General Descriptions
Have a sour taste. Vinegar is a solution of acetic acid. Citrus
fruits contain citric acid.
React with certain metals to produce hydrogen gas.
React with carbonates and bicarbonates to produce carbon
dioxide gas
Bases
Have a bitter taste.
Feel slippery. Many soaps contain bases.
 Produce H+ (as H3O+) ions in water (the hydronium ion is a hydrogen ion
attached to a water molecule) see next slide
 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”
Lone Hydrogen ions do not exist by themselves in
solution. H+ is always bound to a water molecule
to form a hydronium ion
H3O+ (hydronium ion) also can be thought of as H+ (hydrogen ion) for our
purposes

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”
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
aluminum hydroxide
Maalox (antacid)
Al(OH)3
 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 water
Arrhenius base is a substance that produces OH- in water

Definition #2: Brønsted – Lowry
Acids – proton donor
Bases – proton acceptor
A “proton” is really just a hydrogen atom
that has lost it’s electron!
Brønsted-Lowry Theory of Acids & Bases
Conjugate Acid-Base Pairs
General Equation
Brønsted-Lowry Theory of Acids & Bases
Conjugate Acid-Base Pairs
base
Acid



Conjugate acid
conjugate base
In this example: HA donates a proton to water
which would make HA the acid and water
accepts the proton which make it a base.
The conjugate acid is H₃O⁺ the hydronium ion.
The conjugate base is A⁻, which is everything
that remains after the acid donates its proton
Conjugate Acid-Base Pairs
base
Acid

Conjugate acid
conjugate base
More specific example of acid-base conjugate
pairs




Which of the following represent conjugate
acid base pairs
A. HF and FB. NH4+ , NH3
C. HCl, H2O





Which of the following represent conjugate
acid base pairs
A. HF and FYES
B. NH4+ , NH3
YES
C. HCl, H2O
NO
* you can identify the pairs because the two
species differ by one H+




Write the conjugate base for each:
A. HClO4
B. H3PO4
C. CH3NH3+







Write the conjugate base for each:
A. HClO4
B. H3PO4
C. CH3NH3+
A. HClO4
B. H3PO4
C. CH3NH3+
H + + ClO4H + + H2PO4H + + CH3NH2
Electrolytes are species which conducts electricity when
dissolved in water. Acids, Bases, and Salts are all electrolytes.
Salts and strong Acids or Bases form Strong Electrolytes. Salt
and strong acids (and bases) are fully dissociated therefore all
of the ions present are available to conduct electricity.
HCl(s) + H2O  H3O+ + ClWeak Acids and Weak Bases for Weak Electrolytes. Weaks
electrolytes are partially dissociated therefore not all species
in solution are ions, some of the molecular form is present.
Weak electrolytes have less ions avalible to conduct
electricity.
NH3 + H2O  NH4+ + OH-
STRONG
vs
_ completely ionized
_ strong electrolyte
_ ionic/very polar bonds
Strong Acids:
HClO4
H2SO4
HI
HBr
HCl
HNO3
WEAK
_ partially ionized
_ weak electrolyte
_ some covalent bonds
Strong Bases:
LiOH
NaOH
KOH
Ca(OH)2
Sr(OH)2
Ba(OH)2
P 477 (8, 10, 12, 14, 17 and 18)
pH Scale

The scale for measuring the hydronium ion concentration [H3O+]
in any solution must be able to cover a large range. A
logarithmic scale covers factors of 10. The “p” in pH stands for
log.



A solution with a pH of 1 has [H3O+] of 0.1 mol/L or 10-1
A solution with a pH of 3 has [H3O+] of 0.001 mol/L or 10-3
A solution with a pH of 7 has [H3O+] of 0.0000001 mol/L or 10-7
pH = - log [H3O+]
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
If the pH of Coke is 3.12, [H+] = ???
Because pH = - log [H+] then
- pH = log [H+]
Take antilog (10x) of both
sides and get
10-pH = [H+]
[H+] = 10-3.12 = 7.6 x 10-4 M
*** to find antilog on your calculator, look
for “Shift” or “2nd function” and then the log
button
 A solution has a pH of 8.5.
What is the
Molarity of hydrogen ions in the solution?
 A solution has a pH of 8.5.
What is the
Molarity of hydrogen ions in the solution?
pH = - log [H+]
8.5 = - log [H+]
-8.5 = log [H+]
Antilog -8.5 = antilog (log [H+])
10-8.5 = [H+]
3.16 X 10-9 = [H+]
Find the pH of these:
1.) A 0.15 M solution of Hydrochloric acid
2) A 3.00 X 10-7 M solution of Nitric acid
 Since acids and bases are opposites,
pH and pOH are opposites!
 pOH does not really exist, but it 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
Kw = [H+] [OH-] = 1.0 x 10-14
Equilibrium constant for water
 Water or water solutions in which [H+] = [OH-] = 10-7 M
are neutral solutions.
 A solution in which [H+] > [OH-] is acidic
 A solution in which [H+] < [OH-] is basic
What is the pH of the
0.0010 M NaOH solution?
[OH-] = 0.0010 (or 1.0 X 10-3 M)
pOH = - log 0.0010
pOH = 3
pH = 14 – 3 = 11
OR
Kw = [H3O+] [OH-]
[H3O+] = 1.0 x 10-11 M
pH = - log (1.0 x 10-11) = 11.00


Acids can react with bases to form a salt and
water.
Salt definition: sometimes 'salt' simply
refers to table salt, which is sodium chloride.
Usually the term is applied to
an ionic compound produced
by reacting an acid with a base.
 Examples: NaCl, KCl, CuSO4
P 478 (32, 34, 36, 37, 40, 42, 44, 46, 48)
complete only part a and b of each question
Titration of a strong acid with a strong base
ENDPOINT = POINT OF NEUTRALIZATION =
EQUIVALENCE POINT
At the end point for the titration of a strong acid with a strong
base, the moles of acid (H+) equals the moles of base (OH-) to
produce the neutral species water (H2O). If the mole ratio in
the balanced chemical equation is 1:1 then the following
equation can be used.
MOLES OF ACID = MOLES OF BASE
Since M=n/V
nacid = nbase
MAVA = MBVB
MAVA = MBVB
1. Suppose 75.00 mL of hydrochloric acid was required to
neutralize 22.50 mLof 0.52 M NaOH. What is the molarity of
the acid?
HCl + NaOH  H2O + NaCl
Ma Va = Mb Vb rearranges to Ma = Mb Vb / Va
so Ma = (0.52 M) (22.50 mL) / (75.00 mL)
= 0.16 M
Now you try:
2. If 37.12 mL of 0.843 M HNO3 neutralized 40.50 mL of KOH,
what is the molarity of the base?
Mb = 0.773 mol/L
Molarity and Titration
Titration of a strong acid with a strong base
ENDPOINT = POINT OF NEUTRALIZATION =
EQUIVALENCE POINT
At the end point for the titration of a strong acid with a strong
base, the moles of acid (H+) equals the moles of base (OH-)
to produce the neutral species water (H2O). If the mole
ratio in the balanced chemical equation is NOT 1:1 then you
must rely on the mole relationship and handle the problem
like any other stoichiometry problem.
MOLES OF ACID = MOLES OF BASE
nacid = nbase
1. If 37.12 mL of 0.543 M LiOH neutralized 40.50 mL of
H2SO4, what is the molarity of the acid?
2 LiOH + H2SO4  Li2SO4 + 2 H2O
First calculate the moles of base:
0.03712 L LiOH (0.543 mol/1 L) = 0.0202 mol LiOH
Next calculate the moles of acid:
0.0202 mol LiOH (1 mol H2SO4 / 2 mol LiOH)= 0.0101 mol H2SO4
Last calculate the Molarity:
Ma = n/V = 0.010 mol H2SO4 / 0.4050 L = 0.248 M
2. If 20.42 mL of Ba(OH)2 solution was used to
titrate29.26 mL of 0.430 M HCl, what is the molarity
of the barium hydroxide solution?
Mb = 0.308 mol/L

A student finds that 23.54 mL of a 0.122 M NaOH
solution is required to titrate a 30.00-mL sample of
hydr acid solution. What is the molarity of the acid?
A student finds that 37.80 mL of a 0.4052 M
NaHCO3 solution is required to titrate a 20.00-mL
sample of sulfuric acid solution. What is the
molarity of the acid?
 The reaction equation is:

H2SO4 + 2 NaHCO3 → Na2SO4 + 2 H2O + 2 CO2
1. How many milliliters of 1.25 M LiOH must be added
to neutralize 34.7 mL of 10.8
0.389
M HNO3?
mL
2. What mass of Sr(OH)2 will be required to neutralize
19.54 mL of 0.00850 M HBr solution?
0.0101 g
3. How many mL of 0.998 M H2SO4 must be added to
neutralize 47.9 mL of 1.233 M KOH?
29.6 mL
4. What is the molar concentration of hydronium ion in
a solution of pH 8.25?
5.623 x 10-9 M
5. What is the pH of a solution that has a molar
concentration of hydronium ion of 9.15 x 10-5?
pH = 4.0
6. What is the pOH of a solution that has a molar
concentration of hydronium ion of 8.55 x 10-10?
pOH = 4.9
______1. How many milliliters of 0.75 M KOH must be
added to neutralize 50.0 mL of 2.50 M HCl?
______2. What mass of Ca(OH)2 will be required to
neutralize 100 mL of 0.170 M HCl solution?
______3. How many mL of 0.554 M H2SO4 must be added to
neutralize 25.0 mL of 0.9855 M NaOH?
______ 4. What is the molar concentration of hydronium ion
in a solution of pH 2.45?
______ 5. What is the pH of a solution that has a molar
concentration of hydronium ion of 3.75 x 10-9?
______ 6. What is the pOH of a solution that has a molar
concentration of hydronium ion of 4.99 x 10-4?