The Chemistry of
Acids and Bases
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Acid and Bases
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Acid and Bases
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Acid and Bases
Some Properties of Acids
 Produce H+ (as H3O+) ions in water (the hydronium ion is a
hydrogen ion attached to a water molecule)
 Taste sour
 Corrode metals
 Electrolytes
 React with bases to form a salt and water
 pH is less than 7
 Turns blue litmus paper to red “Blue to Red A-CID”
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Some Properties of Bases
 Produce OH- ions in water
 Taste bitter, chalky
 Are electrolytes
 Feel soapy, slippery
 React with acids to form salts and water
 pH greater than 7
 Turns red litmus paper to blue
“Basic Blue”
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Some Common Bases
Formula
Name
Common Name
NaOH
sodium hydroxide
lye
KOH
potassium hydroxide
liquid soap
Ba(OH)2
barium hydroxide
stabilizer for plastics
Mg(OH)2
magnesium hydroxide
“MOM”
Milk of magnesia
Al(OH)3
aluminum hydroxide
Maalox (antacid)
Acid Nomenclature
Anion
Ending
Binary 
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Acid Name
-ide
hydro-(stem)-ic acid
-ate
(stem)-ic acid
-ite
(stem)-ous acid
Ternary
An easy way to remember which goes with which…
“In the cafeteria, you ATE something ICky”
Acid Nomenclature Flowchart
ACIDS
start with 'H'
2 elements
3 elements
hydro- prefix
-ic ending
no hydro- prefix
-ate ending
becomes
-ic ending
-ite ending
becomes
-ous ending
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Acid Nomenclature Review
• HBr (aq)
• H2CO3
• H2SO3

hydrobromic acid

carbonic acid

sulfurous acid
Try these
• HI
• HCl
(aq)
(aq)
•Hydroiodic Acid
•Hydrochloric Acid
• H2SO4
•Sulfuric Acid
• HNO3
•Nitric Acid
• H2CO3
•Carbonic Acid
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Try these
• HBr
• HCN
(aq)
(aq)
•Hydrobromic Acid
•Cyanic Acid
• H2SO3
•Sulfurous Acid
• HClO4
•Perchloric Acid
• CH3OOH •Acetic Acid
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Acid/Base definitions
• Definition #1: Arrhenius (traditional)
Acids – produce H+ ions (or hydronium ions H3O+)
Bases – produce OH- ions
(problem: some bases don’t have hydroxide ions!)
Arrhenius acid is a substance that produces H+ (H3O+) in water14
Arrhenius base is a substance that produces OH- in water
NaOH + H2O  OH- + Na+ + H2O
Acid & Base Definitions
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• Definition #2: Brønsted–Lowry
Acids – proton donor
PDA (proton donated acid)
BAD (Brønsted–Lowry acid donates)
Bases – proton acceptor
A “proton” is really just a hydrogen atom that has
lost it’s electron!
ACID-BASE THEORIES
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A Brønsted-Lowry Acid is a proton donor
A Brønsted-Lowry Base is a proton acceptor
The Brønsted definition means NH3 is a Base in
water — and water is itself an Acid
Base
Acid
Conjugate
Acid
Conjugate
Base
Conjugate Pairs
Conjugate Acid – the remaining ion or molecule that can re-accept the
proton and act as a base
Conjugate Base – the species that is formed when a Brønsted-Lowry base
gains a proton.
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Acids & Base Definitions
Definition #3 – Lewis
Lewis acid - a
substance that
accepts an electron
pair
Lewis base - a
substance that
donates an electron
pair
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Lewis Acids & Bases
Formation of hydronium ion is an excellent
example.
H
+
ACID
•• ••
O—H
H
BASE
••
H O—H
H
•Electron pair of the new O-H bond
originates on the Lewis base.
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Acids & Base Definitions
Summary
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SOME DEFINITIONS
• Amphoteric - A substance that can act as
either an acid or a base, e.g., H2O, HCO3-
• Polyprotic acid or base - An acid or
base that can donate or accept more than one
proton or hydroxide, e.g., H3PO4, H2CO3,
H4EDTA…Ba(OH)2, Al(OH)3
Strong and Weak Acids/Bases
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The strength of an acid (or base) is
determined by the amount of IONIZATION.
•Generally divide acids and bases into STRONG or WEAK
Strong and Weak Acids/Bases
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STRONG ACID:
HNO3 (aq) +
H2O(l) --> H3O+(aq) + NO3-(aq)
HNO3 is about 100% dissociated in water.
HNO3, HCl, HBr, HI, H2SO4 and HClO4 are among
the only known strong acids.
Strong and Weak Acids/Bases
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• Weak acids are much less than 100% ionized in
water.
One of the best known is acetic acid = CH3COOH
Strong and Weak Acids/Bases
• Strong Base: 100% dissociated in
water.
NaOH (aq)
---> Na+ (aq) +
OH- (aq)
Other common strong
bases include KOH and
Ca(OH)2.
CaO (lime) + H2O -->
Ca(OH)2 (slaked lime)
CaO
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Strong and Weak Acids/Bases
• Weak base: less than 100% ionized in
water
One of the best known weak bases is ammonia
NH3(aq) + H2O(l)  NH4+(aq) + OH-(aq)
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Other Weak Bases
The pH scale is a way
of expressing the
strength of acids and
bases. Instead of
using very small
numbers, we just use
the NEGATIVE power
of 10 on the Molarity
of the H+ (or OH-) ion.
Under 7 = acid
7 = neutral
Over 7 = base
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pH of Common Substances
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Determine the pH of the following
H3O+ concentrations
1) Solution A [H3O+] = 5.89 x 10-7
pH = ?
2) Solution B [H3O+] = 4.365 x 10-12
pH = ?
3) Solution C [H3O+] = 1.05 x 10-4
pH = ?
4) Solution D [H3O+] = 1.00 x 10-6
pH = ?
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Determine the pH of the following
H3O+ concentrations
1) Solution A [H3O+] = 5.89 x 10-7
pH = 6.23
2) Solution B [H3O+] = 4.365 x 10-12
pH = 11.36
3) Solution C [H3O+] = 1.05 x 10-4
pH = 3.98
4) Solution D [H3O+] = 1.00 x 10-6
pH = 6.00
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pH meter
• Tests the voltage of
the electrolyte
• Converts the voltage
to pH
• Very cheap, accurate
• Must be calibrated
with a buffer solution
Calculating the pH
pH = - log [H+]
(Remember that the [ ] mean Molarity)
Example: If [H+] = 1 X 10-10
pH = - log 1 X 10-10
pH = - (- 10)
pH = 10
Example: If [H+] = 1.8 X 10-5
pH = - log 1.8 X 10-5
pH = - (- 4.74)
pH = 4.74
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pH calculations – Solving for H+
If the pH of Coke is 3.12, [H3O+] = ???
Because pH = - log [H3O+] then
- pH = log [H3O+]
Take antilog (10x) of both
sides and get
10-pH = [H3O+]
[H3O+] = 10-3.12 = 7.59 x 10-4 M
*** to find antilog on your calculator, look for “Shift” or
“2nd function” and then the log button
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pOH
• Since acids and bases are opposites,
pH and pOH are opposites!
• pOH is useful for changing bases to
pH.
• pOH looks at the perspective of a
base
pOH = - log [OH-]
Since pH and pOH are on opposite
ends,
pH + pOH = 14
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pH
[H+]
[OH-]
pOH
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The pH of rainwater collected in a certain region of the
northeastern United States on a particular day was
4.82. What is the H+ ion concentration of the
rainwater?
pH testing
• There are several ways to test pH
– Blue litmus paper (red = acid)
– Red litmus paper (blue = basic)
– pH paper (multi-colored)
– pH meter (7 is neutral, <7 acid, >7
base)
– Universal indicator (multi-colored)
– Indicators like phenolphthalein
– Natural indicators like red
cabbage, radishes
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pH indicators
• Indicators are dyes that can be
added that will change color in
the presence of an acid or base.
• Some indicators only work in a
specific range of pH
• Once the drops are added, the
sample is ruined
• Some dyes are natural, like
radish skin or red cabbage
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The pigment in red cabbage juice is anthocyanin,
which changes color from red in acid solution to
purplish to green in mildly alkaline solution to
yellow in very alkaline solution.
Other pH Indicators
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Universal Indicator
1. Determine the pH to the tenth and ion concentration of the
following unknown solution given the color of each indicator:
q
Phenolphthalein - colorless
q
Universal indicator – red-orange
q
Bromcresol Green – yellowish green
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Universal Indicator
2. Determine the pH to the tenth and ion concentration of the
following unknown solution given the color of each indicator:
q
Alzarin Yellow R – yellow
q
Phenolphthalein - pink
q
Thymolphthalein –blue
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Universal Indicator
3. Determine the pH to the tenth and ion concentration of the
following unknown solution given the color of each indicator:
q
Methyl Orange – yellowish orange
q
Chlorphenol Red – yellow
q
Bromthymol Blue – yellow
q
Bromcresol Green – blue
ACID-BASE REACTIONS
Titrations
H2C2O4(aq)
acid
+
2 NaOH(aq) --->
base
Na2C2O4(aq) + 2 H2O(liq)
Carry out this reaction using a TITRATION.
Oxalic acid,
H2C2O4
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Setup for titrating an acid with a base
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Titration
1. Add solution from the buret.
2. Reagent (base) reacts with
compound (acid) in solution in the
flask.
3. Indicator shows when exact
stoichiometric reaction has
occurred. (Acid = Base)
This is called NEUTRALIZATION.