CHE 157-
PHYSICAL CHEMISTRY 1
ACIDS, BASES AND SALTS
BY
BABATUDE B. ADELEKE
RECOMMENDED TEXT BOOKS
FUNDAMENTAL PHYSICAL CHEMISTRY
Edited by: Iweibo I., Okonjo K. and Obi-Egbedi N.
Chapter 2
UNIVERSITY CHEMISTRY
Bruce H. Mahan
Chapter 6
COURSE OUTLINE
Qualitative and Quantitative Aspect of Acids,
bases and salts
What are acids, bases and salts?
Different definitions and concepts
How do they ionize or dissociate in water
Quantitative description of aqueous equilibria of
solutions of acids, bases and salts
Relationships between acid or base strength
and chemical structure
Buffer solution
Hydrolysis of salt
Acid-base indicators
INTRODUCTION
What acid is in the human stomach and what is
its use?
Some citrus fruits taste sour, why?
CITRUS FRUIT
\
LEMONS
Ascorbic acid
(Vitamin c)
GRAPE FRUITS
Citric acid
•Vinegar contains acetic acids
•Aspirin is acetylsalicylic acid
• Fluid inside a car battery is a solution of sulphuric
acid
• Vitamin C- Ascorbic acid
• Milk of magnesia which is taken to soothe an
upset stomach is a base-magnesium hydroxide
• Sodium hydroxide is sold in supermarket under
the name lye and is also present in oven cleaners
and drain cleaners such as Drano
• Ammonia is present/found in cleaning products
around the house
Phosphoric acid, H3 PO4 is used in the
manufacture of fertilizers, detergents and
various soft drinks e.g. colas
Acidic anhydride
CO2 + H2O
H2CO3
SO2 + H2O
H2SO3
SO3 + H2O
H2SO4
2NO2 + H2O
2HNO3
Air pollutants- Acid rain
PROPERTIES OF ACIDS
 Give food a sour taste e.g. vinegar, lemon juice
 Dissolve metals e.g. iron and zinc
 Aqueous solution conducts electricity
Electrolyte------some strong e.g. HCl(aq), H2SO4(aq)
Others weak e.g. acetic acid, citric acid
 Turn damp blue litmus red or pink
PROPERTIES OF BASES
 React with acids to form salt and water only
 Aqueous solution- bitter taste and feel slippery
to the touch
 Aqueous solution conducts electricity
Electrolyte------some strong e.g. NaOH(aq)
Ca(OH)2.
Others weak e.g. NH3 (aq)
 Turn damp red litmus blue
Neutralization
 Reaction between an acid and a base
 Reaction between a strong acid and a strong
base in an aqueous solution takes the form of the
net equation:
H3O+(aq)+ OH- (aq)
2H2O
 Reaction
between
sodium
hydroxide
hydrochloric acid
NaOH(aq) + HCl(aq))
OH- + H+
NaCl(aq) + H2O
2H2O
and
The products of a neutralization reaction in
aqueous solution are salt and water only
 the tendency for the formation of water in a
neutralization reaction is very strong so much so
that acids also react with insoluble metal
hydroxides and oxides e.g.
Mg(OH)2(s)+ 2H+(aq)
Fe2 O3(s)+ 6H+(aq)
Mg+ + 2H2O
2Fe3+ + 3H2O
ACID SALT
The reaction between NaOH
(aq)
and HCl (aq) can
produce only one possible salt, NaCl
NaOH (aq) + HCl (aq)
NaCl(aq) + H2O
Reaction between a base and a polyproptic acid can
produce two or more salts, depending on the extent
of neutralization of the acid. Add 2mol of NaOH to
1mol of H2SO4. Complete neutralization
2NaOH (aq) + H2SO4 (aq)
Na2 SO4(aq) + 2H2O
On the other hand, add 1mol of NaOH to 1mol of
H2SO4.
NaOH (aq) + H2SO4 (aq)
NaHSO4(aq) + 2H2O
The products are NaHSO4(aq) and water. This
corresponds to partial neutralization of the diprotic
acid. Similarly with tripotic acid H3PO4 you can
produce
i.Na3PO4–Trisodium phosphate (sodium phosphate
ii. Na2 HPO4–Disodium hydrogen phosphate
iii. NaH2PO4–Sodium dihydrogen phosphate
 Salts that are the result of the partial
neutralization of polyprotic acid are called acid
salts
 acid salts such as Na2HPO4 are acidic and can
undergo further neutralization. Thus Na2HPO4
is able to react with NaOH.
Na2HPO4 (aq) + NaOH (aq)
Na2 SO4(aq) + H2O
DIFFERENT CONCEPTS OF ACID AND
BASES
The different concepts
• Arrhenius concept
• Lowry-Bronsted concept
• Lewis concept
Arrhenius concept: An acid is a compound which
dissociates in aqueous solution to produce
hydrogen ions H+ (H+) exists in aqueous solution
Examples of Arrhenius acids are HCl (aq), HNO3 (aq),
HC2H3O2 (aq), or CH3COOH (aq), H2SO4(aq), H3PO4(aq)
not all compounds that contain hydrogen are
arrhenius acids e.g. CH3CH2OH
H
O
H
C
H
C
O
H
H
Ethanoic acid
Ionisable hydrogen bonded to highly
electronegative element, oxygen
H
H
C
C OH
H
H
Ethanol
CLASSIFICATION OF ARRHENIUS ACIDS
The types of classifications
A. STRUCTURAL CLASSIFICATION
i.
Non-oxygen acids- HCl(aq) HCN(aq) and HF(aq)
ii. Oxy acids or oxygen-containing acids
HClO(aq), HNO3(aq), H2SO4 (aq) one or more oxygen
in a molecule of the acids
B. PROTIC CLASSIFICATION
i.
Monoprotic Acids- can release one hydrogen
ion per molecule e.g. HCl(aq), HF(aq), CH3COOH
ii. Diprotic Acids-can release two hydrogen ions
per molecule of acids H2SO4(aq), H2S(aq)
iii. Triprotic Acids-can release three hydrogen ions
per molecule of acids H3PO4(aq)
Note:(ii) and (iii) are referred to as Polyprotic acids
C. STRENGTH OF ACID CLASSIFICATION
i. Strong Arrhenius acids-acids in which all
ionisable hydrogen atoms in the molecule react
with
water
to
produce
hydroxonium
ions/hydrogen ions
The reaction
HCl(aq) + H2O(L)
goes to completion
HCl(aq) is a strong acid
H3O+(aq) + Cl-(aq)
ii. Weak Arrhenius acids- only a relatively small
number of the hydrogen atoms react with water to
form the hydroxonium ion, H3O
In the reaction
HF(aq) + H2O
H3O+(aq) + F-(aq)
The equilibrium lies essentially to the left.
Arrhenius concept of bases- a compound that
contains the hydroxyl group OH and dissociates in
aqueous solution to produce the hydroxide ion (OH-)
is a base according to Arrhenius concept.
NaOH(aq)
Na+(aq) + OH-(aq)
Mg(OH)2
Mg+(aq) + 2OH-(aq)
A strong Arrhenius base ionizes completely in water.
A weak base releases a relatively few of its
hydroxide ions Al(OH)3 which is insoluble in water.
Al(OH)3(s)
Al3+(aq) + 3OH-(aq)
Equilibrium lies essentially to the left
SUMMARY OF ARRHENIUS CONCEPT
The proton H+ is responsible for acidic properties
The hydroxide ion, OH- is responsible for basic
properties. However
i. What is the nature of the proton in aqueous
solution?
ii.Substances which did not contain hydroxide ion,
OH- were capable of acting as bases or neutralize
acids e.g. NH3(L) + HCl
NH4+Cl-(s)
Pure liquid Ammonia
Ammonia is acting, in this reaction as a base.
When sodium carbonate, Na2CO3, is dissolved in
water a solution that will neutralize a acid results.
Sodium carbonate cannot by itself dissociate
directly to produce hydroxide ions, but its reaction
suggests it must be a base
To accommodate the deficiencies of the Arrhenius
concept,
a
new
concept
called
LOWRY-
BRONSTED CONCEPT was introduced
DEFINITION
An acid is any substance, molecular or ionic,
capable of giving up (donating) a proton.
A base is any substance molecular or ionic, capable
of accepting a proton.
 An acid is a proton donor
 A base is a proton acceptor
ACID = BASE + PROTON (H+)
Examples
Ionization of HCl in water
HCl(aq) + H2O
H3O+(aq) + Cl-(aq)
Acid
Base
In the reverse reaction
HCl and Cl- differ only by a proton
HCl and Cl- are called a CONJUATE ACID-BASE
PAIR
H3O and H2O are CONJUGATE ACID-BASE
PAIR
HCl + H2O
Acid 1
Base 2
H3O+ + ClAcid 2
Base 1
Acid1 and base 1 form Acid-Base conjugate pair
Acid 2 and Base 2 form Acid-Base conjugate pair
CARBONATE ION AS A BASE
CO32- + H2O
HCO32- + OHBase 2
Acid 1
Acid 2
Base 1
The Lowry-Bronsted definition extends the terms
acid and base to include substances besides H+ and
OH- with the resulting advantage that a large number
of reactions can be discussed in the same language
and treated mathematically by the same methods
Examples
HCl(g) + NH3(aq)
Acid 1
Base 2
H2 O(aq) + H2O(aq)
Base 2
Acid 1
NH4+ + ClAcid 2
Base 1
H3O+(aq) + OH-(aq)
Acid 2
Base 1
Note: Amphoteric nature of water
Auto-ionization or self-ionization of water
CH3COOH(aq) + H2O(aq)
CH3COO-(aq) + H3O+(aq)
Acid 1
Base 2
HS- + H3O+
H2 S(aq) + H2O
Acid 1
Base 2
HS- + H2O
Acid 1
Base1
Base1
HA(aq) + H2O(aq)
Base 2
Acid 2
S2- + H3O
Base1
HA + B
A-+ HB+
Acid 1
Base1
Base 2
Acid 2
S2- + H3O
Base 2
Acid 1
Base 1
Acid 2
Acid 2
Acid 2
B = H2O, NH3, OH-,CH3CH2NH2,
(CH3CH2)2 NH, R3N, etc
SRENGTH OF ACIDS AND BASES
Strong acid- a Large tendency to transfer a proton to
another molecule
Strong base- a large tendency to accept a proton
from another
If we measure strength quantitatively by the degree
to which reactants are converted to products in a
reaction such as
HA + B
A-+ HB+
Acid 1
Base 2
Base1
Acid 2
Extent of reaction depends on the strength of both
HA and B
Therefore we can only compare the strength of
individual acids by measuring their tendencies to
transfer a proton to the same base
In aqueous solution, water is the same base
Quantitative measure of acid strength is Ka, the
acid dissociation constant
HA + H2O
A- + H3O+
K= equilibrium constant
Ka is an equilibrium constant whose value
Is temperature dependent
Auto-ionization of water
H2 O + H2O
OH- + H3O+
Kb {H2O} = {OH-} {H3O+}
Kw = {OH-} {H3O+}
Ion product of water
Table of Acid dissociation constant of some weak
acids at 250c
In
HA/A- conjugate pair
If HA is a strong acid
A- must be a weak base
HB+/B conjugate acid-base pair
If B is a strong base
HB+ must be a weak acid
In general, if an acid is a strong acid, its conjugate
base is a weak base and vice versa.
For a base, water is the common acid
B + H2O
BH+ + OH-
Kb = Base dissociation constant
If B≡A-, the conjugate base of acid HA
A - + H 2O
HA + OH-
Ka X Kb = {H3O} {OH-}=Kw
Note that Ka is inversely proportional to Kb
Therefore the larger the value of Ka the smaller is
Kb similarly the larger the value of Kb the smaller
is Ka
i.e the stronger the acid, the weaker the conjugate
base or the stronger the base, the weaker the
conjugate acid
Table of Dissociation constants of some weak base
at 250c
Note: Absolute values of Ka and Kb are very small
Therefore introduce another scale of measurement
of acid or base strength
pKa = - log Ka
pKb = - log Kb
Hence,the stronger the acid the smaller the value
of pKa the weaker the base the larger the value of
Kb and vice-versa.
Similarly we define
pKw = - log Kw
pH = - log{H3O+}
pOH = - log{OH-}
Note that
Kw = {H3O+} {OH-}
-logKw = - log{H3O+} - log{OH-}
pKw = pH + pOH
True in all aqueous
At 250c, Kw = 10-14
At 250c, pH + pOH= 14
SAMPLE CALCULATIONS
If Ka = 1.8 X 10-4
Calculate the pKa of the acid
pKa = - log Ka = - log 1.8 X 10-4
= - log 1.8 - log10-4
= - log 1.8 - x -4
= 4 – log1.8 = 3.74 (log1.8=0.2553)
Ka = 1 x 10-12
pKa = -log1 x –log 10-12
= 0 + 12 = 12
The stronger the acid (Ka = 1.8 x 10-4 ) the smaller
the pKa = 3.74
THE LEWIS THEORY OF ACID AND BASES
In the Lewis theory
• A base is defined as a substance that can furnish
an electron pair i.e a base is a electron donor
• An acid is a substance that can accept electron
pairs i.e an acid is an electron acceptor
• A coordinate covalent bond is formed between
the acid and the base.
Two conditions are necessary
(a) Existence of a vacant orbital of appropiate
energy on the acid
(b) existence of a pair of electrons of appropiate
energy on the base
Examples
BF3 + F- = BF4Acid
base
..
-:
:F
..
F
B
F
F
F
F
B
F
F
Coordinate covalent bond
Ag+ + 2CN- = [Ag(CN)2]Acid
Base
[:C≡N:]..
H+ + NH3
Acid
base
..
Ag+ + 2NH3
Acid
Base
NH4+
[Ag(NH3)2]-
H+ + H2O:
H3O+
All Lowry-Bronsted bases are electron donors
Therefore All Lowry-Bronsted bases are lewis bases
H+ is a lewis Acid because it has vacant valence
orbital. Other Lewis acids such as Ag+, Cu2+, Fe3+,
BCl3, AlCl3 do not contain protons
Lewis acid-base theory extended concept of acid-
base reactions
Ligands are Lewis bases attached to metal ions,
e.g. H2O in [Cu(H2O)4]2
CN- in [ Fe(CN)6]3(a) Monodentate Ligands-one donor atoms
involved in bonding e.g H2O, NH3, Cl-, F-, CN(b) Bidentate Ligands- two atoms in a molecule
involved in covalent bond formation
..
..
e.g HN-CH2-CH2- NH2 Ethylene diamine
..N
..N
N
Bipyridyl
N
1, 10- phenanthroline
CH2
CH2
H2N
H2N
M
..
[:O
..
O
HO
.. 2C
O:]
..
oxalate ion
O
C
O
C
from
O
C
OH
oxalic acid
O
C
..
:O
O
C
..
:
O
M
C. Polydentaate ligands- many atoms involved in
covalent bond formation, most common
example-Ethylene Diamine Tetra Acetic Acid
(EDTA)
HO2CCH2
CH2COOH
N
CH2
CH2
N
HO2CCH2
CH2COOH
CH2COO-
CH2COOH
ionizes
N
N
CH2COOH
CH2COONitrilotriacetate
CH2COOH
..
-:O
..
O
O
CH2 C O:-
C CH2
..
N
-:O
..
C CH2
O
CH2COO-
CH2
CH2
..N
CH2 C
O
..
O
..:-
Biologically significant Lewis acids are metal ions
in biological molecules such as hemoglobin (fe),
Chlorophyll(Mg), the vitamin B12 group and
cytochrome
Hydrogen ion concentration in a saturated aqueous
solution of CO2 is 1.3 x 10-4M. What is the pH of
the solution
ACID BASE EQUILIBRIA IN AQUEOUS
SOLUTIONS
A.
The pH Scale
pH= -log[H3O+]
Hydrogen ion concentration in a saturated aqueous
solution of CO2 is 1.3 x 10-4M. What is the pH
of the solution?
pH= -log[H3O+]
[H3O+] = 1.3x 10-4M
pH = -log(1.3 x 10-4)
= -log(1.3 + log10-4)
= -log1.3 - log10-4 = = -log1.3 + 4
= 4-0.11
= 3.89
If a solution has a pH of 4.5, what is the hydrogen
ion concentration of the solution?
pH= -log[H3O+]
4.5 = -log[H3O+]
-4.5 = log[H3O+]
i.e log[H3O+] = (0.5-5)
[H3O+] = 100.5 x 10-5M
= 3.2 x 10-5M
B. THE IONIZATION OF WATER
Water as a weak electrolyte
H2O + H2O
H3O+(aq) + OH-(aq)
Kw = [H3O+] [OH-]
pH + pOH = pKw
At 250C, Kw = 10-14
For a neutral solution
[H3O+] = [OH-]
in any aqueous solution and
at any temperature
Acidic solution
[H3O+] ˃ [OH-]
Basic solution
[H3O+] ˂ [OH-]
In a neutral solution at 250C
[H3O+] = [OH-] = 10-7M
pH = pOH=7
SOLUTION OF STRONG ACIDS AND BASES
Hcl and HNO3 are the most common monoprotic
acids. Acids are strong because they are essentially
100% dissociated in aqueous solution
HNO3 + H2O
HNO3
H3O + NO3-
or
H+ + NO3-
Contribution of a strong acid to the H+ in solution
is determined by the concentration of the strong
acid.
Thus 0.05M HNO3 will contribute to the solution
0.05M H+ and 0.05M NO3Strong bases are usually ionic compounds e.g
soluble metal hydroxides, such as NaOH, KOH and
ca(OH)2. They dissociate completely in aqueous
solution
NaOH
H2O
Na+ + OH-
H2O
Ca(OH)2
Ca2+ + 2 OH-
Note that for every mole of Ca(OH)2 that dissolves
2mol of OH- are added to the solution
pH of solution of strong acids
Assume that essentially all the H+ in the solution
comes from the dissolved acid. We assume that
contribution of ionization of water to H+ in the
solution is negligible. This is because the presence
of H+ from an acid (HNO3) shifts the equilibrium
H2O
H+ + OH- to the left
H+ Coming from the dissociation of water is less
than 10-7M
pOH of solution of strong bases
OH- in a solution of a strong base is assumed to
come from the dissociated strong base i.e
contribution of ionization of water to OH- in
solution is negligible
H2O
H+ + OH-
Common ion Effect
What is the pH of 0.1M solution of Hcl?
[H3O+] = 0.1M = 10-1M
pH = -log [H3O+] = -log 10-1 = -x-1 = 1
What is the pOH of the same solution at 250C
At 250C, pKw = 14
pH+ pOH = 14
pOH = 14-pH
pOH = 14-1
= 13
Note that the only source of OH- is an acid solution
is the auto ionization of water
H2O
H+ + OH-
Thus is 0.1M Hcl solution, [OH-] = 10-13M
In pure water [neutral] [OH-] = 10-7M
Therefore since 10-13M ˂˂ 10-7M, the selfionization of water has been suppressed by the
common ion, H+ from both the acid and selfionization of water and the contribution of H+ from
self-ionization of water (10-13M) is negligible
compared with [H+]acid = 10-1M.
What is the pOH of a 0.01M NaOH solution? In
0.01M NaOH solution, [OH-] = 10-2 M
pOH = -log [OH-] =
= -log 10-2 = -1x-2
=2
What is the pH of the same solution at 250C
pH + pOH = 14
pH= 14-pOH = 14-2 = 12
[H+] = 10-12M
In alkaline solution, the only source of H+ is the
auto-ionization of water
H2O
H+ + OH-
In pure water, [OH- = [H+] = 10-7M
But 10-12M ˂˂ 10-7M i.e auto-ionization of water
has been suppressed.
EQUILIBRIA IN AQUEOUS SOLUTION OF
WEAK ACID
HA + H2O
H3O + A-
Acid ionization constant
Degree of ionization, α =
If C is the total concentration of the acid
{H3O+}= αc
{A-} = αc
{HA} = c- α c = (1- α) c
α is small cared with 1, since HA is a weak acid
Therefore Ka = α2c
α=
pH of solution of a weak acid
HA + H2O
H3O+ + A-
C is the total concentration of the acid
C = [HA] + [A-]
mass balance
Two assumptions are made
1. HA is a weak acid hence fraction of ionized
acid is small, [A-] ˂ [HA]
Then C = [HA]
2. Contribution of self-ionization of water to the
hydrogen concentration in solution is negligible,
hence
[HA]= [H3O+]
Substituting (1) and (2) into expression of Ka
[H3O] =
= Ka C
-log[H3O] = -logKa.C
= - logKa - logC
pH = pKa - log C
pH of aqueous solution of a weak base
B + H2O
Base ionization
Constant
BH+ + OH-
If c mol/dm3 is the total concentration of the base in
the solution then
C = [B] + [BH+]
Two assumption are made
1. The base is a weak base then,
[BH+] ˂ [B]
Therefore C = [B]
mass balance
2. Assume that the contribution of self-ionization
of water to the total OH- in solution is negligible
[BH+] = [OH-]
Subtituting the results of the two assumptions in
equation for Kb
Taking –log of both side of equation
-log[OH-] = - logKb - logC
pOH = pKb - log C
Remember that pH = pOH
Substitute the equation for pOH
pH = pKw - logKb + logC
Note that α. The degree of ionization of the base is
but
HYDROLYSIS
Hydrolysis is the interaction of the ions of salts with
the ions of water.
Classification of salts according to hydrolysis
reactions
1) Salts of strong acid and strong bases e.g NaCl,
KNO3
2) Salts of weak acids and strong bases e.g
CH3COONa, KCN
3) Salts of strong acids and weak bases e.g NH4Cl
4) Salts of weak acids and weak bases e.g
CH3COONH4
A. Salts of strong acids and strong bases, e.g NaCl
2H2O
H3O+ + OH-
NaCl
Cl- + Na+
No interaction between the ions of salts and water
hence [H3O+] = [OH-]. Solution is neutral
B. Salts of strong base and weak acids, NaA
2H2O
NaA
H3O+ + OHA- + Na+
HA + H2 O
Removal of H3O from solution hence OH- is in
excess and solution is basic
C. Salts of strong acid and weak base, BH+ ClBH+ Cl-
Cl- + BH+
B + H2O
Removal of OH- from solution hence H3O is in
excess and solution is basic.
D. Salts of weak acids and weak bases, BH+A-
2H2O
H3O+ + OH-
BH+A-
A- + BH+
HA + H2O
B + H2O
Both H3O and OH- are removed form solution. The
solution may be acidic, basic or neutral depending
on the relative values of Ka and Kb
Calculation of pH or pOH of aqueous solution of
salts
A. Salts o a strong base and a weak acid, \na\a
A - + H 2O
HA + OH-
Ionization of the conjugate base of a weak acid
Hydrolysis
Constant
.
=
Therefore
Since
Ka = Acid ionization constant
Kh is inversely proportional to Ka
Therefore the smaller the value of Ka i.e the weaker
the acid, the bigger the value of Kh i.e the more
extensively hydrolsed is the conjugate base.
To calculate OH- concentration in the solution if
the slat where C = total concentration of salt, apply
the usual two assumptions
1. A- is a weak base
c = [A-] + [HA]
[HA] ˂ [A-]
Therefore C = [A-]
2. Contribution of self ionization of water to OHis negligible. Therefore
[HA] = [OH-]
Substitute into Kh
Solution is basic, excess OH-
(B) Salt of a strong acid and a weak base BH+Cle.g NH4+
BH+ + H2O
B + H3O+
Hydrolysis
constant
Since
Kb = Base ionization costant
To calculate [H3O+] is solution
Apply the two assumption
C= [BH+] + [B]
[B] ˂ [BH+] since BH+ is a weak acid
C = [BH+]
Neglect the contribution of self-ionization of water
to {H3O+]
Therefore [B] = [H3O
substituting into Kh
Solution is acidic.
BUFFER SOLUTION
A buffer solution is defined as one that resists
change in hydrogen-ion concentration or pH when
small amount of a strong acid or a strong base is
added to the solution
A buffer solution is used to control the pH of a
medium and to keep the pH within a fairly narrow
limits. This is necessary in many biological and
chemical processes.
A buffer solution is made from:
• A weak acid and one of its salts or
• A weak base and one of its salts.
A number of such buffers with their useful pH range
are given in Table below.
COMPONENTS
USEFUL PH RANGE
Glycine and glycine hydrochloride
1.0—3.7
Phthalic acid and potassium phthalate
2.2—3.8
Acetic acid and sodium acetate
3.7—5.6
Monosodium phosphate and disodium phosphate
5.8—8.0
Boric acid and borax
6.8—9.2
Borax and sodium hydroxide
9.2—11.0
Disodium phosphate and trisodium phosphate
11.0—12.0
BUFFER ACTION
Consider a buffer solution made from acetic and
sodium acetic mixture
Note that acetic ion is the conjugate base (weak) of
the weak acid, acetic acid. The solution contains a
high concentration of acetic acid which is almost
completely unionized and a high concentration of
acetate ions from the salt.
If a small amount of a strong acid is added to the
buffer solution, hydrogen ions of the strong, acid
react with acetate ions to form undissociated acetic
acid until equilibrium for acetic acid is satisfied
according to the equation
H3O+ (from added acid) + A-
HA + H2O
Because of this reaction, there is little change in the
hydrogen-ion concentration of the solution
If a small amount of a strong base is added,
hydroxide ions will react with undissociated acetic
acid until the constancy of Ka is satisfied. The
reaction is
HA + OH (from added base)
H2O + A-
Again the hydrogen-ion concentration remain
essentially unchanged.
Buffers with a pH higher than 7 can be prepared by
using a base----- is stronger than its conjugate acid.
A common basic buffer is formed by mixing
ammonia with an ammonium salt such as NH4Cl
and contains the conjugate acid-base pair NH4+ and
NH3. if a strong acid is added to the buffer, it reacts
as follows,
H3O + NH3
NH4+ + H2O
If a strong base is added, the reaction is
OH- + NH4+
NH3 + H2O
The calculation of pH of Buffer solutions
Consider a buffer solution containing the weak acid
HA and its highly ionized salt Na+A-. Its is
necessary that Ka for the acid be satisfied, thus
Because of the common ion action of A- ions, [HA]
is almost equal to the total concentration of the
acid, Also [A-] is essentially equal to the
concentration of the completely ionized salt, Na+AThe above equation becomes
Herdersen-Hasselbalch
equation
Use the equation to :
a) Calculate pH of a buffer solution
b) Calculate
ratio at a given value of pH
c) Calculate pKa of an acid
Note that the buffer solution has its maximum
capacity when
= 1. At this point, pH = pKa.
Therefore in choosing a buffer mixture of a
specified pH with the maximum capacity,
select an acid whose pKa value is equal or
nearly equal to that of the desired pH of the
solution.
Also, in order to be most effective, the numbers of
moles of weak acid and its salt (base) used to
prepare the buffer must be considerably greater than
the numbers of moles of acid or base that may later
be added to the buffer solution.
Buffers find many important applications. Living
systems employ buffers to maintain nearly constant
PH so that biochemical reactions can follow their
correct paths.
For example, blood contains among other things, a
H2CO3/HCO3- buffer system that helps maintain
the PH at 7.4.
In the laboratory many inorganic and chemical
reactions are performed in buffered solutions to
minimize an adverse effects caused by acids or
bases that might be consumed or produced during
reaction.
Sodium bicarbonate, NaHCO3 (baking soda) is
added to a swimming pool because it acts as a
buffer and controls the PH of the water in the pool.
This is because of the ability of HCO3 ion to react
with both acids and bases
HCO3- + H+
HCO3- + OH-
H2CO3
CO32- + H2O
ACID – BASE INDICATORS
Judge acidity or basicity of a solution – use some
substances called – base indicators.
Acid-base indicators – organic compounds whose
color depends on the PH of the solution in which
they are dissolved.
Example is litmus
Litmus is pink (red) in an acid solution
Litmus is blue in a basic solution
Litmus paper - strips of absorbent paper impregnated
with the litmus dye.
- Use it to test whether a solution is
acidic or basic
- Acidic solution will turn blue litmus
paper red
- Basic solution will turn pink litmus
paper blue
Other PH test papers are available that contain
mixtures of indicators dyes and can actually be used
to estimate the approximate value of the PH of a
solution.
Today, usually instrument called PH meters, which
re electronic devises are used to measure the PH of
a solution with a high degree of precision and
accuracy.
One principal use of acid-base indicators is in the
detection of the equivalence point in acid-base
titrations.
Indicators – usually weak organic acids or bases
that change color over a range of PH values.
Not all indicators change color at the same PH
The choice of indicator for a particular titration
depends on the PH at which the equivalence point
is expected to occur.
The table below shows a list of indicators with
their color change and the PH ranges over which
the color changes are observed.
SOME COMMON INDICATORS
INDICATOR
COLOR CHANGE
PH RANGE
(for color change)
Thymol blue
Red to yellow
1.2 – 2.8
Bromophenol blue
Yellow to blue
3.0 – 4.6
Congo Red
Blue to Red
3.0 – 5.0
Methyl Orange
Red to yellow
3.2 – 4.4
Brumocresol green
Yellow to blue
3.8 – 5.4
Methyl Red
Red to yellow
4.8 – 6.0
Bromocresol purple
yellow to purple
5.2 – 6.8
Bromothymol blue
yellow to blue
6.0 – 7.6
Cresol red
yellow to red
7.0 – 8.8
Thymol blue
yellow to blue
8.0 – 9.6
Phenolphthalein
colorless to pink
8.2 – 10.0
Alizarin yellow
yellow to red
10.1 – 12.0
HOW INDICATOR WORK
Denote an indicator as HIn
Dissociation reaction HIn
H+ + In-
Apply Le chatelier’s principle:
In acidic solution (excess of H+), the predominat
species is HIn
In basic solution, equilibrium shifts to the right and
predominant species is InHIn is acid form of the indicator
In- is the basic form of the indicator
The color of the acid from HIn is different from the
color of the base In- . The ability of HIn to function
as an indicator as an indicator is based on the
difference in color between acid and basic forms.
For example, with litmus the acid form, (HIn) is
pink while the basic form (In-) is blue
The dissociation constant Ka for the indicator is
The pH changes very rapidly as we pass through
the equivalence point is acid-base titration
Consider NaOH– HCl titration
pH changes from 4.7 to 9.3 with the addition of
only 0.02ml of base. This is only about one-half of
a drop of solution
When pH = 4.7, [H+] = 2 x 10-5M
When pH =9.3, [H+] = 5 x 10-10M
How does this rapid change of affect the ratio
Assume Ka of the indicator to be 1x10-7, then
before equivalence point
=
=
We have 200 times as much HIn as InTherefore color observed is that of HIn
After the equivalence point, [H+] = 5 x 10-10M
=
=
?
Now we have 200 times as much of In- as HIn
Therefore color observed is that if InThus as we pass through the equivalence point
there is a sudden change in the relative amounts of
the acid and basic forms of the indicator, which we
notice as color change.
Relationship between acid or base strength and
chemical structure.
See previous slides– tables of Ka and Kb of acids
and bases respectively.