Acid bases and salts.part 2

Chapter #14
Acids, Bases, and Salts

Acids, Bases and Salts Topics
The Arrhenius Theory
The Brønsted Theory
Naming Acids (See Nomenclature Notes)
The Self-Ionization of Water
The pH Concept
Properties of Acids
Properties of Bases
Salts
The Strengths of Acids and Bases
Analyzing Acids and Bases
Titration Calculations
Hydrolysis Reactions of Salts
Buffers

History of Acids and Bases
In the early days of chemistry chemists were organizing
physical and chemical properties of substances. They
discovered that many substances could be placed in two
different property categories:
Substance A
1. Sour taste
2. Reacts with carbonates to make CO2
3. Reacts with metals to produce H2
4. Turns blue litmus pink
5. Reacts with B substances to make
salt water
Substance B
1. Bitter taste
2. Reacts with fats to make soaps
3. Do not react with metals
4. Turns red litmus blue
5. Reacts with A substances make
salt and water
Arrhenius was the first person to suggest a reason why
substances are in A or B due to their ionization in water.

The Swedish chemist Svante Arrhenius proposed the first
definition of acids and bases.
(Substances A and B became
known as acids and bases)
According to the Arrhenius model:
“acids are substances that dissociate in water to
produce H+
ions and bases are substances that
dissociate in water to produce OH-
ions”
NaOH (aq)  Na+
(aq) + OH-
(aq) Base
HCl (aq)  H+
(aq) + Cl-
(aq) Acid
Arrhenius Theory

What is H+
?
+
e-
+
Hydrogen (H) Proton (H+
)

Unknown to Arrhenius free H+
ions do not exist in water. They
covalently react with water to produce hydronium ions, H3O+
.
or:
H+
(aq) + H2O (l)  H3O+
(aq)
This new bond is called a coordinate covalent bond since
both new bonding electrons come from the same atom
Hydronium Ion

Hydronium ion is the name for H3O+
and is often times
abbreviated as H+
(aq) they both mean the same thing.
What is the difference between a strong acid and a weak
acid?
Hydronium Ion

and is often times
abbreviated as H+
acid? Strong acids ionize 100% and weak ones do not!
Hydronium Ion

and is often times
abbreviated as H+
A single arrow is used to represent the ionization of a strong
acid. Double arrows (Equilibrium) are used to represent
weak acids.
For example: HCl (g) H+
(aq) + Cl -
(aq)
HF (g) H+
(aq) + F -
Hydronium Ion

and is often times
abbreviated as H+
weak acids.
(aq) + Cl -
(aq)
HF (g) H+
(aq) + F -
According to Arrhenius, is water an acid or base?
HOH (l) H+
(aq) + OH –
(aq)
Hydronium Ion

and is often times
abbreviated as H+
weak acids.
(aq) + Cl -
(aq)
HF (g) H+
(aq) + F -
According to Arrhenius, is water an acid or base?
HOH (l) H+
(aq) + OH –
(aq)
Neither, he called it Neutral (same amount of OH-
and H+
Hydronium Ion

Strong Acids and Bases
How can we identify strong acids or bases?

Easy memorize them!

Easy, memorize them!
Memorized Strong Acids
1. HClO4
2. H2SO4
3. HI
4. HBr
5. HCl
6. HNO3
Memorized Strong Bases
Hydroxides of group 1 and 2
metals, excluding Be and Mg

Johannes Brønsted and Thomas Lowry revised
Arrhenius’s acid-base theory to include this behavior.
They defined acids and bases as follows:
“An acid is a hydrogen containing species that
donates a proton. A base is any substance that
accepts a proton”
HCl (aq) + H2O (l)  Cl-
(aq) + H3O+
(aq)
In the above example what is the Brønsted acid? What is
Bronsted Lowry
Bronsted Lowry Theory

HCl (aq) + H2O (l)  Cl -
( aq) + H3O+
(aq)
In reality, the reaction of HCl with H2O is an equilibrium
and occurs in both directions, although in this case the
equilibrium lies far to the right.
For the reverse reaction Cl -
behaves as a Brønsted
base and H3O+
behaves as a Brønsted acid.
The Cl-
is called the conjugate base of HCl. Brønsted
acids and bases always exist as conjugate acid-base
pairs.
Bronsted Lowry Theory

In pure water (no solute) water molecules behave as both an
acid and base!!
e.g.
H2O (l) + H2O (l)  H3O+
(aq) + OH-
(aq)
This is called the self-ionization (autoionizaion) of water.
Although the equilibrium lies far to the left it is very important to
take into consideration, especially for living systems.
Does anyone know how we write the equilibrium constant for
this reaction?
Autoionization of Water

The auto-ionization of water is described by the
equation:
H2O (l) + H2O (l)  H3O+
(aq) + OH-
(aq)
The equilibrium constant for this reaction is given by:
]OH][O3H[2]O2H[K
2]O2H[
]OH][O3H[
]O2H][O2H[
]OH][O3H[
K
−+=
−+
=
−+
=
Kw = K[H2O]2
= 10-14
This equilibrium lies very much
to the left i.e. mostly water. For pure water [OH-
] =
[H+
] = 1 x 10-7
M

]OH][O3H[wK2]O2H[K
2]O2H[
]OH][O3H[
]O2H][O2H[
]OH][O3H[
K
−+==
−+
=
−+
=
As [OH-
] and [H+
] are so small the [H2O] is not affected by their
formation. It is useful to define a new constant Kw such that:
Kw is called the ion product of water.
What is the value for the ion product of water?
1.00 g
ml 18.0 g
mole ml
10-3
L
= 55.5 M

]OH][O3H[wK2]O2H[K
2]O2H[
]OH][O3H[
]O2H][O2H[
]OH][O3H[
K
−+==
−+
=
−+
=
As [OH-
] and [H+
] are so small the [H2O] is not affected by their
formation. It is useful to define a new constant Kw such that:
Kw is called the ion product of water.
What is the value for the ion product of water?
[H+
][OH-
] = 10-14
1.00 g
ml 18.0 g
mole ml
10-3
L
= 55.5 M

We define an aqueous solution as being neutral when the
[H+
] = [OH-
].
We define an aqueous solution as being acidic when
[H+
] > [OH-
].
We define an aqueous solution as being basic when
[H+
] < [OH-
].
However, in each case Kw = 1 x 10-14
M2
[H+
] = 0.0000001 = 10-7
(how can this be abbreviated further?)

[H+
] = [OH-
].
[H+
] > [OH-
].
[H+
] < [OH-
].
M2
[H+
] = 0.0000001 = 10-7
By just describing the power

[H+
] = [OH-
].
[H+
] > [OH-
].
[H+
] < [OH-
].
M2
[H+
] = 0.0000001 = 10-7
By just describing the power Called the power of H, or pH.

[H+
] = [OH-
].
[H+
] > [OH-
].
[H+
] < [OH-
].
M2
[H+
] = 0.0000001 = 10-7
By just describing the power Called the power of H, or pH.
pH = 7
Our math departments tells us that log means
power too.

The mathematical definition of pH using [H+
] for [H3O+
] is
listed below:
pH = -log [H+
], or [H+
]= 1x10-pH
(both are mathematically
equivalent)
How about the power for the OH -
, what should this be
called?

The mathematical definition of pH using [H+] for [H3O+] is
listed below:
pH = -log [H+
], or [H+
] = 1x10-pH
equivalent)
called? Would you believe pOH?

listed below:
pH = -log [H+
], or [H+
]= 1x10-pH
equivalent)
Have you heard of pOH before?

listed below:
pH = -log [H+
], or [H+
]= 1x10-pH
equivalent)
Have you heard of pOH before?
pH + pOH = 14 for water solutions.

Now for some examples
1. Find the pH and pOH, when [H+
] = 10-4

Now for an example
1. Find the pH and pOH, when [H+
] = 10-4
pH = 4 and pOH = 10, since they must add to 14
using the calculator pH = -log [H+
], type in 10-4
, push
the log button and pH = -(-4) = 4. Same for pOH

A pH Number line
Number lines have been used in history and math classes,
so to keep up we use them in chemistry classes.
pH = 16
pH = 12
pH = 7
pH = 2
[H+
] = 10-2

A pH Number line
pH = 16
pH = 12
pH = 7
pH = 2
[H+
] = 10-2
[OH -
] = 10-12

A pH Number line
pH = 16
pH = 12
pH = 7
pH = 2
[H+
] = 10-2
[OH -
] = 10-12
[H+] > [OH -]
acidic

A pH Number line
pH = 16
pH = 12
pH = 7
pH = 2
[H+
] = 10-2
[OH -
] = 10-12
[H+
] = 10-7
[OH -
] = 10-7
[H+] > [OH -]
acidic

A pH Number line
pH = 16
pH = 12
pH = 7
pH = 2
[H+
] = 10-2
[OH -
] = 10-12
[H+
] = 10-7
[OH -
] = 10-7
[H+
] = [OH -
]
neutral
[H+
] > [OH -
]
acidic

A pH Number line
pH = 16
pH = 12
pH = 7
pH = 2
[H+
] = 10-2
[OH -
] = 10-12
[H+
] = 10-7
[OH -
] = 10-7
[H+
] =10-12
[OH -
] = 10-2
[H+
] = [OH -
]
neutral
[H+
] > [OH -
]
acidic

A pH Number line
pH = 16
pH = 12
pH = 7
pH = 2
[H+
] = 10-2
[OH -
] = 10-12
[H+
] = 10-7
[OH -
] = 10-7
[H+
] =10-12
[OH -
] = 10-2
[H+
] < [OH -
]
basic
[H+
] = [OH -
]
neutral
[H+
] > [OH -
]
acidic

A pH Number line
pH = 16
pH = 12
pH = 7
pH = 2
[H+
] = 10-2
[OH -
] = 10-12
[H+
] = 10-7
[OH -
] = 10-7
[H+
] =10-12
[OH -
] = 10-2
[H+
] =10-16
[OH -
] =
[H+
] < [OH -
]
basic
[H+
] = [OH -
]
neutral
[H+
] > [OH -
]
acidic

A pH Number line
pH = 16
pH = 12
pH = 7
pH = 2
[H+
] = 10-2
[OH -
] = 10-12
[H+
] = 10-7
[OH -
] = 10-7
[H+
] =10-12
[OH -
] = 10-2
[H+
] =10-16
[OH -
] = 102
[H+
] < [OH -
]
basic
[H+] = [OH -]
neutral
[H+] > [OH -]
acidic

A pH Number line
pH = 16
pH = 12
pH = 7
pH = 2
[H+
] = 10-2
[OH -
] = 10-12
[H+
] = 10-7
[OH -
] = 10-7
[H+
] =10-12
[OH -
] = 10-2
[H+
] =10-1
[OH -
] = 102
[H+
] < [OH -
]
basic
[H+
] < [OH -
]
basic
[H+] = [OH -
]
neutral
[H+] > [OH -
]
acidic

A pH Number line
pH = 16
pH = 12
pH = 7
pH = 2
[H+
] = 10-2
[OH -
] = 10-12
[H+
] = 10-7
[OH -
] = 10-7
[H+
] =10-12
[OH -
] = 10-2
[H+
] =10-16
[OH -
] = 102
[H+
] < [OH -
]
basic
[H+
] < [OH -
]
basic
[H+] = [OH -
]
neutral
[H+
] > [OH -
]
acidic
acidic

A pH Number line
pH = 16
pH = 12
pH = 7
pH = 2
[H+
] = 10-2
[OH -
] = 10-12
[H+
] = 10-7
[OH -
] = 10-7
[H+
] =10-12
[OH -
] = 10-2
[H+
] =10-16
[OH -
] = 102
[H+
] < [OH -
]
basic
[H+
] < [OH -
]
basic
[H+
] = [OH -
]
neutral
[H+
] > [OH -
]
acidic
acidic
basic

Acids undergo characteristic double replacement reactions
with oxides, hydroxides, carbonates and bicarbonates.
e.g.
2HCl (aq) + CuO (s)  CuCl2 (aq) + H2O (l)
2HCl (aq) + Ca(OH)2 (aq)  CaCl2 (aq) + 2H2O (l)
2HCl (aq) + CaCO3 (aq)  CaCl2 (aq) + H2O (l) + CO2 (g)
2HC l (aq) + Sr(HCO3)2 (aq)  SrCl2 (aq) + 2H2O (l) + 2CO2 (g)

Bases undergo a double replacement reaction with acids
called neutralization:
NaOH (aq) + HCl (aq)  H2O (l) + NaC l (aq)
In words this well known reaction is often described as:
“acid plus base = salt plus water”
We previously discussed this reaction when describing types of
reactions.

We have discussed the double replacement reactions and ionic
equations before. Since the acids and bases undergo double
replacement reactions called neutralization reactions, then they
can have ionic equations too.
e.g.
Molecular equation:
HCl (aq) + NaOH (aq)  NaCl (aq) + H2O (l)
Total ionic equation:
H+
(aq) + Cl-
(aq) + Na+
(aq) + OH-
(aq)  Na+
(aq) + Cl-
(aq) + H2O (l)
Net ionic equation:
H+
(aq) + OH-
(aq)  H2O (l)

Another property of acids is their reaction with certain metals to
produce hydrogen gas, H2 (g).
Zn (s) + 2HC l (aq)  H2 (g) + ZnCl2 (aq)
This is an example of a single replacement reaction and is a
redox reaction.
Total ionic equation:
Zn (s) + 2H+
(aq) + 2Cl-
(aq)  H2 (g) + Zn2+
(aq) + 2Cl-
(aq)
Net ionic equation:
Zn (s) + 2H+
(aq)  H2 (g) + Zn2+
(aq)

Salts
Salts are the ionic product of an acid base neutralization
reaction.
Acidic Salts are formed from a strong acid and a weak base.
Neutral salts are formed from a strong acid and strong base.
Basic salts are formed from a strong base and a weak acid.
Give the acid and base the following salts were formed from
and label the salts as acidic, basic, or neutral.
1. NaCl
2. NaC2H3O2
3. NH4Cl

Salts
reaction.
1. NaCl
1. NaC2H3O2
1. NH4Cl
NaCl + HOHReactants are?

Salts
reaction.
1. NaCl
2. NaC2H3O2
3. NH4Cl
NaCl + HOHHCl + NaOH
NaC2H3O2 + HOH
S.A. s.b.

Salts
reaction.
1. NaCl
2. NaC2H3O2
3. NH4Cl
NaC2H3O2 + HOH
S.A. s.b.Neutral salt

Salts
reaction.
1. NaCl
2. NaC2H3O2
3. NH4Cl
NaC2H3O2 + HOHHC2H3O2 + NaOH
Neutral salt s.a. s.b.

Salts
reaction.
1. NaCl
2. NaC2H3O2
3. NH4Cl
w.a. s.b.

Salts
reaction.
1. NaCl
2. NaC2H3O2
3. NH4Cl
w.a. s.b.basic salt

Salts
reaction.
1. NaCl
2. NaC2H3O2
3. NH4Cl
w.a. s.b.basic salt
NH4Cl + HOH

Salts
reaction.
1. NaCl
2. NaC2H3O2
3. NH4Cl
w.a. s.b.basic salt
NH4Cl + HOHNH4OHHCl +

Salts
reaction.
1. NaCl
2. NaC2H3O2
3. NH4Cl
w.a. s.b.basic salt
s.a. w.b.

Salts
reaction.
1. NaCl
2. NaC2H3O2
3. NH4Cl
w.a. s.b.basic salt
neutral salt s.a. s.b.
s.a. w.b.acidic salt

Acid, Base, and Salt Hydrolysis
HBr (aq)

HBr (aq) H+
(aq) + Br -
(aq)

HBr (aq) H+
(aq) + Br -
(aq) Acidic, because H+
(aq)

HBr (aq) H+
(aq) + Br -
(aq)
0.1 Initial concentration

HBr (aq) H+
(aq) + Br -
(aq)
0.1 Initial concentration0.0

HBr (aq) H+
(aq) + Br -
(aq)
0.1 Initial concentration0.0 ?

HBr (aq) H+
(aq) + Br -
(aq)
0.1 Initial concentration0.0 0.0

HBr (aq) H+
(aq) + Br -
(aq)
Final concentration?

HBr (aq) H+
(aq) + Br -
(aq)
Final concentration0.0

HBr (aq) H+
(aq) + Br -
(aq)
Final concentration0.0 ?

HBr (aq) H+
(aq) + Br -
(aq)
Final concentration0.0 0.1

HBr (aq) H+
(aq) + Br -
(aq)
Final concentration0.0 0.1 ?

HBr (aq) H+
(aq) + Br -
(aq)
Final concentration0.0 0.1 0.1

HBr (aq) H+
(aq) + Br -
(aq)
pH = ?

HBr (aq) H+
(aq) + Br -
(aq)
pH = ?
[H+
] = ?

HBr (aq) H+
(aq) + Br -
(aq)
pH = ?
[H+
] = 0.1 = 10-1

HBr (aq) H+
(aq) + Br -
(aq)
pH = 1
Ca(OH)2 (aq)

HBr (aq) H+
(aq) + Br -
(aq)
pH = 1
Ca(OH)2 (aq) Ca2+
(aq) + 2 OH-
(aq)

HBr (aq) H+
(aq) + Br -
(aq)
pH = 1
Ca(OH)2 (aq) Ca2+
(aq) + 2 OH-
(aq) acidic?

HBr (aq) H+
(aq) + Br -
(aq)
pH = 1
Ca(OH)2 (aq) Ca2+
(aq) + 2 OH-
(aq) No, basic OH-

HBr (aq) H+
(aq) + Br -
(aq)
pH = 1
Ca(OH)2 (aq) Ca2+
(aq) + 2 OH-
(aq)
Initial concentration0.1

HBr (aq) H+
(aq) + Br -
(aq)
pH = 1
Ca(OH)2 (aq) Ca2+
(aq) + 2 OH-
(aq)
Initial concentration0.1 0.0

HBr (aq) H+
(aq) + Br -
(aq)
pH = 1
Ca(OH)2 (aq) Ca2+
(aq) + 2 OH-
(aq)
Initial concentration0.1 0.0 ?

HBr (aq) H+
(aq) + Br -
(aq)
pH = 1
Ca(OH)2 (aq) Ca2+
(aq) + 2 OH-
(aq)
Initial concentration0.1 0.0 0.0

HBr (aq) H+
(aq) + Br -
(aq)
pH = 1
Ca(OH)2 (aq) Ca2+
(aq) + 2 OH-
(aq)
Final concentration?

HBr (aq) H+
(aq) + Br -
(aq)
pH = 1
Ca(OH)2 (aq) Ca2+
(aq) + 2 OH-
(aq)
Final concentration0.0 ?

HBr (aq) H+
(aq) + Br -
(aq)
pH = 1
Ca(OH)2 (aq) Ca2+
(aq) + 2 OH-
(aq)
Final concentration0.0 0.1

HBr (aq) H+
(aq) + Br -
(aq)
pH = 1
Ca(OH)2 (aq) Ca2+
(aq) + 2 OH-
(aq)
Final concentration0.0 0.1 ?

HBr (aq) H+
(aq) + Br -
(aq)
pH = 1
Ca(OH)2 (aq) Ca2+
(aq) + 2 OH-
(aq)

HBr (aq) H+
(aq) + Br -
(aq)
pH = 1
Ca(OH)2 (aq) Ca2+
(aq) + 2 OH-
(aq)
pH = ?

HBr (aq) H+
(aq) + Br -
(aq)
pH = 1
Ca(OH)2 (aq) Ca2+
(aq) + 2 OH-
(aq)
pH = ?
[OH -
] = 0.2

HBr (aq) H+
(aq) + Br -
(aq)
pH = 1
Ca(OH)2 (aq) Ca2+
(aq) + 2 OH-
(aq)
pH = ?
[OH -
] = 0.2 [H+
] = ? [H] [OH -
] = 10-14
[H+
] = ?

HBr (aq) H+
(aq) + Br -
(aq)
pH = 1
Ca(OH)2 (aq) Ca2+
(aq) + 2 OH-
(aq)
final concentration0.0 0.1 0.2
pH = - log[H+
] pH = -log[5X10-14
] = 13.30103 = 13.3
[OH -
] = 0.2 [H+
] = 5X10-14
[H] [OH -
] = 10-14
[H+
] =
10-14
0.2

HBr (aq) H+
(aq) + Br -
(aq)
pH = 1
Ca(OH)2 (aq) Ca2+
(aq) + 2 OH-
(aq)
pH = 13.3

HBr (aq) H+
(aq) + Br -
(aq)
pH = 1
Ca(OH)2 (aq) Ca2+
(aq) + 2 OH-
(aq)
pH = 13.3
NaF (aq)

HBr (aq) H+
(aq) + Br -
(aq)
pH = 1
Ca(OH)2 (aq) Ca2+
(aq) + 2 OH-
(aq)
pH = 13.3
NaF (aq) Na+
(aq) + F –
(aq)

HBr (aq) H+
(aq) + Br -
(aq)
pH = 1
Ca(OH)2 (aq) Ca2+
(aq) + 2 OH-
(aq)
pH = 13.3
NaF (aq) Na+
(aq) + F –
(aq) Acidic, basic, or neutral?

HBr (aq) H+
(aq) + Br -
(aq)
pH = 1
Ca(OH)2 (aq) Ca2+
(aq) + 2 OH-
(aq)
pH = 13.3
NaF (aq) Na+
(aq) + F –
(aq)
Basic, since HF is w.a. and
NaOH is s.b.
Will sodium and fluorine ions react with water?

HBr (aq) H+
(aq) + Br -
(aq)
pH = 1
Ca(OH)2 (aq) Ca2+
(aq) + 2 OH-
(aq)
pH = 13.3
NaF (aq) Na+
(aq) + F –
(aq)
NaOH is s.b.
Na+
+ HOH NaOH + H+

HBr (aq) H+
(aq) + Br -
(aq)
pH = 1
Ca(OH)2 (aq) Ca2+
(aq) + 2 OH-
(aq)
pH = 13.3
NaF (aq) Na+
(aq) + F –
(aq)
NaOH is s.b.
Na+
+ HOH NaOH + H+
Cannot make strong
acids or bases from weak
oness.b.

HBr (aq) H+
(aq) + Br -
(aq)
pH = 1
Ca(OH)2 (aq) Ca2+
(aq) + 2 OH-
(aq)
pH = 13.3
NaF (aq) Na+
(aq) + F –
(aq)
NaOH is s.b.
Na+
+ HOH NaOH + H+
Cannot make strong
ones
F -
+ HOH HF + OH-
w.a.

HBr (aq) H+
(aq) + Br -
(aq)
pH = 1
Ca(OH)2 (aq) Ca2+
(aq) + 2 OH-
(aq)
pH = 13.3
NaF (aq) Na+
(aq) + F –
(aq)
NaOH is s.b.
Na+
+ HOH NaOH + H+
Cannot make strong
ones
F -
+ HOH HF + OH- Yes, HF weak acid and
OH-
is formed, thus basic
salt!
w.a.

NH4Cl (aq) NH4
+
(aq) + Cl-
(aq)

NH4Cl (aq) NH4
+
(aq) + Cl-
(aq) acidic, basic, or neutral?

NH4Cl (aq) NH4
+
(aq) + Cl-
HCl + NH4OH NH4Cl + HOH

NH4Cl (aq) NH4
+
(aq) + Cl-
s.a. w.b.

NH4Cl (aq) NH4
+
(aq) + Cl-
(aq) Acidic!
s.a. w.b.
Will the ions from the salt combine with water?
NH4
+
+ HOH NH4OH + H+

NH4Cl (aq) NH4
+
(aq) + Cl-
(aq) Acidic!
s.a. w.b.
NH4
+
+ HOH NH4OH + H+
w.b.

NH4Cl (aq) NH4
+
(aq) + Cl-
(aq) Acidic!
s.a. w.b.
NH4
+
+ HOH NH4OH + H+
w.b.
This reaction is OK,
since a w.b. is formed

NH4Cl (aq) NH4
+
(aq) + Cl-
(aq) Acidic!
s.a. w.b.
NH4
+
+ HOH NH4OH + H+
w.b.
Cl-
+ HOH HCl + OH-

NH4Cl (aq) NH4
+
(aq) + Cl-
(aq) Acidic!
s.a. w.b.
NH4
+
+ HOH NH4OH + H+
w.b.
Cl-
+ HOH HCl + OH-
s.a.

NH4Cl (aq) NH4
+
(aq) + Cl-
(aq) Acidic!
s.a. w.b.
NH4
+
+ HOH NH4OH + H+
w.b.
Cl-
+ HOH HCl + OH-
s.a.
Cannot form s.a. from
weaker reactants, thus
N.R.

NH4Cl (aq) NH4
+
(aq) + Cl-
(aq) Acidic!
s.a. w.b.
NH4
+
+ HOH NH4OH + H+
w.b.
Cl-
+ HOH HCl + OH-
s.a.
Cannot form s.a. from
weaker reactants, thus
N.R.
Since H+
was formed in the first reaction, then [H+
] is now
greater than [OH-
] making the solution acidic

NaCl (aq)

NaCl (aq) Na+
(aq) + Cl-
(aq)

NaCl (aq) Na+
(aq) + Cl-

NaCl (aq) Na+
(aq) + Cl-
HCl + NaOH NaCl + HOH

NaCl (aq) Na+
(aq) + Cl-
s.a. s.b.

NaCl (aq) Na+
(aq) + Cl-
(aq) Acidic!
s.a. s.b.

NaCl (aq) Na+
(aq) + Cl-
(aq) Acidic!
s.a. s.b.
Now react each of the ions with water.
Na+
+ HOH NaOH + H+

NaCl (aq) Na+
(aq) + Cl-
(aq) Acidic!
s.a. s.b.
Na+
+ HOH NaOH + H+
s.b.

NaCl (aq) Na+
(aq) + Cl-
(aq) Acidic!
s.a. s.b.
Na+
+ HOH NaOH + H+
s.b.
Cannot form strong
bases from weaker
ones, thus N.R.

NaCl (aq) Na+
(aq) + Cl-
(aq) Acidic!
s.a. s.b.
Na+
+ HOH NaOH + H+
s.b.
Cannot form strong
bases from weaker
ones, thus N.R.
Cl-
+ HOH HCl + OH-

NaCl (aq) Na+
(aq) + Cl-
(aq) Acidic!
s.a. s.b.
Na+
+ HOH NaOH + H+
s.b.
Cannot form strong
bases from weaker
ones, thus N.R.
Cl-
+ HOH HCl + OH-
s.a.

NaCl (aq) Na+
(aq) + Cl-
(aq) Acidic!
s.a. s.b.
Na+
+ HOH NaOH + H+
s.b.
Cannot form strong
bases from weaker
ones, thus N.R.
Cl-
+ HOH HCl + OH-
s.a.
Cannot form strong
acids from weaker
ones, thus N.R.

Buffers
Buffers are extremely important in chemistry and biology. They
maintain a nearly consistent pH in various solutions.

Buffers
Our blood must maintain a pH around 7.35-7.45. If the pH is
above 7.45 you would have a condition called alkalosis. If the pH
is below 7.35, then one would suffer from acidosis.

Buffers
Acidosis leads to depression of the nervous system. Mild acidosis
can result in dizziness, disorientation, or fainting; a more severe
case can cause coma, or death.

Buffers
What would happen to the pH of our blood if we were to eat
acidic foods, such as apples, oranges, or limes? What might
happen to the pH of our blood if some of the hydrochloric acid
from our stomach were to seep into our blood?

Buffers
What would happen to the pH of our blood if we were to eat
acidic foods, such as apples, oranges, or limes? What might
happen to the pH of our blood if some of the hydrochloric acid
from our stomach were to seep into our blood? The pH would
be lower in both

Despite the possibility of pH increases or decreases, the body
maintains a nearly constant pH of 7.4. The body uses buffers to
maintain this remarkable feat.
What is a buffer and how does it work?

A buffer consists of a weak acid and the salt of its conjugate base,
or a weak base and the salt of its conjugate acid.
Examples:
HF + NaOH NaF + HOH
w.a. c.b.

A buffer consists of a weak acid and the salt of its conjugate base,
or a weak base and the salt of its conjugate acid.
Examples:
HF + NaOH NaF + HOH
w.a. c.b.
NH3 + HCl NH4Cl
w.b. c.a.

1.0 L
HF (g) NaF (s)
Buffer preparation: Add
0.10 mole HF (g) and NaF (s)
to 1.0 L of water.

1.0 L
HF (g) NaF (s)
to 1.0 L of water.
HF (g)  H+
+ F-
NaF (s) Na+
+ F-
H+
Na+
F-
HF
large small

1.0 L
to 1.0 L of water.
HF (g)  H+
+ F-
NaF (s) Na+
+ F-
HF
H+
Na+
F-
Now add the strong acid HCl
HCl
large small

1.0 L
to 1.0 L of water.
HF (g)  H+
+ F-
NaF (s) Na+
+ F-
HF
H+
Na+
F-
HCl
HCl H+
+ Cl-
H+ Cl-
What will the pH be if just water and no
buffer?
Large small

1.0 L
to 1.0 L of water.
HF (g)  H+
+ F-
NaF (s) Na+
+ F-
H+
HF
Na+
F-
HCl
HCl H+
+ Cl-
H+ Cl-
buffer? pH = 1, dead if this is your blood.
Large small

1.0 L
to 1.0 L of water.
HF (g)  H+
+ F-
NaF (s) Na+
+ F-
H+
HF
Na+
F-
HCl
HCl H+
+ Cl-
H+ Cl-
Large small
What removes the H+
to keep the pH near 7?

1.0 L
to 1.0 L of water.
HF (g)  H+
+ F-
NaF (s) Na+
+ F-
H+
HF
Na+
F-
HCl
HCl H+
+ Cl-
H+ Cl-
Large small
What removes the H+
to keep the pH near 7? The conjugate base, F-
H+
+ F-
HF (a weak acid, low H+
)

1.0 L
to 1.0 L of water.
HF (g)  H+
+ F-
NaF (s) Na+
+ F-
H+
HF
Na+
F-
Now add the strong base NaOH
NaOH
Na+ OH-Large small
What will the pH be if just water
and no buffer?
NaOH Na+
+ OH-

1.0 L
to 1.0 L of water.
HF (g)  H+
+ F-
NaF (s) Na+
+ F-
H+
HF
Na+
F-
Now add the strong base NaOH
NaOH
Na+ OH-Large small
What will the pH be if just water and
no buffer? PH = 13, dead again
NaOH Na+
+ OH-
What removes the OH-
to keep the pH near 7? The acid HF
HF + OH-
F-
+ HOH

Titration
Titration is an experimental procedure to
determine the concentration of an
unknown acid or base.
The figure on the left shows the
glassware for a titration experiment. A
buret clamp holds the buret to a ring
stand and below the buret is a flask
containing the solution to be titrated,
which includes an indicator. The
purpose of the indicator is to indicate
the point of neutralization by a color
change.

The picture on the left shows the
tip of a buret, with air bubble,
which is not good, and also shows
the stop-cock. Note the position
of the stop-cock is in the “off”
position. This picture shows the
color of the phenolphthalein
indicator at the end-point. In this
experiment a 23.00 mL aliquot of
0.1000 M NaOH titrant is added to
5.00 mL of an unknown HCL
solution. The acid solution in the
beaker starts out clear and
becomes pink when all of the HCL
has been consumed.
NaOH + HCl  NaCl + HOH

How can we calculate the concentration of
acid in the beaker?
Titration

How can we calculate the concentration of acid in the
beaker?
Normal procedure, yes, a conversion. Steps 1-4, again!
Titration

How can we calculate the concentration of acid
in the beaker?
Normal procedure, yes, a conversion. Steps 1-4,
again!
0.100 mole NaOH
L NaOH solution

in the beaker?
0.100 mole NaOH
L NaOH solution
10-3
L solution
mL solution

in the beaker?
again!
0.100 mole NaOH
L NaOH solution
10-3
L solution
mL solution
23.00 mL soln

in the beaker?
again!
0.100 mole NaOH
L NaOH solution
10-3
L solution
mL solution
23.00 mL soln
mole NaOH
mole HCl

in the beaker?
0.100 mole NaOH
L NaOH solution
10-3
L solution
mL solution
23.00 mL soln
mole NaOH
mole HCl

in the beaker?
0.100 mole NaOH
L NaOH solution
10-3
L solution
mL solution
23.00 mL soln
mole NaOH
mole HCl
10-3
L HCl soln.
mL HCl soln.

in the beaker?
0.100 mole NaOH
L NaOH solution
10-3
L solution
mL solution
23.00 mL soln
mole NaOH
mole HCl
10-3
L HCl soln.
mL HCl soln.
5.00 mL

in the beaker?
0.100 mole NaOH
L NaOH solution
10-3
L solution
mL solution
23.00 mL soln
mole NaOH
mole HCl
10-3
L HCl soln.
mL HCl soln.
5.00 mL
=
0.460 M HCl

Indicators
Indicators are weak organic (carbon containing) acids of
various colors depending on the formula of the acid.
Below is a generic acid.
HA  H+
+ A-
colorless pink
1. Describe the color change when a strong acid is added?

Indicators
HA  H+
+ A-
colorless pink
Less pink

Indicators
HA  H+
+ A-
colorless pink
2. Describe the color change when a strong base is added?
Less pink

Indicators
various colors depending on the formula of the acid. Below is
a generic acid.
HA  H+
+ A-
colorless pink
Less pink
Darker pink

Indicators
HA  H+
+ A-
colorless pink
3. Describe the color change when the pH is lowered?
Less pink
Darker pink

Indicators
HA  H+
+ A-
colorless pink
Less pink
Darker pink
Less pink

Indicators
HA  H+
+ A-
colorless pink
4. Describe the color change when the pH is raised?
Less pink
Darker pink
Less pink

Indicators
HA  H+
+ A-
colorless pink
4. Describe the color change when the pH is raised?
Less pink
Darker pink
Less pink
Darker pink

Color versus pH of Many Different indicators

How can we make an indicator?
Step One
Red Cabbage
Step Two
Cook the Cabbage
Step Three
Filter the Juice

What color is the juice after filtering?

What color is the juice after filtering? The color of pH 6, 7, or
8
Colors of cabbage juice at various pH values

ACIDS BASES AND SALTS
The End Ch#14

Acid bases and salts.part 2

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