1.
Take out HW and Stamp Sheet
 #’s 20-22,24,32a,34a,35,36
#’s 31a,33a,37,39,40, 48-50,56,58ab
2.
3.
Solve this problem:
What is the pH of a solution with a pOH of
11.09?
HW = pH problem set due Mon &
study for quiz
More acidic
More basic
pH
NaOH, 0.1 M
Household bleach
Household ammonia
Lime water
Milk of magnesia
Borax
Baking soda
Egg white, seawater
Human blood, tears
Milk
Saliva
Rain
Black coffee
Banana
Tomatoes
Wine
Cola, vinegar
Lemon juice
Gastric juice
14
13
12
11
10
9
8
7
6
5
4
3
2
1
0
[H1+]
[OH1-]
1 x 10-14
1 x 10-13
1 x 10-12
1 x 10-11
1 x 10-10
1 x 10-9
1 x 10-8
1 x 10-7
1 x 10-6
1 x 10-5
1 x 10-4
1 x 10-3
1 x 10-2
1 x 10-1
1 x 100
1 x 10-0
1 x 10-1
1 x 10-2
1 x 10-3
1 x 10-4
1 x 10-5
1 x 10-6
1 x 10-7
1 x 10-8
1 x 10-9
1 x 10-10
1 x 10-11
1 x 10-12
1 x 10-13
1 x 10-14
pOH
0
1
2
3
4
5
6
7
8
9
10
11
12
13
14
pH
pH = -log[H3O+]
[H3O+]
[H3O+] = 10-pH
pH + pOH = 14
pOH
Kw = [H3O+] [OH-] = 1 x10-14
pOH = -log[OH-]
[OH-]
[OH-] = 10-pOH
pH + pOH = 14
pH + 11.09 = 14
pH = 14 – 11.09 = 2.91
Solution is acidic 
Kw = [H3O+] [OH-] = 1.0 x10-14
1.0 x10-14 = [3.3 x 10-10] [OH-]
[OH-] = 3.0 x 10-5 M
Solution is basic 
pH
pH = -log[H3O+]
[H3O+]
[H3O+] = 10-pH
pH + pOH = 14
pOH
Kw = [H3O+] [OH-] = 1 x10-14
pOH = -log[OH-]
[OH-]
[OH-] = 10-pOH
concentration (moles/L)
10-1
pH = 3
pH = 11
OH-
H3O+
pH = 7
10-7
H3O+
OH-
OH-
H3O+
10-14
[H3O+] > [OH-]
Timberlake, Chemistry 7th Edition, page 332
acidic
solution
[H3O+] = [OH-]
neutral
solution
[H3O+] < [OH-]
basic
solution
7
6
5
.
4
3 2 1
0.120 M H2SO4
X
2 mol H3O+ =
0.240 M H3O+
1 mol H2SO4
pH = - log [H3O+] = -log [0.240M H3O+]
pH = 0.620
pH
pH = -log[H3O+]
[H3O+]
[H3O+] = 10-pH
pH + pOH = 14
pOH
Kw = [H3O+] [OH-] = 1 x10-14
pOH = -log[OH-]
[OH-]
[OH-] = 10-pOH
pH = - log [H3O+]
4.60 = -log [H3O+]
2nd
log
-4.6
[H3O+] = 2.51 x 10-5 M
10- 4.60
You can check your answer by working backwards.
pH = - log [2.51x10-5 M]
pH = 4.60
pH
pH = -log[H3O+]
[H3O+]
[H3O+] = 10-pH
pH + pOH = 14
pOH
Kw = [H3O+] [OH-] = 1 x10-14
pOH = -log[OH-]
[OH-]
[OH-] = 10-pOH
standard solution

Definition
◦ Analytical method
in which a standard
solution is used to
determine the
concentration of an
unknown solution.
unknown solution
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Buret
stopcock
Erlenmeyer flask

Titrant

Analyte

Equivalence point
◦ The substance added to the analyte
in a titration (a standard solution)
◦ The substance being analyzed
◦ The point in a titration at which the
quantity of titrant is exactly
sufficient for stoichiometric reaction
with the analyte.
If the concentration of the titrant is
known, then the unknown concentration
of the analyte can be determined.
Titrant
Analyte
Quantitative analysis —
used to determine the
amounts or concentrations
of substances present in a
sample by using a
combination of chemical
reactions and
stoichiometric calculations
 The
“perfect pink”
for a titration with
phenolphthalein
Indicator changes color to
indicate pH
change
pH
Example…
7
phenolphthalein
is colorless in
acid and pink in
basic solution
point at
which
exactly
enough
reactant
pink has been
added for
Endpoint =
the
solution to
be
neutralized
and no
Volume base added more
◦ Point at which equal
amounts of H3O+ and
OH- have been added.
◦ Determined by…
 indicator color change
 dramatic change in pH
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
moles H3
+
O
= moles
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
OH

25.5 mL of 1.3M KOH are required to neutralize
50.0 mL of HCl. Find the molarity of HCl.
Acid (H3O+) Base (OH-)
M=?
M = 1.3M
V = 50.0 mL V = 25.5 mL
MA = 0.66M HCl
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Titration Data
pH
0.00
10.00
20.00
22.00
24.00
25.00
26.00
28.00
30.00
40.00
50.00
1.00
1.37
1.95
2.19
2.70
7.00
11.30
11.75
11.96
12.36
12.52
Solution
of NaOH
Na+
OH-
OHNa+ Na+
OH-
Solution
of HCl
H+
Cl-
25 mL
ClH+
H+
Cl-
H+
Cl-
14.0
12.0
10.0
8.0
equivalence point
6.0
4.0
2.0
OHNa+
Titration of an Acid With an Base
pH
NaOH added
(mL)
0.0
0.0
10.0
20.0
30.0
Volume of 0.100 M NaOH added
(mL)
40.0
•
Arrhenius - In aqueous solution…
Acids increase hydrogen ion
concentration [H+]
HCl(aq) 
+
H (aq)+
H
H
Cl
O
H
Cl-
(aq)
O
H
H
–
+
Cl
H
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
A Brønsted-Lowry acid is a proton donor
A Brønsted-Lowry base is a proton acceptor
base
acid
conjugate
acid
conjugate
base


From the Latin word conjugare, meaning “to
join together.”
Reactions between acids and bases always
yield their conjugate bases and acids.
Strong and Weak Acids/Bases
The strength of an acid (or base) is
determined by the amount of
IONIZATION.
HNO3, HCl, H2SO4 and HClO4 are among the
only known strong acids.
Strong and Weak Acids/Bases
Weak acids are much less than 100%
ionized in water.
One of the best known is acetic acid =
CH3CO2H

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
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)


The more polar the H-X bond and/or the weaker
the H-X bond, the more acidic the compound.
Acidity increases from left to right across a row
and from top to bottom down a group.
H2O can function as both an ACID and a
BASE.
Equilibrium constant for water = Kw
Kw = [H3O+] [OH-] = 1.00 x 10-14 at 25 oC
gastric
juice
1.6
vinegar
2.8
carbonated
beverage
3.0
0
1
2
acidic
Timberlake, Chemistry 7th Edition, page 335
urine
6.0
4
5
bile
8.0
6
7
neutral
[H+] = [OH-]
8
ammonia
11.0
bleach
12.0
seawater
8.5
9
1.0 M
NaOH
(lye)
14.0
milk of
magnesia
10.5
detergents
8.0 - 9.0
milk
6.4
tomato
4.2
coffee
5.0
3
blood
7.4
potato
5.8
apple juice
3.8
lemon
juice
2.2
drinking water
7.2
bread
5.5
orange
3.5
1.0 M
HCl
0
water (pure)
7.0
soil
5.5
10
11
basic
12
13
14