A id Bases
Acids,
B
andd B
Buffers
ff
Chapter 9
1
The Acids and Bases Around Us
„ Classify each of the following as acid or base:
„ Windex
Wi d
„ Oven cleaner
„ Vitamin
Vit i C
„ Vinegar
„ Baking soda
„ Milk of magnesia
Acid and Base
„
„
„
„
„
„
„
„
„
„
„
„
Acids –
Can donate an H+
Produces H3 O+ when dissolved in water
Sour taste
Will dissolve some metals
Turn litmus paper red
Base –
B
Accepts the H+
Bitter taste
P d
Produces
OH- when
h di
dissolved
l d iin water
t
Turns litmus paper blue
3
Ionization of an Acid or Base in Water
„ Predict the product:
p
„ HNO2
„ CH3NH2
+
H2O Æ
base
+ H2O Æ
acid
„ Weak bases or acids partially ionize in water and
reach equilibrium.
q
„ Strong acids ionize 100 % in water.
4
Strong Acids that Ionize 100%
Hydrochloric
y
Acid
HCl
Hydrobromic Acid
HBr
Hydroiodic Acid
HI
Sulfuric Acid
H2SO4
Nitric Acid
HNO3
Perchloric Acid
HClO4
Chloric
C
o c Acid
cd
HClO
C O3
5
Strong Soluble Bases dissociate 100 %
Lithium hydroxide
y
LiOH
Sodium hydroxide
NaOH
Potassium hydroxide
KOH
Rubidium hydroxide
y
RbOH
Calcium hydroxide
Ca(OH)2
Strontium hydroxide
Sr(OH)2
Barium
a u hydroxide
yd o de
Ba(OH)
a(O )2
6
Acid and Base
„ H2CO3 + H2O ÆH3O+ + HCO
CO3-1
„ acid
base
acid
base
„ Conjugate acid-base pair - Chemical species
whose formulas differ only by one proton. They
have an important role in physiological equilibria.
Conjugate Acid
Acid-Base
Base Pair
Acid
Base
CH3COOH
NO2-1
HSO3HSO3HPO4-2
Amphoterism
„ Substance that can behave as an acid (donates H+) or
a base (accepts H+)
„ Which of the following substances are amphoteric?
„ a. H3 PO4
„ B. HCO3„ C. CrO4-2
„
d. CH3NH2
Acid Base Reactions
„ Identify each substance as acid or base and
relate to their conjugate pair:
HSO4- + HPO42- ÅÆ
H2SO4 + PO43-
10
Autoionization of Water
• Water can act as an acid and a base.
base
•
H2O +
acid
H2O
base
•A very small number of water molecules react to form a
hydronium ion and a hydroxide ion.
ion
11
Weak Acids and Bases
„ These substances do not ionize 100 % because the
conjugate pair is stronger than water and drives the
reaction back to the product side.
„ Predict
P di t th
the products:
d t
„
H2 CO3 + H2O
„ The conjugate reacts to form back the carbonic acid.
„ An equilibrium
q
is established.
12
Substances Present at the end
before
end
before
end
More acid is present at the
end due to the formation
of an equilibrium.
13
Chemical Equilibrium
„ Dynamic process at which the rate of forward
and reverse reaction are the same.
„
H2O (l)
H2O (g)
„ Applies to many reactions.
14
Equilibrium Constant (K)
„
„
aA + bB Æ cC + dD
K=
[C]c [D]d
[A]a [B]b
„ [ A] = molarity = mol/L
„ Equilibrium constant depends on temperature.
15
Equilibrium Constant
„ Solvents, solids, and liquids do not affect the
equilibrium constant.
„ Magnitude of K is related to stability of
products or reactants.
p
16
Equilibrium Constant
K = [products]p
[reactants]r
„ K>>103
[Products]>>[Reactants]
„ Products more stable than reactants
„ K<<10-3
„ [Products]<<[Reactants]
[
] [
]
„ Reactants more stable than products
„ K~1 (Intermediate)
„ [Products] ~ [Reactants]
„
17
Le Chatelier’s
Chatelier s Principle
„ If an external stress is applied to a system at equilibrium, the system
adjusts in such a way that the stress is partially relief.
„ Stress is related to changes in:
„ Concentration
„ Pressure – applies
pp
to g
gases only
y
„ Volume
„ Temperature
„ Competing Reactions
18
Le Chatelier’s Principle
„ Concentration Changes:
„
The concentration stress of an added reactant or product is
relieved by reaction in the direction that consumes the added
substance.
b t
„
The concentration stress of a removed reactant or product is
relieved
li
db
by reaction
i iin the
h di
direction
i that
h replenishes
l i h (makes
( k
more) the removed substance.
19
Le Chatelier’s Principle
„ For the follo
following
ing reaction:
reaction
„
Fe2O3(s) + 3 CO(g) <-> 2 Fe(l) + 3 CO2(g)
„ Use Le Châtelier’s principle to predict the direction of reaction
when an equilibrium mixture is disturbed by:
(a) Adding Fe2O3 (b) Removing CO2
(c) Removing CO
20
Le Chatelier’s
Chatelier s Principle
„ Volume and Pressure Changes: Only reactions containing gases are
affected by changes in volume and pressure.
„ PV = nRT
or
„
P = n RT
V
Increasing pressure = Decreasing volume
„ PV = nRT tells us that increasing pressure or decreasing volume
increases concentration.
„
Pressure
„ Volume
Concentration
Concentration
21
Le Chatelier’s Principle
„ Temperature Changes: Changes in temperature can change
the equilibrium constant.
„ Consider
C
heat as a product ((exothermic)) or a reactant
(endothermic) in the equilibrium equation.
„
aA
A <->
< > bB
∆ H = 234 kJ/mol
kJ/ l
„ What will an increase in temp. do to the above equilibrium?
„
aA <-> bB
∆ H = -234 kJ/mol
„ Le
L Ch
Chatelier’s
t li ’ principle
i i l iis IImportant
t t iin th
the ffunction
ti off b
buffers.
ff
22
Kw: Ionization Constant for Water
„ What is the equilibrium
q
expression
p
for water?
H2O + H2O <--> H3O+ (aq) + OH- (aq)
„
14
Kw = [H3O+ ][OH-] = 1.0 x 10-14
„ [ ] represents concentration in molarity (M)
„ What
What’ss the hydroxide concentration when [H3O+] = 0.054
0 054 M?
23
[H+ ] and [OH-]
„ Calculate the hydroxide concentration when
„ a. [H3O+] = 0.050 M
„ b. [H3O+] = 2.56 x 10-6 M
„ c. [[H3O+] = 7.4 x 10-11 M
„
As [H+] ↓, the [OH-] ↑
24
The pH Scale (0 - 14)
pH = -log
log [H3O+ ]
pH)
or [H3O+ ] = 10 ((-pH)
Calculate the pH of the previous solutions.
When th
Wh
the solution
l ti is
i neutral,
t l [H3O+ ] = [OH-] = 10-77 M and
d th
the
pH is 7.
Acidic solutions have pH < 7 and basic pH > 7
As [H+ ] ↑ pH ↓
25
Practice Problems
„ What is the hydrogen ion concentration in a
solution with pH 8.3? Is this solution acidic or
basic?
„ What is the pH of a solution that has a [OH-] =
3.4x10-4?
„
pH + pOH = 14
26
Physiological pH
„ Between 7.35 – 7.45
„ Affected by stress placed in the body such as kidney
disease, stroke, respiratory failure
„ Disruption of cell functions could occur
27
Buffers
„ Buffers are solutions that resist changes in pH
when an acid or base is added to them.
„ The reaction cannot g
go 100 % to p
products.
„ Weak acids or weak bases could be used under
certain conditions.
CH3COOH + H2O < -->
acid
base
28
Buffers
„ Equilibrium
q
((Two way
y road)) is required
q
for a buffer.
100 %
„
„ HCl + H2O
„ CH3COOH + H2O
H3O+
+ ClH3O+
+ CH3COO-
„ A buffer will contain large
g q
quantities of both the CH3COOH
and its conjugate base CH3COO- to be able to control pH
changes.
29
Buffers
„ Which of the following mixtures could be used
as buffers?
„ H2SO4/HSO4-1
„ HPO4-2/H3PO4
„ NH3/NH4+
30
Buffers
H3O+
„ CH3COOH + H2O
+ CH3COO-
Weak Acid
+ H+
+ OHWeak Base
31
Strong acids have weak
conjugate bases
bases.
Weak acids have conjugate
bases that are stronger than
water.
32
Strength
g of acids
„ The conjugate bases of weak acids
„
HA + H2O
H3O+ + A-1
„ The conjugate base of weak acids are stronger than water
and can grab a hydrogen to establish an equilibrium.
„ What’s the equilibrium expression for the above?
„ Ka = equilibrium for weak acids
33
Ionization Constant (Ka)
„ HA + H2O
„
Ka =
H3O+ + A-1
[H3O+] [A-]
[HA]
„ What happens when [HA] = [A-]?
„
Ka = [H3O+]
„
pH = p
p
pKa
or
Important to indicate
region in which a buffer
works
34
Buffers
Buffers are most resistant to pH
changes
g when the pH
p = pK
p a of the weak
acid and are effective when the pH is
within one unit of the pKa (pH = pKa ± 1).
35
36
„ The larger the Ka for an acid, the stronger an acid it is.
„ The lower the pKa (pKa = -logKa), the stronger the acid.
„
An acid solution has a pKa = 8.45. What is the Ka of
the acid?
37
Buffers in the Blood
„ The pH of your blood normally ranges
between 7.35 and 7.45.
„ A blood pH below the normal range is called
acidosis, while a blood pH above this range
i called
is
ll d alkalosis,
lk l i either
ith one off which
hi h iis
potentially fatal.
38
Buffers in the Blood
• Blood is kept in this narrow range
(pH = 7.35 – 7.45) with the help of buffers.
• The most important buffer system in blood is formed
from carbonic acid (H2CO3) and its conjugate base
base,
hydrogen carbonate ion (HCO3-):
CO2 (g) + H2O
H2CO3 + H2O
Blood Buffer
HCO3- + H3O+
39