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©2007 Pearson Prentice Hall 1
The Chemistry of Everything
Kimberley Waldron
Chapter 4 Salt
Behavior of ions, acids and bases
and the notion of equilibrium
©2007 Pearson Prentice Hall 2
Chapter Topics
• Ionic liquids, ionic interactions, delocalized
electronic charge.
• Polyatomic ions.
• Polar molecules, dipoles, ion–dipole interactions,
solubility.
• Electrolytes, molarity.
• Osmosis.
• Autoionization of water, acids and bases, pH.
• Acid rain.
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©2007 Pearson Prentice Hall 3
Salt: The Staff of Life
• Solid salt is a hard white powder.
• In solution in water, salt becomes a
collection of ions.
• Ions in solution can have powerful effects
on matter.
• Ionic solutions are essential for life.
• Acids and bases are special cases of ionic
solutions.
©2007 Pearson Prentice Hall 4
More than Morton’s
• A salt is any ionic solid
formed by neutralization
of an acid by a base.
• It contains a positive ion
(cation) and negative ion
(anion).
• Formation of salt from
elements involves.
• Transfer of electrons from
metal to non-metal.
3
©2007 Pearson Prentice Hall 5
Ionic Liquids?
• Densely packed ionic
lattices make traditional
salts high melting point
solids.
• New ionic compounds are
liquids at room
temperature.
– Ions are very large and
don’t pack together.
– Charge density is low.
– Attractions between ions
are weak.
©2007 Pearson Prentice Hall 6
Applications of Ionic Liquids
• Environmentally friendly
alternative to organic
solvents.
• Wide liquid range.
• Millions of variations by
changing the cation.
• Separating heavy metal
ion contaminants (lead
etc.) from water. The
lead dissolves better in
the ionic liquid.
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©2007 Pearson Prentice Hall 7
Mummies and Salts
• Questions remain
about how mummies
were made.
• Embalming solutions
contain salts like
natron.
©2007 Pearson Prentice Hall 8
Natron
• Natron is a collection of
salts used in the
mummification process.
• Some contain simple ions
like chloride.
• Other anions are
poly(many)atomic:
– CO32-, SO4
2-
• Polyatomic ions have
several resonance dot
structures.
5
©2007 Pearson Prentice Hall 9
Dot Structures for Ions
©2007 Pearson Prentice Hall 10
Dot Structures for Ions don’t
Follow the Rules of Neutral
Molecules
• Addition of charge (+ or -) changes the
bonding.
• Negative ion bond number decreases:
– C usually makes 4 bonds but in CN- makes 3.
– O usually makes 2 bonds but in OH- makes 1.
• Positive ion bond number increases:
– N usually makes 3 bonds but in NH4+ makes
4.
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©2007 Pearson Prentice Hall 11
Expanding the Octet Rule
• Maximum number of bonds formed by
carbon is 4 – equivalent to 8 electrons.
• Sulfate ion SO42- contains 6 bonds to
sulfur.
• Sulfur is larger than carbon and can
accommodate more atoms.
• Rule 1 for valence electrons is followed.
• Rule 2 for number of bonds is not.
©2007 Pearson Prentice Hall 12
Formulas of Ionic Compounds
• Salts are neutral but ions are charged.
• Charges of the ions must cancel out:• In MgCl2 Mg2+ charge is cancelled by 2 x Cl-
• In general, the formula of the salt can bepredicted using:
– y+ is charge on cation
– x- is charge on anion
– If x = y then both are given value of 1 (exceptperoxide which is O2
2-)
A Bx y
x-y+
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©2007 Pearson Prentice Hall 13
Common Polyatomic Ions
©2007 Pearson Prentice Hall 14
Salts and Desiccation: Polarity
of Water
• Water is a bent
molecule.
• O is more
electronegative than
H and attracts
electrons.
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©2007 Pearson Prentice Hall 15
Water is a Polar Molecule
• The charge
distribution is uneven.
• Water is a polar
molecule.
©2007 Pearson Prentice Hall 16
Water Molecules Hydrate
Ions
• Why do tightly bounds ionic crystal lattices dissolve?
• Ion – dipole interactions aid the solvation process:– The negative O atoms on water attach to the positive ions.
– The positive H atoms on water attach to the negative ions.
• Like dissolves like:– Ionic compounds dissolve in polar solvents but not in non-polar
solvents.
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©2007 Pearson Prentice Hall 17
Saturation and Limits on
Solubility
• Water molecules are
needed to solvate the
ions.
• Solubility of salt is
limited by availability
of water molecules.
• SSaturation is when
solution has reached
the solubility limit.
©2007 Pearson Prentice Hall 18
The Dynamic Equilibrium
• In a saturated solution thesolid salt is in equilibriumwith the dissolved salt:
• There is constantexchange between theions in the solid and inthe solution:– Rate of ions entering
solution = rate of ionsentering solid.
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©2007 Pearson Prentice Hall 19
Equilibrium Everywhere
• The state of equilibrium is everywhere in chemistry:
– Solid in equilibrium with liquid at melting point.
– Gas in equilibrium with liquid at boiling point.
– Reactants in equilibrium with products in a reaction.
©2007 Pearson Prentice Hall 20
What’s This to do with
Mummies?
• Water is required by bacteria to
decompose the body.
• The preserving salt (natron) absorbs the
water.
• The excess salt means equilibrium is
never reached and no water is left for the
bacteria.
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©2007 Pearson Prentice Hall 21
Chemistry and the Crumbling
Temple
• The temple at Luxor
is crumbling into dust.
• Salt becomes lodged
in crevices.
• Salt absorbs water.
• Expands and breaks
the stone structure.
©2007 Pearson Prentice Hall 22
Salts are Electrolytes
• Pure water does not
conduct electricity.
• Salts in solution
contain ions.
• Ions conduct
electricity.
• Salts are eelectrolytes.
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©2007 Pearson Prentice Hall 23
Electrolytes and Body
Function• Ions in the body are essential for blood pressure control and neural
function.
• Exercise depletes the body of ions.
• Hence the need for electrolytes.
• The history of Gatorade derives from the recognition of the role ofions in bodily function.
• Na+, K+, Mg2+ and Cl- are most important ions
©2007 Pearson Prentice Hall 24
Measuring Concentration
• In chemistry, the most common way of
measuring concentration is molarity:
• Unit is M. mM means millimolar.
• Example: What is concentration of solution
that contains 0.520 mol NaCl in 6.00 L?
_
_
MolessoluteMolarity
Litersolution=
0.5200.0867
6.00
molNaClMolarityM
Lsolution= =
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©2007 Pearson Prentice Hall 25
Salt and Growing Tomatoes
• Salt content of soil is critical for growing a
good tomato.
• Salt level in soil controls moisture level.
• Too much salt in salt draws moisture out
of the plant.
©2007 Pearson Prentice Hall 26
Factors Affecting Fluid Flow
• Cell membranes regulate
flow cell contents
• Semi-permeable
membrane prevents flow
of ions but allows
passage of water
• Osmosis is the flow of
water through a semi-
permeable membrane
from a dilute to a more
concentrated solution
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©2007 Pearson Prentice Hall 27
Osmotic Pressure
• The equilibration of the solution concentrations by
osmosis leads to a height difference in the solutions.
• This is tantamount to a pressure.
• The osmotic pressure is the pressure required to prevent
osmosis occurring.
©2007 Pearson Prentice Hall 28
Osmosis Dictates Water
Transport in the Body
• If cell salt
concentration is
higher water flows
into the cell – cell
expands.
• If cell salt
concentration is
lower, water flows out
from the cell – cell
shrinks.
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©2007 Pearson Prentice Hall 29
Growing Tomatoes in Salty
Soils• Excess salt in the soil
causes passage ofwater from plant toequalizeconcentrations.
• Lower soil saltconcentrationmaintains desirablelevel of water in plant.
©2007 Pearson Prentice Hall 30
Genetic Engineering and Salty
Soils
• Genetic engineering
makes plants with
higher natural salt
contents to permit
growth in salty soils.
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©2007 Pearson Prentice Hall 31
Hard Water and Bad Hair
• Hard ions make hardwater:– Ca2+, Mg2+, Fe#+
• Small highly chargedions form precipitateswith soap.
• Water softeningworks by ionexchange – hard forsoft ions.
©2007 Pearson Prentice Hall 32
Mechanics of Ion-Exchange
• Ion-exchange resins
contain soft ions (Na+)
attracted to negative
tethers.
• Exchange involves
the Na+ ion trading
places with the Ca2+
or Mg2+ ions.
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©2007 Pearson Prentice Hall 33
Ions in Pure Water
• Even a sample of pure
water contains ions.
• This is auto-ionization:
– H+ ions are responsible for
acid behavior.
– OH- ions are responsible
for basic behavior.
• Forward reaction:
– H2O H+ + OH-
• Reverse reaction:
– H+ + OH- H2O
©2007 Pearson Prentice Hall 34
Hydrogen Ions and the pH
Scale
• At equilibrium concentration of H+ and OH-
ions = 1 x 10-7 M (1 in 550,000,000)
• pH is a simple way of reporting the H+ ion
concentration:
pH = -log10[H+]
• Example:
[H+] = 1.0 x 10-7M
pH = -log10(1.0 x 10-7) = 7.00
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©2007 Pearson Prentice Hall 35
Significance of p
• In general:
p ≡ -log10
• pOH = -log10[OH-]
• In pure water [H+] =[OH-] = 1.0 x 10-7 M
• pH = pOH = 7 = neutral
• pH < 7 means [H+] > [OH-]
• pH > 7 means [H+] < [OH-]
©2007 Pearson Prentice Hall 36
Acids and Ions
• Acids are substancesthat produce H+ ionsin solution.
• Water solvates theions.
• In a typical acidsolution the H+ ionsprovided by the waterare negligible– IN 0.001 M HCl, [H+] =
10-3 M, pH = 3
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©2007 Pearson Prentice Hall 37
Slippery when Basic
• Basic solutions contain OH-
• Basic solutions have pH > 7– Example:
– Solution 0.055 M OH- ion has pOH = 1.26
– pH + pOH = 14
– pH = 12.74
• Tartness = acidity; caustic = basic
©2007 Pearson Prentice Hall 38
Significance of pH units
• What is more acidic: tomato juice, pH =
4.5 or orange juice pH = 3.5?
– Tomato juice: [H+] = 10-4.5 = 3.2 x 10-5 M
– Orange juice: [H+] = 10-3.5 = 3.2 x 10-4 M
– Orange juice is 10 times more acidic
• One pH unit = factor of ten change in [H+]
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©2007 Pearson Prentice Hall 39
Measuring pH
• A pH meter contains
a small tube with acid.
• The pH meter is able
to detect the
difference in
concentration
between H+ ions
inside the tube and
outside.
©2007 Pearson Prentice Hall 40
Detecting Acidity
• Indicators aresubstances thatreveal pH by color.
• When pH changesfrom acid to basis, theindicator colorchanges.
• Example:– Phenolphthalein is
colorless in acid butpink in base.
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©2007 Pearson Prentice Hall 41
An Indicator for all Occasions
• Individual indicators can only indicate possible ranges ofpH.
• Different indicators can be selected for different pHranges.
• A universal indicator has several colors that give moreprecise measurements of pH over the whole range.
©2007 Pearson Prentice Hall 42
Strong Acids and Bases
• SStrong acids and
bases dissolve
completely in water.
• All of the molecules
are ionized.
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©2007 Pearson Prentice Hall 43
Interaction of Air and Water
• The pH of water is influenced by the air above it.
• Natural rainwater is slightly acidic because of CO2.
• Some gases, especially the products of industry are
strongly acidic.
• SO2 and NO2 are the molecules responsible for acid rain.
©2007 Pearson Prentice Hall 44
Sources of Acid Rain
• SO2 comes from
burning fossil fuels
(especially coal)
– In air becomes SO3
– In water becomes
H2SO4 (strong acid)
• NO2 comes from
automobiles:
– In water becomes
HNO3 (strong acid)
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©2007 Pearson Prentice Hall 45
Acid Rain Stresses the
Environment
• Aquatic life cannot tolerate low pH.
• Penetration depth for light changes which
affects growth of plants.
• Situation is reversible although full
recovery can take years.
• Environmental legislation has been
effective in reducing acid rain.
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