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Chemistry, The Central Science, 10th edition
Theodore L. Brown; H. Eugene LeMay, Jr.;
and Bruce E. Bursten
Chapter 16
Acids and Bases
John D. Bookstaver
St. Charles Community College
St. Peters, MO
 2006, Prentice Hall, Inc.
Acids
and
Bases
Properties of Acids
• sour taste
• react with “active” metals
 recall the activity series (any metals above H are “active”)
 2 Al (s) + 6 HCl (aq)  2 AlCl3 (aq) + 3 H2 (g)
• corrosive
 has the ability to eat away/destroy other substances
• react with carbonates (containing CO32-)
 produce CO2 gas
 marble, baking soda, chalk, limestone
 CaCO3 (A) + 2 HCl (aq)  CaCl2 (aq) + CO2 (g) + H2O (l)
• change color of vegetable dyes
 blue litmus paper turns red
• react with bases to form ionic salts (ionic compounds)
Properties of Bases
• also known as alkalis
• taste bitter
• solutions feel slippery
Many soaps are manufactured with the use of
bases
• change color of vegetable dyes
different color than acid
red litmus paper turns blue
• react with acids to form ionic salts (ionic
compounds)
This is known as “neutralization”
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What are Acids and Bases?
• There are several different ways to
define the terms acid and base
• Each definition comes from a different
acid/base theory:
 Arrhenius Theory
 Brønsted-Lowry Theory
 Lewis Theory
Acids
and
Bases
Arrhenius Theory
• Acid:
 when dissolved in water, the hydrogen ion concentration
increases (basically, an Arrhenius acid produces H+ ions)
Ex) HCl (aq) → H+ (aq) + Cl– (aq)
HCl is an acid b/c it is ionized (molecular acid that is
pulled apart by water) to form H+
• Base:
 when dissolved in water, the hydroxide ion concentration
increases (basically, an Arrhenius base produces OH–
ions)
Ex) NaOH (aq) → Na+ (aq) + OH– (aq)
NaOH is a base b/c it dissociates (ionic compound
breaking up into its ions) to form OH–
Problems with Arrhenius Theory…
• Does not explain why molecular substances,
like NH3, and some ionic compounds, like
Na2CO3 or Na2O, are bases
since they do not contain OH– ions, they can’t
be considered Arrhenius bases
• Does not explain why molecular substances,
like CO2, dissolve in water to form acidic
solutions
since it doesn’t even contain H+ ions it can’t be
considered an Arrhenius acid
• Does not explain acid-base reactions that
take place outside of aqueous solutions
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Hydronium Ion: H3O+
• the H+ ions produced by the acid are so
reactive they cannot exist in water
H+ ions are just protons w/ no valence electrons!!!
• To be more stable, H+ ions react with water
molecule(s) to produce hydronium ions, H3O+
• Note: for our purposes, H+ and H3O+ will be used
interchangeably
Brønsted-Lowry Theory
• The Brønsted-Lowry theory looks at
the entire acid-base chemical reaction
• In an acid-base reaction, an H+ is
transferred from the acid to the base
 Broader, more general, definition than
Arrhenius
 does not have to take place in aqueous
solution
Brønsted-Lowry Acids & Bases
• Brønsted-Lowry acids are H+ donors (aka:
proton donors)
Must have a removable (acidic) hydrogen
because of the molecular structure, often one H
in the molecule is easier to transfer than others
(this is known as the “ionizable hydrogen”)
• Brønsted-Lowry bases are H+ acceptors
(aka: proton acceptors)
Must have a lone pair/ nonbonding pair of
electrons
because of the molecular structure, often one
atom in the molecule is more willing to accept
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H+ transfer than others
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Brønsted-Lowry Acid/Base Rxns
Ex) HCl (aq) + H2O (l) → Cl– (aq) + H3O+ (aq)
acid
base
• HCl transfers an H+ to H2O, forming H3O+ ions
 water acts as base, accepting H+
Ex) NH3 (aq) + H2O (l)  NH4+ (aq) + OH– (aq)
base
acid
• NH3 (aq) is a Brønsted-Lowry base because NH3
accepts an H+ from H2O (it is also an Arrhenius
base because it forms OH–)
water acts as acid, donating H+
Amphiprotic Molecules
• Amphiprotic (or amphoteric) means have qualities
of both…
 substance can act as both an acid and a base
depending what it is reacting with
After
Donated
H+
OH–
Amphoteric
Molecule
H2O
After
Accepted
H+
H3O+
CO32−
HCO3−
H2CO3
SO42−
HSO4−
H2SO4
Brønsted-Lowry Reversible
Acid-Base Reactions
• Brønsted-Lowry theory allows A-B rxns to be
reversible (shown w/ double-sided arrows)
General A-B rxn: H–A + :B  :A– + H–B+
• the original base has an extra H+ after the
reaction, so it will act as an acid in the
reverse process
• and the original acid has a lone pair of
electrons after the reaction, so it will act as a
base in the reverse process
:A– + H–B+  H–A + :B
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Conjugate Pairs
• In a Brønsted-Lowry 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
are called a conjugate pair
• the original base becomes the conjugate
acid; and the original acid becomes the
conjugate base
Acids
and
Bases
Conjugate Pairs
Acids
and
Bases
Brønsted-Lowry
Acid-Base Reactions
H–A
acid
HCHO2
acid
+ :B

:A–
+
base
conjugate
base
+ H2O
base

H–B+
conjugate
acid
CHO2– +
H3O+
conjugate conjugate
base
acid
+
NH4+
H2O + NH3  OH–
acid
base
conjugate
conjugate
base
acid
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Conjugate Pairs Practice
• Determine the acid, base, conjugate acid,
and conjugate base in the following reaction:
H2O
base
+
H2SO4
acid

HSO4–
H3O+
+
conjugate
base
conjugate
acid
Acids
and
Bases
Arrow Conventions
• chemists commonly use two kinds of arrows
in reactions to indicate the degree of
completion of the reactions
• a single arrow indicates all the reactant
molecules are converted to product
molecules at the end
• a double arrow indicates the reaction is in
equillibrium. The reaction stops when some
of the reactant molecules have been
converted into products. Some reactant
molecules remain and some products
molecules have formed
Looks like  in these notes
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 % of the molecules ionize ()
• a weak base is a weak electrolyte
only a small % of the base molecules form
OH– ions ()
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Autoionization of Water
• As we have seen, water is amphoteric.
In pure water, a few molecules act as bases and a
few act as acids.
about 1 out of every 10 million water molecules
form ions through a process called autoionization
H2O(l) + H2O(l)  H3O+(aq) + OH−(aq)
• all aqueous solutions contain both H3O+ and
OH–
the concentration of H3O+ and OH– are equal in
pure water (which is neutral)
Ion Product of Water
• the product of the H3O+ and OH–
concentrations is a constant number
• the number is called the ion product constant
of water (Kw)
At 25ºC,
[H3O+] [OH−] = [H+] [OH−] = Kw = 1 x 10-14
if you measure one of the concentrations
(H+ or OH–) , you can calculate the other
• as [H3O+] increases the [OH–] must decrease
so the product stays constant
inversely proportional
Acidic and Basic Solutions
• all aqueous solutions contain both H+
Be careful!
and OH– ions
We are
• neutral solutions
[H+] = [OH–] = 1 x 10-7
• acidic solutions
[H+] is greater than [OH–]
[H+] > 1 x 10-7; [OH–] < 1 x 10-7
• basic solutions
[OH–] is greater than [3O+]
[OH–] > 1 x 10-7; [H+] < 1 x 10-7
dealing with
NEGATIVE
exponents.
the larger
the negative
exponent,
the
SMALLER
the number
actually is!
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Concentration Calculation
Practice
• Calculate the [OH] at 25ºC when the
[H+] = 1.5 x 10-9 M, and determine if the
solution is acidic, basic, or neutral
1.0 x 10-14 M = [OH] [H+]
1.0 x 10-14 M = [OH] [1.5 x 10-9 M]
[OH] = 6.67 x10-6 M
The solution is basic… [OH] > [H+]
Acids
and
Bases
pH
• pH is a measure of how acidic or basic
a solution is
pH = –log [H+];
[H+] = 10–pH
Ex) pHwater = –log [1.0 x 10-7 M] = 7
Similarly, [H+]water = 10–7 = 1.0 x 10-7 M
• need to know the [H+] concentration to
find pH
pH
• An acid has a higher [H+] than pure
water, so its pH is < 7
• A base has a lower [H3O+] than
pure water, so its pH is >7
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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)
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
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Acids
and
Bases
Other “p” Scales
• The “p” in pH tells us to take the
negative log of the quantity (in this case,
hydrogen ions).
• Some similar examples are
pOH = −log [OH−]
pKw = −log Kw
Acids
and
Bases
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pOH
• another way of expressing the
acidity/basicity of a solution is pOH
pOH = -log [OH];
[OH] = 10-pOH
• pOH measures are the exact opposite of pH:
• pOH < 7 is basic, pOH > 7 is acidic
• pOH = 7 is neutral
Watch This!
Because
[H3O+] [OH−] = Kw = 1.0  10−14,
we know that
−log [H3O+] + −log [OH−] = −log Kw = 14.00
or, in other words,
pH + pOH = pKw = 14.00
Acids
and
Bases
The
Acid/
Base
Square
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Relationships:
pH, pOH, [H+], and [OH–]
pH: 0
[H+]:
100
1
10-1
H+
OH-
3
5
10-3
10-5
H+
9
10-7
10-9
[OH-]: 10-14
10-13 10-11
pOH: 14
13
11
13
10-11
10-13
14
10-14
H+
H+
H+
OH- OH-
OH-
OH-
11
7
10-9
10-7
10-5
10-3
10-1
100
9
7
5
3
1
0
EQUATIONS TO REMEMBER!
[H+][OH–] = 1 x 10-14 = kw
pH = -log[H+]
[H+] = 10-pH
pOH = -log[OH-]
[OH–] = 10-pOH
pH + pOH = 14
How Do We Measure pH?
• For less accurate
measurements, one
can use
 Litmus paper
• “Red” paper turns
blue above ~pH = 8
• “Blue” paper turns
red below ~pH = 5
 An indicator
Acids
and
Bases
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How Do We Measure pH?
For more accurate
measurements, one
uses a pH meter,
which measures the
voltage in the
solution.
Acids
and
Bases
Strong Acids
• You will recall that the seven strong acids are
HCl, HBr, HI, HNO3, H2SO4, HClO3, and
HClO4.
• These are, by definition, strong electrolytes
and exist totally as ions in aqueous solution.
• For the monoprotic strong acids,
[H3O+] = [acid].
Acids
and
Bases
Strong Bases
• Strong bases are the soluble hydroxides,
which are the alkali metal and heavier
alkaline earth metal hydroxides (Ca2+, Sr2+,
and Ba2+).
• Again, these substances dissociate
completely in aqueous solution.
Acids
and
Bases
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Dissociation Constants
• For a generalized acid dissociation,
HA(aq) + H2O(l)
A−(aq) + H3O+(aq)
the equilibrium expression would be
Kc =
[H3O+] [A−]
[HA]
• This equilibrium constant is called the
acid-dissociation constant, Ka.
Acids
and
Bases
Dissociation Constants
The greater the value of Ka, the stronger
the acid.
Acids
and
Bases
Calculating Ka from the pH
• The pH of a 0.10 M solution of formic acid,
HCOOH, at 25°C is 2.38. Calculate Ka for
formic acid at this temperature.
• We know that
Ka =
[H3O+] [COO−]
[HCOOH]
Acids
and
Bases
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Calculating Ka from the pH
• The pH of a 0.10 M solution of formic acid,
HCOOH, at 25°C is 2.38. Calculate Ka for
formic acid at this temperature.
• To calculate Ka, we need the equilibrium
concentrations of all three things.
• We can find [H3O+], which is the same as
[HCOO−], from the pH.
Acids
and
Bases
Calculating Ka from the pH
pH = −log [H3O+]
2.38 = −log [H3O+]
−2.38 = log [H3O+]
10−2.38 = 10log [H3O+] = [H3O+]
4.2  10−3 = [H3O+] = [HCOO−]
Acids
and
Bases
Calculating Ka from pH
Now we can set up a table…
[HCOOH], M
Initially
0.10
Change
−4.2  10-3
At
Equilibrium
0.10 − 4.2  10−3
= 0.0958 = 0.10
[H3O+], M [HCOO−], M
0
0
+4.2  10-3 +4.2  10−3
4.2  10−3
4.2  10−3
Acids
and
Bases
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Calculating Ka from pH
Ka =
[4.2  10−3] [4.2  10−3]
[0.10]
= 1.8  10−4
Acids
and
Bases
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