Transcript Chapter 15 - Chemistry

Chapter 15
Acids and Bases
.
Acid-Base Theories
In defining what is considered to be an acid and
what is considered to be a base, three theories
have been proposed:
Arrhenius acid-base theory
Brønsted-Lowry acid-base theory
Lewis acid-base theory
We will see that each subsequent
theory builds upon what was stated
in the previous theory.
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The Arrhenius Theory
In the Arrhenius theory of acids, an
acid dissolved in water increases
the concentration of hydronium ions
H3O+ in the solution:
Arrhenius acid HA:
HA (aq) + H2O (l)  H3O+ (aq) + A- (aq)
In this reaction we see all Arrhenius acids
contain protons (H+) that are donated to
water
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Arrhenius Theory – Acid Strength
In the Arrhenius theory of
acids, a strong acid
COMPLETELY reacts with
water, so there is no HA left at
the end of the reaction:
HA (aq) + H2O (l)
→H O
3
+
(aq) + A- (aq)
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Arrhenius Theory – Acid Strength
In the Arrhenius theory of acids, a
weak acid reacts with water until
an equilibrium is reached where
HA is still present in the
equilibrium mixture:
HA (aq) + H2O (l)
⇌H O
3
+
(aq) + A- (aq)
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The Arrhenius Theory
In the Arrhenius theory of bases, a
base dissolved in water increases
the concentration of hydroxide ions
OH- in the solution:
Arrhenius base M(OH)x:
M(OH)x (aq)
M
x+
(aq) + x OH- (aq)
In this reaction we see all Arrhenius bases
contain hydroxide (OH-)
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Arrhenius Theory – Base Strength
In the Arrhenius theory of
acids, a strong base
COMPLETELY dissociates in
water, so there is no M(OH)x left
at the end of the reaction:
M(OH)x (aq)
→M
x+
(aq) + x OH- (aq)
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Arrhenius Theory – Acid Strength
In the Arrhenius theory of bases,
a weak base only partially
dissociates in water until an
equilibrium is reached where
M(OH)x is still present in the
equilibrium mixture:
M(OH)x (aq)
M
x+
(aq) + x OH- (aq)
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Common strong acids and bases
.
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Why do we need to improve on
Arrhenius theory?
The Arrhenius theory has a
drawback!
Certain compounds that DO NOT
contain hydroxide can still increase
the hydroxide concentration when
placed in water.
Arrhenius theory does not explain
this!
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The Brønsted-Lowry Theory
The Brønsted-Lowry Theory: an acid
Brønsted-Lowry Theory: an acid is
any substance that donates protons
(H+) while a base is any substance that
can accept protons.
This means that Brønsted-Lowry
acid-base reactions are proton
transfer reactions.
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Proton transfer reactions
Pairs of compounds are related to
each other through Brønsted-Lowry
acid-base reactions. These are
conjugate acid-base pairs.
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Proton transfer reactions
Generally, an acid HA has a conjugate
base A- (an H+ has transferred away from
the acid). Conversely, a base B has a
conjugate acid BH+ (an H+ has transferred
toward the base).
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Water in BL acid-base reactions
When a Brønsted-Lowry acid is placed in
water, it donates a proton to the water
(which acts as a base) and establishes an
acid-base equilibrium.
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Water in BL acid-base reactions
In the reverse reaction of the
equilibrium, the acid H3O+ donates a
proton to the base A- to give back
water and HA.
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Water in BL acid-base reactions
When a Brønsted-Lowry base is placed in
water, it accepts a proton from water
(which acts as an acid) and establishes an
acid-base equilibrium.
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Water in BL acid-base reactions
In the reverse reaction of the
equilibrium, the acid BH+ donates a
proton to the base OH- to give back
water and B.
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Brønsted-Lowry Bases
To accept a proton (to act as a B-L base)
requires a molecule to have an unshared
pair of electrons which can then be used
to create a bond to the H+.
All Brønsted-Lowry bases have
at least one lone pair of
electrons.
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Brønsted-Lowry Bases
In the previous reactions
we’ve seen NH3 has a lone
pair of electrons and can
act as a B-L base.
Also, water has two lone
pairs of electrons, and can
act as a B-L base.
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Amphiprotic substances
Some substances, like water, have protons that
can be donated (BL acid), and lone pairs of
electrons that can accept protons (BL base).
This is why it can act like an acid OR a base
DEPENDING on the other species present.
Such substances are said to be
amphiprotic.
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Problem
Write a balanced equation for the
dissociation of each of the following
Brønsted-Lowry acids in water:
a) H2SO4
b) HSO4c) H3O+
d) NH4+
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Problem
What is the conjugate acid of each of
the following Brønsted- Lowry bases?
a) HCO3b) CO32c) OHd) H2PO422
Why do we need to improve on
Brønsted-Lowry theory?
There are many reactions
that behave VERY MUCH
LIKE proton transfer
reactions that DO NOT
involve protons!
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Lewis Acids and Bases
A Lewis acid is an electron pair
acceptor, while a Lewis base is
an electron pair donor.
These definitions are more general than the
Brønsted-Lowry definitions because
protons DO NOT need to be
involved in Lewis acid-base
reactions.
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In general LA + :LB  LA-LB
BF3 is a Lewis acid where the B atom can
accept an electron pair. The N of the NH3
has a lone pair that can be donated,
making NH3 a Lewis base.
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Metal ions as Lewis acids
Many metal ions have
the ability to act as Lewis
acids. The ions are
willing to accept electron
pairs from LIGANDS
(which act as Lewis
bases) because this often
stabilizes the ion in
solution. The result is
often called a complex
ion.
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Problem
.
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Comparing theories
Since Arrhenius acids must
contain protons, then ALL
Arrhenius acids ARE ALSO
Brønsted-Lowry acids.
We’ve already seen that NOT ALL
Brønsted-Lowry bases are
Arrhenius bases.
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Comparing theories
There are Lewis acids (like metal ions)
that ARE NOT
Brønsted-Lowry acids.
ALL Brønsted-Lowry bases must all have at
least one lone pair of electrons, so
ALL Brønsted-Lowry bases
MUST ALSO BE
Lewis bases
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BL acid and base strength
Brønsted-Lowry acid-base
equilibria are competitions!
The equilibrium is the result of
a tug-of war between the two
bases in the system as they
fight for protons given away
by the two acids.
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Acid Strength and Base Strength
The acid that is “better at donating
protons”
OR
the base that is “better at accepting
protons”
will be found
in lesser amounts
at equilibrium
compared to the other acid (or base).
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Strong BL acids in water
A strong acid (HA) is one that
almost completely dissociates in
water (which acts as a base).
The conjugate base A- will be a
very weak base.


HA  H 2 O  A  H 3 O

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Strong BL acids in water


HA  H 2 O  A  H 3 O

At equilibrium, there will be very little to
no HA present in the system, and the
concentration of A- will essentially be the
same as the initial concentration of HA.
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Weak BL acids in water
A weak acid (HA) is one that
partially dissociates in water
(which acts as a base). The acid
is not as good at donating protons
to the water.
The conjugate base (A-) will be a
weak base. Overall
HA  H 2 O



A  H 3O

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Weak BL acids in water
HA  H 2 O



A  H 3O

At equilibrium, there will be some Aand H3O+ present in the system.
However, the concentration of HA will
still be significant at equilibrium.
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.
36
.
Notice that the
strongest
acids have the
weakest
conjugate
bases, and the
strongest
bases have the
weakest
conjugate
acids!
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Hydrated Protons and Hydronium Ions
The ultimate proton-donor is a
proton itself!
In water there is no such thing as H+.

H  H 2O 
 H 3 O

Often more than one water molecule will crowd around
the proton to give hydrates with the formula H(H2O)n+
where n is 1 to 4.
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Hydrated Protons and Hydronium Ions
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Dissociation of Water
It is possible for one water molecule to act as an acid
while another water molecule acts as a base at the same
time. This leads to the self-ionization of water
equilibrium:
H2O (l) + H2O (l)  H3O+ (aq) + OH- (aq)
The equilibrium constant for this reaction is called the
ion-product constant for water, Kw.
Kw = [H3O+][OH-]
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At 25 °C, Kw = 1.0 x 10-14
so [H3O+] = [OH-] = 1.0 x 10-7 mol/L
Relatively few water molecules are
dissociated at equilibrium at room
temperature!
We will always assume that
[H3O+] [OH-] = 1.0 x 10-14 at 25 °C.
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At 25 C
Acidic
[H3O+] > 1.0 x 10-7 M
or [OH-] < 1.0 x 10-7 M
Basic
[OH-] > 1.0 x 10-7 M
or [H3O+] < 1.0 x 10-7 M
Neutral
[H3O+] = [OH-] = 1.0 x 10-7 M
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At 25 C
We also find, since
[H3O+] [OH-] = 1.0 x 10-14 = Kw
then
[H3O+] = 1.0 x 10-14 / [OH-]
and
[OH-] = 1.0 x 10-14 / [H3O+]
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.
At 25 C
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Problem
OH
The concentration of
in a
sample of seawater is
-6
5.0 x 10 mol/L.
Calculate the concentration of
+
H3O ions, and classify the
solution as acidic, neutral, or
basic.
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Problem
At 50 °C the value of
-14
Kw is 5.5 x 10 .
What are the
+
[H3O ] and [OH ]
in a neutral solution
at 50 °C?
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The pH Scale
[H3O+] in water can range from
very small (strongly basic)
to very large (strongly acidic)
it is sometimes easier to use a
negative logarithmic (power of 10)
scale to express [H3O+] with a term we
call the pH of a solution.
pH = - log [H3O+]
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pH and acidity
At 25 C
pH > 7 is basic
pH < 7 is acidic
pH = 7 is neutral
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pOH and acidity
At 25 C
pOH = - log [OH-]
Or [OH-] = 10-pOH
pOH < 7 is basic
pOH > 7 is acidic
pOH = 7 is neutral
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pH scale
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.
Kw = 1.0 x 10-14 = [H3O+] [OH-]
pKw = - log (1.0 x 10-14) = 14.00
(2 sigfigs! The 14 is not significant!)
pKw = - log ([H3O+] [OH-])
= (- log [H3O+]) + (- log [OH-])
= pH + pOH
so 14.00 = pH + pOH (at 25 C)!
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Problem
Calculate the pH of each of the following
solutions:
a) A sample of seawater that has an
OH- concentration of 1.58 x 10-6 mol/L
b) A sample of acid rain that has an
H3O+ concentration of 6.0 x 10-5 mol/L
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Problem
Calculate [H3O+] and [OH-] in each of
the following solutions:
a) Human blood (pH 7.40)
b) A cola beverage (pH 2.8)
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Measuring pH
We often measure the pH of a solution with
a chemical acid-base indicator.
Indicators are B-L acids (symbolized HIn)
where the acid form has a different
colour than the conjugate base form (In-)
HIn (aq) + H2O (l)  H3O+ (aq) + In- (aq)
colour A
colour B
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.
Indicators
tend to
change
colour only
in small pH
ranges of
about
2 units.
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Measuring pH
To make a universal indicator
that covers the pH range from
about 1 to 12,
a mixture of several different
indicators with different pH
ranges is used.
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.
Methyl violet
Phenolphthalein
Bromothymol blue
Bromocresol green
Universal indicator
Methyl orange
pH
1
2
3
4
5
6
7
8
9
10 11
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Universal indicator
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