Acids and Bases
http://www.unit5.org/chemistry/AcidBase.html
Guiding Questions
What is an acid?
What makes acids dangerous?
Is acid rain an issue for us?
What does pH balanced mean?
Table of Contents
‘Acids, Bases, and Salts’
Definitions
pH Scale
Acid Rain
Common Acids
Common Bases
LeChatelier’s Principle
Maintaining Blood pH
Conjugate Acid-Base Pairs
pKa
Concentration vs. Strength
pH Indicators
Buffers
Titration
Swimming Pools
Acids, Bases, and Salts
You should be able to
Understand the acid-base theories of Arrhenius, Brønsted-Lowry,
and Lewis.
Identify strong acids and bases and calculate their pH’s.
Calculate the pH of a weak acid or base.
Calculate the concentration of a strong or weak acid or base from
its pH.
Calculate the pH and ion concentration in a polyprotic acid.
Predict the pH of a salt from its formula and then calculate the pH
of the salt.
Be familiar with titration curves and selection of an acid-base indicator.
Lecture Outline – Acids and Bases
Lecture Outline – Acids and Bases
student notes outline
textbook questions
Lecture Outline – Acids and Bases
textbook questions
Keys
text
http://www.unit5.org/chemistry/AcidBase.html
pH scale: measures acidity/basicity
ACID
BASE
10x10x
100x 10x
0
1
2
3
4
5
6
7
8
9
10
11
12
13
14
NEUTRAL
Each step on pH scale represents a factor of 10.
pH 5 vs. pH 6
pH 3 vs. pH 5
pH 8 vs. pH 13
(10X more acidic)
(100X different)
(100,000X different)
pH scale: measures acidity/basicity
Søren Sorensen
(1868 - 1939)
ACID
BASE
10x10x
100x 10x
0
1
2
3
4
5
6
7
8
9
10
11
12
13
14
NEUTRAL
Each step on pH scale represents a factor of 10.
pH 5 vs. pH 6
pH 3 vs. pH 5
pH 8 vs. pH 13
(10X more acidic)
(100X different)
(100,000X different)
Richter Scale - Earthquakes
7
6
5
.
4
3
2
1
pH = -log [H1+]
Acid
Base
[H+]
Acidic
[H+] = [OH-]
Neutral
pH = 7
Basic
[OH-]
radius = 112.8 cm
radius = 35.7 cm
.
radius = 11.3 cm
radius = 3.6 cm
radius = 1.1 cm
radius = 0.1128 cm
pH = 1
pH = 2
pH = 3
.
pH = 4
5
6
Acids and Bases
pH < 7
taste sour
react w/bases
proton (H1+) donor
turn litmus red
lots of H1+/H3O1+
react w/metals
pH > 7
taste bitter
react w/acids
proton (H1+) acceptor
turn litmus blue
lots of OH1–
don’t react w/metals
Both are electrolytes.
Acid vs. Base
Different
Alike
pH < 7
Affects pH
and
litmus paper
Topic
sour taste
react with
metals
Acid
Different
pH > 7
Topic
Related to
H+ (proton)
concentration
pH + pOH = 14
Base
bitter taste
does not
react with
metals
Properties
electrolytes
electrolytes
sour taste
bitter taste
turn litmus red
turn litmus blue
react with metals to
form H2 gas
slippery feel
vinegar, milk, soda,
apples, citrus fruits
ammonia, lye, antacid,
baking soda
ChemASAP
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.
Common Acids and Bases
Strong Acids (strong electrolytes)
HCl
HNO3
HClO4
H2SO4
hydrochloric acid
nitric acid
perchloric acid
sulfuric acid
Weak Acids (weak electrolytes)
CH3COOH
H2CO3
acetic acid
carbonic
Strong Bases (strong electrolytes)
NaOH
KOH
Ca(OH)2
sodium hydroxide
potassium hydroxide
calcium hydroxide
Weak
Weak Base
Base (weak
(weak electrolyte)
electrolyte)
NH
NH43OH
ammonia
NH3 + H2O  NH4OH
Kotz, Purcell, Chemistry & Chemical Reactivity 1991, page 145
Acid + Base  Salt + Water
• Orange juice + milk  bad taste
• Evergreen shrub + concrete  dead bush
• Under a pine tree + fertilizer  white powder
HCl + NaOH  NaCl + HOH
salt
water
Acid-Base Neutralization
1-
1+
+
+
H32O+
OH
H2O-
H2O
H2O
Water ion
Hydronium
Hydroxide
Water ion
Water
Water
Dorin, Demmin, Gabel, Chemistry The Study of Matter 3rd Edition, page 584
Acid-Base Neutralization
1-
1+
+
+
H3O+
OH-
H2O
H2O
Hydronium ion
Hydroxide ion
Water
Water
Dorin, Demmin, Gabel, Chemistry The Study of Matter 3rd Edition, page 584
Acid Precipitation
http://nadp.sws.uiuc.edu/amaps2/
Click
Here
Formation of Sulfuric Acid
+
+
SO2(g) + H2O(l)
H2SO3(aq)
2SO2(g) + O2(g)
2SO3(g)
SO3(g) + H2O(l)
H2SO4(aq)
Sulfuric acid
Catalyzed by atmospheric dust
SO2(g) + H2O2(l)
Kelter, Carr, Scott, Chemistry A World of Choices 1999, page 302
H2SO4(aq)
Hydrogen ion concentration as pH from measurements
made
at the
field laboratories
during 2003
made
at the
Central
Analytical Laboratory,
1999
The progressively darker red areas on the map indicate the lowest pH levels
and areas most prone to problems from acid rain.
National Atmospheric Deposition Program/National Trends Network
http://nadp.sws.uiuc.edu
Figure courtesy of the National Atmospheric Deposition Program, Champaign, Ill.
Acid Rain
Estimated sulfate ion deposition, 1999
Smoke stacks pollute SO2
into the atmosphere. This
combines with water to form
acid rain.
SO4 Levels
http://nadp.sws.uiuc.edu
Coal Burning Power Plant
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.
CO2 (g)
H2O (l)
H2CO3 (aq)
Carbon
dioxide
Water
Carbonic
acid
Weak
acid
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.
Carbon
dioxide
Water
Carbonic
Acid
Carbon
dioxide
Water
Carbonic
Acid
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.
Common Acids
Sulfuric Acid
H2SO4
Battery acid
Nitric Acid
HNO3
Used to make fertilizers
and explosives
Phosphoric Acid
H3PO4
Food flavoring
Hydrochloric Acid
HCl
Stomach acid
Acetic Acid
Carbonic Acid
CH3COOH
H2CO3
Vinegar
Carbonated water
Common Acids
Formula
Name of Acid
Name of Negative
Ion of Salt
HF
HBr
HI
HCl
HClO
HClO2
HClO3
HClO4
H2S
H2SO3
H2SO4
HNO2
HNO3
H2CO3
H3PO3
H3PO4
hydrofluoric
hydrobromic
hydroiodic
hydrochloric
hypochlorous
chlorous
chloric
perchloric
hydrosulfuric
sulfurous
sulfuric
nitrous
nitric
carbonic
phosphorous
phosphoric
fluoride
bromide
iodide
chloride
hypochlorite
chlorite
chlorate
perchlorate
sulfide
sulfite
sulfate
nitrite
nitrate
carbonate
phosphite
phosphate
Formation of Hydronium Ions
1+
1+
1+
+
H+
H2O
hydrogen ion
(a proton)
water
H3O+
hydronium ion
Sulfuric Acid, H2SO4
Sulfuric acid is the most commonly produced industrial chemical in the world.
Uses: petroleum refining, metallurgy, manufacture of fertilizer,
many industrial processes: metals, paper, paint, dyes, detergents
Sulfuric acid is used in
automobile batteries.
H2SO4
“oil of vitriol”
Nitric Acid, HNO3
Nitric acid stains proteins yellow (like your skin).
Uses: make explosives, fertilizers, rubber, plastics, dyes, and pharmaceuticals.
HNO3
“aqua fortis”
Hydrochloric Acid, HCl
The stomach produces HCl to aid in the digestion of food.
Uses: For ‘pickling’ iron and steel.
Pickling is the immersion of metals in acid solution to remove
surface impurities.
A dilute solution of HCl is called muriatic acid (available in many hardware
stores). Muriatic acid is commonly used to adjust pH in swimming pools
and in the cleaning of masonry.
HCl(g) + H2O(l)
hydrogen chloride
water
HCl(aq)
hydrochloric acid
Common Bases
Name
Formula
Common Name
Sodium hydroxide
NaOH
lye or caustic soda
Potassium hydroxide
KOH
lye or caustic potash
Magnesium hydroxide
Mg(OH)2
milk of magnesia
Calcium hydroxide
Ca(OH) 2
slaked lime
Ammonia water
NH
H 2O
NH43.OH
household ammonia
NH41+ + OH1ammonium hydroxide
Relative Strengths of Acids and Bases
perchloric
hydrogen chloride
nitric
sulfuric
hydronium ion
hydrogen sulfate ion
phosphoric
acetic
carbonic
hydrogen sulfide
ammonium ion
hydrogen carbonate ion
water
ammonia
hydrogen
Metcalfe, Williams, Catska, Modern Chemistry 1966, page 229
Formula
HClO4
HCl
HNO3
H2SO4
H3O+
HSO4H3PO4
HC2H3O2
H2CO3
H2S
NH4+
HCO3H2O
NH3
H2
acid
Conjugate base
Formula
perchlorate ion
chloride ion
nitrate ion
hydrogen sulfate ion
water
sulfate ion
dihydrogen phosphate ion
acetate ion
hydrogen carbonate ion
hydro sulfide ion
ammonia
carbonate ion
hydroxide ion
amide ion
hydride ion
conjugate base + H+
ClO4ClNO3HSO4H2O
SO42H2PO4C2H3O2HCO3HSNH3
CO32OHNH2H-
Decreasing Base Strength
Decreasing Acid Strength
Acid
Binary Hydrogen Compounds
of Nonmetals When Dissolved in Water
(These compounds are commonly called acids.)
The prefix hydro- is used to represent hydrogen, followed by the name
of the nonmetal with its ending replaced by the suffix –ic and the word
acid added.
Examples:
*HCl
Hydrochloric acid
HBr
Hydrobromic acid
*The name of this compound would be hydrogen chloride if it was NOT dissolved in water.
Naming Simple Chemical Compounds
Ionic (metal and nonmetal)
Metal
Forms
only one
positive
ion
Use the
name of
element
Forms
more than
one positive
ion
Covalent (2 nonmetals)
Nonmetal
Single
Negative
Ion
Use element
Use the name
name followed
of the
by a Roman
element, but
numeral to
end with ide
show the charge
First
nonmetal
Second
nonmetal
Before
element name
use a prefix
to match
subscript
Use a prefix
before
element name
and end
with ide
Polyatomic
Ion
Use the
name of
polyatomic
ion (ate or
Ite)
Naming Ternary Compounds
from Oxyacids
The following table lists the most common families of oxy acids.
one more
oxygen atom
HClO4
perchloric acid
most
“common”
HClO3
chloric acid
H2SO4
sulfuric acid
H3PO4
phosphoric acid
HNO3
nitric acid
one less
oxygen
HClO2
chlorous acid
H2SO3
sulfurous acid
H3PO3
phosphorous acid
HNO2
nitrous acid
two less
oxygen
HClO
hypochlorous acid
H3PO2
hypophosphorous acid
(HNO)2
hyponitrous acid
An acid with a
name ending in
A salt with a
name ending in
-ous
forms
-ite
-ic
forms
-ate
Hill, Petrucci, General Chemistry An Integrated Approach 1999, page 60
Oxyacids  Oxysalts
If you replace hydrogen with a metal, you have formed an oxysalt.
A salt is a compound consisting of a metal and a non-metal. If the
salt consists of a metal, a nonmetal, and oxygen it is called an
oxysalt. NaClO4, sodium perchlorate, is an oxysalt.
OXYACID
OXYSALT
HClO4
perchloric acid
NaClO4
sodium perchlorate
HClO3
chloric acid
NaClO3
sodium chlorate
HClO2
chlorous acid
NaClO2
sodium chlorite
HClO
hypochlorous acid
NaClO
sodium hypochlorite
ACID
SALT
per stem ic
changes to
per stem ate
stem ic
changes to
stem ate
stem ous
changes to
stem ite
hyper stem ous
changes to
hypo stem ite
HClO3
acid
+
Na1+
cation
NaClO3 + H1+
salt
Acid Definitions
Lewis Acid
Brønsted-Lowry
Arrhenius
acids
Arrhenius Acids and Bases
Acids release hydrogen ions in water.
Bases release hydroxide ions in water.
An acid is a substance that produces hydronium ions, H3O+,
when dissolved in water.
Brønsted-Lowry Definitions
A Brønsted-Lowry acid is a proton donor; it donates a hydrogen ion, H+.
A Brønsted-Lowry base is a proton acceptor; it accepts a hydrogen ion, H+.
Lewis Definitions
A Lewis acid is a substance than can accept (and share) an electron pair.
A Lewis base is a substance than can donate (and share) an electron pair.
Acid Definitions
Lewis acids
The Arrhenius model of acids
and bases was broadened by
the Brønsted-Lowry model.
Brønsted-Lowry
The Lewis acid-base model is
the most general in scope.
Arrhenius
acids
Ralph A. Burns, Fundamentals of Chemistry 1999, page 483
The Lewis definition of an acid
includes any substance that
is an electron pair acceptor;
a Lewis base is any substance
that can act as an electron pair
donor.
Acid Definitions
Lewis acids
The Arrhenius model of acids
and bases was broadened by
the Brønsted-Lowry model.
Brønsted-Lowry
The Lewis acid-base model is
the most general in scope.
Arrhenius
acids
Ralph A. Burns, Fundamentals of Chemistry 1999, page 483
The Lewis definition of an acid
includes any substance that
is an electron pair acceptor;
a Lewis base is any substance
that can act as an electron pair
donor.
Acid – Base Systems
Type
Acid
Base
Arrhenius
H+ or H3O +
producer
OH - producer
BrønstedLowry
Lewis
Proton (H +)
donor
Proton (H +)
acceptor
Electron-pair
acceptor
Electron-pair
donor
Arrhenius Acid
Any substance that releases H+ ions as the
only positive ion in the aqueous solution.
1-
1+
+
+
HCl
H2O
H3O+
Cl-
hydrogen chloride
(an Arrhenius acid)
water
hydronium ion
chloride ion
Definitions
 Arrhenius
- In aqueous solution…
• Acids form hydronium ions (H3O+)
HCl + H2O  H3O+ + Cl–
H
H
Cl
O
H
O
H
H
–
+
Cl
H
acid
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Definitions
 Arrhenius
- In aqueous solution…
• Bases form hydroxide ions (OH-)
NH3 + H2O  NH4 +
+
H
H
H
N
H
O
H
–
+
O
N
H
H
OH
H
H
base
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H
Arrhenius Bases and Their Properties
According to the definition of Arrhenius a:
Base - "a substance whose water solution yields...
hydroxide ions (OH-) as the only negative ions."
Are NaOH and NH3 considered to be Arrhenius bases? YES
1) Bases are electroytes
Dissociation equation for NaOH
NaOH(s)
Na1+(aq) + OH1-(aq)
Dissociation equation for NH3
NH3(g) + H2O(l)
NH41+(aq) + OH1-(aq)
2) Bases cause indicators to turn a characteristic color
3) Bases neutralize acids
NaOH(aq) + HCl(aq)
NaCl(aq) + H2O(l)
4) Water solutions of bases tasted bitter and feel slippery.
Neutralization
Neutralization is a chemical reaction between an acid and a base
to produce a salt (an ionic compound) and water.
NaOH(aq) + HCl(aq)
base
acid
NaCl(aq) + H2O(l)
salt
water
Some neutralization reactions:
H2SO4(aq) + 2 NaOH(aq)
sulfuric acid
2 HC2H3O2(aq) +
acetic acid
sodium hydroxide
Ca(OH)2(aq)
calcium hydroxide
Na2SO4 +
sodium sulfate
2 HOH
water
Ca(C2H3O2)2 + 2 HOH
calcium acetate
water
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.
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
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.
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Salts
NaCl
Salts - Ionic compounds containing a positive ion other than the
hydrogen
ion andand
a negative
than salts
the hydroxide ion.
Formulas
namesion
of other
common
i.e., a metal and a non-metal
SALT
FORMULA
Common Name
Under what conditions do salts conduct current?
sodium chloride
NaCl(s) + H2O(l)
sodium nitrate
NaCl
(table) salt
Na1+(aq) + Cl1-(aq)
NaNO3
Chile saltpeter
sodium bicarbonate
NaHCO3
baking soda
potassium carbonate
K2CO3
potash
ammonium chloride
NH4Cl
sal ammoniac
Salt Formation
strong
base
strong
acid
NaOH
HCl
salt of a strong base and a strong acid
NaCl
NaOH + HCl  NaCl + H2O
strong
base
weak
acid
NaOH
HC2H3O2
NaC2H3O2
salt of a strong base and a weak acid
NaOH + HC2H3O2  NaC2H3O2 + H2O
Note: that in each case H-OH (water) is formed
Salt Formation
weak
base
strong
acid
NH3
H2SO4
salt of a weak base and a strong acid
(NH4) 2 SO4
H2SO4
NH4OH
NH4OH + H2SO4  (NH4)2SO4 + H2O
weak
base
weak
acid
NH3
HC2H3O2
NH4 C2H3O2
salt of a weak base and a weak acid
NH4OH + HC2H3O2  NH4C2H3O2 + H2O
Note: that in each case H-OH (water) is also formed
weak
base
strong
acid
NH3
H2SO4
(NH4) 2 SO4
2 NH4OH
ammonium
ion
+
hydroxide
ion
salt of a weak base and a strong acid
H2SO4
(NH4)2SO4
sulfuric
acid
ammonium sulfate
+
2 HOH
water
sulfate ion
1-
1+
H2SO4
NH4OH
21+
1-
1+
2 NH4OH
NH4+
NH4+
1+
OHOH-
+
H2SO4
H2SO4
(NH4)2SO4
(NH4)2SO4
+
2 H 2O
HOH
HOH
Reactions that produce salt
acid
+
base
salt
+
water
H3PO4
+
NH4OH
(NH4)3PO4
+
H2O
phosphoric acid and ammonium hydroxide
HNO3
nitric acid
H2CO3
Mg(OH)2
magnesium hydroxide
KOH
yields ammonium phosphate and water
Mg(NO3)2
magnesium nitrate
K2CO3
carbonic acid
potassium hydroxide
potassium carbonate
HC2H3O2
Al(OH)3
Al(C2H3O2)3
acetic acid
aluminum hydroxide
aluminum acetate
HClO4
perchloric acid
Ba(OH)2
barium hydroxide
H2O
Ba(ClO4)2
barium perchlorate
H2O
H2O
H2O
Brønsted-Lowry Acids and Bases
Acid = any substance that donates a proton.
Base = any substance that accepts a proton.
d+
1-
1+
d-
+
HCl
H2O
H3O+
Cl-
(acid)
(base)
hydronium ion
chloride ion
Brønsted-Lowry Acids and Bases
Acid = any substance that donates a proton.
Base = any substance that accepts a proton.
d+
1-
1+
d-
+
HCl
H2O
H3O+
Cl-
(acid)
(base)
hydronium ion
chloride ion
Brønsted-Lowry Acids and Bases
d-
1-
1+
d+
+
NH3
H2O
(base)
(acid)
NH4+
ammonium ion
OH-
hydroxide ion
Brønsted-Lowry Acids and Bases
d-
1-
1+
d+
+
NH3
H2O
(base)
(acid)
NH4+
ammonium ion
OH-
hydroxide ion
Brønsted-Lowry Acids and Bases
d-
1-
1+
d+
+
NH3
H2O
(base)
(acid)
NH4+
ammonium ion
OH-
hydroxide ion
Brønsted-Lowry Acids and Bases
d-
1-
1+
d+
+
NH3
H2O
(base)
(acid)
NH4+
ammonium ion
OH-
hydroxide ion
Definitions
 Brønsted-Lowry
• Acids are proton (H+) donors.
• Bases are proton (H+) acceptors.
HCl + H2O 
acid
–
Cl
+
+
H3O
base
conjugate base
conjugate acid
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Definitions
H2O + HNO3  H3O+ + NO3–
B
A
CA
Base
Acid
O
H O
H
N
H O
O
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
CB
Definitions
NH3 + H2O 
B
A
+
NH4
CA
Base
+
OH
CB
Acid
H
H N
H
 Amphoteric
H O
H
- can be an acid or a base.
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Definitions

Give the conjugate base for each of the following:
-
HF
F
H3PO4
H2PO4
+
H3O
H2O
 Polyprotic
- an acid with more than one H+
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Definitions

Give the conjugate acid for each of the following:
Br
-
HBr
HSO4
H2SO4
2CO3
HCO3
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Definitions
 Lewis
• Acids are electron pair acceptors.
• Bases are electron pair donors.
Lewis
base
Lewis
acid
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Basic
7
Acid
14
Neutral
pH Scale
Acidic
0
Base
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 515
[H+]
pH
10-14
14
10-13
13
10-12
12
10-11
11
10-10
10
10-9
9
10-8
8
10-7
7
10-6
6
10-5
5
10-4
4
10-3
3
10-2
2
10-1
1
100
0
1 M NaOH
Ammonia
(household
cleaner)
Blood
Pure water
Milk
Vinegar
Lemon juice
Stomach acid
1 M HCl
pH of Common Substances
gastric
juice
1.6
vinegar
2.8
carbonated
beverage
3.0
0
1
2
acidic
Timberlake, Chemistry 7th Edition, page 335
urine
6.0
4
5
bile
8.0
6
7
neutral
[H+] = [OH-]
8
ammonia
11.0
bleach
12.0
seawater
8.5
9
1.0 M
NaOH
(lye)
14.0
milk of
magnesia
10.5
detergents
8.0 - 9.0
milk
6.4
tomato
4.2
coffee
5.0
3
blood
7.4
potato
5.8
apple juice
3.8
lemon
juice
2.2
drinking water
7.2
bread
5.5
orange
3.5
1.0 M
HCl
0
water (pure)
7.0
soil
5.5
10
11
basic
12
13
14
pH of Common Substance
More acidic
More basic
pH
NaOH, 0.1 M
Household bleach
Household ammonia
Lime water
Milk of magnesia
Borax
Baking soda
Egg white, seawater
Human blood, tears
Milk
Saliva
Rain
Black coffee
Banana
Tomatoes
Wine
Cola, vinegar
Lemon juice
Gastric juice
14
13
12
11
10
9
8
7
76
5
4
3
2
1
0
[H1+]
[OH1-]
pOH
1 x 10-14
1 x 10-13
1 x 10-12
1 x 10-11
1 x 10-10
1 x 10-9
1 x 10-8
1 x 10-7
1 x 10-6
1 x 10-5
1 x 10-4
1 x 10-3
1 x 10-2
1 x 10-1
1 x 100
1 x 10-0
1 x 10-1
1 x 10-2
1 x 10-3
1 x 10-4
1 x 10-5
1 x 10-6
1 x 10-7
1 x 10-8
1 x 10-9
1 x 10-10
1 x 10-11
1 x 10-12
1 x 10-13
1 x 10-14
0
1
2
3
4
5
6
8
9
10
11
12
13
14
Acid – Base Concentrations
concentration (moles/L)
10-1
pH = 3
pH = 11
OH-
H3O+
pH = 7
10-7
H3O+
OH-
OH-
H3O+
10-14
Timberlake, Chemistry 7th Edition, page 332
[H3O+] > [OH-]
[H3O+] = [OH-]
acidic
solution
neutral
solution
[H3O+] < [OH-]
basic
solution
pH
pH = -log [H1+]
Kelter, Carr, Scott, Chemistry A World of Choices 1999, page 285
pH Calculations
pH
pH = -log[H3O+]
[H3O+]
[H3O+] = 10-pH
[H3O+] [OH-] = 1 x10-14
pH + pOH = 14
pOH
pOH = -log[OH-]
[OH-]
[OH-] = 10-pOH
pH = - log [H+]
Given: pH = 4.6
pH = - log [H+]
choose proper equation
4.6 = - log [H+]
substitute pH value in equation
- 4.6 =
2nd
log
determine the [hydronium ion]
- 4.6 =
log [H+]
log [H+]
[H+] = 2.51x10-5 M
multiply both sides by -1
take antilog of both sides
Recall, [H+] = [H3O+]
10x
antilog
You can check your answer by working backwards.
pH = - log [H+]
pH = - log [2.51x10-5 M]
pH = 4.6
Acid Dissociation
monoprotic
e.g. HCl, HNO3
HA(aq)
0.03 M
H1+(aq) + A1-(aq)
0.03 M
0.03 M
pH = ?
pH = - log [H+]
pH = - log [0.03M]
pH = 1.52
diprotic
e.g. H2SO4
H2A(aq)
0.3 M
2 H1+(aq) + A2-(aq)
0.6 M
0.3 M
pH = - log [H+]
pH = - log [0.6M]
pH = 0.22
polyprotic
e.g. H3PO4
H3PO4(aq)
?M
3 H1+(aq) + PO43-(aq)
xM
Given: pH = 2.1
find [H3PO4]
assume 100%
dissociation
Given: pH = 2.1
3 H1+(aq) + PO43-(aq)
0.00794 M
H3PO4(aq)
XM
find [H3PO4]
assume 100%
dissociation
Step 1) Write the dissociation of phosphoric acid
Step 2) Calculate the [H+] concentration
pH = - log [H+]
[H+] = 10-pH
2.1 = - log [H+]
[H+] = 10-2.1
- 2.1 = log [H+]
[H+] = 0.00794 M
2nd
7.94 x10-3 M
log
- 2.1 =
2nd
log log [H+]
[H+] = 7.94 x10-3 M
Step 3) Calculate [H3PO4] concentration
Note: coefficients (1:3) for (H3PO4 : H+)
7.94 x10-3 M = 0.00265 M H PO
3
4
3
How many grams of magnesium hydroxide are needed to add to 500 mL of H2O
to yield a pH of 10.0?
Step 1) Write out the dissociation of magnesium hydroxide
Mg(OH)2(aq)
-4 M
0.5
5 x10-5
Step 2) Calculate the pOH
Step 3) Calculate the [OH1-]
Mg2+(aq) + 2 OH1-(aq)
1 x10-4 M
Mg2+ OH1Mg(OH)2
pH + pOH = 14
10.0 + pOH = 14
pOH = 4.0
pOH = - log [OH1-]
[OH1-] = 10-OH
[OH1-] = 1 x10-4 M
Step 4) Solve for moles of Mg(OH)2
x mol
mol
5 x10 -5 M 
0.5 L
L
Step 5) Solve for grams of Mg(OH)2
M
x g Mg(OH)2 = 2.5 x 10-5 mol Mg(OH)2
x = 2.5 x 10-5 mol Mg(OH)2
58 g Mg(OH)2
= 0.00145 g Mg(OH)2
1 mol Mg(OH)2
I WISH I HAD
SWEAT GLANDS.
Equilibrium
• LeChatelier’s Principle
CO2 +
CaO
CaCO3
“chicken
breath”
“food”
“egg shell”
As temperature increases, chickens “pant” more.
This stresses the system and shifts the equilibrium to the
This makes the egg shells THIN and fragile.
LEFT
I wish I had
sweat glands.
In a chicken… CaO
+
CO2
In summer, [ CO2 ] in a chicken’s blood
-- shift
CaCO3
(eggshells)
due to panting.
; eggshells are thinner
How could we increase eggshell thickness in summer?
[ CO2 ]
, shift
-- give chickens carbonated water
[ CaO ]
, shift
-- put CaO additives in chicken feed
-- air condition the chicken house
TOO much $$$
-- pump CO2 gas into the chicken house
would kill all the chickens!
LeChatelier’s Principle
N2 + 3 H2
2 NH3 + heat
Raising the temperature…
Increasing the pressure…
Decreasing the concentration
of NH3…
Dorin, Demmin, Gabel, Chemistry The Study of Matter 3rd Edition, page 532
…favors the endothermic reaction (the reverse
reaction) in which the rise in temperature is
counteracted by the absorption of heat.
…favors the forward reaction in which 4 mol
of gas molecules is converted to 2 mol.
…favors the forward reaction in order to
replace the NH3 that has been removed.
Animation by Raymond Chang
All rights reserved
Equilibrium Expression
Haber Process
N2 + 3 H2

products 
K eq 
reactants 
Dorin, Demmin, Gabel, Chemistry The Study of Matter 3rd Edition, page 532
2 NH3 + heat

NH3 
K eq 
3
N2 H2 
2
reversible reaction:
Reactant
 Product
Reactant
and
P  R
Product
Acid dissociation is a reversible reaction.
H2SO4
equilibrium:
2 H1+ + SO42–
Rate at which
Rate at which
R  P
P  R
=
looks like nothing is happening, however…
system is dynamic, NOT static
Le Chatelier’s principle
Le Chatelier’s principle:
When a system at equilibrium is disturbed, it shifts to a
new equilibrium that counteracts the disturbance.
N2(g) + 3 H2(g)
Disturbance
2 NH3(g)
Equilibrium Shift
Add more N2…………………..
“
“
H2…………………..
“
“
NH3…………………
Remove NH3…………………..
Add a catalyst…………………
Increase pressure…………….
Fritz Haber
no shift
Light-Darkening Eyeglasses
AgCl + energy
(clear)
Go outside…
Ago + Clo
(dark)
Sunlight more intense than inside light;
“energy”
shift
to a new equilibrium: GLASSES DARKEN
Then go inside…
“energy”
shift
to a new equilibrium: GLASSES LIGHTEN
Sensitive Sunglasses
Oxidation-reduction reactions are the basis for many interesting and useful applications of technology.
One such application is photochromic glass, which is used for the lenses in light sensitive glasses.
Lenses manufactured by the Corning Glass Company can change from transmitting 85% of light to only
transmitting 22% of light when exposed to bright sunlight.
Photochromic glass is composed of linked tetrahedrons of silicon and oxygen atoms jumbled together
in a disorderly array, with crystals of silver chloride caught in between the silica tetrahedrons. When the
glass is clear, the visible light passes right through the molecules. The glass absorbs ultraviolet light,
however, and this energy triggers an oxidation-reduction reaction
between Ag+ and Cl-:
Ag+ + Cl- --> Ag0 + Cl0
To prevent the reaction from reversing itself immediately, a few ions of Cu+ are incorporated into the
silver chloride crystal. These Cu+ ions react with the newly formed chlorine atoms:
Cu+ + Cl0 --> Cu2+ + ClThe silver atoms move to the surface of the crystal and form small colloidal clusters of silver metal.
This metallic silver absorbs visible light, making the lens appear dark (colored).
As the glass s removed from the light, the Cu2+ ions slowly move to the surface of the crystal where
they interact with
the silver metal:
Cu2+ + Ag0 --> Cu+ + Ag+
The glass clears as the silver ions rejoin chloride ions in the crystals.
Maintaining Blood pH
Carbon dioxide is exhaled
Acid entering the blood stream
HCO31-
+
H+
H2CO3
H2O + CO2
Bicarbonate ion circulates in the blood stream where it is in equilibrium with H+ and OH-.
In the lungs, bicarbonate ions combine with a hydrogen ion and lose a water molecule
to form carbon dioxide, which is exhaled.
Kelter, Carr, Scott, Chemistry A World of Choices 1999, page 291
Alkalosis
If our breathing becomes too fast (hyperventilation)…
Carbon dioxide is removed from the blood too quickly.
This accelerates the rate of degradation of carbonic acid into carbon dioxide and water.
The lower level of carbonic acid encourages the combination of hydrogen ions and
bicarbonate ions to make more carbonic acid. The final result is a fall in blood H1+
levels that raises blood pH which can result in over-excitability or death.
Kelter, Carr, Scott, Chemistry A World of Choices 1999, page 291
Acidosis
If breathing becomes too slow (hypoventilation)…
…free up acid, pH of blood drops, with associated health risks such as depression
of the central nervous system or death.
The normal pH of blood is between 7.2 – 7.4.
This pH is maintained by the bicarbonate ion and other buffers.
Kelter, Carr, Scott, Chemistry A World of Choices 1999, page 291
Acids: Concentration vs. Strength
WEAK
STRONG
CONCENTRATED
H+ A- H+ A- H+ A- H+ A- HA
A- H+ A- H+ A- H+ A- H+ A H+ A- HA H+ A- H+ A- H+ AA- H+ A- H+ A- H+ A- H+ A- H+
H+ A - H + A - H + A - HA H + A A- H+ A- H+ A- H+ A- H+ A–
H+ A- H+ A- H+ A- H+ A- H+
A- H+ A- H+ A- H+ A- H+ AHA A- H+ A- H+ A- H+ A- H+
HA HA H+ A- HA HA
HA HA HA HA HA
H+ A- HA HA HA HA
HA HA H+ A- HA HA
HA HA HA H+ A- HA
H+ A- HA HA HA HA
HA HA HA H+ A- HA
H+ A- HA HA HA HA
HA HA H+ A- HA HA
HA
H+ AHA
HA
HA
HA
HA
HA
HA
DILUTE
H+
A-
H+
A-
HA
A-
H+
A-
H+
A–
H+
A-
H+
A-
H+
A-
H+
HA
H+
A-
HA
H+
HA
A-
HA
HA
HA
H+ A -
HA
HA
HA
HA
H+A–
A-
Dissociate nearly 100%
HA
H1+
+
A-
+
A-
H+
HA
HA
HA
H+A–
HA
STRONG ACIDS
HA
WEAK ACIDS
Dissociate very little
HA
H1+
Comparison of Strong and Weak Acids
Type of acid, HA
Reversibility
of reaction
Ka value
Ions existing when acid,
HA, dissociates in H2O
Strong
Not
reversible
Ka value very large
H+ and A-, only.
No HA present.
Weak
reversible
Ka is small
H+, A-, and HA
HA(aq) + H2O(l)
H3O+(aq) + A-(aq)
The equilibrium expression for the reaction is
Ka =
[H3O+] [A-]
[HA]
Note: H3O+ = H+
Strong vs. Weak Acid
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 508
Concentrated vs. Dilute
0.3 M HCl
10.0 M CH3COOH
Dilute, strong acid
Concentrated, weak acid
2.0 M HCl
Concentrated, strong acid
OR Dilute, strong, acid
12.0 M HCl
Concentrated, strong acid
Naming Acids
Anion
Acid
_________ ide
(chloride, Cl1-)
add H+
_________ ate
(chlorate, ClO3-)
(perchlorate, ClO4-)
add H+
_________ite
(chlorite, ClO2-)
(hypochlorite, ClO-)
add H+
ions
ions
ions
Hydro____ ic acid
(hydrochloric acid, HCl)
_________ic acid
(chloric acid, HClO3)
(perchloric acid, HClO4)
______ous acid
(chlorous acid, HClO2)
(hypochlorous acid, HClO)
7A
HF
No acid or
base
properties
Weak base
---
Weak acid
SiH4
PH3
H2S
HCl
No acid or
base
properties
Weak base
Weak acid
Strong acid
Increasing acid Strength
Increasing base Strength
Brown, LeMay, Bursten, Chemistry 2000, page 625
Increasing acid Strength
Period 3
6A
H2O
4A
CH4
Increasing base Strength
Period 2
Group
5A
NH3
Equilibrium and pH Calculations
Weak acid
Strong acid
H3O+ + A-
HA + H2O
HA + H2O
H3O+ + A-
acid-dissociation
constant calculations
Ka =
[A-] [H3O+]
[HA]
[HA] = [H3O+]
[H3O+]
+
antilog(-pH)
7
[OH-]
-log [H3O+] =
pH
0
1 x 10-14
=
[OH-]
14
Tocci, Viehland, Holt Chemistry Visualizing Matter 1996, page 525
1 x 10-14
[H3O+]
=
-
Strengths of Conjugate Acid-Base Pairs
Acid strength increases
strong
HCl H2SO4
medium
weak
very weak
HNO3 H3O+ HSO4- H3PO4 HC2H3O2 H2CO3 H2S H2PO4- NH4+ HCO3- HPO42- H2O
Cl- HSO4-
negligible
NO3
H2O
SO42-
H2PO4-
very weak
C2H3O2- HCO3- HS- HPO42-
weak
Base strength increases
NH3
medium
CO32- PO43- OH-
strong
Conjugate
Acid Strength
Relative
acid
strength
Relative
conjugate
base
strength
Very
strong
Very
weak
Strong
HA
H+ + A-
Weak
[H+] [A-]
pKa =
[HA]
Weak
Strong
Very
weak
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 508
Very
strong
Solutions of Acids and Bases: The Leveling Effect
• No acid stronger than H3O+ and no base stronger than OH– can exist
in aqueous solution, leading to the phenomenon known as the
leveling effect.
• Any species that is a stronger acid than the conjugate acid of water
(H3O+) is leveled to the strength of H3O+ in aqueous solution
because H3O+ is the strongest acid that can exist in equilibrium with
water.
• In aqueous solution, any base stronger than OH– is leveled to the
strength of OH– because OH– is the strongest base that can exist in
equilibrium with water
• Any substance whose anion is the conjugate base of a compound
that is a weaker acid than water is a strong base that reacts
quantitatively with water to form hydroxide ion
Weak Acids (pKa)
Weak Acids – dissociate incompletely (~20%)
Strong Acids – dissociate completely (~100%)
A(g) + 2 B(g)
3 C(g) + D(g)
Equilibrium constant (Keq) =
Keq =
[Products]
[Reactants]
[C ]3 [D]
[ A][B ]2
LeChatelier’s Principle
(lu-SHAT-el-YAY’s)
H+(aq) + C2H3O21-(aq)
HC2H3O2(aq)
1
Equilibrium constant
Keq =
[H  ][C2H3O2 ]
[HC2H3O2 ]
= Ka = Acid dissociation constant
Ka = 1.8 x 10-5 @ 25 oC for acetic acid
1
[H  ][C2H3O2 ]
Ka =
[HC2H3O2 ]
1
1.8 x 10-5
[H  ][C2H3O2 ]
=
[HC2H3O2 ]
Ionization Constants for Acids
Ka
HCl
H+ + Cl1-
very large
HNO3
H+ + NO31-
very large
H2SO4
H+ + HSO41-
HC2H3O2
H2S
large
H+ + C2H3O21-
1.8 x 10-5
H+
9.5 x 10-8
+
HS1-
Ionization of Acids
Acid
Hydrochloric
Ionization Equation
HCl
Sulfuric
H2SO4
Acetic
HC2H3O2
Ionization Constant, pKa
H1+ + Cl1H1+ + HSO41H1+ + C2H3O21-
very large
large
1.8 x 10-5
Formula
Name
Value of Ka*
HSO4HClO2
HC2H2ClO2
HF
HNO2
HC2H3O2
HOCl
HCN
NH4+
HOC6H5
hydrogen sulfate ion
chlorous acid
monochloracetic acid
hydrofluoric acid
nitrous acid
acetic acid
hypochlorous acid
hydrocyanic acid
ammonium ion
phenol
1.2 x 10-2
1.2 x 10-2
1.35 x 10-3
7.2 x 10-4
4.0 x 10-4
1.8 x 10-5
3.5 x 10-8
6.2 x 10-10
5.6 x 10-10
1.6 x 10-10
*The units of Ka are mol/L but are customarily omitted.
Increasing acid strength
Values of Ka for Some Common Monoprotic Acids
Sample 1)
One gram of concentrated sulfuric acid (H2SO4) is diluted to a 1.0 dm3 volume
with water. What is the molar concentration of the hydrogen ion in this solution?
What is the pH?
Solution)
First determine the number of moles of H2SO4
x mol H2SO4 = 1 g H2SO4
H2SO4
H+ + HSO41-
1 mol H2SO4
98 g H2SO4
&
= 0.010 mol H2SO4
HSO41-
H+ + SO42-
OVERALL:
H2SO4
0.010 M
2 H+ + SO42-
in dilute solutions...occurs ~100%
0.020 M
pH = - log [H+]
substitute into equation
pH = - log [0.020 M]
pH = 1.69
A volume of 5.71 cm3 of pure acetic acid, HC2H3O2, is diluted with water at
25 oC to form a solution with a volume of 1.0 dm3.
What is the molar concentration of the hydrogen ion, H+, in this solution?
(The density of pure acetic acid is 1.05 g/cm3.)
Step 1) Find the mass of the acid
Mass of acid = density of acid x volume of acid
= 1.05 g/cm3 x 5.71 cm3
= 6.00 g
Step 2) Find the number of moles of acid. (From the formula of acetic acid,
you can calculate that the molar mass of acetic acid is 60 g / mol).
 1 mol HC 2H3O 2 
= 0.10 mol acetic acid (in 1 L)
x mol acetic acid = 6.00 g HC2H3O2 
60
g
HC
H
O

2 3 2 
Molarity: M = mol / L
Substitute into equation
M = 0.10 mol / 1 L
M = 0.1 molar HC2H3O2
Step 3) Find the [H+]
Ka =
HC2H3O2
Step 3) Find the
0.1 M
[H+]
weak acid
H+ + C2H3O21-
0.1? M
Ka = 1.8 x 10-5 @ 25 oC for acetic acid
1
[H  ][C2H3O2 ]
Ka =
[HC2H3O2 ]
1
1.8 x
10-5
[H  ][C2H3O2 ]
=
[HC2H3O2 ]
How do the concentrations of
H+ and C2H3O21- compare?
[x][x]
[HC2H3O2 ]
Substitute into equation:
1.8 x 10-5 
1.8 x 10
-5
x2

[0.10 M]
pH = - log[H+]
x2 = 1.8 x 10-6 M
x = 1.3 x 10-3 molar
pH = - log [1.3 x10-3 M]
= [H+]
pH = 2.9
H+ Concentrations
…Strong vs. Weak Acid
Moles of Acid used to make
1 L of solution
H+
pH
0.010 mol H2SO4
0.0200 M
1.7
Strong acid
0.100 mol HC2H3O2
0.0013 M
2.9
Weak acid
Note: although the sulfuric acid is 10x less
concentrated than the acetic acid...
…it produces > 10x more H+
pH = - log[H+]
Practice Problems:
1a) What is the molar hydrogen ion concentration in a 2.00 dm3 solution
of hydrogen chloride in which 3.65 g of HCl is dissolved?
1b) pH
2a) What is the molar concentration of hydrogen ions in a solution
containing 3.20 g of HNO3 in 250 cm3 of solution?
2b) pH
3a) An acetic acid solution is 0.25 M. What is its molar concentration of
hydrogen ions?
3b) pH
4) A solution of acetic acid contains 12.0 g of HC2H3O2 in 500 cm3
of solution. What is the molar concentration of hydrogen ions?
1a) 0.0500 M
1b) pH = 1.3
2a) 0.203 M
2b) pH = 0.7
3a) 2.1 x 10-3 M
3b) pH = 2.7
4) 2.7 x 10-3 M
Weak Acids
Cyanic acid is a weak monoprotic acid. If the initial concentration of cyanic
+ is 4.8 x 10-3 M,
acid is 0.150 M and the equilibrium concentrationHof3O
H3+O(aq)
calculate Ka for cyanic acid.
4.8 x 10-3 M
0.150 M
Ka =
+ CN1-(aq)
H+(aq)
HCN(aq)
[Products]
[Reactants]
Ka =
4.8 x 10-3 M
[H3O+] [CN1-]
[HCN]
[4.8 x 10-3 M][4.8
[CNx1-10
] -3 M]
Ka =
[0.150 M]
Ka = 1.54 x 10-4
How is [H3O+] determined?
Measure pH of solution and work backwards
Acid Dissociation
+
H
1-
HCl
Cl
Acid
Conjugate base
Conjugate pair
Kelter, Carr, Scott, Chemistry A World of Choices 1999, page 280
Conjugate Acid-Base Pairs
conjugates
HCl
+
base
acid
H 2O
H3O+
+
acid
Clbase
conjugates
HCl
acid
+
H 2O
H3O+
base
CA
+
ClCB
Conjugate Acid-Base Pairs
conjugates
acid
NH3
+
H2O
base
base
NH41+
+
OH-
acid
conjugates
NH3
base
+
H2O
acid
NH41+
CA
+ OHCB
Water is Amphoteric
Amphoteric or Amphiprotic substances:
Substances which can act as either proton donors (acids) or
proton acceptors (bases) depending on what substances are present.
HCl
+
acid
NH3
base
+
H2O
H3O+
base
CA
H2O
acid
NH41+
CA
+
ClCB
+ OHCB
Amphoteric
A substance that can act as either an acid or a base.
11+
+
H3O+
hydronium ion
+
HSO4hydrogen sulfate
ion
1-
H2O
water
H2SO4
sulfuric acid
2-
1-
+
HSO4hydrogen sulfate
ion
OHhydroxide ion
+
SO42-
H2O
sulfate ion
water
Amphoteric
A substance that can act as either an acid or a base.
11+
+
H3O+
hydronium ion
+
HSO4hydrogen sulfate
ion
(HSO4- as a base)
H2SO4
sulfuric acid
H2O
water
Amphoteric
A substance that can act as either an acid or a base.
1-
21-
+
+
HSO4-
OH-
SO42-
H2O
hydrogen sulfate
ion
hydroxide ion
sulfate ion
water
(HSO4- as an acid)
Conjugate Acid-Base Pairs
conjugates
base2
HC2H3O2 + H2O
acid2
H3O1+
+
acid1
C2H3O2OHbase1
conjugates
HC2H3O2 + H2O
acid
base
H3O1+
CA
+
C2H3O2OHCB
The reaction proceeds in the direction such that the stronger acid
donates its proton to the stronger base.
Litmus Paper
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.
pH Paper
pH 0
1
2
3
4
5
6
pH 7
8
9
10
11
12
13
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.
Desired Features of Sensors
pH paper
1904
pH 0
1
2
3
4
5
pH 7
8
9
10
11
12
6
13
Detection limit
Low deflection
High sensitivity
High selectivity
Wide dynamic
range
Simple to use
Cost-effective
Range and Color Changes of Some
Common Acid-Base Indicators
pH Scale
1
2
3
4
5
6
7
8
9
10
11
12
13
14
Indicators
Methyl orange
Methyl red
Bromthymol blue
Neutral red
Phenolphthalein
3.1 – 4.4
red
red
4.4
yellow
yellow
6.2
6.2
red
colorless
6.8
yellow
7.6
8.0
8.0
blue
yellow
10.0
red
Bromthymol blue indicator would be used in titrating a strong acid with a strong base.
Phenolpthalein indicator would be used in titrating a weak acid with a strong base.
Methyl orange indicator would be used in titrating a strong acid with a weak base.
colorless beyond 13.0
pH
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.
INDICATOR COLORS IN TITRATION
Indicator
Acid color
Transition color
Base color
STRONG ACID – STRONG BASE
Litmus
2
pH
Bromthymol blue
3
4
5
6
7
8
9
10
11
12
INDICATOR COLORS IN TITRATION
Indicator
Acid color
Transition color
Base color
WEAK ACID – STRONG BASE
Phenolphthalein
2
pH
Phenol red
3
4
5
6
7
8
9
10
11
12
INDICATOR COLORS IN TITRATION
Indicator
Acid color
Transition color
Base color
STRONG ACID – WEAK BASE
Methyl orange
2
pH
Bromphenol blue
3
4
5
6
7
8
9
10
11
12
Indicator
1
2
Orange IV
4
5
6
Colorless
Phenolphthalein
Methyl Red
3
Red
Orange
Peach
Orange
7
8
9 10 11 12
Pink
Yellow
Yellow
Red
Common pH Indicators
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 520
Edible Acid-Base Indicators
COLOR CHANGES AS A FUNCTION OF pH
INDICATOR
pH
2
3
4
5
6
7
8
9
10
11
12
RED APPLE SKIN
BEETS
BLUEBERRIES
RED CABBAGE
*
CHERRIES
GRAPE JUICE
RED ONION
YELLOW ONION
PEACH SKIN
PEAR SKIN
PLUM SKIN
RADISH SKIN
RHUBARB SKIN
TOMATO
TURNIP SKIN
*YELLOW at pH 12 and above
Red Cabbage Indicator
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.
H+
Phenolphthalein
Indicator
Colorless = Acidic pH
Pink = Basic pH
-O
O
C
C
OH
HO
C
O-
O
(Colorless acid form, HIn)
C
OH
O-
O
(Pink base form, In-)
Preparation of an Ester
Acetylsalicylic Acid (Aspirin)
OBJECTIVE:
Aspirin Synthesis
To become familiar with the techniques and principle of esterification.
DISCUSSION:
Aspirin is a drug widely used as an antipyretic agent (to reduce fever), as an analgesic agent (to
reduce pain), and/or as an anti-inflammatory agent (to reduce redness, heat or swelling in tissues).
Chemically, aspirin is an ester. Esters are the products of reaction of acids with alcohols, as shown
in the following equation using type formulas:
O
O
R – C – OH
+
ACID
R’ – OH
R – C – O – R’
ALCOHOL
ESTER
+
H2O
WATER
The symbol R refers to the hydrocarbon portion (radical) of the molecules aside from the
O
functional group. In an organic acid, R – C – OH, the functional group is the carboxyl
O
group (-COOH) or –C-OH. The type of formula for an alcohol is R-OH, where the functional group is
the hydroxyl group (-OH). The symbol R’ indicates that the two R-groups in the ester formula need
not be the same. It has been shown by radioactive tracer methods that in the mechanism of the
esterification reaction, the –OH group is split from the acid and the –H from the alcohol.
Aspirin can be made as follows:
C – OH
CH3 – C – OH
+
Acetic acid
HO –
C – OH
CH3 – C – O–
+
Salicylic acid
Aspirin
(containing an
-OH group)
(acetylsalicylic acid,
an ester)
H2O
The use of acetic anhydride instead of acetic acid, however, is a better preparative method, because
the anhydride with the water to form acetic acid tends to drive the reaction to the right as shown
below. An acid catalyst also is used to speed up the reaction.
CH3
O
C
C
CH3
C – OH
O
O
Acetic anhydride
+
HO –
C – OH
CH3 – C – O–
Salicylic acid
Aspirin
(138.12 g/mol)
(180.15 g/mol)
+
CH3 – C – OH
Acetic acid
How to read a buret volume
23
24.55 mL?
23.45 mL
(not 24.55 mL)
24
Titration
standard solution
 Titration
• Analytical method in
which a standard
solution is used to
determine the
concentration of an
unknown solution.
unknown solution
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
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
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Titration
+
O
moles H3 = moles
MVn = MVn
M: Molarity
V: volume
n: # of H+ ions in the acid
or OH- ions in the base
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
OH
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
MV# = MV#
M(50.0mL)(2)
=(1.3M)(42.5mL)(1)
M = 0.55M H2SO4
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Acid-Base Titration
Data Table
0.10 M HCl
Base (mL)
Calibration Curve
0.00 mL
1.00 mL
2.00 mL
4.00 mL
9.00 mL
17.00 mL
27.00 mL
48.00 mL
? M NaOH
1.00 mL
1.00 mL
2.00 mL
5.00 mL
8.00 mL
10.0 mL
15.0 mL
Acid (mL)
1)
2)
3)
Solution
Solution
of
of NaOH
NaOH
Create calibration curve of six data points
Using [HCl], determine concentration of NH3
Determine vinegar concentration using [NaOH]
determined earlier in lab
Solution
of HCl
5 mL
Titration
Curve
Titration
indicator -changes color
to indicate pH change
e.g. phenolpthalein is colorless in acid
and pink in basic solution
endpoint
pink
equivalence
point
pH
7
Pirate…”Walk the plank”
once in water, shark eats and
water changes to pink color
base
Calibration Curve
endpoint
Base (mL)
pink
pH
7
equivalence
point
Acid (mL)
Pirate…”Walk the plank”
once in water, shark eats and
water changes to pink color
base
indicator - changes color to indicate pH change
e.g. phenolphthalein is colorless in acid
and pink in basic solution
Calibration Curve
endpoint
Base (mL)
pink
pH
7
equivalence
point
Acid (mL)
Pirate…”Walk the plank”
once in water, shark eats and
water changes to pink color
base
indicator - changes color to indicate pH change
e.g. phenolphthalein is colorless in acid
and pink in basic solution
Titration Curve
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 527
Acid-Base Titrations
Titration of a Strong Acid With a Strong Base
14.0
12.0
Solution
of NaOH
10.0
OHNa+
Na+
pH
-
OH- OH
Na+ Na+
OH-
8.0
equivalence point
6.0
4.0
Solution
of HCl
H+
Cl-
2.0
Cl
H+
H+
Cl-
H+
Cl-
0.0
0.0
10.0
20.0
30.0
40.0
Volume of 0.100 M NaOH added
(mL)
Additional
Adding
additional
NaOH
NaOH
from
isNaOH
added.
the buret
is added.
pH
to increases
hydrochloric
pH rises
and
as
acid
theninlevels
the flask,
off as
the
NaOH
a strong
equivalence
is acid.
addedInbeyond
point
the beginning
is the
approached.
equivalence
the pH increases
point. very slowly.
Titration Data
pH
0.00
10.00
20.00
22.00
24.00
25.00
26.00
28.00
30.00
40.00
50.00
1.00
1.37
1.95
2.19
2.70
7.00
11.30
11.75
11.96
12.36
12.52
Solution
of NaOH
Na+
OH-
OHNa+ Na+
OH-
Solution
of HCl
H+
Cl-
25 mL
14.0
12.0
10.0
8.0
equivalence point
6.0
4.0
2.0
OHNa+
Titration of a Strong Acid With a Strong Base
pH
NaOH added
(mL)
0.0
0.0
10.0
20.0
30.0
40.0
Volume of 0.100 M NaOH added
(mL)
ClH+
H+
Yellow
Blue
Cl-
H+
Cl-
Bromthymol blue is best indicator: pH change 6.0 - 7.6
Titration of a Strong Acid With a Strong Base
(20.00 mL of 0.500 M HCl by 0.500 M NaOH)
14.0
12.0
Color change
alizarin yellow R
10.0
Color change
phenolpthalein
pH
8.0
Color change
bromthymol blue
equivalence
point
6.0
Color change
bromphenol blue
4.0
Color change
methyl violet
2.0
0.0
0.0
10.0
20.0
30.0
Volume of 0.500 M NaOH added
(mL)
Hill, Petrucci, General Chemistry An Integrated Approach 2nd Edition, page 680
Titration of a Weak Acid With a Strong Base
Titration of a Weak Acid With
a Strong Base
Titration Data
14.0
NaOH added
(mL)
12.0
10.0
pH
equivalence point
8.0
6.0
4.0
2.0
0.0
0.0
10.0
20.0
30.0
Volume of 0.100 M NaOH added
(mL)
40.0
0.00
5.00
10.00
12.50
15.00
20.00
24.00
25.00
26.00
30.00
40.00
pH
2.89
4.14
4.57
4.74
4.92
5.35
6.12
8.72
11.30
11.96
12.36
Phenolphthalein is best indicator: pH change 8.0 - 9.6
Titration of a Weak Base With a Strong Acid
Titration of a Weak Base With a Strong Acid
Titration Data
14.0
HCl added
(mL)
pH
0.00
10.00
20.00
30.00
40.00
45.00
47.00
48.00
49.00
50.00
51.00
11.24
9.91
9.47
8.93
8.61
8.30
7.92
7.70
7.47
5.85
3.34
12.0
pH
10.0
8.0
6.0
equivalence point
4.0
2.0
0.0
0.0
10.0
20.0
30.0
40.0
Volume of 0.100 M HCl added
(mL)
50.0
7. What is the pH of a solution made by dissolving 2.5 g NaOH in 400 mL water?
Determine number of moles of NaOH
 1 mol NaOH 
x mol NaOH = 2.5 g NaOH 
  0.0625 mol NaOH
 40 g NaOH 
Calculate the molarity of the solution
M
mol
L

0.0625 mol NaOH
0.4 L
[Recall 1000 mL = 1 L]
MNaOH = 0.15625 molar
NaOH
0.15625 molar
Na1+ +
0.15625 molar
pOH = -log [OH-]
OH10.15625 molar
or
kW = [H+] [OH-]
pOH = -log [0.15625 M]
1 x 10-14 = [H+] [0.15625 M]
pOH = 0.8
[H+] = 6.4 x 10-14 M
pOH + pH = 14
pH = -log [H+]
0.8 + pH = 14
pH = 13.2
pH = -log [6.4 x 10-14 M]
What volume of 0.5 M HCl is required to titrate 100 mL of 3.0 M Ca(OH)2?
2 HCl
x mL
0.5 M
+
"6.0 M"
Ca(OH)2
100 mL
3.0 M
CaCl2
M1V1 = M2V2
(0.5 M) (x mL) = (3.0 M) (100 mL)
+ 2 HOH
M1V1 = M2V2
(0.5 M) (x mL) = (6.0 M) (100 mL)
x = 600 mL of 0.5 M HCl
mol
M
L
HCl
0.3 mol
x = 1200 mL of 0.5 M HCl
HCl
molHCl = M x L
Ca(OH)2
mol Ca(OH)2 = M x L
mol = (0.5 M)(0.6 L)
mol = (3.0 M)(0.1 L)
mol = 0.3 mol HCl
mol = 0.3 mol Ca(OH)2
H1+ +
0.3 mol
Cl1-
Ca(OH)2
0.3 mol
0.3 mol
[H+] = [OH-]
Ca2+ +
2OH1-
0.3 mol
0.6 mol
6.
10.0 grams vinegar titrated with 65.40 mL of 0.150 M NaOH
(acetic acid + water)
moles HC2H3O2 =
A)
moles NaOH
mol
M
L
NaOH
molNaOH = M x L
therefore, you have ...
0.00981 mol HC2H3O2
mol = (0.150 M)(0.0654 L)
mol = 0.00981 mol NaOH
 60 g HC3H2O2 
  0.59 g HC2H3O2
 1 mol HC3H2O2 
B) x g HC2H3O2 = 0.00981 mol HC2H3O2 

C)
part 
% = 
 x 100%
 whole 
 0.59 g acetic acid 
% = 
 x 100%
 10.0 g vinegar 
% =
5.9 % acetic acid
Commercial vinegar is sold as 3 - 5 % acetic acid
Carboxylic Acid
HC2H3O2
H
= acetic acid
O
H C C
1O
H
H
H+
CH3COOH
R - COOH
carboxylic acid
C2H4O2
O
H
O
H
C
O
H
H
C
H
H
Lactic Acid
OH
H3C
C
CO2H
H
Lactic acid
C3H6O3
Titration
?
1.0 M HCl
titrate with
? M NaOH
2.00 mL
1.00 mL
M1 V1 = M2 V2
(1.0 M)(1.00 mL) = (x M)(2.00 mL)
X = 0.5 M NaOH
1+
2.0 M H
1.0 M H2SO4
1.00 mL
titrate with
? M NaOH
2.00 mL
M1 V1 = M2 V2
(1.0 M)(1.00 mL) = (x M)(2.00 mL)
X = 0.5 M NaOH
HCl
Hydrochloric acid
stomach acid, pickling metal
H2SO4
Sulfuric acid
battery acid, # 1 selling chemical
H3PO4
Phosphoric acid
food flavoring
HNO3
Nitric acid
fertilizer, explosives
CH3COOH
Acetic acid
vinegar
HF
Hydrofluoric acid
etch glass
NaOH
Ca(OH)2
NH4OH
sodium hydroxide
calcium hydroxide
ammonium hydroxide
Soren Sorenson developed pH scale
7
neutral
pH scale
0
[H+] = [OH-]
acid
14
base
(alkalinity)
Arnold Beckman invented the pH meter
pH = -log [H+]
pOH = -log [OH-]
pH + pOH = 14
kW =
[H+]
[OH-]
kw = 1 x 10-14
H+ + H2O
proton
H3O+
hydronium ion
Concentrated vs. Dilute
Concentration:
Molarity
molality
mol
M= L
mol
m = kg
H2SO4
2 H1+
3M
“6 M”
+
SO42-
Normality
Strong / Weak Acid
Strong
HA
H+
+
A-
(~100% dissociation)
Weak
HA
H+
+
A-
(~20% dissociation)
H2A
2 H+
[Product]
Ka = [Reactant]
+
Ka =
A-
[H+]2 [A-]
[H2A]
acid dissociation constant
Ka
0.8
0.0021
H3PO4
HF
3H+ + PO43H + + F-
Acid
+
Base
Salt
+
Water
How would you make calcium sulfate in the lab?
H2?SO4
+
Ca(OH)
?
2
BASE
ACID
Sour taste, litmus
CaSO4 + 2 H2O
red
bitter taste, litmus
blue
Arrhenius – H+ as only ion in water
Arrhenius – OH- as only ion in water
Brønsted-Lowry – proton donor
Brønsted-Lowry – proton acceptor
Indicators
colorless
weak acid
phenolphthalein
yellow
strong acid
bromthymol blue
universal indicator
&
blue
strong base
R O Y G B I V
pH 4
litmus paper
pink
strong base
pH paper
7
12
Buffers - salts of weak acids and weak bases that maintain a pH
e.g. Aspirin (acetyl salicylic acid) vs. Bufferin
low pH upsets stomach
LeChatelier’s Principle
- acidosis & alkalosis (bicarbonate ion acts as buffer)
- darkening glasses
- egg shells thinner in summer (warm)
Acid – Base Concentrations
concentration (moles/L)
10-1
pH = 3
pH = 11
OH-
H3O+
pH = 7
10-7
H3O+
OH-
OH-
H3O+
10-14
Timberlake, Chemistry 7th Edition, page 332
[H3O+] > [OH-]
[H3O+] = [OH-]
acidic
solution
neutral
solution
[H3O+] < [OH-]
basic
solution
Amino Acids – Functional Groups
Amine
Base Pair
Carboxylic Acid
R- COOH
NH21lose H+
NH21-
H+
NH3
amine
ammonia
NH41+
ammonium ion
1+
:
1-
N
N
N
H
H
:
:
H
H
H
H
H
H
H
Amino Acids – Functional Groups
Amine
NH21-
Base Pair
Carboxylic Acid
R- COOH
+
d+ H
H+ O2- d
http://fig.cox.miami.edu/~cmallery/150/chemistry/sf3x17b.jpg
Water – Amphiprotic
lose H+
OH1-
H+
hydroxide
H2O
H3O1+
water
hydronium ion
:
N
H
H
:
:
N
H
1+
H
1-
N
H
H
H
H
H
Water – Also Amphiprotic
Amphiprotic – Act as an acid (proton donor) or base (proton acceptor)
OH1hydroxide
lose H+
H+
H2O
H3O1+
water
hydronium ion
+
d+ H
H+ O2- d
Amino Acids – Functional Groups
Amphiprotic – Act as an acid (proton donor) or base (proton acceptor)
OH1hydroxide
lose H+
H+
H2O
H3O1+
water
hydronium ion
+
d+ H
H+ O2- d
http://fig.cox.miami.edu/~cmallery/150/chemistry/sf3x17b.jpg
Range and Color Changes of Some
Common Acid-Base Indicators
pH Scale
1
2
3
4
5
6
7
8
9
10
11
12
13
14
Indicators
Methyl orange
Methyl red
Bromthymol blue
3.1 – 4.4
red
red
4.4
yellow
Neutral red
red
Phenolphthalein
colorless
yellow
6.2
6.2
6.8
yellow
7.6
blue
8.0
8.0
yellow
10.0
red
colorless beyond 13.0
Range and Color Changes of Some
Common Acid-Base Indicators
pH Scale
1
2
3
4
5
6
7
8
9
10
11
12
13
14
Indicators
Methyl orange
Methyl red
Bromthymol blue
3.1 – 4.4
red
red
4.4
yellow
Neutral red
red
Phenolphthalein
colorless
yellow
6.2
6.2
6.8
yellow
7.6
blue
8.0
8.0
yellow
10.0
red
colorless beyond 13.0
pH Paper
pH Paper
pH 0
1
2
3
4
5
pH 7
8
9
10
11
12
6
pH 0
1
2
3
4
5
13
pH 7
8
9
10
11
12
13
6
pH Paper
6
pH Paper
pH 0
1
2
3
4
5
pH 7
8
9
10
11
12
6
pH 0
1
2
3
4
5
13
pH 7
8
9
10
11
12
13
Neutralization of Bug Bites
Wasp - stings with base
Red Ant - bites with acid
(neutralize with lemon juice or vinegar)
(neutralize with baking soda)
Strength
 Strong
Acid/Base
• 100% ionized in water
• strong electrolyte
HCl
HNO3
H2SO4
HBr
HI
HClO4
-
+
NaOH
KOH
Ca(OH)2
Ba(OH)2
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Strength
 Weak
Acid/Base
• does not ionize completely
• weak electrolyte
HF
CH3COOH
H3PO4
H2CO3
HCN
-
+
NH3
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Ionization of Water
H 2O + H 2 O
Kw =
+
[H3O ][OH ]
H3
+
O
+
= 1.0 
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
OH
-14
10
Why is pure water pH = 7?
1
in 500,000,000 water molecules will
autoionize.
 H2O + H2O  H3O+ + OH1 This yields a hydronium ion concentration
of 1 x 10-7 M H3O+ per liter of solution
 pH = -log[H3O+]
 pH = -log[1 x 10-7] or pH = 7
H
H
O
1-
H
O
H
H
H
H
H
1+
O
H
O
O
O
H
O
H
H
H
H
H
O
H
H
Ionization of Water
 Find
the hydroxide ion concentration of
3.0  10-2 M HCl.
[H3O+][OH-] = 1.0  10-14
[3.0  10-2][OH-] = 1.0  10-14
[OH-] = 3.3  10-13 M
Acidic or basic? Acidic
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
pH Scale
14
0
7
INCREASING
ACIDITY
pH =
NEUTRAL
INCREASING
BASICITY
+
-log[H3O ]
pouvoir hydrogène (Fr.)
“hydrogen power”
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
pH Scale
pH =
+
-log[H3O ]
pOH =
-log[OH ]
pH + pOH = 14
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
pH Scale
 What
is the pH of 0.050 M HNO3?
pH = -log[H3O+]
pH = -log[0.050]
pH = 1.3
Acidic or basic? Acidic
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
pH Scale
 What
is the molarity of HBr in a solution
that has a pOH of 9.6?
pH + pOH = 14
pH = -log[H3O+]
pH + 9.6 = 14
4.4 = -log[H3O+]
pH = 4.4
-4.4 = log[H3O+]
Acidic
[H3O+] = 4.0  10-5 M HBr
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Resources - Acids and Bases
Objectives
Episode 16 – The Proton in Chemistry
Worksheet - vocabulary
Video (VHS) - future of the past
Worksheet - pH and pOH calculations
Outline -
Worksheet - practice problems (key)
Worksheet -
Textbook - text ?'s chemical equilibrium
Worksheet -
Worksheet - weak acid, pKa
Worksheet -
Article - aspirin
Worksheet -
Lab - synthesis of aspirin
Worksheet - aqueous acids and bases titration
Lab - titration
Textbook - questions
Outline (general)