Chapter 14 Lecture

Report
Chapter 14
Acids
and
2006, Prentice hall
Bases
CHAPTER OUTLINE






General Properties
Arrhenius Acids & Bases
Strength of Acids & Bases
Brønsted-Lowery Acids & Bases
Ionization of Water
pH and pOH Scales
2
GENRAL PROPERTIES
OF ACIDS & BASES
 Many common substances in our daily lives are
acids and bases.
 Oranges, lemons and vinegar are examples of
acids. In addition, our stomachs contain acids
that help digest foods.
 Antacid tablets taken for heartburn and
ammonia cleaning solutions are examples of
bases.
3
GENRAL PROPERTIES
OF ACIDS
 General properties associated with acids include
the following:
 sour taste
 change color of litmus from blue to red
 react with metals to produce H2 gas
 react with bases to produce salt & water
4
Structure of Acids
• binary acids have acid
hydrogens attached to
a nonmetal atom
– HCl, HF
Hydrofluoric acid
Structure of Acids
• oxy acids have acid
hydrogens attached
to an oxygen atom
– H2SO4, HNO3
Structure of Acids
• carboxylic acids
have COOH group
– HC2H3O2, H3C6H5O3
• only the first H in
the formula is acidic
– the H is on the COOH
GENRAL PROPERTIES
OF BASES
 General properties associated with bases include
the following:
 bitter taste
 slippery, soapy feeling
 change color of litmus from red to blue
 react with acids to produce salt & water
8
Structure of Bases
• most ionic bases
contain OH ions
– NaOH, Ca(OH)2
• some contain CO32ions
– CaCO3 NaHCO3
• molecular bases
contain structures
that react with H+
– mostly amine groups
amine groups
Ammonia
H–N–H
H
+ H+
H
H – N+ – H
H
Ammonium ion
ARRHENIUS
ACIDS & BASES
According
 The most
to the
common
Arrhenius
definition
definition,
of acids and bases
was formulated
by the
Swedish
chemist
Svante
 Acids
are substances
that
produce
hydronium
Arrhenius
ions
(H3O+)inin1884.
aqueous solution.
HCl (g) + H2O (l)
H3O+ (aq) + Cl– (aq)
Commonly written as
HCl (g)
H2O
H+ (aq) + Cl– (aq)
11
ARRHENIUS
ACIDS & BASES
According to the Arrhenius definition,
 Bases are substances that produce hydroxide
ion (OH-) in aqueous solution.
NaOH (s)
H2O
NH3 (aq) + H2O (l)
Na+ (aq) + OH– (aq)
NH4+ (aq) + OH – (aq)
12
BRØNSTED-LOWRY
ACIDS & BASES
 The
According
Arrhenius
to Brønsted-Lowry
definition of acids
definition,
and bases
anis
limited
acid
is atoproton
aqueous
donor,
solutions.
and a base is a proton
A substance that can act as a Brønsted-Lowry acid
acceptor.
 base
A broader
definition
acids amphiprotic.
and bases was
and
(such as
water) isofcalled
developed by Brønsted and Lowry in the early
20th century.
Base
Acid
Acid
Base
++ (aq) + OH
– –(aq)
NH
HCl
(aq)
(g)
+
H
O
(l)
→
NH
H
O
Cl
(aq)
3
2
3 4
13
BRØNSTED-LOWRY
ACIDS & BASES
 In Brønsted-Lowry definition, any pair of
molecules or ions that can be inter-converted
by transfer of a proton is called conjugate
acid-base pair.
HCl (g) + H2O (l) → H3O+ (aq) + Cl– (aq)
Acid
Base
Conjugate
acid
Conjugate
base
14
BRØNSTED-LOWRY
ACIDS & BASES
NH3 (aq) + H2O (l) → NH4+ (aq) + OH– (aq)
Base
Acid
Conjugate
acid
Conjugate
base
15
Example 1:
Identify the conjugate acid-base pairs for each
reaction shown below:
H2O +
Acid
Cl 
Base
HCl +
Conjugate
acid
OH
Conjugate
base
16
Example 1:
Identify the conjugate acid-base pairs for each
reaction shown below:
C6H5OH + C2H5O  C6H5O + C2H5OH
Acid
Base
Conjugate
base
Conjugate
acid
17
Example 2:
Write the formula for the conjugate acid for each
base shown:
HS + H+  H2S
NH3 + H+  NH4+
CO32 + H+  HCO3
18
Example 3:
Write the formula for the conjugate base for each
acid shown:
HI - H+  I
CH3OH - H+  CH3O
HNO3 - H+  NO3
19
ACID & BASE
STRENGTH
 According
Strong acids
to and
the Arrhenius
bases are those
definition,
that ionize
the
strength
completely
of acids
in water.
and bases is based on the
of their
water.electrolytes.
 amount
Strong acids
andionization
bases arein
strong
H 2O
++
-
NaO
H (aq)
(s) ¾ ¾¾ ¾
¾
Na (aq)
(aq)++Cl
O H(aq)
(aq)
H Cl
® H
100
100
100
1M
1M
1M
20
ACID & BASE
STRENGTH
 Weak acids and bases are those that ionize
partially in water.
 Weak acids and bases are weak electrolytes.
H O
+
2
+
¾
¾
®
H
C
H
O
H
(aq)
+
H 3+OO2 H(aq)(aq)
¾
®
¬2 O
¾ (l)
¾
¾ ¬ ¾¾ NH 4 C
NH 3 (aq)
(aq)
2
3 +2 H
2
100
~1
~1
1M
~0.01M
~0.01M
21
IONIZATION OF
STRONG vs. WEAK ACIDS
Ionizes
completely
Ionizes
partially
22
COMMON
ACIDS
Strong Acids
Weak Acids
HCl
Hydrochloric acid
HC2H3O2
Acetic acid
HBr
Hydrobromic acid
H2CO3
Carbonic acid
HI
Hydroiodic acid
H3PO4
Phosphoric acid
HNO3
Nitric acid
HF
Hydrofluoric acid
H2SO4
Sulfuric acid
HCN
Hydrocyanic acid
H2S
Hydrosulfuric acid
23
COMMON
BASES
Strong Bases
Weak Bases
LiOH
Lithium hydroxide
NH3
Ammonia
NaOH
Sodium hydroxide
CO(NH2)2
Urea
Ca(OH)2
Calcium hydroxide
KOH
Potassium hydroxide
Ba(OH)2
Barium hydroxide
24
COMPARISON OF
ACIDS & BASES
25
Titration
• using reaction stoichiometry
to determine the
concentration of an unknown
solution
• Titrant (unknown solution)
added from a buret
• indicators are chemicals
added to help determine
when a reaction is complete
• the endpoint of the titration
occurs when the reaction is
complete
Titration
Titration
The base solution is the
titrant in the buret.
As the base is added to
the acid, the H+ reacts with
the OH– to form water.
But there is still excess
acid present so the color
does not change.
At the titration’s endpoint,
just enough base has been
added to neutralize all the
acid. At this point the
indicator changes color.
Example 1
•The titration of 10.00 mL HCl solution of unknown
concentration requires 12.54 mL of 0.100 M NaOH
solution to reach the end point. What is the
concentration of the unknown HCl solution?
Find: concentration HCl, M
Collect Needed Equations and Conversion Factors:
HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)
 1 mole HCl = 1 mole NaOH
0.100 M NaOH 0.100 mol NaOH  1 L sol’n
Molarity 
moles solute
liters solution
12.54 mL NaOH 
0.001 L

0.100 mol NaOH
1 mL

1L
1 mol HCl
1 mole NaOH
= 1.25 x 10-3 mol HCl
10.00 mL HCl 
0.001 L
 0.01000 L HCl
1 mL
Molarity 
1.25 x 10
-3
moles HCl
 0.125 M
0.01000 L HCl
molarity = 0.125 M HCl
IONIZATION
OF WATER
 Water can act both as an acid and a base.
 In pure water, one water molecule donates a
proton to another water molecule to produce
ions.
H2O
Acid
+
H2O 
Base
H3O+
Conjugate
acid
+ OH–
Conjugate
base
31
IONIZATION
OF WATER
+ + andOH
 When
Allpure
In
aqueous
the
water,
concentrations
solutions
the transfer
haveofof
HH3protons
O
between
OH– –ions.
are
3O and
water
molecules
together,
produces
the ion-product
equal numbers
constant
 multiplied
An increase
in the
concentration
of one
of of
the
+isand
– ions.
H
(K
O
)
formed.
OH
ions
3w will cause an equilibrium shift that causes
 a
The
decrease
number
in of
theions
other
produced
one. in pure water is
very small, as indicated below:
[H3O+] = [OH–] = 1.0 x 10-7 M
Kw=[H3O+][OH–] =[1.0x10-7][1.0x10-7]=1.0x10-14
32
ACIDIC & BASIC
SOLUTIONS
 When [H3O+] is
and
greater
[OH–]than
are equal
[OH–]ininaa
solution, it is acidic.
neutral.
 For example, if [H3O+] is 1.0 x 10–4 M, then
[OH–] would be 1.0 x 10–10 M.
[O H ]=
Kw
+
[H 3 O ]
=
1.0 x 10
14
1.0 x 10
4
=1.0 x 10
10
M
33
ACIDIC & BASIC
SOLUTIONS
 When [OH-] is greater than [H3O+] in a
solution, it is basic.
 For example, if [OH-] is 1.0 x 10–6 M, then
[H3O+] would be 1.0 x 10–8 M.
+
[H 3 O ]=
Kw
[O H ]
=
1.0 x 10
-14
1.0 x 10
-6
=1.0 x 10
-8
M
34
ACIDIC & BASIC
SOLUTIONS
Basic
Acidic
Neutral
[H3O+]>[OH-]
[H3O+]=[OH-]
[H3O+]<[OH-]
35
pH SCALE
 The acidity of a solution is commonly
measured on a pH scale.
 The pH scale ranges from 0-14, where acidic
solutions are less than 7 and basic solutions
are greater than 7.
pH = -log [H3O+]
36
pH SCALE
Acidic solutions
pH < 7
H3O+ > 1x10-7
Neutral solutions
pH = 7
H3O+ = 1x10-7
Basic solutions
pH > 7
H3O+ < 1x10-7
37
Example 1:
The [H3O+] of a liquid detergent is 1.4x10–9 M.
Calculate its pH.
Solution
is basic
pH = -log [H3O+] = -log [1.4x10-9] = -(-8.85)
pH = 8.85
2 significant
figures
The number of decimal places in a logarithm is equal to
the number of significant figures in the measurement.
38
Example 2:
The pH of black coffee is 5.3. Calculate its
Solution
[H3O+].
is acidic
[H3O+] = antilog (-pH) = 10 –pH = 10 -5.3
[H3O+] = 5 x 10-6
1 significant
figure
39
Example 3:
The [H3O+] of a solution is 3.5 x 10–3 M. Calculate
its pH.
Solution
is acidic
pH = -log [H3O+] = -log [3.5x10-3] = -(-2.46)
pH = 2.46
2 significant
figures
40
Example 4:
The pH of tomato juice is 4.1. Calculate its
[H3O+].Solution
is acidic
[H3O+] = antilog (-pH) = 10 –pH = 10 -4.1
[H3O+] = 8 x 10-5
1 significant
figure
41
Example 5:
The [OH] of a cleaning solution is 1.0 x 105 M.
What is the pH of this solution?
2 sig
figs
Kw=[H3O+][OH–Solution
]
is basic
[H3O+] =
Kw
[O H ]
pH = log[H3O+]
=
1.0 x 10
14
1.0 x 10
5
= 1.0 x 109 M
= log(1.0x109)
= 9.00
42
Example 6:
The pH of a solution is 11.50. Calculate the
[H3O+] Solution
for this solution.
is basic
[H3O+] = antilog (-pH) = 10 –pH = 10 -11.50
[H3O+] = 3.2 x 10-12
43
THE END
44

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