Chapter 17 Acids and Bases
Necessary Terminology to
Begin
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Hydronium = H3O+ (aq) or H+ (aq)
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Hydroxide = OH- (aq)
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Neutral Water = Equal amounts of
above
General Info Bases
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Taste Bitter to the taste
React with our skin to form soap (so it
feels soapy)
React with oils and greases (so often
used in cleaning products (ammonia)
Can cause dyes to change color.
Electrolytes
Corrosive
General Info about Acids
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Taste sour
Corrosive
Electrolytes
Attack skin by dissolving fatty acids.
Common Acids and Bases
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Can you name some?
Common
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Sulfuric (Batteries in Car)
Nitric (Explosives and Fertilizers
Production)
Hydrochloric (Steel Industry (Pickling))
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Also called Muriatic
Common
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Aluminum Hydroxide (Deodorant)
Magnesium Hydroxide (Laxative)
Ammonia (Cleaner)
Cough Syrups (taste awful without
flavoring)
What is an Acid and Base?
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3 different definitions that describe
what Acids and Bases are
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Arrhenius Acids and Bases
Bronsted-Lowry Acids and bases
Lewis Acids and Bases
Arrhenius Background
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His theory defined an acid as any
substance that when added to water
increases the hydronium ion
concentration
Bases as any substance when added
to water that increase the hydroxide
ion concentration.
Bronsted Acid and Bases
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The Bronsted definition of an acid
comes from a man from Denmark who
made his proposal in 1923. His theory
helped to overcome the shortcomings
of the Arrhenius definition by allowing
us to describe solutions which were
not aqueous.
Bronsted Acids and Bases
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His definition has acids as hydrogen
donors and bases as hydrogen
acceptors.
In Summary (Acids)
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An Arrhenius acid generates
hydronium ions in water.
A Bronsted acid donates hydrogens
In Summary (Bases)
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An Arrhenius base generates
hydroxide ions in water
A Bronsted base accepts protons
Practice Identifying
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On the following slides, identify the
acid and base (forward reaction) and
then whether each acid/base definition
works
Questions
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CH3COOH (aq) + H2O (L)  CH3COO- (aq)
+ H3O+ (aq)
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HCl (aq) + NH3 (aq)  NH4+ (aq) + Cl-
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NH3 (aq) + H2O (L)  NH4+ (aq) +
OH- (aq)
Amphoteric
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A substance that can act as an acid or
a base
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Water and Ammonia are common
examples
Nomenclature
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Throw as many hydrogen’s onto the
anion/polyatomic ion and change the
ending from ate to ic and ite to ous.
With the exception the halogens (Add
hydrochloric acid = HCl)
Perchlorate (ClO4-1) becomes
Perchloric Acid (HClO4)
Nomenclature
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Nitrate (NO3-1) becomes Nitric Acid
(HNO3)
Sulfite (SO3-2) becomes Sulfurous
Acid (H2SO3)
Also: When writing the formula, if it is
an acid, the H is placed at the
beginning to denote that the chemical
is an acid.
Major Ideas
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Strong and Weak Acids and Bases
Conjugate Acids and Bases
Bond Strength
Acid/Base Equilibrium
What are Strong Acids/Bases?
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Strong acids and bases fully ionize
when placed in water.
The Unionized from of the acid is not
present. There is no equilibrium,
Strong acids and bases are completion
reactions in water.
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HCl (aq) + H2O (l)  H3O+ (aq) + Cl- (aq)
If you take more chemistry,
you will need to know these
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There are 6 strong acids
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Nitric Acid (HNO3)
Sulfuric Acid (H2SO4)
PerChloric Acid (HClO4)
HydroBromic Acid (HBr)
HydroChloric Acid (HCl)
HydroIodic Acid (HI)
Strong Acids
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Every one of those, when placed in
water, will ionize and all you will have
is Hydronium and the Anion.
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HCl (aq) + H2O (l)  H3O+ (aq) + Cl- (aq)
Acid + Water  Hydronium + Anion
Strong Bases
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A strong base is any Alkali Metal with
a Hydroxide.
Such as: Sodium Hydroxide (NaOH),
Potassium Hydroxide (KOH),
Big Idea behind Strong Acids
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If 1,000,000 moleculesof HCl are
placed into 1.0 L of water, the
Concentration of Hydronium is equal to
the concentration of the Strong Acid
placed into the solution.
The concentration of Hydronium is
1,000,000 molecules in 1.0 L in this
example.
Conjugate Acids/Bases
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When an Acid gives up its proton, the
molecule left (minus a hydrogen) is
called a conjugate base.
The conjugate base has a negative
charge and, being negative, has the
ability to attract a nearby hydrogen
(which is positive) to bond to it.
Conjugate Acids/Bases
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When a Base receives a proton, the
molecule (plus a hydrogen) is called a
conjugate acid.
The conjugate acid has a positive
charge and is looking to give up its
positive charge to another molecule.
Conjugate Acids/Bases
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Summarized
Conjugate Acid  An acid that forms
when a base gains a proton
Conjugate Base  A base that forms
when an acid loses a proton
Example
Acid
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Base
HC2H3O2 (aq) + H2O 
C2H3O2- (aq) + H3O+ (aq)
Conj. Base
Conj. Acid
Example
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Base
Acid
NH3 (aq) + H2O (l) 
NH4+ (aq) + OH- (aq)
Conj. Acid Conj. Base
Practice
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Questions
What is the conjugate Base of the
following Acids?
HCl
H2SO4
Hydronium
Ammonium (NH4+)
Practice
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Label the following as
Acid/Base/CB/CA
HF + H2O  F- + H3O+
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CH3O- +H2O  CH3OH +OH-
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What are Weak Acids/Bases?
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Weak acids and bases partially ionize
when placed into water. An
equilibrium is established where the
Hydrogens are fought over (who gets
to have the hydrogen?).
Acetic acid (HC2H3O2) when placed
into water partially ionizes.
Weak Acid/Base Example
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Weak Acid
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HC2H3O2 (aq) + H2O 
C2H3O2- (aq) + H3O+ (aq)
Weak Base
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NH3 (aq) + H2O (l) 
NH4+ (aq) + OH- (aq)
How weak is weak?
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If 100 weak acid molecules were put
into a solution of water, only about 5
would react.
Most weak acids are found with their
hydrogen
New Ideas
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The pH Scale
Calculating the Hydronium
Concentration
pH Scale: Before we start…
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Quick Math: The Logarithm Scale
On a logarithmic scale, a change in 1
represents a change in 10, a change in
2 represents a change of 100.
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The Richter Scale (Earthquakes) An
earthquake that registers a 5.0 is 1000x
stronger than an earthquake that
registers 2.0
The pH Scale
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Is a measure of how many
Hydroniums are in the water.
The pH means: Powers of Hydrogen
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On the Log Scale
A pH of 7 means the concentration of
Hydronium is 0.0000001 or 1x10-7
pH Scale
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pH
0
1
2
3
4
5
pOH
14
13
12
11
10
9
[H3O+]
1
0.1
0.01
1x10-3
1 x10-4
1 x10-5
[OH-]
1 x10-14
1x10-13
1x10-12
1x10-11
1 x10-10
1 x10-9
pH Scale
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pH
6
7
8
9
10
11
pOH
8
7
6
5
4
3
[H3O+]
1 x10-6
1 x10-7
1 x10-8
1 x10-9
1 x10-10
1 x10-11
[OH-]
1 x10-8
1 x10-7
1 x10-6
1 x10-5
1 x10-4
1 x10-3
Hydronium in Distilled Water
pH: Why it adds up to 14
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Water self ionizes (to a very small
extent)
 H O+ (aq) + OH- (aq)

2 H2O (l) 
3
Kw = [H3O+][ OH-]
Kw = 1.0 x10-14 (This is a constant)
The pH scale
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When acids are added to water, they
add to the hydronium concentration
(decreasing the hydroxide
concentration)
When bases are added to water, they
add to the hydroxide concentration
(decreasing the hydronium
concentration).
The pH Scale at neutral
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Kw = [OH][H]
Kw = (1E-7)(1E-7) = 1E-14
pH = 7
A pH of 7 is considered Neutral (it
contains just as much base as acid)
Question
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Can you have a pH less than 0 and
greater than 14?
Yes, a 10.0 Molar HCl solution has a
pH of -1
What is the pH?
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A solution has a hydronium
concentration of 0.001?
A solution has a hydroxide
concentration of 0.00001?
A solution has equal concentrations of
hydronium and hydroxide?
Carbonic Acid Our Blood
Buffer
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Carbon Dioxide in blood  Reacts
with water forming Carbonic Acid
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Same as pop
More Carbon Dioxide = More Acid =
Lower pH
After a marathon, many runners have
lower pH’s then 7.4. This is because
of high carbon dioxide concentrations
in the blood.
Breathing Rate
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When you hyperventilate, what
happens to the blood pH?
When you breath into a brown bag,
what happens to the blood pH?
Neutralization
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Acids and Bases react to neutralize
each other
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Produce a salt and water
HCl + NaOH  Water and NaCl
What is a Titration?
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A way to determine the acidity or
baseness of a solution
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You have an unknown solution, you can
use a known concentration of acid to find
out what the base concentration is
What is a Titration?
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Let’s say a student puts an unknown
amount of NaOH into some water.
You can determine how much NaOH
was added by titrating the solution
against a known concentration of acid
Equivalence Point
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At the equivalence point, equal amounts of
acid and base have been added together,
so there are no reactants left.
In the case of adding HCl (a strong acid) to
NaOH (a strong base), at the equivalence
point, there is only water and salt.
HCl + NaOH  NaCl + H2O
Indicators
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Indicators change color at different
pH’s and are useful for knowing what
the pH of the solution is (they won’t
give exact, but they will let you know
more or less acidic than some number)
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During a titration, you place indicators
into the solution to let you know when
the pH has crossed a certain point.
Different titrations have different
equivalent points, so choosing an
appropriate one is important.
For a SA and SB reaction, indicator isn’t important.
For WA and WB reactions, it is more important
Phenolphthalein
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The two main differences are the extra
H (in acidic solution) in the top left of
the molecule, and the bonding of the
central carbon