as strong
as the covalent bonds that connect atoms together within
molecules.
H–bonds are intermolecular bonds.
Covalent bonds are intramolecular bonds.
5.2
Intermolecular Forces



The dipole/dipole interaction is like a weak
ionic bond.
For this reason, water is also able to
dissolve many ionic compounds.
Several water molecules can surround ions
and bring them into solution.
Substances that will dissociate in
solution are called electrolytes.
Ions are simply charged
particles – atoms or groups of
atoms.
They may be positively
charged – cations.
Or negatively charged –
anions.
Dissolution of NaCl in Water
NaCl(s)
H2O
Na+(aq) + Cl–(aq)
The polar water molecules stabilize the ions
as they break apart (dissociate).
5.8
Intermolecular Forces



Polar bonds can result in polar molecules.
For molecules like CO2, the polar bonds
cancel each other out.
For other molecules, like water, the polar
bonds cause slight positive and negative
ends on each molecule.
Water on Earth

Only 3% of all water is fresh (potable)

Of this:

68% is in glaciers

30% is underground

1% in the atmosphere

only 0.3% in lakes, rivers, streams
Water Footprint


The average person needs 1E6 (1 million)
liters per year.

This is equivalent to 250000 gallons,
half of an Olympic sized pool.
Some of this water is used directly, and
some indirectly.
Directly used water includes:





Drinking water
Bathing water
Water used for washing dishes
Water used for washing clothes
Toilet water
Indirectly used water includes:




Water for crops
Water for livestock
Water needed for services,

Electrical power

Waste treatment
Water needed for industry

Production of consumer goods

Construction
Water Footprint
Water is necessary to produce food:
5.3
Water Footprint
Water is necessary for products:
5.3
Fresh Water


Surface water – Lakes, rivers, streams

Easily accessible

Not abundant enough to meet our
needs

May need to be filtered/treated to drink
Ground water – Underground in aquifers
(trapped in geological formations)

Harder to access

More abundant

Often drunk without treatment
The average American uses
almost 100 gallons of water a day.
Nearly ¾ of the water entering
our homes goes down the drain.
Much of our clean water comes from
underground aquifers.
The Ogallala Aquifer is shown in dark blue.
While normally free of pollutants, groundwater can be contaminated
by a number of sources:
Abandoned mines
Runoff from fertilized fields
Poorly constructed landfills and septic systems
Household chemicals poured down the drain or on the ground
5.4
Salt Water

Very abundant (97% of all water on Earth)

Easily accessible, at least near coasts

Not potable due to high salt content

Difficult to purify for human consumption
Access to safe drinking water varies widely across the world.
5.4
Solutions




Solutions are one pure substance
(compound or element) dissolved in
another
Solute – what is being dissolved (minor
component)
Solvent – what the solute is dissolved in
(major component)
Solution – solute/solvent mixture




Solutions
Usually, solvents dissolve solutes of similar
chemical structure (“like dissolves like”)
Non-polar solvents dissolve non-polar
solutes: oil and gasoline
Polar solvents dissolve polar solutes: water
and sugar
But oil and water will not mix
Covalent molecules in solution
A sucrose molecule – when dissolved in water,
sugar molecules interact with and become
surrounded by water molecules, but the sucrose
molecules do not dissociate like ionic
compounds do; covalent molecules remain
intact when dissolved in solution.
They will not conduct electricity; they are
nonelectrolytes.
5.9
Like dissolves like
5.9
Solutions


Water is also able to dissolve many ionic
compounds due to it's strong polarity
Even 'insoluble' solutes may be very slightly
soluble
When ions (charged particles) are in aqueous solutions,
the solutions are able to conduct electricity.
(a) Pure distilled water (nonconducting)
(b) Sugar dissolved in water (nonconducting): a nonelectrolyte
(c) NaCl dissolved in water (conducting): an electrolyte
5.6
Pure water – very poor electrical conductor
Water – very good solvent
Therefore, pure water is hard to obtain
Dissolved ions make water a good
conductor
Any soluble ionic compound splits apart
into ions in solution
Simple generalizations about ionic compounds
allow us to predict their water solubility.
*Insoluble means that the compounds have extremely low solubility in water (less than 0.01 M).
All ionic compounds have at least a very small solubility in water.
5.8
Concentration




A ratio of solute to solvent
Independent of the amount of solution
For very small concentrations, often
measured in ppm or ppb
For aqueous solutions, we often use
Molarity
Concentration Terms
Parts per hundred (percent)
20 g of NaCl in 100 g of water is a 20% NaCl solution
Parts per million (ppm)
Parts per billion (ppb)
2 g Hg
2 10-6 g Hg 2 g Hg
2 ppb Hg


9
3
110 g H 2O 110 g H2O 1 L H2O
5.5
Molarity (M)
moles of solute
Molarity=
liters of solution
n
M=
V
Molarity makes conversion between moles
and volume easy to calculate.
How to prepare a 1.00 M NaCl solution:
solute
M = Lmol
of solution
Note – you do NOT add
58.5 g NaCl to 1.00 L of
water.
The 58.5 g will take up
some volume, resulting in
slightly more than 1.00 L
of solution – and the
molarity would be lower.
5.5
Maximum Contaminant Level (MCL)



Maximum allowed concentration for
minimal risk to humans
Depends on the contaminant:

heavy metals, solvents, other
chemicals
The MCL for a contaminant may be set by
the federal government, state, county, or
city.
Maximum Contaminant Level Goal (MCLG)
and Maximum Contaminant Level (MCL)
5.10
Nitrate concentrations from California domestic
groundwater wells and agricultural irrigation
5.10
Typical steps used to treat fresh water from
rivers/streams/aquifers for drinking:



Filter particulates
Use flocculating agents to remove smaller
particles (traps particles together for
filtration)
Treat to remove/kill bacteria

Chlorination

or Ozone

or UV irradiation
Schematic drawing of a typical municipal water treatment facility.
5.11
Chlorination




Most common method to treat water
Uses hypochlorous acid (HClO), the same
compound used in swimming pools
Kills microorganisms
Keeps treating from facility, through pipes,
up to the point of use (home/factory)
Ozone




Works well to kill microorganisms
Expensive
Short lifetime, quickly decomposes
Does not protect water after leaving the
treatment facility
UV Irradiation




UV light used to kill microorganisms
Easy to implement
Cheaper than O3
Only treats water at exposure, so offers no
protection after leaving treatment facility
Purification of water – removing all particulates,
organisms, dissolved ions/compounds


Distillation – boil water and condense steam
to obtain pure water (very minor impurities)
Purifiers/filters – series of filters/membranes

Each filter removes certain types of
contaminants

Expensive, not practical for large volumes
Making freshwater from saltwater
Desalinization – a process that removes ions from saltwater
5.12
Making freshwater from saltwater – continued
Distillation – a separation process in which a liquid solution is
heated and the vapors are condensed and collected
Either perform distillation in laboratory (left) or use solar power (right).
5.12

Reverse osmosis – produces very pure
water

Water is forced through a special
membrane that only allows very small
molecules to penetrate

Requires a lot of energy

Not currently practical for large
volumes
Making freshwater from saltwater – continued
Osmosis – the passage of water through a semipermeable membrane from a
solution that is less concentrated to a solution that is more concentrated
Reverse Osmosis – uses pressure to force the movement of water through a
semipermeable membrane from a solution that is more concentrated to a
solution that is less concentrated
5.12
LifeStraw – created for developing countries to remove bacteria,
viruses, and parasites from water to use for drinking
5.12
Water, water, every where,
And all the boards did shrink;
Water, water, every where,
Nor any drop to drink.
And every tongue, through utter drought,
Was withered at the root;
We could not speak, no more than if
We had been choked with soot.
The Rime of the Ancient Mariner, excerpt
Samuel Taylor Coleridge