Acids And Bases
Section 19.1
Acid-Base Theories
• OBJECTIVES:
–Define the properties of
acids and bases.
Section 19.1
Acid-Base Theories
• OBJECTIVES:
–Compare and contrast
acids and bases as defined
by the theories of:
a) Arrhenius,
b) Brønsted-Lowry, and
c) Lewis.
Class question
• Where can acids
be found?
–
–
–
–
Sodas
Stomach
Vinegar
Citrus fruits
• Where can bases
be found?
–
–
–
–
–
Soap
Drano
Antacid tablets
Windex
detergent
Properties of Acids
• Taste sour
• React with bases
• Litmus paper test – turn blue litmus
paper red
• Electrolytic – conduct electricity
– Can be strong or weak electrolytes
in aqueous solution
Properties of Acids
• They have a pH of less than 7 (more
on this concept of pH in a later
lesson)
• How do you know if a chemical is an
acid?
– It usually starts with Hydrogen.
– HCl, H2SO4, HNO3, etc. (but not water!)
Acids Affect Indicators, by
changing their color
Blue litmus paper turns red in
contact with an acid (and red paper
stays red).
Acids
have a
pH
less
than 7
Properties of Bases
•
•
•
•
Taste bitter
Feels slippery
React with acids
Litmus paper test – turn red litmus
paper blue
• electrolytic
Bases Affect Indicators
Red litmus paper
turns blue in contact
with a base (and blue
paper stays blue).
Phenolphthalein
turns purple in a
base.
Bases
have a
pH
greater
than 7
Acid Nomenclature
Nomenclature of Acids
• Acids are composed of a(n)
Hydrogen ion (H+)
________________ followed by
anion
a(n)
_______
Ex:
H+ + Cl1H+ + SO42-
HCl
H2SO4
Binary Acids
H+ + anion
• H+ + anion with –ide ending 
Hydro _____ic acid
acid name is __________________
HCl
chloride
anion? _______
Hydrochloric
acid
acid name ________________
Binary Acids
H+ + anion
• H+ + anion with –ide ending 
Hydro _____ic acid
acid name is __________________
HF
fluoride
anion? _______
Hydrofluoric
acid
acid name ________________
Oxyacids
H+ + anion
• H+ + anion with –ate ending 
_____ic acid
acid name is __________________
HNO3
nitrate
anion? _______
nitric
acid
acid name ________________
Oxyacids
H+ + anion
• H+ + anion with –ate ending 
_____ic acid
acid name is __________________
H2SO4
sulfate
anion? _______
sulfuric
acid
acid name ________________
Oxyacids
H+ + anion
• H+ + anion with –ite ending 
_____ous acid
acid name is __________________
HNO2
nitrite
anion? _______
nitrous
acid
acid name ________________
Oxyacids
H+ + anion
• H+ + anion with –ite ending 
_____ous acid
acid name is __________________
HClO2
chlorite
anion? _______
chlorous
acid
acid name ________________
Writing acid formulas
• Hydrobromic acid
1-)
HBr
Bromide
(Br
anion? ___________ formula ______
• Acetic acid
1-)
HC
H
O
acetate(C
H
O
2
3
2
3
2
anion? ___________ formula ______ 2
• Nitrous acid
1-)
HNO
nitrite
(NO
2
2
anion? ___________ formula ______
Base Nomenclature
Nomenclature of Bases
• Bases are composed of a(n)
cation followed by
_______
hydroxide (OH1-)
a(n) ________________
Writing Base Names
• Rule: name the cation and add
“hydroxide”
• NaOH
• Mg(OH)2
• Fe(OH)3
sodium hydroxide
magnesium hydroxide
Iron (III) hydroxide
Memorize: NH3 = ammonia
Writing base formulas
• potassium hydroxide
+
KOH
K
cation? ______ formula ______
• Calcium hydroxide
2+
Ca(OH)
Ca
cation? ______ formula ______ 2
• Aluminum hydroxide
Al(OH)
3+
Al
cation? ______ formula ______ 3
Ions In Solution
• Why are some solutions acidic, basic,
or neutral?
It depends on number of H+ and OHions present.
Ions In Solution
• Acidic solution – contain more H+ ions
than OH- ions
4000 H+ and 0 OH- is acidic
1000 H+ and 500 OH- is acidic
5 H+ and 3 OH- is acidic
Ions In Solution
• Basic Solution – contain more OHions than H+ ions
4000 OH- and 0 H+ is basic
1000 OH- and 500 H+ is basic
5 OH- and 3 H+ is basic
Ions In Solution
• Neutral Solution – equal amounts of
H+ and OH- ions
4000 OH- and 4000 H+ is neutral
1000 OH- and 1000 H+ is neutral
5 OH- and 5 H+ is neutral
Self Ionization of Water
• Proper ionization
H2O + H2O  H3O+ + OHhydronium
ion
H
O
+
O
H
H
H
H

+
O
H
H
O
H
Self Ionization of Water
• simplified version
H2 O
H
+
+ OH-
Acid-Base Theories
Types of Acids/Bases
• Arrhenius Model
• Bronsted-Lowry Model
• Lewis Model
Svante Arrhenius
• He was a Swedish chemist (1859-1927),
and a Nobel prize winner in chemistry
(1903)
• one of the first chemists to explain the
chemical theory of the behavior of acids
and bases
• Dr. Hubert Alyea (professor emeritus at
Princeton University) was the last
graduate student of Arrhenius.
Hubert N. Alyea (1903-1996)
Svante Arrhenius (1859-1927)
Arrhenius Model of
Acids and Bases
• Arrhenius Acids
– Defn: contain H+ and ionizes to form H+
– Examples
+ + ClH
HCl 
+ + NO H
3
HNO3 
makes
solution
ACIDIC
Arrhenius Model of
Acids and Bases
• Arrhenius Bases
– Defn:
– contain OH- and ionizes to produce OH- ions
makes
– Examples
solution
NaOH 
Na+ + OH-
Ca(OH)2 
Ca2+ + 2 OH-
BASIC
Flaw with Arrhenius model
• Not all bases contain hydroxide
• Ex: ammonia (NH3) is basic
– According to Arrhenius, since ammonia can
NOT produce OH- it is NOT a base
• Therefore a new type of acid/base must
be determined
Gilbert Lewis (1875-1946)
Lewis Acids and Bases
• Gilbert Lewis focused on the
donation or acceptance of a pair of
electrons during a reaction
• Most general of all 3 definitions;
acids don’t even need hydrogen!
Lewis Model
• Lewis acid
– an atom, ion, or molecule that accepts an
electron pair to form a covalent bond
• Lewis base
– An atom, ion, or molecule that donates
an electron pair to form a covalent bond
Lewis Model
• Lewis acid-base reaction
– The formation of one or more covalent
bonds between an electron-pair donor
and an electron-pair acceptor
Example
– Acids are electron pair acceptors.
– Bases are electron pair donors.
Lewis
base
Lewis
acid
Johannes Brønsted
(1879-1947)
Denmark
Thomas Lowry
(1874-1936)
England
Bronsted-Lowry Model
• Bronsted-Lowry Acid
– Defn: proton/H+ donor
• can give H+ to another species
• Bronsted-Lowry Base
– Defn: proton/H+ acceptor
• can take H+ from another species
**Acids and bases always come in pairs.**
Bronsted-Lowry Model
• REMEMBER!!!! REMEMBER!!!!
acids donate, bases accept protons
Ashley does boys always
Bronsted-Lowry Model
• Examples
HCl + H2O  Cl- + H3O+
Acid
Base
(donates (accepts
proton) proton)
HCl is an acid —
when it dissolves in
water, it gives it’s
proton to water.
What is
happening
here?
Which is the
acid? base?
Water is a base—
when the HCl gives up
the proton, water
accepts it to form
the hydronium ion
Bronsted-Lowry Model
• Examples
NH3 + H2O  NH4+ +
Base
(accepts
proton)
Acid
(donates
proton)
What is
happening
OH here?
Which is the
acid? base?
Why Ammonia is a Base
• Ammonia can be explained as a
base by using Brønsted-Lowry:
NH3(aq) + H2O(l) ↔ NH41+(aq) + OH1(aq)
Ammonia is the hydrogen ion
acceptor (base), and water is the
hydrogen ion donor (acid).
Conjugate Acid/Base Pairs
• Conjugate acid – new species
produced when base gains H+ ion
• Conjugate base – new species
produced when acid donates H+ ion
Thus, a conjugate acid-base pair
is related by the loss or gain of a
single hydrogen ion.
Conjugate Acid/Base
Pairs
• general Bronsted-Lowry reaction
conj. acid/base pair
acid + base  conj. acid + conj. base
conj. acid/base pair
Every acid has a conjugate base.
Every base has a conjugate acid.
Conjugate Acid/Base
Pairs
• Examples
conj. acid/base pair
HNO3 + H2O  H3O+ + NO3acid
base
C.A.
conj. acid/base pair
C.B.
What is
the acid?
base?
What is
the
conjugate
acid/base?
Conjugate Acid/Base
Pairs
• Examples
conj. acid/base pair
NH3 + H2O  NH4+ + OHbase
acid
C.A.
conj. acid/base pair
C.B.
What is
the acid?
base?
What is
the
conjugate
acid/base?
Conjugate acid-base
pairs
• What is the conjugate base of:
1HSO
H2O
4
H2SO4 _________
H3O+ ________
• What is the conjugate acid of:
1H
PO
H2O
22
4
HPO4 _________
OH1- ________
How can H2O be both acid and base?
What is the conjugate
base?
Acid
H2SO4
HPO42NH41+
H3O1+
H2O
Conjugate Base
HSO41PO43NH3
H2O
OH-
How can H2O be both acid and base?
Amphoteric
• Defn – substance that can act as
both acids and bases
HNO3 + H2O  H3O+ + NO3base
NH3 + H2O  NH4+ + OHacid
Is H2O a
base or acid?
Is H2O a
base or acid?
Water is amphoteric b/c it acts as a
base in one reaction and acts as an acid
in the second
Mono-, Di-, Triprotic Acids
• Defns
– monoprotic (HA) – one ionizable proton
ex: HF, HCl, HBr (= normality is 1)
– diprotic (H2A) – two ionizable protons
ex: H2SO4, H2CO3 (= normality is 2)
– triprotic (H3A) – three ionizable protons
ex: H3PO4, H3BO3 (= normality is 3)
Polyprotic Acid
Ionization
• Always forms ONE H+
H3PO4 
H2PO41- + H+
H2PO41- 
HPO42- + H+
HPO42- 
PO43- + H+
Strength of Acids and
Bases
• Acid/base strength is based on 
the degree to which they ionize
1) strong ()
2) weak ()
Strong Acid/Base
• Defn – acid or base that completely
ionizes
HA
XOH
100%
ionization
100%
ionization
H+ + AX+ + OH-
every single HA molecule ionizes into
H+ and A-
Strength
-
+
• Strong Acid/Base
– 100% ionized in water
– strong electrolyte
HCl
HNO3
H2SO4
HBr
HI
HClO4
NaOH
KOH
Ca(OH)2
Ba(OH)2
Strong Acid
• Illustration
H
H
H
A
A
A
+
+
+
+
-
+
-
+
-
All break into ions
6 Strong Acids
•
•
•
•
•
•
HCl – hydrochloric acid
HBr – hydrobromic acid
HI – hydroiodic acid
HClO4 – perchloric acid
H2SO4 – sulfuric acid
HNO3 – nitric acid
Strong Bases
• Group I and II metal hydroxides
LiOH
Mg(OH)2
NaOH
Ca(OH)2
KOH
Sr(OH)2
RbOH
Ba(OH)2
No need to
memorize
exact ones
Weak Acid/Base
• Defn – acid or base that partially
ionizes
HA
XOH
partial
ionization
H+ + AX+ + OH-
not all will ionize; the weaker it is
the less it ionizes
Strength
-
+
• Weak Acid/Base
– does not ionize completely
– weak electrolyte
– Dissociates into both ions and molecules
HF
CH3COOH
H3PO4
H2CO3
HCN
NH3
Weak Acid
• Illustration
H
A
H
A
H
A
+
+
Only some break into ions
-
What are the weak acids
and bases?
• The ones that are NOT strong
Strong or weak,
concentrated or diluted
• For acids and bases, it is important
to distinguish between concentrated
and dilute from strong and weak. The
strong
concentrated
words _________
and __________
have different meanings. Similarly,
weak
dilute
___________
and ___________
are not the same either.
Strong or weak,
concentrated or diluted
• Strong and weak refer to
how much substance ionizes
____________________________
• Concentrated and dilute refer to
how much solute is present
____________________________
Example
• 1 M HCl
Strong and dilute
• 12 M HCl
Strong and concentrated
• 1 M H2CO3
weak and dilute
• 12 M H2CO3
weak and concentrated
Ion Product Constant for
Water (Kw)
• Defn: equilibrium value for self
ionization of water (H2O  H+ + OH-)
• Formula
Kw = [H+][OH-] = 1 x 10-14
ALWAYS
ALWAYS
ALWAYS
Ion Product Constant for
Water (Kw)
• in pure water 
[H+] = [OH-] = 1 x 10-7
Remember pure
water is neutral
• in non pure water 
(acidic/basic conditions), value of
[H+] and [OH-] differ
But still [H+][OH-] = 1 x 10-14
Is solution acidic, basic,
or neutral?
• Acidic
[H+] > [OH-]
• Basic
[H+] < [OH-]
• neutral
[H+] = [OH-]
The pH Scale
14
0
7
INCREASING
ACIDITY
pH = -log[
NEUTRAL
+
H]
pouvoir hydrogène (Fr.)
“hydrogen power”
INCREASING
BASICITY
The pH Scale
pH of Common Substances
Relation of pH and pOH
• pH + pOH = 14
If given one variable, subtract to find the other
The pH Scale
pH =
pOH =
+
-log[H ]
-log[OH ]
pH + pOH = 14
pH
c) change one pH unit 
represents a ten fold change in
strength
- ex: pH = 3 vs pH = 4
pH 3 is 101 or 10 times more acidic
- ex: pH = 7 vs pH = 10
pH 7 is 103 or 1000 times more acidic
Overall Relationship
pH
pOH
[H+]
[OH-]
Overall Relationship
The pH Scale
• What is the pH of 0.050 M HNO3?
pH = -log[H+]
pH = -log[0.050]
pH = 1.3
Acidic or basic? Acidic
Calculating [H+] and [OH-]
• Find the hydroxide ion concentration of
3.0  10-2 M HCl.
[H+][OH-] = 1.0  10-14
[3.0  10-2][OH-] = 1.0  10-14
[OH-] = 3.3  10-13 M
Acidic or basic?
Acidic
Calculating [H+] and [OH-]
• What is the molarity of HBr in a solution
that has a pOH of 9.6?
pH + pOH = 14
pH = -log[H+]
pH + 9.6 = 14
4.4 = -log[H+]
pH = 4.4
-4.4 = log[H+]
Acidic
[H+] = 4.0  10-5 M HBr
Calculating [H+] and [OH-]
• A Ca(OH)2 solution has a pH of 8.0.
Determine the [H+], [OH-], and [Ca(OH)2]
for the solution.
pH = - log [H+]
log [H+] = -pH
Antilog is the same
thing as 10^(x)
[H+] = antilog (-pH) = antilog (-8.0)
= 1 x 10-8 M H3O+
Calculating [H+] and [OH-]
[H+] [OH-] = 1 x 10-14 M2
[OH-] = 1 x 10-14 M2 = 1 x 10-14 M2
[H+]
1 x 10-8 M
[OH-] = 1 x 10-6 M
Calculating [H+] and [OH-]
• Ca(OH)2
Ca2+ +
2OH( 1 x 10-6 M)
1 x 10-6 mol OH- 1 mol Ca(OH)2 = 5 x 10-7 mol/L Ca(OH)2
Liter
2 mol OH-
[Ca(OH)2] = 5 x 10-7 M
Sample problem #1
• Calculate the pH of a solution with
[H+] = 3.0 x 10-6 M.
pH = -log [H+]
= - log [3.0 x 10-6]
= 5.52
Sample problem #2
• Calculate the pH of a solution with
[OH-] = 8.2 x 10-6. [OH-]  pOH  pH
pOH = -log[OH-]
= -log [8.2 x 10-6]
= 5.09
pH + pOH = 14
pH + 5.09 = 14
pH = 8.91
Sample problem #2
• Calculate the pH of a solution with
[OH-] = 8.2 x 10-6. [OH-]  [H+]  pH
[H+][OH-] = 1 x 10-14
[H+][8.2 x 10-6] = 1 x 10-14
[H+] = 1.22 x 10-9
pH = -log[1.22 x 10-9] = 8.91
Sample problem #3
• What is the [H+] of a solution with
pH = 2?
pH = -log[H+]
[H+] = 10-pH
= 10-2
= 0.01 M
Sample problem #4
• (i) What is the [OH-] of a solution
with pOH = 3.7?
[OH-] = 10-pOH
[OH-] = 10-3.7
= 2 x 10-4 M
Sample problem #4
•(ii) What is the pH and the [H+] if the
pOH is 3.7?
pOH  pH  [H+]
pH + pOH = 14
pH + 3.7 = 14
pH = 10.3
[H+] = 10-pH
= 10-10.3
= 5 x 10-11 M
Reaction between acids
and bases
• Neutralization (defn) – reaction of
acid and base to form a salt and
water
– The reaction is a double replacement
• Salt (defn) – ionic compound made of
– cation from base and
– anion from acid
Reaction between acids
and bases
• Ex reaction
Mg(OH)2 + HCl 
base
Mg2+ + OH-
acid
H+ + Cl-
MgCl2 + H2O
salt
water
Mg2+ - cation from base
Cl- - anion from acid
Neutralization
ACID + BASE  SALT + WATER
HCl + NaOH  NaCl + H2O
strong
strong
neutral
HC2H3O2 + NaOH  NaC2H3O2 + H2O
weak
strong
basic
– Salts can be neutral, acidic, or basic.
– Neutralization does not mean pH = 7.
Ex problems
• i) What is the salt formed from
sulfuric acid (H2SO4) and
potassium hydroxide (KOH)?
base cation? K+
2SO
acid anion?
4
What is salt?
K2SO4
Ex problems
• ii) What is the salt formed when
Al(OH)3 and HBr react?
base cation?
acid anion?
Al3+
Br-
What is salt? AlBr3
Is salt solution acidic,
basic, or neutral?
• a) strong acid + strong base 
Neutral salt
• b) strong acid + weak base 
acidic salt
• c) weak acid + strong base 
basic salt
Ex problem
•
Determine if salt solution is acidic,
basic, or neutral.
a) LiBr  Li+ + Bracid? HBr (strong acid)
base? LiOH (strong base)
NEUTRAL
Salt is ____________
Ex problem
•
Determine if salt solution is acidic,
basic, or neutral.
b) Fe(NO3)3  Fe3+ + NO3(strong acid)
acid? HNO3
base? Fe(OH)3 (weak base)
ACIDIC
Salt is ____________
B. Titration
• Titration
standard solution
– Analytical method in
which a standard
solution is used to
determine the
concentration of an
unknown solution.
unknown solution
B. Titration
• Equivalence point (endpoint)
– Point at which equal amounts of
H3O+ and OH- have been added.
– Determined by…
• indicator color change
• dramatic change in pH
Titration
• dramatic change in pH
B. Titration
+
O
moles H3 = moles
MVn = MVn
OH
M: Molarity
V: volume
n: # of H+ ions in the acid
or OH- ions in the base
Titration
• 42.5 mL of 1.3M KOH are required to
neutralize 50.0 mL of H2SO4. Find the
molarity of H2SO4.
H3O+
OH-
M=?
M = 1.3M
V = 50.0 mL
n=2
V = 42.5 mL
n=1
Subscript of H or
OH from formulas!
MV# = MV#
M(50.0mL)(2)
=(1.3M)(42.5mL)(1)
M = 0.55M H2SO4