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Chapter 15: Acids and Bases
Mrs. Brayfield
Heart Burn
Why do people take milk of magnesia or other
medication for heart burn?
Well, what is heart burn?
It is when the acid from your stomach irritates your esophagus
This is why we take medicine that is basic to “neutralize”
the acid
15.2: The Nature of Acids and Bases
Acids have a few common properties:
Sour taste (citric acid in lemons/limes)
Ability to dissolve metals (remember Mg and HCl lab?)
Turns litmus paper red
Neutralize bases
Bases also have a few common properties:
Bitter taste (antacid tablets)
Slippery feel (like the lye in soap)
Turns litmus paper blue
Neutralizes acids
Nature of Acids and Bases
Acids are much more common in food because of the
bitter taste of bases
This is because organic bases in nature are usually poisonous
Common Acids Common Bases
Hydrochloric Acid HCl Sodium Hydroxide NaOH
Sulfuric Acid H2SO4 Potassium Hydroxide KOH
Nitric Acid HNO3 Sodium Bicarbonate NaHCO3
Acetic Acid HC2H3O2 Sodium Carbonate Na2CO3
Citric Acid H3C6H5O7 Ammonia NH3
Carbonic Acid H2CO3
Hydrofluoric Acid HF
Phosphoric Acid H3PO4 Tables 15.1 & 15.2
15.3: Definitions of Acids and Bases
Swedish chemist Svante Arrhenius (1880s) developed the
first definition for acids and bases:
Acid: A substance that produces H+ ions in aqueous solns
Base: A substance that produces OH- ions in aq solns
HCl is an acid:
𝐻𝐶𝑙 𝑎𝑞 → 𝐻+ 𝑎𝑞 + 𝐶𝑙−(𝑎𝑞)
NaOH is a base:
𝑁𝑎𝑂𝐻 𝑎𝑞 → 𝑁𝑎+ 𝑎𝑞 + 𝑂𝐻−(𝑎𝑞)
Definitions of Acids and Bases
When acids and bases are put in water they dissociate
(break up) into ions
These ions then interact with water:
As well as:
𝐻+ 𝑎𝑞 + 𝑂𝐻−(𝑎𝑞) → 𝐻2𝑂(𝑙)
Definitions of Acids and Bases
The widely accepted definition of acids and bases is called
the Brønsted-Lowry definition (1923)
Acid: Proton (H+) donor
Base: Proton (H+) acceptor
Well what does this mean?
HCl is an acid because it donates a proton:
𝐻𝐶𝑙 𝑎𝑞 + 𝐻2𝑂 𝑙 → 𝐻3𝑂+ 𝑎𝑞 + 𝐶𝑙−(𝑎𝑞)
NH3 is a base because it accepts a proton:
𝑁𝐻3 𝑎𝑞 + 𝐻2𝑂 𝑙 → 𝑁𝐻4+ 𝑎𝑞 + 𝑂𝐻−(𝑎𝑞)
Definitions of Acids and Bases
According to the Brønsted-Lowry definition, acids and
bases always occur together in the same reaction:
𝐻𝐶𝑙 𝑎𝑞 + 𝐻2𝑂 𝑙 → 𝐻3𝑂+ 𝑎𝑞 + 𝐶𝑙−(𝑎𝑞)
𝑁𝐻3 𝑎𝑞 + 𝐻2𝑂 𝑙 → 𝑁𝐻4+ 𝑎𝑞 + 𝑂𝐻−(𝑎𝑞)
acid
(proton donor)
acid
(proton donor)
base
(proton acceptor)
base
(proton acceptor)
Definitions of Acids and Bases
A substance that can act as a acid or base are called
amphoteric
Look when reactions are reversed:
𝑁𝐻4+ 𝑎𝑞 + 𝑂𝐻−(𝑎𝑞) → 𝑁𝐻3 𝑎𝑞 + 𝐻2𝑂 𝑙
What was the base (NH3) has become the acid (NH4+)
and visa versa
So we call these pairs conjugate acid-base pairs
Or two substances related to each other by the transfer of a proton
acid
(proton donor)
base
(proton acceptor)
Definitions of Acids and Bases
In any acid-base reaction,
A base accepts a proton and becomes a conjugate acid
An acid donates a proton and becomes a conjugate base
Definitions of Acids and Bases Ex 1
Write conjugate acid for the following:
Write the conjugate base for the following:
Base Conjugate Acid
H2O
NH3
CO32-
H2PO4-
Acid Conjugate Base
H2O
NH3
CO32-
H2PO4-
H3O+
NH4+
HCO3-
H3PO4
OH-
NH2-
Cannot be an acid
HPO42-
Definitions of Acids and Bases Ex 2
Identify the Brønsted-Lowry acid, base, conjugate acid,
and conjugate base:
𝐶5𝐻5𝑁 𝑎𝑞 + 𝐻2𝑂 𝑙 → 𝐶5𝐻5𝑁𝐻+ 𝑎𝑞 + 𝑂𝐻−(𝑎𝑞)
𝐻𝑁𝑂3 𝑎𝑞 + 𝐻2𝑂 𝑙 ↔ 𝐻3𝑂+ 𝑎𝑞 + 𝑁𝑂3
−(𝑎𝑞)
Acid
Acid
Base
Base
Conjugate
Base
Conjugate
Base
Conjugate
Acid
Conjugate
Acid
Definitions of Acids and Bases
Which of the following is not a conjugate acid-base pair?
A. (CH3)3N; (CH3)3NH+
B. H2SO4; H2SO3
C. HNO2; NO2-
Homework Problems: #1, 2, 4, 6, 8
B – two acids, not pairs
15.4: Acid Strength and Ka
Remember from chapter 4 that the strength of an
electrolyte is determined by the amount of dissociation
The more it dissociates, the stronger the electrolyte is
The same goes for acids:
A strong acid completely dissociates (or ionizes) in solution
A weak acid only partially dissociates in solution
Acid Strength
Hydrochloric acid is a strong acid:
There are only six strong acids (you MUST memorize
these):
Table 15.3: Strong Acids
Hydrochloric Acid HCl Nitric Acid HNO3
Hydrobromic Acid HBr Perchloric Acid HClO4
Hydriodic Acid HI Sulfuric Acid H2SO4
Acid Strength
HF is a weak acid:
For the strength, if the attraction between the conjugate
acid and conjugate base are weak, the acid is strong (the
forward reaction is favored)
Acid Strength
Acid Strength
If you notice there are some acids with multiple
hydrogens
This means that they have more than one ionizable protons
If the acid has two hydrogens (for example H2SO4) we
call that diprotic
If the acid has three hydrogens (for example H3PO4) we
call that triprotic
Whenever the acid undergoes a second (or third) dissociation,
the acid strength greatly decreases (we’ll see in section 6)
Acid Ionization Constant (Ka)
The strengths of weak acids are quantified with the acid
ionization constant (Ka), which is the equilibrium constant
for the ionization reaction of a weak acid:
So for the following reaction:
𝐻𝐴 𝑎𝑞 + 𝐻2𝑂 𝑙 ↔ 𝐻3𝑂+ 𝑎𝑞 + 𝐴− 𝑎𝑞
𝐻𝐴 𝑎𝑞 ↔ 𝐻+ 𝑎𝑞 + 𝐴− 𝑎𝑞
The equilibrium constant is:
𝐾𝑎 =𝐻3𝑂
+ [𝐴−]
[𝐻𝐴]=
𝐻+ [𝐴−]
[𝐻𝐴]
See page 560 (table 15.5) for values
Ka Example
Which acid is stronger:
HF Ka = 3.5x10-4
HClO Ka = 2.9x10-8
Which conjugate base is stronger?
HClO (the stronger the acid, the weaker the conjugate
base)
Homework Problems: #10, 11, 12, 14
HF
15.5: Autoionization of Water and pH
We know from previous sections that water is
amphoteric (can be an acid or a base)
Even in pure water, water acts as an acid and a base with
itself (called ionization):
𝐾𝑤 = 𝐻3𝑂+ 𝑂𝐻− = 𝐻+ [𝑂𝐻−]
Autoionization of Water
Kw is called the dissociation constant for water
Its value at 25°C is 1.0x10-14
In a neutral solution, [H+] = [OH-]
Or when those concentrations equal 1.0x10-7 (pg. 561)
In acidic solution, [H+] > [OH-]
In basic solution, [H+] < [OH-]
Autoionization of Water Example
Calculate [H+] at 25°C for each of the following solutions
and determine if the solution is basic, acidic, or neutral:
1. [OH-] = 1.5 x 10-2 M
2. [OH-] = 1.0 x 10-7 M
3. [OH-] = 8.2 x 10-10 M
[H+] = 6.74 x 10-13 M, basic
[H+] = 1.0 x 10-7 M, neutral
[H+] = 1.2 x 10-5 M, acidic
pH
The pH scale allows us to specify the acidity of a solution
𝑝𝐻 = −log[𝐻3𝑂+]
pH is reported to 2 decimal places
In general:
pH < 7, the solution is acidic
pH > 7, the solution is basic
pH = 7, the solution is neutral
pH Scale
pH scale is from 0-14
The higher the number,
the more basic it is
The pH changes by 1 for
every power of 10 change
in [H+]
pH Example 1
Calculate the pH of each solution and indicate whether
the solution is acidic or basic:
[H3O+] = 9.5 x 10-9 M
𝑝𝐻 = − log 𝐻3𝑂+ = −log(9.5 × 10−9)
𝑝𝐻 = 8.02
[OH-] = 7.1 x 10-3 M
𝐻3𝑂+ 𝑂𝐻− = 𝐾𝑤
𝐻3𝑂+ = 1.41 × 10−12𝑀
𝑝𝐻 = − log 1.41 × 10−12 = 11.85
pH Example 2
Calculate the H3O+ concentration for a solution with a
pH of 8.37.
𝑝𝐻 = − log 𝐻3𝑂+ = 8.37
𝐻3𝑂+ = 10−8.37 = 4.27 × 10−9𝑀
pOH
The pOH scale is the opposite of the pH scale:
𝑝𝑂𝐻 = −log[𝑂𝐻−]
𝑝𝐻 + 𝑝𝑂𝐻 = 14.00
The sum of pH and pOH is always equal to 14.00 at 25°C
Other p Scales
The other p scale is pKa; which is just another way to
quantify weak acid strength
The smaller pKa is, the stronger the acid
𝑝𝐾𝑎 = −log(𝐾𝑎)
Homework Problems: #16, 18, 20, 22
15.6: Finding [H+] and pH of Strong and
Weak Acids
In every solution of acid, we have two sources of H+:
𝐻𝐴 𝑎𝑞 + 𝐻2𝑂 𝑙 ↔ 𝐻3𝑂+ 𝑎𝑞 + 𝐴− 𝑎𝑞
𝐻2𝑂 𝑙 + 𝐻2𝑂 𝑙 ↔ 𝐻3𝑂+ 𝑎𝑞 + 𝑂𝐻−(𝑎𝑞)
Except in VERY dilute acid solutions, the ionization of
water is negligible
AKA doesn’t matter
So the only contribution of [H+] is from the acid itself
Strong and Weak Acids
Because a strong acid is completely dissociated, we can
say that the concentration of the strong acid is equal to
the concentration of H+
For example:
0.10𝑀𝐻𝐶𝑙 ⇒ 𝐻+ = 0.10𝑀 ⇒ 𝑝𝐻 = − log 0.10 = 1.00
However, we cannot make the same assumption for weak
acids
If we measured the pH of two solutions:
0.10M HCl pH = 1.00
0.10M HC2H3O2 pH = 2.87
We see that the weak acid does not fully dissociate (smaller K)
Weak Acids
To find the concentration of [H+], we must go back to
chapter 14 (equilibrium):
𝐻𝐴 𝑎𝑞 + 𝐻2𝑂 𝑙 ↔ 𝐻3𝑂+ 𝑎𝑞 + 𝐴− 𝑎𝑞
For the reaction above, we can summarize the
concentrations at equilibrium:
𝐾𝑎 =𝐻3𝑂
+ [𝐴−]
[𝐻𝐴]=
𝑥2
0.10 − 𝑥
[HA] [H3O+] [A-]
I 0.10 ≈ 0 0
C -x +x +x
E 0.10-x x x
Weak Acid Example 1
Find the concentration of H3O+ of a 0.250M HF solution
(Ka = 3.5 x 10-4)
𝐻𝐹(𝑎𝑞) + 𝐻2𝑂 𝑙 ↔ 𝐻3𝑂+ 𝑎𝑞 + 𝐹−(𝑎𝑞)
𝐾𝑎 =𝐻3𝑂
+ [𝐹−]
[𝐻𝐹]=
𝑥2
0.250 − 𝑥= 3.5 × 10−4
[HF] [H3O+] [F-]
I 0.250 ≈ 0 0
C -x +x +x
E 0.250 - x x x
Weak Acid Example 2
Find the pH of a 0.0150M HC2H3O2 solution (Ka = 1.8 x
10-5)
𝐻𝐶2𝐻3𝑂2(𝑎𝑞) + 𝐻2𝑂 𝑙 ↔ 𝐻3𝑂+ 𝑎𝑞 + 𝐶2𝐻3𝑂2
−(𝑎𝑞)
𝐾𝑎 =𝐻3𝑂
+ [𝐶2𝐻3𝑂2−]
[𝐻𝐶2𝐻3𝑂2]=
𝑥2
0.0150 − 𝑥= 1.8 × 10−5
[𝐇𝑪𝟐𝑯𝟑𝑶𝟐] [H3O+] [𝑪𝟐𝑯𝟑𝑶𝟐
-]
I 0.0150 ≈ 0 0
C -x +x +x
E 0.0150 - x x x
Weak Acid Example 3
Find the pH of a 0.010M HNO2 solution. (Ka = 4.6 x 10-4)
𝐻𝑁𝑂2(𝑎𝑞) + 𝐻2𝑂 𝑙 ↔ 𝐻3𝑂+ 𝑎𝑞 + 𝑁𝑂2
−(𝑎𝑞)
𝐾𝑎 =𝐻3𝑂
+ [𝑁𝑂2−]
[𝐻𝑁𝑂2]=
𝑥2
0.010 − 𝑥= 4.6 × 10−4
[𝐇𝑵𝑶𝟐] [H3O+] [𝑵𝑶𝟐
-]
I 0.010 ≈ 0 0
C -x +x +x
E 0.010 - x x x
Weak Acid Example 4
A 0.175M weak acid solution has a pH of 3.25. Find Ka for
the acid.
𝑝𝐻 = − log 𝐻3𝑂+ = 3.25 ⇒ 𝐻3𝑂
+ = 5.62 × 10−4𝑀
𝐾𝑎 =𝑥2
0.175 − 𝑥=
(5.62 × 10−4)2
0.175 − 5.62 × 10−4= 1.8 × 10−6
[HA] [H3O+] [A-]
I 0.175 ≈ 0 0
C -x +x +x
E 0.175 - x x x
Polyprotic Acids
Remember that polyprotic acids are acids with more than
one ionizable protons
In general a polyprotic acid ionizes in successive steps,
each with its own Ka; for example:
𝐻2𝑆𝑂4 𝑎𝑞 ↔ 𝐻+ 𝑎𝑞 + 𝐻𝑆𝑂4− 𝑎𝑞 𝐾𝑎1 = 1.6 × 10−2
𝐻𝑆𝑂4− 𝑎𝑞 ↔ 𝐻+ 𝑎𝑞 + 𝑆𝑂4
2− 𝑎𝑞 𝐾𝑎2 = 6.4 × 10−8
Notice that Ka2 is smaller than Ka1
This is because it is harder to remove a proton from a charged
molecule
Polyprotic Acids
Homework Problems: #24, 26, 28, 30, 32, 34, 36, 38, 42, 44,
46
15.7: Base Solutions
Just like with strong acids, strong bases also completely
dissociates in solution
There are only 6 strong bases:
As you may notice, these are group 1A and 2A metal
hydroxides
Table 15.8 Strong Bases
Lithium hydroxide LiOH Strontium hydroxide Sr(OH)2
Sodium hydroxide NaOH Calcium hydroxide Ca(OH)2
Potassium hydroxide KOH Barium hydroxide Ba(OH)2
Base Solutions
Unlike diprotic acids, bases with more than one OH-
dissociate in one step:
𝑆𝑟(𝑂𝐻)2 𝑎𝑞 → 𝑆𝑟2+ 𝑎𝑞 + 2𝑂𝐻−(𝑎𝑞)
Weak Bases
We can write a weak base dissociation just like a weak
acid dissociation:
𝐵 𝑎𝑞 + 𝐻2𝑂 𝑙 ↔ 𝐵𝐻+ 𝑎𝑞 + 𝑂𝐻−(𝑎𝑞)
𝐾𝑏 =𝐵𝐻+ [𝑂𝐻−]
[𝐵]
The p scale also applies here:
𝑝𝐾𝑏 = −log𝐾𝑏
See table 15.9 (page 575) for common Kb values
pH From Strong Base Example
Find the [OH-] and pH of a 0.010M Ba(OH)2 solution.
𝐵𝑎 𝑂𝐻 2 𝑎𝑞 → 𝐵𝑎2+ 𝑎𝑞 + 2𝑂𝐻−(𝑎𝑞)
𝑂𝐻− = 2 0.010 = 0.020𝑀
𝑝𝑂𝐻 = − log 0.020 = 1.70
𝑝𝐻 = 14.00 − 𝑝𝑂𝐻 = 12.30
pH of a Weak Base Example
Find the [OH-] and pH of a 0.33M methylamine solution.
(Kb = 4.4 x 10-4)
𝐶𝐻3𝑁𝐻2 𝑎𝑞 + 𝐻2𝑂 𝑙 ↔ 𝐶𝐻3𝑁𝐻3+ 𝑎𝑞 + 𝑂𝐻−(𝑎𝑞)
𝐾𝑏 =𝐶𝐻3𝑁𝐻3
+ [𝑂𝐻−]
[𝐶𝐻3𝑁𝐻2]=
𝑥2
.33 − 𝑥= 4.4 × 10−4
[𝑪𝑯𝟑𝑵𝑯𝟐] [𝑪𝑯𝟑𝑵𝑯𝟑+] [OH-]
I 0.33 ≈ 0 0
C -x +x +x
E 0.33 - x x x
Homework Problems: #48, 52, 54, 58
15.8: Acid-Base Properties of Salts
We know that ions can act as an acid or base
For example:
𝐻𝐶𝑂3− 𝑎𝑞 + 𝐻2𝑂 𝑙 ↔ 𝐻2𝐶𝑂3 𝑎𝑞 + 𝑂𝐻−(𝑎𝑞)
We also know that ions are not stable by themselves
This is why we would see the bicarbonate ion bonded to
sodium in a salt:
𝑁𝑎𝐻𝐶𝑂3 𝑠 ↔ 𝑁𝑎+ 𝑎𝑞 + 𝐻𝐶𝑂3−(𝑎𝑞)
Acid-Base Properties of Salts
We can think of any anion as the conjugate base of an
acid:
And any cation as the conjugate acid of a base:
This anion is the conjugate base of this acid
Cl– HCl
F– HF
NO3– HNO3
C2H3O2– HC2H3O2
This cation is the conjugate acid of this base
NH4+ NH3
C2H5NH3+ C2H5NH2
CH3NH3+ CH3NH2
Acid-Base Properties of Salts Example
Classify the following as a weak base, weak acid, or pH-
neutral:
1. CHO2 –
2. CH3NH3+
3. F –
4. Na+
Weak base
Weak acid
Weak base
pH-neutral
Acid-Base Properties of Salts
Every anion may not act as a base – it depends on the
strength of the corresponding acid
An anion that is the cb of a weak acid is itself a weak base
An anion that is the cb of a strong acid is pH-neutral
Every cation may not act as a acid – it depends on the
strength of the corresponding base
A cation that is the ca of a weak base is itself a weak acid
A cation that is the ca of a strong base is pH-neutral
The stronger the starting compd, the weaker the conj. is
Acid-Base Properties of Salts
There are 4 rules to follow when determining if a salt is
acidic, basic, or neutral:
1. Salts in which neither the cation nor the anion acts as an acid
or base form pH-neutral solutions
2. Salts in which the cation does not act as an acid and the
anion acts as a base form basic solutions
3. Salts in which the cation acts as an acid and the anion does
not acts as a base form acidic solutions
4. Salts in which the cation acts as an acid and the anion acts as
a base form solutions in which the pH depends on the
relative strengths of the acid and the base
1. Salts with pH-neutral solutions
Some examples include:
NaCl
Na+ is neutral
Cl – is from a strong acid, so it is neutral
Ca(NO3)2
Ca2+ is neutral
NO3 – is from a strong acid, so it is neutral
KBr
K+ is neutral
Br – is from a strong acid, so it is neutral
2. Salts that form basic solutions
Some examples include:
NaF
Na+ is neutral
F – is from a weak acid, so it is basic
Ca(C2H3O2)2
Ca2+ is neutral
C2H3O2 – is from a weak acid, so it is basic
KNO2
K+ is neutral
NO2 – is from a weak acid, so it is basic
3. Salts that form acidic solutions
Some examples include:
FeCl3
Fe3+ is from a weak base, so it is acidic
Cl – is from a strong acid, so it is neutral
Al(NO3)3
Al3+ is from a weak base, so it is acidic
NO3 – is from a strong acid, so it is neutral
NH4Br
NH4+ is from a weak base, so it is acidic
Br – is from a strong acid, so it is neutral
*Small highly charged metal cations (like Al3+ and Fe3+) form weakly acidic
solutions
4. Salts in which it depends on strength
Some examples include:
FeF3
Fe3+ is weakly acidic
F – is weakly basic
Al(C2H3O2)3
Al3+ is weakly acidic
C2H3O2 – is weakly basic
NH4NO2
NH4+ is weakly acidic
NO2 – is weakly basic
So are these acidic or basic?????
Strength Dependent
For those salts, we need to compare the Ka of the acid to
the Kb of the base
The ion with the higher K value dominates and determines if
the solution is acidic or basic
Salts: Acidic, Basic, or Neutral Example
Determine whether the solutions formed by each of the
following will be acidic, basic, or neutral:
NaHCO3
Basic
CH3CH2NH3Cl
Acidic
KNO3
Neutral
Fe(NO3)3
Acidic
NH4NO2
Can’t tell….
Homework Problems: #60, 62, 64, 65, 66, 68, 70
Section 9 is beyond the scope of this course
15.10: Lewis Acids and Bases
Under the Brønsted-Lowry definition of acids and bases,
ammonia (NH3) is a base because it accepts a proton
There is a better definition of acids and bases, called
Lewis acids and bases
A Lewis acid is an electron pair acceptor
A Lewis base is an electron pair donator
Lewis Acids and Bases
This new definition is why dry ice, when placed in water,
will acidify water:
A Lewis acid has an empty orbital (or can rearrange
electrons to create an empty orbital) that can accept an
electron pair
Homework Problems: #80, 82
Review Problems: #88, 94(a-c), 99
http://www.youtube.com/watch?v=ANi709MYnWg&index
=9&list=PL8dPuuaLjXtPHzzYuWy6fYEaX9mQQ8oGr
http://www.youtube.com/watch?v=LS67vS10O5Y&index=
31&list=PL8dPuuaLjXtPHzzYuWy6fYEaX9mQQ8oGr