Acids &
Bases
Edward Wen, PhD
Learning Outcomes
• Properties of acids and bases and definitions
• pH scale and calculation of pH
• Completing and balancing Neutralization
reactions
• Titration calculations for neutralization
reactions
• Defining Weak vs. Strong electrolytes (using
the concept of equilibrium)
• Buffers – recognition of a buffer system, how
a buffer works
2
Types of Ionic Compounds
• Acids = form H+ ions in water solution
• Bases = combine with H+ ions in water solution
 increases the OH- concentration
 may either directly release OH- or pull H+ off H2O
• Salts = Ionic compounds formed from Acid and Base.
 all strong electrolytes
 Cation: except H+
 Anion: except OH3
Properties of Acids
• Sour taste
• react with “active” Metals
i.e. Al, Zn, Fe, but NOT w/ Ag, Au
Zn + 2 HCl ZnCl2 + H2
• react with Carbonates, producing CO2
marble, baking soda, limestone
CaCO3 + 2 HCl CaCl2 + CO2 + H2O
• change color of vegetable dyes
blue litmus turns red
• react with Bases to form ionic salts
4
Most food contains acids
• Citric acid (HO2CCH(CO2H)COHCO2H):
citrus fruits, tomato
• Malic acid (HO2CCH2CHOHCO2H):
green apple, tomato, grape
• Ascorbic acid (aka Vitamin C)
• Folic acid
5
Common Acids
Chemical Name
Formula
Uses
Strength
Nitric Acid
HNO3
explosive, fertilizer, dye, glue
Strong
explosive, fertilizer, dye, glue,
batteries
metal cleaning, food prep, ore
refining, stomach acid
fertilizer, plastics & rubber,
food preservation
plastics & rubber, food
preservation, Vinegar
Sulfuric Acid
H2SO4
Strong
Hydrochloric Acid
HCl
Phosphoric Acid
H3PO4
Acetic Acid
HC2H3O2
Hydrofluoric Acid
HF
metal cleaning, glass etching
Weak
Carbonic Acid
H2CO3
soda water
Weak
Boric Acid
H3BO3
eye wash
Weak
Strong
Moderate
Weak
6
Binary acids
• (HmX): acid hydrogens
attached to a nonmetal
atom
HCl, HF, HBr, HI
H2S, H2Se
Hydrofluoric acid
7
Oxyacids
• acid hydrogens (H+)
attached to an oxygen
atom
H2SO4, HNO3, H3PO4
HClO4
8
Carboxylic acids
• Many exist in food like
vinegar, tomato, citrus
fruit
• -COOH group
 HC2H3O2, H3C6H5O3
• only the first H in the
formula is acidic
 the H is on the COOH
9
Properties of Bases
•
•
•
•
also known as alkalis
taste bitter
solutions feel slippery
change color of vegetable dyes
 different color than acid
 red litmus turns blue
• react with acids to form ionic salts
 neutralization
10
Common Bases
Chemical
Name
sodium
hydroxide
potassium
hydroxide
calcium
hydroxide
Aluminum
hydroxide
magnesium
hydroxide
ammonium
hydroxide
Formula
NaOH
Common
Name
lye,
caustic soda
KOH
caustic potash
Ca(OH)2
slaked lime
NH4OH,
{NH3(aq)}
soap, plastic,
petrol refining
soap, cotton,
electroplating
Strength
Strong
Strong
cement
Strong
Antacid
Weak
milk of
magnesia
antacid
Weak
ammonia
water
detergent, Windex
fertilizer,
explosives, fibers
Weak
Al(OH)3
Mg(OH)2
Common Uses
11
Structure of Bases
• most ionic bases contain OH- ions
 Drano clog-remover: NaOH, Ca(OH)2
• some contain CO32- ion: it produces OH- with water
 Baking soda: CaCO3
 Alka-Seltzer: NaHCO3
• molecular bases that react with H+
 Windex: Ammonia (NH3)
12
Acid-Base Reactions
(Neutralization, Double Displacement Reaction)
• H+ (from the acid) + OH- (from the base)  H2O
it is often helpful to think of H2O as H-OH
• Cation (from base) + Anion (from acid)  Salt
acid + base → salt + water
HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)
13
Acid Reactions. I.
Reaction with Metals
• Reaction with many metals: Al,
Zn, Fe, Mg
but not all!! Not for Cu, Au, Ag, etc.
• Producing a Salt and hydrogen
gas H2
3 H2SO4(aq) + 2 Al(s) → Al2(SO4)3(aq) + 3 H2(g)
14
Acid Reactions. II
Reaction with Metal Oxides
• when acids react with metal oxides, they
produce a salt and water
3 H2SO4 + Al2O3 → Al2(SO4)3 + 3 H2O
15
Acid Reactions. III
Gas-evolving Reaction with Salts
• when acids react with metal carbonate,
bicarbonate, sulfide, sulfite, and bisulfite, gas will
be produced along with other products
2 HNO3 + FeCO3 → Fe(NO3)2 + CO2 + H2O
HCl + NaHCO3 → NaCl + CO2 + H2O
ZnS + 2 HBr → ZnBr2 + H2S
CaSO3 + 2 HI → CaI2 + SO2 + H2O
H2SO4 + 2 NH4HSO3 → (NH4)2SO4 + 2SO2 + 2H2O
16
Base Reactions
• Neutralization of acids
• Reaction with Nonmetal oxides, CO2
2 NaOH + CO2 → Na2CO3 + H2O
• Strong bases will react with Al metal to form
sodium aluminate and hydrogen gas
Example: Dissolving recycled aluminum can
with NaOH solution
2 NaOH + 2 Al + 6 H2O → 2 NaAl(OH)4 + 3 H2
17
Relative Strength of Acids
Acid
Strong Acid
Conjugate Base
Hydroiodic acid
HI
I-
Hydrobromic acid
HBr
Br-
Hydrochloric Acid
HCl
Cl-
Sulfuric Acid
H2SO4
HSO4-
Nitric Acid
HNO3
NO3-
Hydronium ion
H3O+
H2O
Phosphoric Acid
H3PO4
H2PO4-
Hydrofluoric Acid
HF
F-
Acetic Acid
HC2H3O2
C2H3O2-
Carbonic Acid
H2CO3
HCO3-
Ammonium ion
NH4+
NH3
18
Strong Acids
• The stronger the acid, the
more willing it is to donate
H+
Stomach acid
HCl  H+ + Cl-
 use water as the standard base
• Strong acids donate
practically all their H+
 100% ionized in water
• [H3O+] = [strong acid]
 [ ] = molarity
19
Strong Acids
Examples:
• Binary Acid: HCl, HBr, HI
• Oxyacid: HNO3, H2SO4, HClO4, HClO3
• Example:
HNO3 = H+ + NO3H2SO4 = 2H+ + SO42-
20
Weak Acids
• Weak acids donate a
small fraction of their
H+
Vinegar
HC2H3O2  H+ + C2H3O2-
most of the weak acid
molecules do not donate
H+ to water
• [H3O+] << [weak acid]
21
Weak Acids
Examples:
• Binary Acid: HF, H2S, H2Se
• Oxyacid: HNO2, H2SO3, H3PO4, HClO
• Most carboxylic acids, such as acetic
acid
22
Strong Bases
• The stronger the base, the more
willing it is to accept H+
DranoTM
 use water as the standard acid
NaOH  Na+ + OH-
• Strong bases: practically all
molecules are dissociated into
OH– or accept H+
 1 mol NaOH = 1 mol OH 1 mol Ca(OH)2 = 2 mol OH-
• [OH–] = [strong base] x (# OH)
23
Weak Bases
• Definition: a small
fraction of molecules
accept H+
WindexTM
NH3 + H2O  NH4+ + OH-
most of the weak base
molecules do not take H+
from water
• [HO–] << [weak base]
24
Autoionization of Water
• Water: extremely Weak electrolyte
therefore there must be a few ions present
• about 2 out of every 1 billion water molecules
form Ions: Autoionization
H2O + H2O  H3O+ + OH–
H2O  H+ + OH–
• ALL aqueous solutions contain both H+ and OH–
the concentration of H+ and OH– are equal in water
@ 25°C: [H+] = [OH–] = 10-7M
25
Ion Product of Water
• [H+] x [OH–] = constant: Ion Product of
water, Kw
• At 25°C, [H+] x [OH–] = 1 x 10-14 = Kw
• as [H+] increases, [OH–] must decrease so the
product stays constant
14
110
[H ] 

[OH ]

14
110
[OH ] 

[H ]

26
Acidic and Basic Solutions
• Neutral solutions have equal [H+] and [OH–]
[H+] = [OH–] = 1 x 10-7 M
• Acidic solutions : [H+] > [OH–]
[H+] > 1 x 10-7 M
[OH–] < 1 x 10-7 M
• Basic solutions: [OH–] > [H+]
[H+] < 1 x 10-7 M
[OH–] > 1 x 10-7 M
27
Practice - Determine the [H+] concentration
and whether the solution is acidic, basic or
neutral for the following
All [H+] compared to 1 x 10-7 M
• [OH–] = 3.50 x 10-8 M
• [NaOH] = 0.000250 M
• [HCl] = 0.50 M
28
Practice - Determine the [H+] concentration
and whether the solution is acidic, basic or
neutral for the following
• [OH–] = 3.50 x 10-8 M
[H+]
-14
1
x
10
-7 M
=
=
2.86
x
10
3.50 x 10-8
[H+] >[OH-], therefore acidic
• NaOH = 0.000250 M
[H+]
-14
1
x
10
=
= 4.00 x 10-11 M [H+] < [OH-], therefore basic
0.000250
• [HCl] = 0.50 M
[H+] = 0.50 M
[H+] > 1.0 x 10-7 M therefore acidic
Acidic/Basic: [H+] vs. [OH-]
[H+] 100 10-1
+
H
OH-
Acid
10-3
10-5
+
H
OH-
[OH-]10-14 10-13 10-11
10-9
10-7
10-9
Base
10-11
H+
+
H
10-13 10-14
H+
OH OH OH
10-7
10-5
10-3
10-1 100
even though it may look like it, neither H+ of OH- will ever be 0
the sizes of the H+ and OH- are not to scale
because the divisions are powers of 10 rather than units
30
pH
• The measure of the acidity/basicity of a solution
• pH = -log[H+], [H+] = 10-pH
 exponent on 10 with a positive sign
 pHwater = -log[10-7] = 7
 need to know the [H+] concentration to find pH
• pH < 7 : Acidic;
• pH = 7 : Neutral
pH > 7 : Basic
31
pH scale
• pH↓, Acidity↑
• pH↑, basicity↑
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
32
pH measurement
pH can be measured by
pH meter:
• The change in [H+]
affects the voltage of a
standard cell
33
pH of Common Substances
Substance
pH
1.0 M HCl
0.0
0.1 M HCl
1.0
stomach acid
1.0 to 3.0
lemons
2.2 to 2.4
soft drinks
2.0 to 4.0
plums
2.8 to 3.0
apples
2.9 to 3.3
cherries
3.2 to 4.0
unpolluted rainwater
5.6
human blood
7.3 to 7.4
egg whites
7.6 to 8.0
milk of magnesia (sat’d Mg(OH)2)
10.5
household ammonia
10.5 to 11.5
1.0 M NaOH
14
34
Example - Calculate the pH of the
following strong acid or base solutions
• 0.0020 M HCl
• 0.010 M NaOH
35
Example - Calculate the pH of the
following strong acid or base solutions
• 0.0020 M HCl
HCl as strong acid, so [H+] = 0.0020 M
pH = - log (2.0 x 10-3) = 2.7
• 0.010 M NaOH
NaOH as strong base, so [OH-] = 0.010 M
[H+]
=
1 x 10-14= 1 x 10-12 M
1 x 10-2
pH = - log (1.0 x 10-12) = 12
36
pH in everyday life
Stomach acid Vinegar Pure water Windex
pH
0
1
[H+] 100 10-1
+
H
OH-
Acid
3
10-3
5
7
9
10-5
10-7
10-9
+
H
OH-
[OH-]10-14 10-13 10-11
10-9
+
H
Base
11
13
10-11
OH
10-5
14
10-13 10-14
H+
H+
OH
10-7
Drano
OH
10-3
10-1 100
37
Example - Calculate the concentration of
[H+] for a solution with pH 3.7
[H+] = 10-pH
[H+] = 10-3.7
= 2 x 10-4 M
= 0.0002 M
38
Find concentration of Acid or
Base? Titration
• Purpose: using Reaction
Stoichiometry to determine the
Concentration of an unknown
solution
• Titrant (solution of known
concentration) added from a Buret
• Indicators: chemicals added to
help determine when a reaction is
complete
• the Endpoint of the titration occurs
when the reaction is complete
39
Titration: Color change w/ Indicator
40
Titration
Start: The base solution
as titrant in the buret.
Titrating: As the Base is
added to the Acid, H+ +
OH–  HOH. But still
excess Acid present so the
color does not change.
Endpoint: just enough
Base to neutralize all the
acid. The indicator
changes color.
41
Calculations in Titration
• At the Endpoint of the titration, acid base
neutralization reaction is complete. The mole ratio
between acid and base in the reaction mixture is the
same as in the balanced equation.
• Given the concentration of titrant, the mole of titrant
can be calculated as: mole = Molarity x Volume (L)
• Then the mole of the other reactant can be calculated
from the mole of titrant and the mole ratio in the
equation (review stoichiometry: mole-to-mole).
• Finally, the molarity of other reactant can be
determined.
Example: Acid-Base Titration
Example:
• The titration of 10.00 mL of HCl solution of unknown
concentration requires 12.51 mL of 0.100 M Ba(OH)2
solution to reach the endpoint. What is the concentration
of the unknown HCl solution?
44
Example:
The titration of 10.00 mL of HCl
solution of unknown concentration
requires 12.54 mL of 0.100 M
Ba(OH)2 solution to reach the end
point. What is the concentration of
the unknown HCl solution?
Information
Given: 10.00 mL HCl
12.54 mL 0.100 M
Ba(OH)2
Find: M HCl
• First, write balanced equation:
2 HCl(aq) + Ba(OH)2(aq) → BaCl2 (aq) + 2H2O(l)
2 mole HCl = 1 mole Ba(OH)2
0.100 M Ba(OH)2 0.100 mol Ba(OH)2  1 L sol’n
moles solute
M olarity 
liters solution
45
Example:
The titration of 10.00 mL of HCl
solution of unknown concentration
requires 12.51 mL of 0.100 M
Ba(OH)2 solution to reach the end
point. What is the concentration of
the unknown HCl solution?
Information
Given: 10.00 mL HCl
12.51 mL Ba(OH)2
Find: M HCl
CF:
2 mol HCl = 1 mol Ba(OH)2
0.100 mol Ba(OH)2 = 1 L
M = mol/L
• Write a Solution Map:
mL
Ba(OH)2
L
Ba(OH)2
0.001 L
1 mL
mL
HCl
0.001 L
1 mL
mol
Ba(OH)2
0.100 mol Ba  O H  2
1 L Ba  OH  2
L
HCl
mol
HCl
2 mol HCl
1 mol Ba(OH) 2
moles HC l
M olarity 
liters HC l
46
Example:
The titration of 10.00 mL of HCl
solution of unknown concentration
requires 12.51 mL of 0.100 M
Ba(OH)2 solution to reach the end
point. What is the concentration of
the unknown HCl solution?
Information
Given: 10.00 mL HCl
12.51 mL Ba(OH)2
Find: M HCl
CF:
2 mol HCl = 1 mol Ba(OH)2
0.100 mol Ba(OH)2 = 1 L
M = mol/L
SM: mL Ba(OH)2 → L Ba(OH)2 →
mol Ba(OH)2 → mol HCl;
mL HCl → L HCl & mol  M
= 2.50 x 10-3 mol HCl
47
Example:
The titration of 10.00 mL of HCl
solution of unknown concentration
requires 12.51 mL of 0.100 M
Ba(OH)2 solution to reach the end
point. What is the concentration of
the unknown HCl solution?
Information
Given: 10.00 mL HCl
12.51 mL Ba(OH)2
Find: M HCl
CF:
2 mol HCl = 1 mol Ba(OH)2
0.100 mol Ba(OH)2 = 1 L
M = mol/L
SM: mL Ba(OH)2 → L Ba(OH)2 →
mol Ba(OH)2 → mol HCl;
mL HCl → L HCl & mol  M
0.001 L
1 0.00 mL H Cl 
 0 . 01000 L H Cl
1 mL
-3
2.50 x 10 moles HCl
Molarity 
 0.250 M
0.01000 L HCl
48
How does pH change?
Initial pH
1 L Pure
water
1 L 0.14 M
K2HPO4 +
0.10 M
KH2PO4
pH after
pH after
adding 1 mL adding 1 mL
1 M HCl
1 M NaOH
4.00
10.00
7.00
6.99
7.01
7.00
49
Buffers
• Definition: solutions that resist changing pH when
small amounts of acid or base are added
• The mixture of 0.14 M K2HPO4 + 0.10 M KH2PO4
solution has much smaller pH change when strong
acid or base is added, thus is called Buffer.
• Ingredient: mixing together a weak acid and its
conjugate base
 or weak base and it conjugate acid
Online demo: https://www.youtube.com/watch?v=P-R-Cqvb5yo
• Human body fluid as buffer: H2CO3/HCO350
Buffer
Composition:
• a weak acid + its salt;
example: HC2H3O2 / NaC2H3O2, HF/KF
When acid is added:
C2H3O2- + H+  HC2H3O2
When base is added:
OH- + HC2H3O2  C2H3O2- + H2O
• OR, a weak base + its salt
example: NH3 / NH4Cl
51
Acetic Acid/Acetate Buffer
52
Treasure Hunt:
Which two can combine into a
Buffer?
HCl
NH4+
C2H3O2ClHCO3CO32HC2H3O2
NH3
H2CO3
53