Transcript Chapter 14 Acids and Bases

Chapter 14
Acids and
Bases
Types of Electrolytes
• salts = water soluble ionic compounds
all strong electrolytes
• acids = form H+1 ions in water solution
• bases = combine with H+1 ions in water solution
increases the OH-1 concentration
may either directly release OH-1 or pull H+1 off H2O
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Properties of Acids
• Sour taste
• react with “active” metals
 i.e. Al, Zn, Fe, but not Cu, Ag or Au
2 Al + 6 HCl AlCl3 + 3 H2
 corrosive
• react with carbonates, producing CO2
 marble, baking soda, chalk, limestone
CaCO3 + 2 HCl CaCl2 + CO2 + H2O
• change color of vegetable dyes
 blue litmus turns red
• react with bases to form ionic salts
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Common Acids
Chemical Name
Formula
Uses
Strength
Nitric Acid
HNO3
explosive, fertilizer, dye, glue
Strong
explosive, fertilizer, dye, glue,
batteries
metal cleaning, food prep, ore
refining, stomach acid
fertilizer, plastics & rubber,
food preservation
plastics & rubber, food
preservation, Vinegar
Sulfuric Acid
H2SO4
Strong
Hydrochloric Acid
HCl
Phosphoric Acid
H3PO4
Acetic Acid
HC2H3O2
Hydrofluoric Acid
HF
metal cleaning, glass etching
Weak
Carbonic Acid
H2CO3
soda water
Weak
Boric Acid
H3BO3
eye wash
Weak
Strong
Moderate
Weak
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Structures of Acids
• binary acids have acid
hydrogens attached to a
nonmetal atom
HCl, HF
Hydrofluoric acid
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Structure of Acids
• oxy acids have acid
hydrogens attached to
an oxygen atom
H2SO4, HNO3
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Structure of Acids
• carboxylic acids have
COOH group
 HC2H3O2, H3C6H5O3
• only the first H in the
formula is acidic
 the H is on the COOH
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Properties of Bases
• also known as alkalis
• taste bitter
 alkaloids = plant product that is alkaline
 often poisonous
• solutions feel slippery
• change color of vegetable dyes
 different color than acid
 red litmus turns blue
• react with acids to form ionic salts
 neutralization
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Common Bases
Chemical
Name
sodium
hydroxide
potassium
hydroxide
calcium
hydroxide
sodium
bicarbonate
magnesium
hydroxide
ammonium
hydroxide
Formula
NaOH
Common
Name
lye,
caustic soda
Uses
soap, plastic,
petrol refining
soap, cotton,
electroplating
Strength
Strong
KOH
caustic potash
Strong
Ca(OH)2
slaked lime
cement
Strong
NaHCO3
baking soda
cooking, antacid
Weak
Mg(OH)2
milk of
magnesia
antacid
Weak
NH4OH,
{NH3(aq)}
ammonia
water
detergent,
fertilizer,
explosives, fibers
Weak
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Structure of Bases
• most ionic bases
contain OH ions
NaOH, Ca(OH)2
• some contain CO32- ions
CaCO3 NaHCO3
• molecular bases contain
structures that react
with H+
mostly amine groups
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Arrhenius Theory
• bases dissociate in water to produce OHions and cations
 ionic substances dissociate in water
NaOH(aq) → Na+(aq) + OH–(aq)
• acids ionize in water to produce H+ ions
and anions
 because molecular acids are not made of ions,
they cannot dissociate
 they must be pulled apart, or ionized, by the
water
HCl(aq) → H+(aq) + Cl–(aq)
 in formula, ionizable H written in front
HC2H3O2(aq) → H+(aq) + C2H3O2–(aq)
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Arrhenius Acid-Base Reactions
• the H+ from the acid combines with the OHfrom the base to make a molecule of H2O
it is often helpful to think of H2O as H-OH
• the cation from the base combines with the
anion from the acid to make a salt
acid + base → salt + water
HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)
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Problems with Arrhenius Theory
• does not explain why molecular substances, like NH3,
dissolve in water to form basic solutions – even though
they do not contain OH– ions
• does not explain acid-base reactions that do not take
place in aqueous solution
• H+ ions do not exist in water. Acid solutions contain
H3O+ ions
 H+ = a proton!
 H3O+ = hydronium ions
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Brønsted-Lowery Theory
•
in a Brønsted-Lowery Acid-Base reaction, an
H+ is transferred
 does not have to take place in aqueous solution
 broader definition than Arrhenius
•
acid is H donor, base is H acceptor
 base structure must contain an atom with an
unshared pair of electrons
•
in the reaction, the acid molecule gives an H+
to the base molecule
H–A + :B  :A– + H–B+
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Amphoteric Substances
• amphoteric substances can act as either an
acid or a base
have both transferable H and atom with lone pair
• HCl(aq) is acidic because HCl transfers an H+
to H2O, forming H3O+ ions
water acts as base, accepting H+
HCl(aq) + H2O(l) → Cl–(aq) + H3O+(aq)
• NH3(aq) is basic because NH3 accepts an H+
from H2O, forming OH–(aq)
water acts as acid, donating H+
NH3(aq) + H2O(l)  NH4+(aq) + OH–(aq)
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Brønsted-Lowery
Acid-Base Reactions
• one of the advantages of Brønsted-Lowery theory
is that it allows reactions to be reversible
H–A + :B → :A– + H–B+
• the original base has an extra H+ after the reaction
– so it could act as an acid in the reverse process
• and the original acid has a lone pair of electrons
after the reaction – so it could act as a base in the
reverse process
:A– + H–B+ → H–A + :B
• a double arrow, , is usually used to indicate a
process that is reversible
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Conjugate Pairs
• In a Brønsted-Lowery Acid-Base reaction, the
original base becomes an acid in the reverse
reaction, and the original acid becomes a base in
the reverse process
• each reactant and the product it becomes is
called a conjugate pair
• the original base becomes the conjugate acid;
and the original acid becomes the conjugate base
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Brønsted-Lowery
Acid-Base Reactions
H–A
acid
+
HCHO2
acid
H2O +
acid
:B
base

+ H2O
base

NH3
base

:A–
conjugate
base
CHO2–
conjugate
base
+
HO–
conjugate
base
+
H–B+
conjugate
acid
+
H3O+
conjugate
acid
NH4+
conjugate
acid
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Conjugate Pairs
In the reaction H2O + NH3  HO– +
NH4+
–
H2O and HO constitute an
Acid/Conjugate Base pair
NH3 and NH4+ constitute a
Base/Conjugate Acid pair
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Practice – Identify the Brønsted-Lowery
Acids and Bases and their Conjugates in
each Reaction
+
H2O

HSO4–
+
H3O+
HCO3– +
H2O

H2CO3
+
HO–
H2SO4
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Neutralization Reactions
• H+ + OH- H2O
• acid + base salt + water
• double displacement reactions
 salt = cation from base + anion from
acid
 cation and anion charges stay constant
H2SO4 + Ca(OH)2 → CaSO4 + 2 H2O
• some neutralization reactions are gas
evolving where H2CO3 decomposes
into CO2 and H2O
H2SO4 + 2 NaHCO3 → Na2SO4 + 2 H2O + 2 CO2
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Nonmetal Oxides are Acidic
• nonmetal oxides react with water to form
acids
• causes acid rain
CO2 (g) + H2O(l) → H2CO3(aq)
2 SO2(g) + O2(g) + 2 H2O(l) → 2 H2SO4(aq)
4 NO2(g) + O2(g) + 2 H2O(l) → 4 HNO3(aq)
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Acid Reactions
Acids React with Metals
• acids react with many metals
but not all!!
• when acids react with metals, they
produce a salt and hydrogen gas
3 H2SO4(aq) + 2 Al(s) → Al2(SO4)3(aq) + 3 H2(g)
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Acid Reactions
Acids React with Metal Oxides
• when acids react with metal oxides, they
produce a salt and water
3 H2SO4 + Al2O3 → Al2(SO4)3 + 3 H2O
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Base Reactions
• the reaction all bases have is common is
neutralization of acids
• strong bases will react with Al metal to form
sodium aluminate and hydrogen gas
2 NaOH + 2 Al + 6 H2O → 2 NaAl(OH)4 + 3 H2
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Titration
• using reaction stoichiometry to
determine the concentration of
an unknown solution
• Titrant (unknown solution)
added from a buret
• indicators are chemicals added
to help determine when a
reaction is complete
• the endpoint of the titration
occurs when the reaction is
complete
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Titration
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Titration
The base solution is the
titrant in the buret.
As the base is added to
the acid, the H+ reacts with
the OH– to form water.
But there is still excess
acid present so the color
does not change.
At the titration’s endpoint,
just enough base has been
added to neutralize all the
acid. At this point the
indicator changes color.
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Example 14.4
Acid-Base Titration
The titration of 10.00 mL of HCl solution of unknown concentration
requires 12.54 mL of 0.100 M NaOH solution to reach the end
point. What is the concentration of the unknown HCl solution?
Strong or Weak
• a strong acid is a strong electrolyte
 practically all the acid molecules ionize, →
• a strong base is a strong electrolyte
 practically all the base molecules form OH– ions, either
through dissociation or reaction with water, →
• a weak acid is a weak electrolyte
 only a small percentage of the molecules ionize, 
• a weak base is a weak electrolyte
 only a small percentage of the base molecules form OH–
ions, either through dissociation or reaction with water, 
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Strong Acids
• The stronger the acid, the
more willing it is to donate H
 use water as the standard base
HCl  H+ + ClHCl + H2O H3O+ + Cl-
• strong acids donate
practically all their H’s
 100% ionized in water
 strong electrolyte
• [H3O+] = [strong acid]
 [ ] = molarity
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Strong Acids
Pure Water
HCl solution
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Weak Acids
• weak acids donate a small
fraction of their H’s
HF  H+ + FHF + H2O  H3O+ + F-
most of the weak acid
molecules do not donate H
to water
much less than 1% ionized
in water
• [H3O+] << [weak acid]
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Weak Acids
Pure Water
HF solution
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Strong Bases
• The stronger the base, the more
willing it is to accept H
NaOH  Na+ + OH-
 use water as the standard acid
• strong bases, practically all
molecules are dissociated into
OH– or accept H’s
 strong electrolyte
 multi-OH bases completely
dissociated
• [HO–] = [strong base] x (# OH)
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Weak Bases
• in weak bases, only a small
fraction of molecules accept
H’s
NH3 + H2O  NH4+ + OH-
 weak electrolyte
 most of the weak base molecules
do not take H from water
 much less than 1% ionization in
water
• [HO–] << [strong base]
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Relationship between Strengths of
Acids and their Conjugate Bases
• the stronger an acid is, the weaker the attraction
of the ionizable H for the rest of the molecule is
• the better the acid is at donating H, the worse its
conjugate base will be at accepting a H
strong acid HCl + H2O → Cl– + H3O+ weak conj. base
weak acid HF + H2O  F– + H3O+ strong conj. base
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Autoionization of Water
• Water is actually an extremely weak electrolyte
therefore there must be a few ions present
• about 1 out of every 10 million water molecules
form ions through a process called
autoionization
H2O  H+ + OH–
H2O + H2O  H3O+ + OH–
• all aqueous solutions contain both H+ and OH–
the concentration of H+ and OH– are equal in water
[H+] = [OH–] = 10-7M @ 25°C
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Ion Product of Water
• the product of the H+ and OH–
concentrations is always the same number
• the number is called the ion product of
water and has the symbol Kw
• [H+] x [OH–] = 1 x 10-14 = Kw
• as [H+] increases the [OH–] must decrease
so the product stays constant
inversely proportional
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Acidic and Basic Solutions
• neutral solutions have equal [H+] and [OH–]
[H+] = [OH–] = 1 x 10-7
• acidic solutions have a larger [H+] than [OH–]
[H+] > 1 x 10-7; [OH–] < 1 x 10-7
• basic solutions have a larger [OH–] than [H+]
[H+] < 1 x 10-7; [OH–] > 1 x 10-7
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Example - Determine the [H+1] for a 0.00020 M
Ba(OH)2 and determine whether the solution is
acidic, basic or neutral
Ba(OH)2 = Ba2+ + 2 OH– therefore
[OH–] = 2 x 0.00020 = 0.00040 = 4.0 x 10-4 M
 
K w  H   OH 
H  
Kw


1 1014 
4 

4
.
0

10

OH

[H+] = 2.5 x 10-11 M
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Practice - Determine the [H+1] concentration
and whether the solution is acidic, basic or
neutral for the following
• [OH–] = 0.000250 M
• [OH–] = 3.50 x 10-8 M
• Ca(OH)2 = 0.20 M
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pH
• the acidity/basicity of a solution is often
expressed as pH
• pH = -log[H+], [H+] = 10-pH
exponent on 10 with a positive sign
pHwater = -log[10-7] = 7
need to know the [H+] concentration to find pH
• pH < 7 is acidic; pH > 7 is basic, pH = 7 is
neutral
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pH
• the lower the pH, the more acidic the solution; the
higher the pH, the more basic the solution
1 pH unit corresponds to a factor of 10 difference
in acidity
• normal range 0 to 14
pH 0 is [H+] = 1 M, pH 14 is [OH–] = 1 M
pH can be negative (very acidic) or larger than 14
(very alkaline)
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pH of Common Substances
Substance
pH
1.0 M HCl
0.0
0.1 M HCl
1.0
stomach acid
1.0 to 3.0
lemons
2.2 to 2.4
soft drinks
2.0 to 4.0
plums
2.8 to 3.0
apples
2.9 to 3.3
cherries
3.2 to 4.0
unpolluted rainwater
5.6
human blood
7.3 to 7.4
egg whites
7.6 to 8.0
milk of magnesia (sat’d Mg(OH)2)
10.5
household ammonia
10.5 to 11.5
1.0 M NaOH
14
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Example - Calculate the pH of a 0.0010 M
Ba(OH)2 solution & determine if is acidic,
basic or neutral
Ba(OH)2 = Ba2+ + 2 OH- therefore
[OH-] = 2 x 0.0010 = 0.0020 = 2.0 x 10-3 M
[H+] =
1 x 10-14
2.0 x 10-3
= 5.0 x 10-12M
pH = -log [H+] = -log (5.0 x 10-12)
pH = 11.3
pH > 7 therefore basic
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Practice - Calculate the pH of the
following strong acid or base solutions
• 0.0020 M HCl
• 0.0050 M Ca(OH)2
• 0.25 M HNO3
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Sample - Calculate the concentration of
[H+] for a solution with pH 3.7
[H+] = 10-pH
[H+] = 10-3.7
means 0.0001 < [H+1] < 0.001
[H+] = 2 x 10-4 M = 0.0002 M
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Practice - Determine the [H+] for each of
the following
• pH = 2.7
• pH = 12
• pH = 0.60
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Buffers
• buffers are solutions that resist changing pH
when small amounts of acid or base are
added
• they resist changing pH by neutralizing
added acid or base
• buffers are made by mixing together a weak
acid and its conjugate base
or weak base and it conjugate acid
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How Buffers Work
• the weak acid present in the buffer mixture
can neutralize added base
• the conjugate base present in the buffer
mixture can neutralize added acid
• the net result is little to no change in the
solution pH
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What is Acid Rain?
• natural rain water has a pH of 5.6
naturally slightly acidic due mainly to CO2
• rain water with a pH lower than 5.6 is called
acid rain
• acid rain is linked to damage in ecosystems
and structures
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What Causes Acid Rain?
• many natural and pollutant gases dissolved in the air
are nonmetal oxides
 CO2, SO2, NO2
• nonmetal oxides are acidic
CO2 + H2O  H2CO3
2 SO2 + O2 + 2 H2O  2 H2SO4
• processes that produce nonmetal oxide gases as waste
increase the acidity of the rain
 natural – volcanoes and some bacterial action
 man-made – combustion of fuel
• weather patterns may cause rain to be acidic in regions
other than where the nonmetal oxide is produced
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Damage from Acid Rain
• acids react with metals, and materials that
contain carbonates
• acid rain damages bridges, cars and other
metallic structures
• acid rain damages buildings and other
structures made of limestone or cement
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Damage from Acid Rain
circa 1935
circa 1995
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