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Electrolytes when dissolved in water splits up into charged particles called ions. The
process is called ionisation or dissociation. Certain electrolytes such as NaCl, KCl, HCl are
almost completely ionised in solutions, whereas electrolytes such as NH 4OH, CH3COOH
etc. are weakly ionised. The electrolytes which are almost completely ionised in their
solutions are called strong electrolytes. On the other hand, electrolytes which are
weakly ionised in their solutions are called weak electrolytes. In case of solutions of
weak electrolytes, the ions produced by dissociation of electrolyte are in equilibrium
with undissociated molecules of the electrolyte. The equations for dissociation of strong
electrolytes are written with only a single arrow directed to the right.
KCl(aq) → K+ (aq) + Cl −(aq)
NH4Cl(aq) → NH4+ (aq) + Cl−(aq)
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On the other hand, equations for the dissociation of weak electrolytes are written with
double arrows
CH3COOH(aq) + H2O(l) ⇌ H3O+ (aq) + CH3COO−(aq)
NH3(aq) + H2O(l) ⇌ NH4+ (aq) + OH −(aq)
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VARIOUS CONCEPTS OF ACIDS AND BASES
ARRHENIUS CONCEPT OF ACIDS AND BASES
According to Arrhenius concept, an acid is a substance which can furnish hydrogen ions
in its aqueous solution.
A base is a substance which can furnish hydroxyl ions in its aqueous solution.
Eg. Substances such as HNO3, HCl, and CH3COOH etc are acids, where as substances
such as NaOH, KOH, NH4OH etc. are bases, according to this concept.
HNO3 ⇌ H+(aq) + NO3−(aq)
HCl(aq) ⇌ H+(aq) + Cl−(aq)
CH3COOH(aq) ⇌ H+(aq) + CH3COO−(aq)
NaOH(aq) ⇌ Na+(aq) + OH−(aq)
KOH(aq) ⇌ K+(aq) + OH−(aq)
NH4OH(aq) ⇌ NH4+(aq) + OH−(aq)
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Acids such as HCl and HNO3 which are almost completely ionised in aqueous solution
are termed as strong acids whereas acids such as CH3COOH which are weakly ionised
are called weak acids .
Similarly, bases which are almost completely ionised in aqueous solution are called
Strong bases, for example, NaOH and KOH. The bases such as NH4OH are only slightly
ionised are called weak bases.
According to Arrhenius theory, neutralisation of acids and bases is basically a reaction
between H+ and OH− ions in solutions.
H + + OH− ⇌ H2O
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Nature of Hydrogen ion in aqueous solutions
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Hydrogen atom contains one proton and one electron. H + ion is formed by loss of this
electron. Therefore H+ ion is simply a proton. Charge density of this unshielded proton is
very high. Therefore, it is not likely to exist as H+ ion. In an aqueous solution, H+ ion is
considered to be present in hydrated form in combination with a water molecule as
H3O+.
H+ + H2 O ⇌ H3O+
(H3O+ is called hydronium ion)
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CLASSIFICATION OF ACIDS AND BASES AS WEAK AND
STRONG AND THEIR SALTS
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Acids, such as hydrochloric acid, and bases, such as potassium hydroxide, that have a
great tendency to dissociate in water are completely ionized in solution; they are called
strong acids or strong bases.
Acids, such as acetic acid, and bases, such as ammonia, that are reluctant to dissociate
in water are only partially ionized in solution; they are called weak acids or weak bases.
Strong acids in solution produce a high concentration of hydrogen ions, and strong
bases in solution produce a high concentration of hydroxide ions and a correspondingly
low concentration of hydrogen ions.
The hydrogen ion concentration is often expressed in terms of its negative logarithm or
pH. Strong acids and strong bases make very good electrolytes, i.e., their solutions
readily conduct electricity. Weak acids and weak bases make poor electrolytes.
ACID-BASE NEUTRALISATION-SALTS
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When aqueous solutions of hydrochloric acid and sodium hydroxide are mixed in the
proper proportion, a reaction takes place to form sodium chloride and water.
HCl(aq) + NaOH(aqr) ⇌ NaCl(aq) + H2O(l)
Sodium chloride
Such a reaction is termed neutralisation because both acidic (H +) and basic (OH-)
properties are eliminated during the reaction. The hydrogen ion, which is responsible
for the acidic properties, has reacted with the hydroxyl ion which is responsible for the
basic properties, producing neutral water. The Na + and CI- ions have undergone no
chemical change and appear in the form of crystalline sodium chloride upon evaporation of the solution. Sodium chloride is an example of the class of compounds called
salts.
HCl(aq) + Cl-(aq) + Na+(aq) ⇌ H2O(l) + Na+(aq) + Cl-(aq)
or
H+(aq) + OH-(aq) ⇌ H2O(l)
Thus, the neutralisation of a base with an acid involves the interaction between OH - and
H+ ions.
or
The reaction between an acid and a base to form salt and water is termed
neutralisation.
The process of neutralization does not produce the resulting solution always neutral; no
doubt it involves the interaction of H+ and OH- ions. The nature of the resulting solution
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depends on the particular acid and a particular base involved in the reaction. The
following examples illustrate this point when equivalent amounts of acids and bases are
reacted in aqueous solution.
(i) A strong acid plus a strong base gives a neutral solution because both are completely
ionised and the reaction goes to completion.
H+ + CI- + Na+ + OH- ⇌ H2O + Na+ + CI(ii) A strong acid plus a weak base gives an acidic solution as the weak base is not
completely ionised. The reaction does not go to completion and there is an excess of
hydrogen ions in solution.
H+ + CI- + NH4OH ⇌ H2O + NH4+ + Cl-
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(iii) A weak acid plus a strong base gives a basic solution as the weak acid is not
completely ionised. The reaction does not go to completion and there is an excess of
hydroxyl ions in solution.
CH3COOH + Na+ + OH- ⇌ H2O + CH3COO- + Na+
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SALTS:
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(iv) A weak acid plus a weak base gives an acidic or a basic or a neutral solution
depending on the relative strength of acid and base. In case both have equal strength,
the resulting solution is neutral in nature.
CH3COOH + NH4OH ⇌ H2O + NH4+ + CH3COOSalts are regarded as compounds made up of positive and negative ions. The
positive part comes from a base while negative part from an acid. Salts are ionic
compounds. Salts may taste salty, sour, bitter, astringent or sweet or tasteless. Solutions
of salts may be acidic, basic or neutral. Fused salts and aqueous solutions of salts
conduct electricity and undergo electrolysis. The properties of salts in aqueous solutions
are the properties of ions. The salts are generally crystalline solids.
The salts are classified into the following classes:
1. SIMPLE SALTS
The salt formed by the neutralization process, i.e., interaction between acid and base, is
termed as simple salt. These are of three types:
a) Normal salts: The salts formed by the loss of all
possible protons (replaceable hydrogen atoms as H+) are called normal salts. Such a
salt does not contain either replaceable hydrogen or a hydroxyl group.
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Examples are: NaCl, NaNO3, K2SO4, Ca3(PO4)2, Na3BO3, Na2HPO3 (one H atom is not
replaceable as H3PO3 is a dibasic acid), NaH2PO2 (both H atoms are not replaceable
as H3PO2 is a monobasic acid), etc.
b) Acid salts: Salts formed by incomplete neutralization of poly-basic acids are called
acid salts. Such salts still contain one or more replaceable hydrogen atoms. These
salts when neutralised by bases form normal salts.
Examples are: NaHCO3, NaHSO4, NaH2PO4, Na2HPO4, etc.,
c) Basic salts: Salts formed by incomplete neutralization of poly acidic bases are called
basic salts. Such salts still contain one or more hydroxyl groups. These salts when
neutralised by acids form normal salts.
Examples are: Zn(OH)Cl, Mg(OH)Cl, Fe(OH)2Cl, Bi(OH)2Cl, etc.
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2. DOUBLE SALTS
The addition compounds formed by the combination of two simple salts are termed
double salts. Such salts are stable in solid state only.
Examples are: Ferrous ammonium sulphate, FeS0 4-(NH4)2SO4.6H2O, Potash alum,
K2SO4Al2(SO4)3.24H2O, and other alums.
Properties:
a) When dissolved in water, it furnishes all the ions present in the simple salts from
which it has been constituted.
b) The solution of double salt shows the properties of the simple salts from which it has
been constituted.
3. COMPLEX SALTS
These are formed by combination of simple salts or molecular compounds. These
are stable in solid state as well as in solutions.
→ K4Fe(CN)6 + K2SO4
Simple salt
Complex salt
CoSO4 + 6NH3 → Co(NH3)6SO4
Simple salt Molecular Complex salt compound
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BRONSTED-LOWRY CONCEPT OF ACIDS AND BASES
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In 1923, Bronsted and Lowry independently proposed new definitions for acids and
bases.
They proposed that: An acid is a substance that can donate a proton.
A base is a substance that can accept a proton.
These definitions are more general because according to these definitions even ions can
behave as acids or bases. Moreover, these definitions are not restricted to reactions
taking place in aqueous solutions only. In order to understand this concept of acids and
bases , consider some specific example :
HCl(aq) + H2 O(l) ⇌ H3O+(aq) + Cl−(aq)
acid
base
+
NH4 (aq) + H2 O(l) ⇌ H3O+(aq) + NH3(aq)
acid
base
H2 O(l) + NH3(aq) ⇌ NH4+(aq) + OH−(aq)
acid
base
H2 O(l) + CO 3(aq)− ⇌ HCO3−(aq) + OH−(aq)
acid
base
From the above equations, it is obvious that acid base reactions according to BronstedLowry concept involve transfer of proton from the acid to a base. A substance can act as
an acid only if another substance capable of accepting a proton is present.
Conjugate acid-base pairs
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An acid after losing a proton becomes a base whereas a base after accepting the proton
becomes an acid. For example, consider the reaction between water and ammonia as
represented by the following equilibrium:
In the forward reaction, water donates a proton to ammonia (base) and acts as acid. In
the reverse reaction NH4+ ions donate a proton to OH− ions (base) and acts as acid. A
base formed by the loss of proton by an acid is called conjugate base of the acid,
whereas an acid formed by gain of a proton by the base is called conjugate acid of the
base. In the above example, OH− ion is the conjugate base of H2O and NH4+ ion is the
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conjugate acid of NH3 . Acid- base pairs such as H2 O − OH− and NH4+ − NH3 which are
formed by loss or gain of proton are called conjugate acid-base pairs.
A strong acid would have large tendency to donate a proton. Thus, conjugate base of a
strong acid would be a weak base. Similarly, a conjugate base of a weak acid would be a
strong base.
Some more conjugate acid - base pairs has been given in the following equations:
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It may be noted that in equation
a) H2O is behaving as a base whereas in equation
b) it is behaving as an acid. Similarly HCO3− ion in equation
c) acts as an acid and in equation
d) it acts as a base. Such substances which act as acids as well as bases are called
amphoteric substances.
In both Arrhenius and Bronsted concepts, acids are sources of protons. Hence all
Arrhenius acids are also Bronsted acids. However, there is a difference in the definition
of bases. Arrhenius theory requires base to the source of OH− ions in aqueous medium,
but Bronsted theory requires base to be a proton acceptor. Hence Arrhenius bases may
not be Bronsted bases. For example, NaOH is a base according to Arrhenius theory
because it gives OH− ions in aqueous solution, but NaOH does not accept proton as such.
Hence it may not be classified as a base according to Bronsted theory.
Strengths of acids and bases
Strength of an acid is measured in terms of its tendency to lose proton whereas strength
of a base is measured in terms of its tendency to accept proton. The conjugate base of a
strong acid is a weak base.
HCl(aq) + H2 O(l)
⇌
H3O+(aq) + Cl−(aq)
strong acid
weak base
On the other hand, conjugate base of a weak acid is a strong base.
CH3COOH(aq) + H2O(l)
⇌
H3O+ (aq) + CH3COO−(aq)
weak acid
strong base
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A base is considered to be strong if it has great tendency to accept a proton. Therefore,
conjugate acid of a strong base has little tendency to lose proton and hence is a weak
acid.
CH3COOH(aq) + OH−(aq)
⇌
H2O(l) + CH3COO−(aq)
strong base
weak acid
On the other hand, conjugate acid of a weak base is a strong acid.
HCl(aq) + H2O(l)
⇌
H3O+ (aq) + Cl−(aq)
weak base
strong acid
The strength of acids or bases is experimentally measured by determining its ionisation
or dissociation constants.
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THE LEWIS ACIDS AND BASES
Although Bronsted-Lowry theory was more general than Arrhenius theory of acids and
bases , but failed to explain the acid base reactions which do not involve transfer of
protons. For example it fails to explain how acidic oxides such as anhydrous CO 2 , SO2 ,
SO3 etc. can neutralise basic oxides such as CaO, BaO etc. even in absence of solvent.
Lewis proposed a more general definition for acids and bases, which do not require the
presence of protons to explain the acid-base behaviour.
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Accoding to Lewis concept :
An acid is a substance which can accept a pair of electrons.
A base is a substance which can donate a pair of electrons .
Acid-base reactions according to this concept involve the donation of electron pair by a
base to an acid to form a co-ordinate bond. Lewis bases can be neutral molecules such
as :
having one or more unshared pairs of electrons. , or anions such as : −CN− , −OH− , −Cl− ,
etc.
Lewis acids are the species having vacant orbitals in the valence shell of one of its
atoms.
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SOME IMPORTANT TERMS
HA ⇌ H+ + A[
][
]
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Ka =
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HA → All acids, specially monobasic acids
BOH → All bases/alkalies
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Higher the Ka, stronger is the acid
Similarly base is represented by BOH
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BOH ⇌ B+ + OH-
Kb =
[
[
][
]
]
Higher the Kb, stronger is the base
For salts
B+A- → Represents salts of any type
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SALT HYDROLYSIS
Pure water is a weak electrolyte and neutral in nature, i.e., H + ion concentration is
exactly equal to OH" ion concentration
[H+] = [OH-]
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When this condition is disturbed by decreasing the concentration of either of the two
ions, the neutral nature changes into acidic or basic.
When [H+] > [OH-], the water becomes acidic and when [H+] < [OH-], the water acquires
basic nature. This is exactly the change which occurs during the phenomenon known as
salt hydrolysis. It is defined as a reaction in which the cation or anion or both of a salt
react with water to produce acidity or alkalinity.
Salts are strong electrolytes. When dissolved in water, they dissociate almost
completely into ions. In some salts, cations are more reactive in comparison to anions
and these react with water to produce H+ ions. Thus, the solution acquires acidic nature.
M+ + H2O ⇌ MOH + H+
Weak base
In other salts, anions may be more reactive in comparison to cations and these react
with water to produce OH- ions. Thus, the solution becomes basic.
A- + H2O ⇌ HA + OHWeak acid
The process of salt hydrolysis is actually the reverse of neutralization.
Salt + Water ⇌ Acid + Base
If acid is stronger than base, the solution is acidic and in case base is stronger than acid,
the solution is alkaline. When both the acid and the base are either strong or weak, the
solution is generally neutral in nature.
As the nature of the cation or the anion of the salt determines whether its solution will
be acidic or basic, it is proper to divide the salts into four categories.
1. Salt of a strong acid and a weak base.
Examples: FeCl3, CuCl2, AlCl3, NH4Cl, CuSO4, etc.
2. Salt of a strong base and a weak acid.
Examples: CH3COONa, NaCN, NaHCO3, Na2CO3, etc.
3. Salt of a weak acid and a weak base.
Examples: CH3COONH4, (NH4)2CO3, NH4HCO3, etc.
4. Salt of a strong acid and a strong base.
Examples: NaCl, K2SO4, NaNO3, NaBr, etc.
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OSTWALD’S DILUTION LAW
According to Arrhenius theory of electrolyte dissociation, the molecules of an
electrolyte in solution are constantly splitting up into ions and the ions are constantly
reuniting to form unionized molecules. Therefore, a dynamic equilibrium exists between
ions and unionized molecules of the electrolyte in solution.
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It was pointed out by Ostwald that like chemical equilibrium, law of mass action van be
applied to such systems also.
Consider a binary electrolyte AB which dissociates into A + and B- ions and the
equilibrium state is represented by the equation:
AB ⇌ A+ + BInitially t = o
C
0 0
At equilibrium C(1-α) Cα Cα
So, dissociation constant may be given as
K = [A+][B-]/[AB] = (Cα * Cα)/C(1-α)
= Cα2 /(1-α)
....... (i)
For very weak electrolytes,
α <<< 1, (1 - α ) = 1
.·.
K = Cα2
α = √K/C
....... (ii)
Concentration of any ion = Cα = √CK .
From equation (ii) it is a clear that degree of ionization increases on dilution.
Thus, degree of dissociation of a weak electrolyte is proportional to the square root of
dilution.
LIMITATIONS OF OSTWALD'S DILUTION LAW:
The law holds good only for weak electrolytes and fails completely in the case of strong
electrolytes. The value of 'α' is determined by conductivity measurements by applying
the formula Λ/Λ∞. The value of 'α' determined at various dilutions of an electrolyte
when substituted in Eq. (i) gives a constant value of K only in the case of weak
electrolytes like CH3COOH, NH4OH, etc. the cause of failure of Ostwald's dilution law in
the case of strong electrolytes is due to the following factors"
I. The law is based on the fact that only a portion of the electrolyte is dissociated
into ions at ordinary dilution and completely at infinite dilution. Strong
electrolytes are almost completely ionized at all dilutions and Λ/Λ ∞ does not give
accurate value of 'α'.
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II.
When concentration of the ions is very high, the presence of charges on the ions
appreciably effects the equilibrium. Hence, law of mass action its simple form
cannot be strictly applied in the case of string electrolytes.
SOME EXAMPLES
Example 1: A 0.01 M solution of acetic is 5% ionized at 25o C. Calculate its dissociation
constant.
According to Ostwald's dilution law
Kα = α2/(1-α)V
Hence,
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α = 0.05, V = 1/0.01 = 100 litres
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Solution:
Ka = 0.05 * 0.05/ (1-0.05)100 = 2.63 * 10-5
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HA ⇌ H+ + A-
Applying Ostwald's dilution law of a weak acid,
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Solution:
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Example 2: Calculate the H+ ion concentration of a 0.02 N weak monobasic acid. The
value of dissociation constant is 4.0 × 10-10.
α = √kaV
Ka= 4.0 ×10-10,
V = 1/0.01 = 100 litres
α = √(4 * 10-10 * 102) = 2 * 10-4
Concentration of hydrogen ions
a/√V = (2*10-4)/100 = 2*10-6 mol L-1
or Concentration of hydrogen ions
= √(CK) = √(0.01 * 4 *10-10) = 2 * 10-6mol L-1
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BUFFER SOLUTIONS
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For several purposes, we need solutions which should have constant pH. Many
reactions, particularly the biochemical reactions, are to be carried out at a constant pH.
But it is observed that solutions and even pure water (pH = 7) cannot retain the constant
pH for long. If the solution comes in contact with air, it will absorb CO 2 and becomes
more acidic. If the solution is stored in a glass bottle, alkaline impurities dissolve from
glass and the solution becomes alkaline.
A solution whose pH is not altered to any great extent by the addition of small
quantities of either an acid (H+ ions) or a base (OH- ions) is called the buffer solution. It
can also be defined as a solution of reserve acidity or alkalinity which resists change of
pH upon the addition of small amount of acid or alkali.
General characteristics of a buffer solution
a. It has a definite pH, i.e., it has reserve acidity or alkalinity.
b. Its pH does not change on standing for long.
c. Its pH does not change on dilution.
d. Its pH is slightly changed by the addition of small quantity of an acid or a base.
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Buffer solutions can be obtained:
1. by mixing a weak acid with its salt with a strong base, eg;
i. CH3COOH + CH3COONa
ii. Boric acid + Borax
iii. Phthalic acid + Potassium acid phthalate
2. by mixing a weak base with its salt with a strong acid, e.g;
i. PNH4OH + NH4Cl
ii. Glycine + Glycine hydrochloride
3. by a solution of ampholyte. The ampholytes or amphoteric electrolytes are the
substances which show properties of both an acid and a base. Proteins and
amino acids are the examples of such electrolytes.
4. by a mixture of an acid salt and a normal salt of a polybasic acid, e.g.,
Na2HPO4 + Na3PO4, or a salt of weak acid and a weak base, such as CH3COONH4.
The first and second types are also called acidic and basic buffers respectively.
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Explanation of buffer action
A. Acidic buffer:
Consider the case of the solution of acetic acid containing sodium acetate. Acetic
acid is feebly ionised while sodium acetate is almost completely ionised. The
mixture thus contains CH3COOH molecules, CH3COO- ions, Na+ ions, H+ ions and
OH- ions. Thus, we have the following equilibria in solution:
CH3COOH ⇌ H+ + CH3COO- (Feebly ionised)
CH3COONa ⇌ Na+ + CH3COO- (Completely ionised)
H2O ⇌ H+ + OH(Very feebly ionised)
When a drop of strong acid, say HCl, is added, the H+ ions furnished by HCl
combine with CH3COO- ions to form feebly ionised CH3COOH whose ionisation is
further suppressed due to common ion effect. Thus, there will be a very slight
effect in the overall H+ ion concentration or pH value.
When a drop of NaOH is added, it will react with free acid to form undissociated
water molecules.
CH3COOH + OH- ⇌ CH3COO- + H2O
Thus, OH- ions furnished by a base are removed and pH of the solution is
practically unaltered.
B. Basic buffer:
Consider the case of the solution containing NH4OH and its salt NH4Cl. The
solution will have NH4OH molecule, ions, Cl- ions, OH- ions and H+ ions
NH4OH ⇌ NH+4 + OH(Feebly ionized)
NH4Cl ⇌ NH4 + Cl
(Completely ionized)
+
H2O ⇌ H + OH
(Very feebly ionized)
When a drop of NaOH is added, the added OH- ion combine with NH4 ions to
form feebly ionised NH4OH whose ionisation is further suppressed due to
common ion effect. Thus pH is not disturbed considerably.
NH4 + OH- ⇌ NH4OH
(From strong base)
When a drop of HCl is added, the added H+ ions combine with NH4OH to form
undissociated water molecules.
NH4OH + H+ ⇌ NH+4 + H2O
(From strong acid)
Thus pH of the buffer is practically unaffected.
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DIAGRAMATIC REPRESENTATION OF BUFFER ACTION
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Base Buffer (NH4OH + NH4Cl)
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Acid Buffer (CH3COOH + CH3COONa)
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