chapter 5: the water we drink. “water has never lost its mystery. after at least two and a half...
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“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?
Where Does Potable (fit for consumption) Drinking Water Come From?
Surface water: from lakes, rivers, reservoirsGround water: pumped from wells drilled into underground aquifers
5.2
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
Poorly constructed landfills and septic systems
Run off from fertilized fields
Household chemicals poured down the drain or on the ground.
The average American usesalmost 100 gallons of water a day.
Nearly ¾ of the water enteringour homes goes down the drain.
5.2
5.3
A solution is a homogeneous mixture of uniform composition.
Solutions are made up of solvents and solutes.
Substances capable of dissolving other substances- usually present in the greater amount.
Substances dissolved in a solvent- usually present in the lesser amount.
When water is the solvent, you have an aqueous solution.
5.4
Concentration TermsParts per hundred (percent)
Parts per million (ppm)
Parts per billion (ppb)
20 g of NaCl in 80 g of water is a 20% NaCl solution
2 ppb Hg 2 g Hg
1109 g H 2O2 10-6 g Hg
1103 g H2O2 g Hg
1 L H 2O
Molarity (M) = moles soluteliter of solution
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
[ ] = “concentration of”
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
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:
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.
mol soluteL of solutionM =
5.5
Different Representations of Water
Lewis structures Space-filling Charge- density
Charge-density
Region of partial negative charge
Regions of partial positive charge
5.5
EN Values assigned by Linus Pauling, winner of TWO Nobel Prizes.
Electronegativity is a measure of an atom’s attraction for the electrons it shares in a covalent bond.
On periodic table, EN increases
Difference in Electronegativity Examples EN equal or greater than 2.0 = ionic bond NaCl EN 0.4-1.9 = polar covalent bond HF, H2OEN 0.0-0.3 = non-polar covalent bond N2, O2
Molecules can have individual polar bonds and still be non-polar overall because the bond dipoles (red arrows) cancel out. Examples: CO2, CH4
O C O
5.5
HH
O
A difference in the electronegativities of the atoms in a bond creates a polar bond.
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 HH2 has a non-polar covalent bond.
NaClNaCl has an ionic bond-look at the EN difference.
Na = 1.0
Cl = 2.9
EN = 1.9
A water molecule is polar – due to polar covalent bonds and the shape of the molecule.
5.6
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.
5.7
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.
Dissolution of NaCl in Water
The polar water molecules stabilize the ions as they break apart (dissociate).
Ions are simply charged particles-atoms or groups of atoms.
They may be positively charged – cations.
Or negatively charged- anions.
NaCl(s) Na+ (aq) + Cl-(aq)H2O
Naming simple ionic compounds is easy-
Name the metallic element (cation) first, followed by the non-metallic element (the anion) second, but with an –ide suffix.
5.7
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+
Sodium ions
and
Sulfate ions
The sulfate group stays together in solution.
Naming polyatomic ionic compounds is also easy-
Name the cation first, followed by the anion second.
5.7
MgOH Mg+ is the cation, OH- is the anion
magnesium hydroxide
NH4Br NH4+is the anion, Br- is the anion
ammonium bromide
5.8
Simple generalizations about ionic compounds allow us to predict their water
solubility.Ions
Solubility of Compounds
Solubility Exceptions Examples
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.
5.9
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.
A pipe with hard-water scale build up
Hard water contains high concentrations of dissolved calcium and magnesium ions.
Soft water contains few of these dissolved ions.
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 CO3
2- 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.
5.12
Getting the lead out:
Schematic of a typical spectrophotometer
Using a plot of absorbance vs. concentration
Calibration graph