Chapter 5: The Water We Drink. “Water has never lost its mystery. After at least two and a half millennia of philosophical and scientific inquiry, the.

Chapter 5: The Water We Drink

“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.

5.3

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.

+ 1 e-Na Na

Na atom Na+ ion

Forming ions

+ 1 e-

Cl atomCl- ion

ClCl

5.7

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

Some atoms form more than one stable ion

5.7

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.

Like dissolves like

5.9

5.10

Maximum Contaminant Level Goal (MCLG)and Maximum Contaminant Level (MCL)

5.10

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.

5.11

Schematic drawing of a typical municipal water treatment facility

5.12

Getting the lead out:

Schematic of a typical spectrophotometer

Using a plot of absorbance vs. concentration

Calibration graph

5.14

Access to safe drinking water varies widely across the world.

5.14

Two water purification techniques:

Distillation Reverse osmosis

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 ifWe had been choked with soot.

The Rime of the Ancient MarinerSamuel Taylor Coleridge