Chapter 05 Water We Drink
Chapter 5:
The Water We Drink
Excessive Water = flld
(Jakarta 2008)
Kevin Carter’s 1994 Pulitzer prize
winning photo of a vulture waiting for
a child to die, so that it will eat it
epitomizes not only the hunger crises
in Sudan but also in the whole of
Africa. (Photo source: Pulitzer)
Water scarcity
(Sudan, Africa)
“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?
While normally free of pollutants, groundwater can be
contaminated by a number of sources:
Abandoned mines
Run off from fertilized fields
Poorly constructed landfills and septic systems
Household chemicals poured
down the drain or on the
ground.
Water distributiln in Indlnesia
http://catharsiscorner.wordpress.com/2009/01/26/peta-airtanah-dunia-sumber-kesejahteraan-dan-potensi-konflik/
Global Water Usage
The average water use in the world:
• 70 % for agricultural needs,
• 8 % for domestic needs and
• 22 % for industry.
• Afghanistan and India >95% of water use for agriculture,
• Britain and Canada > 70% for industry.
• Japan, Indonesia and Brazil 60% of water use for
agriculture,
• the Americans use the 42 per cent for agriculture and 46
percent for industrial use.
Solution
• A solution is a homogeneous mixture of uniform
composition.
• Solutions are made up of solvents and solutes.
– Solvent = Substances capable of dissolving other
substances- usually present in the greater amount.
– Solutes = Substances dissolved in a solvent- usually
present in the lesser amount.
• When water is the solvent, you have an aqueous solution
5.3
The
importance
of water as
a solvent in
our bodies
5.3
Water in the Environment
Concentration Terms
Parts per hundred (percent)
20 g of NaCl in 80 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
110 g H 2O 110 g H2 O 1 L H 2 O
5.4
Molarity (M) = moles solute
liter of solution
[ ] = “concentration of”
1.0 M NaCl solution
[NaCl] = 1.0 M = 1.0 mol NaCl/L solution
Also – this solution is 1.0 M in Na+ and 1.0
M in Cl[Na+] = 1.0 M and [Cl-] = 1.0 M
5.4
What is the concentration (in M and mass
%) of the resulting solution when you add 5
grams of NaOH to 95 mL of water?
95 mL H2O = 95 g H2O mass % : 5 g NaOH/100 g solution
95 mL H2O = .095 L
= 5% NaOH
5 g NaOH = 0.125 moles NaOH
0.125 mole NaOH/0.095 L
= 1.3 M solution of NaOH
5.4
What is the molarity of glucose (C6H12O6) in a
solution containing 126 mg glucose per 100.0 mL
solution?
6.99 x 10-3 M
5.4
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.
5.4
Different Representations of Water
Lewis structures
Space-filling
Chargedensity
Region of partial negative charge
Charge-density
Regions of partial positive charge
5.5
Electronegativity is a measure of an atom’s
attraction for the electrons it shares in a covalent
bond.
On periodic
table, EN
increases
EN Values assigned by Linus Pauling,
winner of TWO Nobel Prizes.
5.5
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.
5.5
H
H
H2 has a non-polar
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
DEN = 1.9
5.5
Polarized bonds
allow hydrogen
bonding to occur.
H–bonds are intermolecular
bonds. Covalent bonds are
intramolecular bonds.
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.
5.6
Forming ions
Na
Na
Na+ ion
Na atom
Cl
Cl atom
+ 1 e-
+ 1 e-
Cl
Cl- ion
5.7
When ions (charged particles) are in aqueous
solutions, the solutions are able to conduct
electricity.
(a) Pure distilled water (non-conducting)
(b) Sugar dissolved in water (non-conducting): a nonelectrolyte
(c) NaCl dissolved in water (conducting): an electrolyte
5.7
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 chargedanions.
Dissolution of NaCl in Water
NaCl(s)
H2O
Na+ (aq) + Cl-(aq)
The polar water molecules stabilize the
ions as they break apart (dissociate).
5.7
Some atoms form more than one stable ion
5.7
Naming simple ionic compounds is easyName the metallic element (cation) first,
followed by the non-metallic element (the
anion) second, but with an –ide suffix.
MgO
Mg is the metal, O is the non-metal
magnesium oxide
NaBr
Na is the metal, Br is the non-metal
sodium bromide
5.7
Ions that are themselves made up of more than
one atom or element are called polyatomic
ions.
NaSO4 (sodium sulfate) dissociates in water to form:
Na+
The sulfate group
stays together in
solution.
and
Sodium ions
Sulfate ions
5.7
Naming polyatomic ionic compounds is also easyName the cation first, followed by the anion second.
MgOH
Mg+ is the cation, OH- is the anion
magnesium
hydroxide
NH4Br
NH4+is the anion, Br- is the anion
ammonium
bromide
5.7
Simple generalizations about ionic
compounds allow us to predict their water
solubility.
Solubility of
Solubility Exceptions
Examples
Ions
Compounds
sodium, potassium,
and ammonium
All soluble
None
NaNO3 is soluble
KBr is soluble
nitrates
All soluble
None
LiNO3 is soluble
Mg(NO3)2 is soluble
chlorides
Most soluble
Silver, some mercury, and
lead chlorides
MgCl2 is soluble
PbCl2 is insoluble
sulfates
Most soluble
Strontium, barium, and lead
sulfate
K2SO4 is soluble
BaSO4 is insoluble
carbonates
Mostly insoluble*
Group IA and NH41
carbonates are soluble
Na2CO3 is soluble CaCO3
is insoluble
hydroxides and
sulfides
Mostly insoluble*
Group IA and NH41
hydroxides and sulfides are
soluble
KOH is soluble Al(OH)3 is
insoluble
*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
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
non-electrolytes.
5.9
Like dissolves like
5.9
Maximum Contaminant Level Goal (MCLG)
and Maximum Contaminant Level (MCL)
5.10
5.10
Hard water contains high concentrations of dissolved
calcium and magnesium ions.
Soft water contains few of these dissolved ions.
A pipe with hard-water scale build up
Not in 6th ed.
Because calcium ions, Ca2+, are generally the largest
contributors to hard water, hardness is usually expressed in
parts per million of calcium carbonate (CaCO3) by mass.
It specifies the mass of solid CaCO3 that could be formed
from the Ca2+ in solution, provided sufficient CO32- ions were
also present:
Ca2+(aq) + CO32–(aq) CaCO3(s)
A hardness of 10 ppm indicates that 10 mg of CaCO3 could be
formed from the Ca2+ ions present in 1 L of water.
Not in 6th ed.
Schematic drawing of a typical municipal water treatment facility
5.11
Getting the lead out:
Schematic of a typical spectrophotometer
Using a plot of
absorbance vs. concentration
Calibration graph
5.12
AAS =
Atomic
absorption
Spectrophotometer
Access to safe drinking water varies widely across the world.
5.14
Two water purification techniques:
Distillation
Reverse osmosis
5.14
The Water We Drink
Excessive Water = flld
(Jakarta 2008)
Kevin Carter’s 1994 Pulitzer prize
winning photo of a vulture waiting for
a child to die, so that it will eat it
epitomizes not only the hunger crises
in Sudan but also in the whole of
Africa. (Photo source: Pulitzer)
Water scarcity
(Sudan, Africa)
“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?
While normally free of pollutants, groundwater can be
contaminated by a number of sources:
Abandoned mines
Run off from fertilized fields
Poorly constructed landfills and septic systems
Household chemicals poured
down the drain or on the
ground.
Water distributiln in Indlnesia
http://catharsiscorner.wordpress.com/2009/01/26/peta-airtanah-dunia-sumber-kesejahteraan-dan-potensi-konflik/
Global Water Usage
The average water use in the world:
• 70 % for agricultural needs,
• 8 % for domestic needs and
• 22 % for industry.
• Afghanistan and India >95% of water use for agriculture,
• Britain and Canada > 70% for industry.
• Japan, Indonesia and Brazil 60% of water use for
agriculture,
• the Americans use the 42 per cent for agriculture and 46
percent for industrial use.
Solution
• A solution is a homogeneous mixture of uniform
composition.
• Solutions are made up of solvents and solutes.
– Solvent = Substances capable of dissolving other
substances- usually present in the greater amount.
– Solutes = Substances dissolved in a solvent- usually
present in the lesser amount.
• When water is the solvent, you have an aqueous solution
5.3
The
importance
of water as
a solvent in
our bodies
5.3
Water in the Environment
Concentration Terms
Parts per hundred (percent)
20 g of NaCl in 80 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
110 g H 2O 110 g H2 O 1 L H 2 O
5.4
Molarity (M) = moles solute
liter of solution
[ ] = “concentration of”
1.0 M NaCl solution
[NaCl] = 1.0 M = 1.0 mol NaCl/L solution
Also – this solution is 1.0 M in Na+ and 1.0
M in Cl[Na+] = 1.0 M and [Cl-] = 1.0 M
5.4
What is the concentration (in M and mass
%) of the resulting solution when you add 5
grams of NaOH to 95 mL of water?
95 mL H2O = 95 g H2O mass % : 5 g NaOH/100 g solution
95 mL H2O = .095 L
= 5% NaOH
5 g NaOH = 0.125 moles NaOH
0.125 mole NaOH/0.095 L
= 1.3 M solution of NaOH
5.4
What is the molarity of glucose (C6H12O6) in a
solution containing 126 mg glucose per 100.0 mL
solution?
6.99 x 10-3 M
5.4
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.
5.4
Different Representations of Water
Lewis structures
Space-filling
Chargedensity
Region of partial negative charge
Charge-density
Regions of partial positive charge
5.5
Electronegativity is a measure of an atom’s
attraction for the electrons it shares in a covalent
bond.
On periodic
table, EN
increases
EN Values assigned by Linus Pauling,
winner of TWO Nobel Prizes.
5.5
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.
5.5
H
H
H2 has a non-polar
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
DEN = 1.9
5.5
Polarized bonds
allow hydrogen
bonding to occur.
H–bonds are intermolecular
bonds. Covalent bonds are
intramolecular bonds.
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.
5.6
Forming ions
Na
Na
Na+ ion
Na atom
Cl
Cl atom
+ 1 e-
+ 1 e-
Cl
Cl- ion
5.7
When ions (charged particles) are in aqueous
solutions, the solutions are able to conduct
electricity.
(a) Pure distilled water (non-conducting)
(b) Sugar dissolved in water (non-conducting): a nonelectrolyte
(c) NaCl dissolved in water (conducting): an electrolyte
5.7
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 chargedanions.
Dissolution of NaCl in Water
NaCl(s)
H2O
Na+ (aq) + Cl-(aq)
The polar water molecules stabilize the
ions as they break apart (dissociate).
5.7
Some atoms form more than one stable ion
5.7
Naming simple ionic compounds is easyName the metallic element (cation) first,
followed by the non-metallic element (the
anion) second, but with an –ide suffix.
MgO
Mg is the metal, O is the non-metal
magnesium oxide
NaBr
Na is the metal, Br is the non-metal
sodium bromide
5.7
Ions that are themselves made up of more than
one atom or element are called polyatomic
ions.
NaSO4 (sodium sulfate) dissociates in water to form:
Na+
The sulfate group
stays together in
solution.
and
Sodium ions
Sulfate ions
5.7
Naming polyatomic ionic compounds is also easyName the cation first, followed by the anion second.
MgOH
Mg+ is the cation, OH- is the anion
magnesium
hydroxide
NH4Br
NH4+is the anion, Br- is the anion
ammonium
bromide
5.7
Simple generalizations about ionic
compounds allow us to predict their water
solubility.
Solubility of
Solubility Exceptions
Examples
Ions
Compounds
sodium, potassium,
and ammonium
All soluble
None
NaNO3 is soluble
KBr is soluble
nitrates
All soluble
None
LiNO3 is soluble
Mg(NO3)2 is soluble
chlorides
Most soluble
Silver, some mercury, and
lead chlorides
MgCl2 is soluble
PbCl2 is insoluble
sulfates
Most soluble
Strontium, barium, and lead
sulfate
K2SO4 is soluble
BaSO4 is insoluble
carbonates
Mostly insoluble*
Group IA and NH41
carbonates are soluble
Na2CO3 is soluble CaCO3
is insoluble
hydroxides and
sulfides
Mostly insoluble*
Group IA and NH41
hydroxides and sulfides are
soluble
KOH is soluble Al(OH)3 is
insoluble
*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
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
non-electrolytes.
5.9
Like dissolves like
5.9
Maximum Contaminant Level Goal (MCLG)
and Maximum Contaminant Level (MCL)
5.10
5.10
Hard water contains high concentrations of dissolved
calcium and magnesium ions.
Soft water contains few of these dissolved ions.
A pipe with hard-water scale build up
Not in 6th ed.
Because calcium ions, Ca2+, are generally the largest
contributors to hard water, hardness is usually expressed in
parts per million of calcium carbonate (CaCO3) by mass.
It specifies the mass of solid CaCO3 that could be formed
from the Ca2+ in solution, provided sufficient CO32- ions were
also present:
Ca2+(aq) + CO32–(aq) CaCO3(s)
A hardness of 10 ppm indicates that 10 mg of CaCO3 could be
formed from the Ca2+ ions present in 1 L of water.
Not in 6th ed.
Schematic drawing of a typical municipal water treatment facility
5.11
Getting the lead out:
Schematic of a typical spectrophotometer
Using a plot of
absorbance vs. concentration
Calibration graph
5.12
AAS =
Atomic
absorption
Spectrophotometer
Access to safe drinking water varies widely across the world.
5.14
Two water purification techniques:
Distillation
Reverse osmosis
5.14