Chapter 15
&
Properties of acids
Taste Sour (kids, don’t try this at home).
 Conduct electricity.
 Some are strong, some are weak
electrolytes.
 React with metals to form hydrogen
gas.
 Change indicators (litmus red).
 React with hydroxides to form water
and a salt.

Properties of bases
React with acids to form water and a
salt.
 Taste bitter.
 Feel slippery (Don’t try this either).
 Can be strong or weak electrolytes.
 Change indicators (litmus blue).

Properties

electrolytes

electrolytes

sour taste

bitter taste

turn litmus red

turn litmus blue

react with metals to
form H2 gas

slippery feel

vinegar, milk, soda,
apples, citrus fruits

ammonia, lye, antacid,
baking soda
Water
Water ionizes- falls apart into ions.
+
 H2O  H + OH
 Called the self ionization of water.
 Only a small amount.
 [H+ ] = [OH-] = 1 x 10-7M
 A neutral solution.
 In water Kw = [H+ ] x [OH-] = 1 x 10-14
 Kw is called the ion product constant.

Ionization of Water
H2O + H2O
Kw =
+
[H3O ][OH ]
H3
+
O
+
OH
= 1.0 
-14
10
Ionization of Water

Find the hydroxide ion concentration of 3.0 
10-2 M HCl.
[H3O+][OH-] = 1.0  10-14
[3.0  10-2][OH-] = 1.0  10-14
[OH-] = 3.3  10-13 M
Acidic or basic?
Acidic
Ion Product Constant
H2O
H+ + OH Kw is constant in every aqueous.
solution [H+] x [OH-] = 1 x 10-14M2
 If [H+] > 10-7 then [OH-] < 10-7
 If [H+] < 10-7 then [OH-] > 10-7
 If we know one, we can determine the
other.
 If [H+] > 10-7 acidic [OH-] < 10-7
 If [H+] < 10-7 basic [OH-] > 10-7

Logarithms
Powers of ten.
 A shorthand for big, or small numbers.
 pH = -log[H+]
 in neutral pH = - log(1 x 10-7) = 7
 in acidic solution [H+] > 10-7
 pH < -log(10-7)
 pH < 7
 in base pH > 7

pH and pOH
pOH = - log [OH-]
+
-14 2
 [H ] x [OH ] = 1 x 10 M
 pH+pOH = 14

[H+]
100 10-1 10-3 10-5 10-7 10-9 10-11 10-13 10-14
pH
0
1
Acidic
14 13
10-14 10-13
3
11
5
7 9
Neutral
9
7 5
11
3
13
14
Basic
1
0
pOH
10-11 10-9Basic
10-7 10-5 10-3 10-1 100
[OH-]
pH Scale
14
0
7
INCREASING
ACIDITY
pH =
NEUTRAL
+
-log[H3O ]
pouvoir hydrogène (Fr.)
“hydrogen power”
INCREASING
BASICITY
pH Scale
pH of Common Substances
pH Scale
pH =
+
-log[H3O ]
pOH =
-log[OH ]
pH + pOH = 14
pH Scale

What is the pH of 0.050 M HNO3?
pH = -log[H3O+]
pH = -log[0.050]
pH = 1.3
Acidic or basic?
Acidic
pH Scale

What is the molarity of HBr in a solution that
has a pOH of 9.6?
pH + pOH = 14
pH = -log[H3O+]
pH + 9.6 = 14
4.4 = -log[H3O+]
pH = 4.4
-4.4 = log[H3O+]
Acidic
[H3O+] = 4.0  10-5 M HBr
Types of Acids and Bases
Several Definitions
Arrhenius Definition
Acids produce hydrogen ions in
aqueous solution.
 Bases produce hydroxide ions when
dissolved in water.
 Limits to aqueous solutions.
 Only one kind of base.
 NH3 ammonia could not be an
Arrhenius base.

Definition

Arrhenius - In aqueous solution…
– Bases form hydroxide ions (OH-)
NH3 + H2O  NH4+ + OHH
H
H
N
H
base
O
H
H
O
N
H
–
+
H
H
H
Polyprotic Acids
Some compounds have more than 1
ionizable hydrogen.
 HNO3 nitric acid - monoprotic



H2SO4 sulfuric acid - diprotic - 2 H+
H3PO4 phosphoric acid - triprotic - 3 H+
Bronsted-Lowry Definitions
And acid is an proton (H+) donor and a
base is a proton acceptor.
 Acids and bases always come in pairs.
 HCl is an acid.
 When it dissolves in water it gives its
proton to water.

H3O+ + Cl-

HCl(g) + H2O(l)

Water is a base makes hydronium ion.
Bronsted-Lowry Definitions

Brønsted-Lowry
– Acids are proton (H+) donors.
– Bases are proton (H+) acceptors.
HCl + H2O  Cl– + H3O+
acid
base
conjugate base
conjugate acid
Bronsted-Lowry Definitions

Give the conjugate base for each of the following:

HF
F-
H3PO4
H2PO4-
H3O+
H2O
Polyprotic - an acid with more than one H+
Bronsted-Lowry Definitions

Give the conjugate acid for each of the following:
Br -
HBr
HSO4-
H2SO4
CO3
HCO3
2-
-
Come in Pairs
General equation
 HA(aq) + H2O(l)
H3O+(aq) + A-(aq)
 Acid + Base
Conjugate acid +
Conjugate base
 This is an equilibrium.
 B(aq) + H2O(l)
BH+(aq) + OH-(aq)
 Base + Acid
Conjugate acid +
Conjugate base
 NH3(aq)+H2O(l)
NH4+(aq)+OH-(aq)

How Strong
Strength
Strong acids and bases are strong
electrolytes
 They fall apart completely.
 Weak acids don’t completely ionize.
 Concentrated much dissolved.
 Strong forms may ions when dissolved.
 Mg(OH)2 is a strong base- it falls
completely apart when dissolved.
 Not much dissolves.

Measuring strength
Ionization is reversible.
 HA
H+ + A makes an equilibrium.
 Equilibrium constant for an acid(acid
dissociation constant.
 Ka = [H+ ][A- ]
[HA]
 Stronger acid- more products.
 larger Ka (pg 450)

What about bases?

Strong bases dissociate completely.

B + H2O
BH+ + OH-
Base dissociation constant.
 Kb = [BH+ ][OH-]
[B] we can ignore the water
 Stronger base more dissociated.
 Larger Kb.

Practice
Write the expression for HNO2
 Write the Kb for NH3

Neutralization reactions
Neutralization Reactions
+ Base  Salt + water
 Salt = an ionic compound
 Water = HOH
 HNO3 + KOH 
 HCl + Mg(OH)2 
 H2SO4 + NaOH 
 Really just double replacement.
 Acid
Reactions Happen in Moles
 How
many moles of HNO3 are need to
neutralize 0.86 moles of KOH?
 How many moles of HCl are needed to
neutralize 3.5 moles of Mg(OH)2 ?
Usually happen in solutions
 If
it takes 87 mL of an HCl solution to
neutralize 0.67 moles of Mg(OH)2 what
is the concentration of the HCl solution?
 If it takes 58 mL of an H2SO4 solution to
neutralize 0.34 moles of NaOH what is
the concentration of the H2SO4 solution?
Titration
Determining an unknown
Neutralization
Chemical reaction between an acid and a
base.
 Products are a salt (ionic compound) and
water.

Neutralization
ACID + BASE  SALT + WATER
HCl + NaOH  NaCl + H2O
strong
strong
neutral
HC2H3O2 + NaOH  NaC2H3O2 + H2O
weak
strong
basic
– Salts can be neutral, acidic, or basic.
– Neutralization does not mean pH = 7.
Titration

Titration
standard solution
– Analytical method in
which a standard
solution is used to
determine the
concentration of an
unknown solution.
unknown solution
Titration

Equivalence point
(endpoint)
– Point at which equal amounts of
H3O+ and OH- have been added.
– Determined by…
• indicator color change
• dramatic change in pH
Titration
+
O
moles H3 = moles
MVn = MVn
M: Molarity
V: volume
n: # of H+ ions in the acid
or OH- ions in the base
OH
Titration

42.5 mL of 1.3M KOH are required to
neutralize 50.0 mL of H2SO4. Find the
molarity of H2SO4.
H3O+
OH-
M=?
M = 1.3M
V = 50.0 mL
n=2
V = 42.5 mL
n=1
MV# = MV#
M(50.0mL)(2)
=(1.3M)(42.5mL)(1)
M = 0.55M H2SO4
Titration
When you add the same number of moles
of acid and base, the solution is neutral.
 By measuring the amount of a base added
you can determine the concentration of
the acid.
 If you know the concentration of the base.
 This is a titration.

Normality
Want moles of H+ and OH molarity x liters = moles of acid or base
 Don’t want moles of acid or base
 Want moles of H+ and OH Moles H+ = Molarity x liters x # of H+
 Normality = Molarity x # of H+


Normality x Liters = Moles of H+

Same process for base yields
Titration equations
Ma x Va x # of H+ = Mb x Vb x # of OH Na x Va = Nb x Vb
 really moles of H+= moles of OH
Practice
What is the normality of the following.
 2.0 M hydrofluoric acid
 0.18 M phosphoric acid
 4.0 M potassium hydroxide
 0.0020 calcium hydroxide

More Practice
If it takes 45 mL of a 1.0 M NaOH solution
to neutralize 57 mL of HCl, what is the
concentration of the HCl ?
 If it takes 67 mL of 0.500 M H2SO4 to
neutralize 15mL of Al(OH)3 what was the
concentration of the Al(OH)3 ?
 How much of a 0.275 M HCl will be
needed to neutralize 25mL of .154 M
NaOH?
