Chapter 14
Acids and Bases
Homework
• Assigned Questions and Problems (odd only)
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Section 14.1
Section 14.2
Section 14.3
Section 14.4
Section 14.5
Section 14.6
Section 14.8
Section 14.9
(optional)
Read this section
Read this section
35, 37, 39
Neutralization Reactions section only, 45
53, 55, 57, 99, 105
67, 69, 71, 111, 121
73, 75, 77, 79, 81, 85, 87, 89, 109, 123
14.2 Acids: Properties and Examples
• Among the most common and important
compounds known
• Aqueous solutions are important in
– biological systems
– chemical industrial processes
• Early known characteristics
– Causes a sour taste (lemons and vinegar)
– Causes litmus (dye) to change from blue to red
– Dissolves some metals generating hydrogen gas
(e.g., zinc and iron)
14.2 Acids: Properties and Examples
• Common Acids (common laboratory
reagents)
– Sulfuric acid (H2SO4): most produced chemical in
U.S.
– Used to produce phosphoric acid and phosphate
fertilizers, rust removal in iron and steel production,
paper production
– Hydrochloric acid (HCl): steel production, organic
compound production, pH control and
neutralization, main component of stomach acid
– Nitric Acid (HNO3): mainly used in the production of
fertilzers, explosives, and used to dissolve metals
– See table 14.1
Section 5.9 Naming Binary Acids
• Use the prefix hydro- before the root
name of the element
• Add the suffix -ic and the word acid to the
root name for the element
• Example: HCl
– hydrochloric acid
• Example: HI
– hydroiodic acid
Section 5.9 Naming Oxyacids
• Produce H+ and a polyatomic ion when dissolved in
water
• Composed of hydrogen, oxygen, and another
nonmetal
• Use the root name of the polyatomic ion
• If it ends in -ate use the suffix -ic acid
• If it ends in -ite use the suffix -ous acid
• Example: H2SO4 (from SO42- , sulfate ion)
– sulfuric acid
• Example: H2SO3 (from SO32- , sulfite ion)
– sulfurous acid
14.3 Bases: Properties and Examples
• Also recognized as an important group of
compounds
• Aqueous solutions are important in
– biological systems
– chemical industrial processes
• Early known characteristics
– Causes a bitter taste (when dissolved in water)
– Causes litmus (dye) to change from red to blue
– Dissolves fats that are placed in base solutions
– Feel slippery like soap
Naming Bases
• Most bases are usually ionic compounds
• The hydroxide ion has a charge of (-1) and is
combined with a positively charged ion (group IA
or IIA metal ion)
• Hydroxides (bases) are named by their cation first,
then the word “hydroxide” is added to the metal
cation
• Most common bases:
–
–
–
–
NaOH (sodium hydroxide)
KOH (potassium hydroxide)
Ca(OH)2 (calcium hydroxide)
NH4OH (ammonium hydroxide)
14.4 The Arrhenius Definition
• Arrhenius (1884) first person to recognize the
essential nature of acids and bases
• Defined them in terms of chemical composition
• Arrhenius defined an acid as a substance that
produces hydrogen ions (H+) in water
• Acids are covalent compounds that ionize in
water to produce H+


 H (aq)  Cl (aq)
• Acids: HCl (g) 
H 2O
14.4 Molecular Definitions of Acids and Bases
• Arrhenius Acids
• When dissolved in water will generate H+ and an
anion
HCl (g) 
 H  (aq)  Cl (aq)
H 2O
– The H+ that is generated will give a sour taste
– Vinegar (acetic acid)
– Lemon (citric acid)
• Most acids are oxyacids (will also gen. H+)


HNO3 (l ) 
 H (aq)  NO3 (aq)
H 2O
Types of Acids
• Multiprotic Acids
–Can donate more than one
proton
–H2SO4
–H3PO4
Strong Acid
Weak Acid
H2 O


H 2SO 4 
HSO 4  H
H 2O


2
HSO 4 
SO 4  H
Types of Acids
• Binary Acids
• Molecular compounds in which hydrogen is attached to a
second nonmetallic element
• Formed from the pure compound which has been dissolved
in water
2O
HCl (g) H

H  (aq)  Cl (aq)
• Oxyacids
• Molecular compounds composed of hydrogen, oxygen, and
a nonmetal (e.g. S, C, N, Cl, or P)
• Molecular compounds which become acids when dissolved
in water

2O
HNO3 (l ) H

H (aq)  NO3 (aq)
• Acid hydrogen is attached to an oxygen
• These acids (formulas) look like acids of polyatomic ions
Types of Acids
• Organic Acids
–Acids with a carbon backbone
(carboxyl group)
–Acetic Acid (CH3COOH or
HC2H3O2)
O
H3C
O
C
H3C
O
Acetic Acid
H
C
Acetate Anion
O
H+
14.3 Bases: Properties and Examples
• Arrhenius Bases
• Arrhenius defined a base as a substance which
contains hydroxide and produces hydroxide
ions (OH-) in water
• Bases are ionic compounds that produce
hydroxide ions (OH-) in water
H 2O


• Bases: NaOH(s) 
Na (aq)  OH (aq)
14.3 Bases: Properties and Examples
• Arrhenius Bases
– When dissolved in water will generate OH- and a metal
ion. Two common examples:
NaOH (s) 
 Na  (aq)  OH  (aq)
H 2O


KOH (s) 
 K (aq)  OH (aq)
H 2O
• Bases are sometimes called alkalis
– The OH- that is generated will give a bitter taste
– Slippery feel (like soap)
– Most bases that are generated are composed from
group 1A and 2A metals
14.4 The Brønsted-Lowry Definition
of Acids and Bases
• Arrhenius definition is widely used, but model has
limitations
– Only for aqueous solutions
– Does not explain how compounds such as
ammonia (does not contain OH-) produce a basic
water solution
• New model (1923) by Brønsted and Lowry
proposed a broader definition
• New model applied to both aqueous and nonaqueous solutions
• Explained how compounds without OH- could
produce a basic solution when added to water
14.4 The Brønsted-Lowry Definition
of Acids and Bases
• Brønsted-Lowry Acid
– Any substance that can donate a proton (H+) to another
substance: a proton donor
• Brønsted-Lowry Base
– Any substance that can accept a proton (H+) from
another substance: a proton acceptor
– Proton donation does not occur unless a proton
acceptor is present
• Arrhenius Acids/Bases vs. Br.-Lowry Acids/Bases
– All acids in the Arrhenius definition are also acids according
to the Bronsted/Lowry model
– The Br.-Lowry definition of bases is quite different from that
of Arrhenius
– H+ ions produced by Br.-Lowry acids also react with water
Water as a Base
• For example: HCl is a Bronsted/Lowry acid because
it produces H+ in water
2O
HCl (g) H

H  (aq)  Cl (aq)
• H+ does not exist in water due to the strong
attraction to the polar water molecule
H+ acceptor
H
O
H
Hydronium ion
H
Cl
H+ donor
H
O
H
H
• In an aqueous solution the Arrhenius H+ ion
generated reacts with water to form H3O+
Cl
14.4 Identifying Brønsted-Lowry Bases
and Their Conjugates



HA (aq)  H 2 O ()  H 3 O (aq)  A (aq)
Acid
Base
Conjugate
Acid
Conjugate
Base
• Any chemical reaction that involves a BronstedLowry acid must also involve a Br.-Lowry base
• The reaction involves a proton transfer
• Acid can lose its proton to form a conjugate base
(something that could accept a proton back again)
• Base accepts a proton to form a conjugate acid
(something that could donate a proton)
14.4 Identifying Brønsted-Lowry Bases
and Their Conjugates


HCl (g)  H 2 O ()  H 3 O (aq)  Cl (aq)
Acid
Base
Conjugate
Acid
Conjugate
Base
• HCl is the Brønsted-Lowry acid
because it is donating a proton (H+)
to the water molecule
• The water molecule is the
Brønsted-Lowry base since it
accepts a proton (H+)
14.4 Conjugate Acid/Base Pairs
• Related to each other by
donating/accepting a single proton
(H+)
–The acid of the pair has the proton
–The base of the pair does not
–Every acid has a conjugate base
–Every base has a conjugate acid
–“Conjugate” is given to the part of the
pair that is produced in the reaction
14.4 Conjugate Acid/Base Pairs


NH3 (aq)  H 2 O (l )  NH4 (aq)  OH (aq)
Base
Acid
Conjugate
Acid
Conjugate
Base
• Each acid is related to a base on
the opposite side
–On the reactants side (left side)
substances are called “acid” or “base”
–On the products side (right side)
substances are called “conjugate acid”
or “conjugate base”
14.4 Conjugate Acid/Base Pairs
• Which of the following
represent conjugate acidbase pairs?
• H2O, H3
• OH-, HNO3
• H2SO4, SO42• HC2H3O2, C2H3O2+
O
14.4 Conjugate Acid/Base Pairs

H2O, H3O+
OH-,
HNO3
H2SO4, SO42-
H 2O  H  H 3O
HNO3  NO3 + H +

Not Pairs
+
H 2SO 4  SO4 +2H Not Pairs
+
H 2SO 4  HSO 4 + H
HC2H3O2, C2H3O2-
2-
-
HC2 H 3O2  C2 H 3O2 + H +
14.5 Reactions of Acids and Bases
Neutralization Reactions
• When Arrhenius acids and bases are mixed,
they will react with each other
• The acidic properties and the basic
properties are destroyed and this means
both substances are neutralized
• A neutralization is the reaction between an
acid and a hydroxide base which forms a
salt and water
14.5 Reactions of Acids and Bases
Neutralization Reactions
• A neutralization is a double-displacement reaction
AB  CD  AD  CB
• Whenever an acid is completely reacted with a
base, a neutralization occurs
HCl (aq)  KOH(aq)  KCl (aq)  H 2 O ()
Acid
Base
Salt
Water
• The net ionic equation is water formation

H (aq)  OH (aq)  H 2 O ()
-
14.5 Reactions of Acids and Bases
Neutralization Reactions
• Neutralization: The reaction between
an acid and a base to form a salt and
water
• H+ from the acid combines with the OHfrom the base to form water
• The properties of both reactants are
neutralized
• The salt contains the positive ion from
the base and the negative ion from the
acid
14.5 Reactions of Acids and Bases
Neutralization Reactions
• The reaction hydrochloric acid and sodium hydroxide:
• A double replacement reaction
HCl (aq)  NaOH (aq)  NaCl (aq)  H 2O (l )
base
acid
salt
water
• Ionic Equation
H (aq)  Cl- (aq)  Na  (aq)  OH- (aq)  Na  (aq)  Cl- (aq)  H2O (l)
• Net Ionic Equation
H (aq)  OH- (aq)  H2O (l)
14.8 Water: Acid and Base in One
(Ionization of Water)
• Substances that can act as acids or bases are
termed amphoteric
• Water can act as an acid or a base
– Amphoteric substance


H 2 O ()  H 2 O ()  H 3 O (aq)  OH (aq)
H+ acceptor
H
O
H
Base
Hydronium ion
H
O
H
Acid
H+ donor
H
O
H
H
Conj. Acid
O
H
Conj. Base
14.8 Water: Acid and Base in One
(Ionization of Water)
• Experiments show that in sample of pure water, a
very small percentage has dissociated to produce
ions
• It involves a proton transfer (a Br/L acid-base rxn)
• The net result is an equal amount of H3O+ and OH


H 2 O ()  H 2 O ()  H 3 O (aq)  OH (aq)
Equal amounts produced
H
O
H
Base
H
O
H
Acid
H
O
H
H
Conj Acid
O
H
Conj Base
14.8 Water: Acid and Base in One
(Ionization of Water)
• The acid-base reaction of water with itself is
called autoionization (self-ionization)
– It is an equilibrium reaction
– Ionization of water to form hydronium ion and
hydroxide ion is balanced by the recombination of
ions to form water
– At 25 ºC, the concentration of each ion: 1.0×10-7 M
– Square brackets around a compound denotes
“concentration” in moles per liter



H 2 O  H 2 O  H 3 O  OH
H O  OH  1 10

3

7
M
14.8 Ion-Product Constant for Water
• At any given temperature, the product of the
concentration of H+ and OH- is always a constant
• This value can be calculated at 25 °C since the
concentration (each ion) is known

-7 



14
(1.00 H
103OM) OH
(1.001
1010
M)  1.00
= Kw 10
-7
-14
M
• Kw (Ion-Product Constant for Water) is the
product of the H3O+ and OH- ion molar
concentrations in water
• Valid in pure water or water with solutes
14.8 Ion-Product Constant for Water
• If the [H3O+] is increased by the addition of
acidic solute, the [OH-] must decrease until the
expression is 1.0 × 10-14 is satisfied
[H3O ]  OH-   1.0  10 14
• Or, if [OH-] is increased by the addition of a
basic solute to the water, the [H3O]+ must
similarly decrease
O]
[H 3
14
 OH   1.0  10
14.8 Ion-Product Constant for Water
• In an aqueous solution, neither [H3O+] or [OH-] is
ever zero
• An acid is a substances that will increase the [H3O+] in
solution
• All acidic solutions have a higher [H3O+] than [OH-]
[H 3O  ] > [OH -]
• An base is a substances that will increase the
[OH-] ions in solution
• All basic solutions have a higher [OH-] than
[H3O+]

-
[H 3O ]  [OH ]
14.8 Ion-Product Constant for Water
• Neutral Solution
• Acidic Solution
[H 3O  ]  [OH -]
• Basic Solution

[H 3O ]  [OH ]
[H 3O  ] > [OH -]
• In all cases: H O OH  1 10

3

14
= Kw
Using Kw in Calculations, Example 1
• Calculate [H3O+] in a solution in which [OH-]
= 2.0x10-2 M. Is this solution acidic, basic or
neutral?

H3O 


OH   1 10 14 = Kw
[H3O ] (2.0 10- 2 )  1.0  10- 14 M
[ H 3O  ] 
1.0  10- 14
= 5.0 10
2
2.0  10
[H 3 O  ]  [OH - ]  basic
13
M
14.9 The pH and pOH Scale
• H3O+ (from H+) concentrations range from
very high values to extremely small valves
• Difficult and inconvenient to work with
numbers over such a large range
[H 3 O  ]  10 M and [H 3 O  ]  1.0 10 14 M
• For example, [H3O+] of 10 M is 1000 trillion
times greater than 10-14 M
• The pH scale of a solution was proposed
as an easier and more practical way to
handle such large numbers
14.9 The pH Scale
• The pH scale is defined as the
negative logarithm of the molar
hydronium ion concentration

pH  log H 3 O


• Logarithms are exponents
• The negative powers of 10 in the
concentrations are converted to
positive numbers
14.9 The pH Scale
• The pH scale uses a common logarithm
based on powers of 10

pH


log
[
H
O
]
3
• Expressed mathematically:
• For a number expressed in scientific
notation with a coefficient of 1, the
logarithm of that number is equal to the
value of the exponent
• Take the log of the number and then
multiply by negative one (change the sign)
14.9 Calculating the pH from [H3O+]
• The negative log of the H3O+ concentration
– To determine the number of significant figures
for logs: The number of decimal places for the log
is equal to the number of sig. figs. in the
concentration (original number)
Given:
H O   1.0  10

3
4
M
2 SF
pH   log [H 3 O  ]   log (1.0  10 4 ) M
pH  (4.00)
 4.00
2 D.P.
14.9 Calculating the pH from [H3O+]
• Calculate the pH for the
following solutions
–A solution in which [H3O+]=1.0x10-3 M
–A solution in which [OH-]= 5.0x10-5 M
Calculating the pH from [H3O+]
Calculating the pH from [OH-]
Given : [ H 3O  ] 1.0  10 3 M
pH   log [H 3 O  ]   log (1.0  10 3 )
pH   (3.00)  3.00
Given: [OH  ]  5.0 10 5 M

5
[H 3 O  ][OH  ] 1.0  10 14
[H 3 O ] (5.0 10 )  1.0  10
[H 3 O  ]  2.0  10 10
14
1.0  10 14
[H 3 O ] 
5.0  10 5

pH   log [H 3 O  ]   log (2.0  10 10 )
pH  (9.70)  9.70
Calculating the [H3O+] from pH
• It is often necessary to calculate the
hydronium ion concentration for a
solution from its pH value
• The logarithm must be undone
• To do this requires determining the
antilog of the pH value
• The antilog can be obtained on a
calculator using the antilog function
inv
log
• Use (-) pH value
Calculating the [H3O+] from pH
• The pH of rainwater in a
polluted area was found to be
3.50. What is the [H3O+] for this
rainwater?

pH
pHlog
log [[HH33OO ]]
inv log (log [H 3 O  ] )  inv log (3.50)
log [H 3 O  ]   3.50
[H 3 O  ]  3.16 10 4 M
The pH Scale
• pH is a log scale
• A change in one unit on the pH
scale means a tenfold increase or
decrease in [H3O+]
–Every time an exponent changes by
one, the pH changes by one
• Lowering the pH increases the
[H3O+]
–low pH value= acidic solution
–high pH value= basic solution
Measuring pH
• Use a pH meter
– Electronic device that measures the pH of the solution
• Use pH paper
– Paper has a chemical in it that changes to different colors
depending on the pH of the solution
• Indicators
– A compound that exhibits different colors depending on the
pH of its surroundings
The pOH Scale
• The negative log is also a way of expressing the
magnitudes of other quantities
• The concentration of OH- can be expressed as pOH
• Same as pH scale, but associated with the [OH-]
– Low pOH value means high [OH-]
– High pOH value means low [OH-]
OH  1.0 10 M
pOH  logOH  log1.0 10 

4

pOH  (4.00)  4.00
4
pH and pOH
• In an aqueous solution, the sum
of the pH and pOH is always 14
pH  pOH  14.00
• The value 14 corresponds to the
-log 1.0 × 10-14 of Kw equation
pH and pOH
[H+] > [OH-]
[H+] = [OH-]
[H+] < [OH-]
Calculating pOH from pH
• A sample of rain in an area with
severe air pollution has a pH of
3.5. What is the pOH of this rain
water?
pH  pOH  14.00
3.5  pOH  14.00
pOH  14.00 - 3.5  10.5
Calculating [OH-] from pOH
• The pOH of a liquid drain
cleaner was found to be 10.50.
What is the [OH-] for this
cleaner?

pOH  log OH

 10.50

log OH

 10.50
OH  inv log(10.50)
OH  3.16 10


11
M
14.6 Acid-Base Titration
• Determining the acid or base concentration
in a solution is a routing laboratory practice
• The pH only determines the amount of
dissociated H+ in solution
• Only dissociated molecules influence the pH
value
HCl (aq)  H  (aq)  Cl  (aq)
• Titration will give info about the total
number of acid or base molecules
present (concentration)
HCl (aq)  H  (aq)  Cl  (aq)
14.6 Acid-Base Titration
• Titration involves the
gradual addition of one
solution to another until
the solute in the first
solution has reacted
completely with the
solute in the second
solution
• First measure a known
volume of acid (unknown
molarity) to the flask
•To determine the concentration of an acid solution, add a
solution of base of known concentration to the flask by a buret
14.6 Acid-Base Titration
• Base addition continues until all the acid has completely reacted with
the added base: endpoint or equivalence point
• To determine endpoint, an indicator is used to detect when the acidbase reaction is complete
• If you know
molesacid = molesbase
– original volume of the acid
– the volume and concentration of the base
– the concentration of the acid can be calculated
At endpoint,
Mbase
Vbase
endpoint
Vacid
14.6 Acid-Base Titration
• As the base (NaOH) is added, OH- will react
with and neutralize HCl (and free H+),
forming water
NaOH (aq)  HCl (aq)  NaCl (aq)  H 2 O ()
• At endpoint all of the acid has completely
reacted with all of the base
At endpoint,
molesacid = molesbase
At endpoint,
Macid Vacid = Mbase Vbase
• end