UNIT 6 – Acids & Bases & Redox Rxns
Chapter 19 – Acids, Bases, and Salts
Chapter 20 – Oxidation-Reduction Rxns
Chapter 19
Acids, Bases, and Salts
Anything in black letters = write it in
your notes (‘knowts’)
19.1 – Acid-Base Theories
Acids
Taste sour
Dissolve active metals to produce hydrogen gas
Turns litmus paper RED
Bases
Taste bitter
Feels slippery on skin (dissolves oils on skin)
Turns litmus paper BLUE
Have you seen the litmus paper yet??
These are experimental definitions, they do not explain
(theory) how an acid is different from a base.
Arrhenius defined an acid
and base theoretically.
Svante Arrhenius
(1857 – 1927)
First, a vocal word…
Dissociate to split or separate from another
Arrhenius Definition (~1887)
ACID – substance that dissociates in water
to form hydrogen ions (H+).
HCl (aq)  H+ (aq) + Cl- (aq)
BASE – substance that dissociates in water
to form hydroxide ions (OH-).
NaOH (aq)  Na+ (aq) + OH- (aq)
When an acid is placed in water, H+ ions are
produced.
Hydrogen ions can also be thought of as H3O+ ions.
H3O+ = hydronium ion
HCl (aq)  H+ (aq) + Cl- (aq)
or equivalently,
HCl
+
H 2O  H 3O +
+
Cl-
Brønsted-Lowry Definition (~1923)
ACID – donates H+
BASE – accepts H+
Johannes Bronstad
(1879 – 1947)
Thomas Lowry
(1874 – 1936)
B-L definition covers more examples than
the Arrhenius definition.
ammonia
ammonium ion
water donates a H+ and so is a B-L acid
ammonia accepts a H+ and so is a B-L base
Conjugate Acid – formed when a base
accepts a H+
Conjugate Base – formed when an acid
donates a H+
Amphoteric – substance that can be an acid or a
base – depending on what it reacts with.
Water is amphoteric
ACIDS donate H+, BASES accept H+
label each reactant as an acid and base,
label the products as conjugate acids or conjugate bases.
HNO3 + H2O  H3O+ + NO3-
CH3COOH + H2O  H3O+ + CH3COONH3 + H2O  NH4+ + OHH2O + CH3COO-  CH3COOH + OH-
Lewis Acids and Bases (not covered)
Assignment:
Chapter 19 #52-57 (p. 684)
Acid/Base Indicators
Litmus
Acid – red, Base – blue, Neutral - colorless
Phenolphthalein
Acid – colorless, Base – pink, Neutral - colorless
Cabbage
Acid – red/pink, Base – yellow/green, Neutral – blue/purple
19.2 – Hydrogen Ions and Acidity
Molarity (M) – unit used to express the concentration of a
solution
mol solute (mol)
Molarity =
liters of soln (L)
anything in [brackets] means the concentration in molarity
[H+] = ‘the hydrogen ion concentration’
[OH-] = ‘the hydroxide ion concentration’
Self-Ionization of Water
Water ionizes to produce a small amount of H+ and
OH- ions.
H2O
 H+ +
OH-
In pure water at 25̊C
[H+] = [OH-] = 1 x 10-7 M
Ion-product constant for water (Kw)
Kw = [H+][OH-] = 1.0 x 10-14
remember…anything in [brackets] represents the concentration
in molarity
A solution is acidic if [H+] > 1.0 x 10-7 M
…or if the pH of the solution is below 7
Just as the mole was used to simplify large numbers of atoms,
pH is used to simplify small concentration numbers
pH = ‘power of the hydrogen ion’
pH = -log[H+]
Instead of writing out numbers like these…
[H+] = 1 x 10-7 M
[H+] = 2.4 x 10-4 M
pH = 7.00
[H+] = 7.3 x 10-10 M
pH = 9.14
pH = 3.62
we can write number like these
Instead of saying “This solution has a hydronium
ion concentration of 2.4 x 10-4 M”.
We can just say “This solution has a pH of 3.62”.
Not only is pH an easier number to talk about, pH
is understood by most people, whereas molarity is
not.
The pH scale is used to describe how acidic or
basic (alkaline) a substance is.
Examples
Pure water has [H+] = 1.00 x 10-7 M
The pH of water would be
pH = -log[H+]
pH = - log [1.00 x 10-7]
pH = 7
Examples
[H+] = 2.3 x 10-5 M. Calculate the pH.
pH = - log [H+]
pH = - log [2.3 x 10-5]
pH = 4.64
Examples
[H+] = 1.0 x 10-5 M. Calculate the pH.
pH = - log [H+]
pH = - log [1.0 x 10-5]
pH = 5.0
Examples
pH = 4.2. Calculate [H+]
pH = - log [H+]
-4.2 = log [H+]
10-4.2 = 10log [H+]
10-4.2 = 6.31 x 10-5 M = [H+]
Summary of pH
The pH scale is used to indicate how acidic or
basic a substance is.
The scale normally ranges from 0 to 14.
Acids have a pH below 7, bases are above 7.
pH = - log [H+]
[H+] = 10-pH
ASSIGNMENT:
Chapter 19 #10-24 (p.655-662)
19.3 – Strengths of Acids and Bases
Coming soon!!
19.4 – Neutralization Reactions
Neutralization Rxn – complete rxn of a strong base
with a strong acid
A neutralization rxn will produce a salt and water.
Acid
HCl
+
Base

+ NaOH 
Salt + H2O
NaCl + H2O
Titration – determining the concentration of an
unknown solution using a solution whose
concentration is known.
Standard – solution of
known concentration.
Equivalence Point –
point where the amount of
acid equals the amount of
base
End Point – point where
the indicator changes
color
EXAMPLE
10.0 mL of 0.5 M HCl solution is added to 20.0
mL of NaOH of unknown concentration. What
is the concentration of the NaOH?
HCl
+ NaOH 
0.5 M
10.0 mL
xM
20.0 mL
NaCl + H2O
Since the reaction of HCl and NaOH is 1:1 and
twice the volume of NaOH was used, the NaOH
must half as strong as HCl; [0.25 M].
EXAMPLE
What volume of 0.10 M KOH is required to
neutralize 20.0 mL of 0.20 M H2SO4 solution?
H2SO4 + 2KOH 
0.20 M
0.10 M
20.0 mL
x mL
K2SO4 + 2H2O
Always adjust if the rxn is NOT 1:1
Since KOH requires twice as many moles as
H2SO4, you should double your answer.
19.5 – Salts in Solution
not covered…
Neutralization Rxn –
Base

Acid
+
HCl
+ NaOH 
Salt + H2O
NaCl + H2O
Write the balanced chemical equation for each
neutralization reaction
Sulfuric acid + magnesium hydroxide
Phosphoric acid + calcium hydroxide
Nitric acid + ammonium hydroxide
Chapter 19 Quiz Questions
1. What color will litmus paper be in an acidic solution?
2. What color will phenolphthalein indicator be in an basic
solution?
3. Define an acid in terms of the Arrhenius definition.
4. What two products are always formed in an acid-base
neutralization reaction?
5. A student titrated 10.0 mL of an HCl solution. The
titration required 23.3 mL of 0.24M NaOH solution.
a. Which solution was the standard?
b. Which solution was more concentrated?
c. Convert the volumes to liters
d. Calculate the number of moles of NaOH that reacted.
e. Calculate the number of moles of HCl that reacted.
f.
Calculate the molarity of the HCl solution.
6. Calculate the pH of solutions with the following
hydrogen ion concentrations.
a. [H+] = 1.23 x 10-4M
b. [H+] = 3.42 x 10-7M
7. Calculate the hydrogen ion concentrations of
solutions with the given pH.
a. pH = 3.14
b. pH = 9.2