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Chapter 19
Introduction To
Acids
and Bases
Properties of Acids
Taste Sour.
 Conduct electricity.
 Some are strong, some are
weak electrolytes.
 React with metals to form hydrogen gas.
movie
 Change indicators (blue litmus turns red).
 React with hydroxides to form water and a
salt.
 React with carbonates such as limestone.

Properties of Bases
React with acids to form water and a
salt.
 Taste bitter.
 Feel slippery.
 Can be strong or weak
electrolytes.
 Change indicators
(red litmus turns blue).

Properties

electrolytes

electrolytes

sour taste

bitter taste

turn blue litmus red

turn red litmus blue

react with metals to
form H2 gas

slippery feel

ammonia, lye, antacid,
baking soda

vinegar, milk, soda,
apples, citrus fruits
ChemASAP
Types of Acids and Bases
Several Definitions are Used
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.

Definitions

Arrhenius - In aqueous solution…
– Acids form hydronium ions (H3O+)
HCl + H2O 
+
H3O
H
H
Cl
acid
O
H
H
–
+
O
H
+
Cl
H
–
Cl
Definitions

Arrhenius - In aqueous solution…
– Bases form hydroxide ions (OH-)
+
NaOH+H2ONa +OH
Base
+H20
Bronsted-Lowry Definitions
An acid is a 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.



HCl(g) + H2O(l) → H3O+ + Cl-
Water is the base because it accepts a
proton and makes the hydronium ion.
Definitions

Brønsted-Lowry
– Acids are proton (H+) donors.
– Bases are proton (H+) acceptors.
HCl + H2O 
acid
–
Cl
+
+
H3O
base
conjugate base conjugate acid
Come in Pairs
General equation for weak acids and base
 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)

Definitions
H2O + HNO3  H3O+ + NO3–
B
A
CA
CB
Definitions
NH3 + H2O
B
A
NH4 +
+
CA
H
H
N
H
H
H

O
H
CB
–
+
O
N
H
OH
H
H
H
Amphoteric - can be an acid or a base.
Definitions


Give the conjugate base for each of the following:
HF
F-
H3PO4
H2PO4
H3O+
H2O
-
Polyprotic - an acid with more than one H+
Mono and Polyprotic Acids
HNO3 nitric acid - monoprotic
 Some compounds have more than 1
ionizable hydrogen and are polyprotic.



H2SO4 sulfuric acid - diprotic - 2 H+
H3PO4 phosphoric acid - triprotic - 3 H+
Definitions

Give the conjugate acid for each of the following:
Br -
HBr
HSO4-
H2SO4
CO3
HCO3
2-
-
How Strong is
an
Acid or Base?
Strength
Strong acids and bases are strong
electrolytes since they fall apart completely.
 Weak acids don’t completely ionize. movie
 Concentrated is how much has dissolved.
Strong and weak tea refers to concentration
not the percent ionization as with acids.
 When referring to acids and bases, use
‘strong and weak’ to indicate ionization not
concentration.

Strength

Strong Acid/Base
-
+
– 100% ionized in water
– strong electrolyte
HCl
HNO3
H2SO4
HBr
HI
HClO4
NaOH
KOH
Ca(OH)2
Ba(OH)2
Strength

Weak Acid/Base
-
+
– does not ionize completely
– weak electrolyte
HF
CH3COOH
H3PO4
H2CO3
HCN
NH3
Common Ionization
of Acids and Bases
Measuring Strength of Weak Acids
Ionization is reversible for weak acids and
bases.
 HA
H+ + A The amounts of H and A can change.
 Equilibrium constant (acid ionization
constant) shows weak acid strength.
 Ka = [H+ ][A- ]
[HA]
 Stronger acids make more products.
 Larger Ka indicates a stronger the acid.

What About Weak Bases?

Weak bases don’t dissociate completely.

B + H2O
BH+ + OH-
Base ionization constant.
 Kb = [BH+ ][OH-]
[B] we can ignore the water
 Stronger base means more dissociation.


The larger Kb is, the stronger the base.
Ka Weak Acids
Kb Weak Bases
Practice

Write the equilibrium expression for
HNO2

Write the Kb for NH3
Water Ionization
Some water ionizes- falls apart into ions.
 H2O  H+ + OH Called the self ionization of water.
 Only a small amount ionizes, but the
amount is constant.
 [H+ ] = [OH-] = 1 x 10-7M
 Pure water is neutral solution.
+
-14
 In water Kw = [H ] x [OH ] = 1 x 10
 Kw is called the water ion product
constant.

Ionization of Water leads to the
pH scale
Kw =
+
[H3O ][OH ]
= 1.0 
-14
10
Ion Product Constant
H2O
H+ + OH Kw is constant in every aqueous solution.
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

[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-]
Ionization Example

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
Now we need a mathematical relationship
between ion concentration and pH.
Logarithms
Powers of ten.
 A shorthand for big, or small numbers.
 pH = -log[H+]
 neutral pH = - log(1 x 10-7) = 7
 acidic solution [H+] > 10-7
 pH < 7 in acids
 in base pH > 7

pH Scale
14
0
7
INCREASING
ACIDITY
NEUTRAL
pH = -log[H3O+]
pouvoir hydrogène (Fr.)
“hydrogen power”
INCREASING
BASICITY
pH Scale
pH =
+
-log[H3O ]
pOH =
-log[OH ]
pH + pOH = 14
+
[H ]
x
[OH ]
=1x
-14
2
10 M
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 = -log[H3O+]
pH + pOH = 14
4.4 = -log[H3O+]
pH + 9.6 = 14
-4.4 = log[H3O+]
pH = 4.4
Antilog (-4.4) = [H3O+]
Acidic
[H3O+] = 4.0  10-5 M HBr
pH Scale
pH of Common Substances
Neutralization Reactions
Neutralization Reactions
+ Base  Salt + Water
 movie
 Salt = an ionic compound
 movie
 Water = HOH
 HNO3 + KOH 
 HCl + Mg(OH)2 
 H2SO4 + NaOH 
 Really just double replacement.
 Acid
Neutralization
ACID + BASE  SALT + WATER
HCl + NaOH  NaCl + H2O
strong
strong
HC2H3O2 + NaOH
weak
strong
neutral
NaC2H3O2 + H2O
basic
– Salts can be neutral, acidic, or basic.
– Neutralization does not always mean pH = 7.
Alka-Seltzer
Tablets
Titration
Determining an Unknown Concentration
Titration

Titration
standard solution
– Analytical method in
which a standard solution
is used to determine the
concentration of an
unknown solution.
– movie
unknown solution
Titration

Equivalence point (endpoint)
– Point at which equal amounts of H+ and
OH- have been added.
– Determined by…
• indicator color change
• dramatic change in pH
Common pH Indicators
Titration
When you add the same number of moles
of acid and base, the solution is neutral.
 By measuring the amount of a known base
added you can determine the
concentration of the unknown acid.
 You must know the concentration of the base,
the standard solution in this case.
 Use Stoichiometry to determine amounts

Normality
We want moles of H+ and OH molarity x liters = moles of acid or base
 Don’t want moles of acid or base
 We really 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 moles of
OH-.

Reactions Happen in Moles
 How
many moles of HNO3 are needed
to neutralize 0.86 moles of KOH?
 How many moles of HCl are needed to
neutralize 3.5 moles of Mg(OH)2 ?
Titration Equations

Ma x Va x # of H+ = Mb x Vb x # of OH-

Na x Va = Nb x Vb
moles H3O+ = moles OHMV n = MV n
M: Molarity
V: Volume
n: # of H+ ions in the acid
or OH- ions in the base
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
More Practice
If it takes 45 mL of a 1.0 M NaOH solution
to neutralize 57 mL of HCl, what is the
concentration (molarity) 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?

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?