Ionic equilibrium part 1

Ionic Equilibrium
Part 1
Dr. Damodar Koirala
Amarsingh Model Secondary School
Pokhara Nepal
1

Content
Part I (This Part)
 Degree of ionization
 Dissociation constant
 Arrhenius concept
 Ostwald dilution law
 Acid-base concept
2 Dr. Damodar Koirala | koirala2059@gmail.com
 Acid-base concept
Part II (Next Part)
 Hydrolysis of salt
 Solubility product
 Common ion effect
 Buffer solution

Terminologies
 Non-electrolyte: a substance which does not split into its
ions when dissolved in water. eg: sugar, benzene, phenol
 Electrolyte:A substance which splits into its ions when
dissolved in water: NaCl, HCl and CH3COOH
 Strong electrolyte: electrolyte that completely ionizes
 Strong electrolyte: electrolyte that completely ionizes
when dissolved in water. Eg: NaCl, KCl, all strong acids
/bases
 Weak electrolyte:An electrolyte which is partially ionized
when dissolved in water is called weak electrolyte. For
examples, all weak acids/ bases.

Ionization of strong electrolyte
 Strong electrolyte undergo complete ionization in aqueous solution
 The concentration of product is calculated accordingly. Consider 3
mole of H SO is added to water solvent then
General eg: AB(aq)  A+(aq) + B– (aq)
Specific eg: NaCl(aq)  Na+(aq) + Cl– (aq)
mole of H2SO4 is added to water solvent then
 Since, H2SO4 is strong electrolyte, the H2SO4 dissociates
completely into two H+ ions and one SO4
– – ion
H2SO4(aq)  2 H+(aq) + SO4
– –(aq)
Before: 3 mol 0 mol 0 mol
After: 0 mol 6 mol 3 mol

Ionization of weak electrolyte
 A weak electrolyte undergoes ionization in aqueous solution
and there exists a dynamic equilibrium between dissociated
and undissociated species of the electrolyte in aqueous
solution at given temperature
 The dissociation reaction is represented by equilibrium arrow
Let AB be the weak electrolyte, then
 Let AB be the weak electrolyte, then
 Since AB is weak electrolyte, only some of it dissociates. The
factor like degree of ionization or ionization constant of AB
helps to determine the concentration of A+ and B-
AB(aq) A+(aq) + B-(aq)

 It is defined as the fraction of total number of molecules
dissolved in the solution that undergoes ionization at a
given temperature
Degree of ionization
Degree of ionization (α) = Number of molecules ionized
 Its value equals to 1 for strong electrolyte since their
ionization is complete
 Its value is less than 1 for weak electrolyte since their
ionization is incomplete
Total number of molecules dissolved

NH4OH(aq) NH4
+(aq) + OH-(aq)
Q: Consider the degree of ionization of NH4OH be 6.0 * 10-3. If 10
mol of NH4OH is dissolved in water then how many moles of
NH4
+ ions are produced.
Solution: Since, NH4OH is weak electrolyte, it dissociates partially
Degree of ionization
We know that,
Degree of ionization (α) = Mole ionized
Total mole dissolved
Mole of NH4OH ionized = α * Total mole dissolved
= 6.0 * 10-3 * 10 = 0.06 mol
Hence, 0.06 mol of NH4
+ ions are produced

Factors affecting degree of ionization
 Nature of solvent: The solvent weaken the electrostatic
force of attraction between the ions, the higher the
power to weaken the electrostatic force of attraction the
more will be the degree of ionization. It is measured in
dielectric constant
 Nature of electrolyte: Strong electrolyte ionizes
completely at all dilution therefore the degree of
ionization is 100%. But in case of weak electrolyte the
degree of ionization is different. It depends upon the
polarity of the molecule. For eg: the degree of ionization
of formic acid is greater than that of acetic acid

Factors affecting degree of ionization
 Temperature: Increasing the temperature increases the
degree of ionization of an electrolyte. At high
temperature the molecular velocity overcomes the force
of attraction between the ions of electrolyte and
consequently the degree of ionization increases
 Concentration of solution:The degree of ionization of
 Concentration of solution:The degree of ionization of
an electrolyte is inversely proportional to the
concentration. It means the higher the concentration, the
lower will be the degree of ionization and vice-versa
 Common ion effect: the common ion from strong
electrolyte reduces the degree of ionization of weak
electrolyte decreases

Ionization constant
 Consider the ionization of a weak binary electrolyte(AB), at
equilibrium,
 Applying, all of mass action,
AB(aq) A+(aq) + B-(aq)
 The ionization constant (K) of a weak electrolyte is the ratio
of the product of the concentration of ionic species to the
concentration of unionized species present at the equilibrium
at the given temperature
K = [A+][B-]
[AB]

Arrhenius theory of ionisation
 Put forward by Svante August Arrhenius in 1887
Postulates:
1. An electrolyte when dissolved in water or any other polar
solvet ionizes to produce ion in solution, these ions are ffree
to move throughout bulk of solution
2. In any solution of any electrolyte, the net positive charge is
equal to net negative charge. Hence, the electrolyte is neutral
3. These ions tends to recombine to form unionized
electrolyte.Therefore, their exist a equilibrium between
ionized ion and unionized molecule. It is dynamic in nature. It
is characterized by ionization consant

Arrhenius theory of ionisation
5. The properties of an electrolyte in the solution are due to
the ions present in it for eg: blue color of CuSO4 solution is
due to cupric ion “Cu2+”. Similarly the green colour of FeSO4
solution is due to “Fe2+” ion
6. The fraction of the total number of moles present as free
ions in the solution is known as degree of ionization. It is
ions in the solution is known as degree of ionization. It is
denoted by alpha (α). It is given by the relation
Degree of ionization (α) = Number of molecules ionized
Total number of molecules dissolved

Ostwald dilution law
 Lets consider a binary electrolyte AB which undergoes
ionization in solution as A+ and B-
 Let α be the degree of ionization of AB and C moles/L be
initial concentration of AB.Then,
AB A+ + B-
 The ionization constant for this system is,
AB A + B
Initial C 0 0
Change - αC +αC +αC
At equilibrium C- αC αC αC
K = [A+][B-]
[AB]

Ostwald dilution law
 Now, K = αC x αC
C- αC
K = α2C
1- α
 This equation gives the relationship between degree of
ionization and ionization constant.
 For weak electrolyte, α <<1 and 1-α is almost equal to zero.
 For weak electrolyte, α <<1 and 1-α is almost equal to zero.
 Hence,
 The law states that an electrolyte in solution gets ionized to a
larger extent with increasing dilution i.e. the degree of
ionization must increase upon dilution
K = α2C or, α = (K/C)1/2

Numerical based on dilution law
 Calculate the degree of ionization and concentration of
H3O+ of 0.1 mole/L solution of acetic acid (Ionization
constant of acetic acid (CH3COOH) is 1.8 * 10-5 M)

Concept of acid and base
 Classical concept
 Acid are substances which react with active metal to give
hydrogen. They turns blue litmus paper into red
 Base are substances which has bitter taste.They turn red
litmus paper into blue
 Modern concept:There is no general definition of acids and
bases.
 Arrhenieus concept (water centered)
 Bronsted-lowry concept (proton centered)
 Lewis concept (electron pair centered)

Arrhenius acid
 An acid is a substance that is capable of producing hydrogen
ion (H+ ) when dissolved in water
 Acids which ionise completely are strong acid and which
ionises partially are weak acid
HCl(aq)  H+(aq) + Cl-(aq)
HNO3(aq)  H+(aq) + NO3-(aq)
H2SO4(aq)  2H+(aq) + SO4--(aq)
CH3COOH(aq)  H+(aq) + CH3COO-(aq)

Arrhenius base
 A base on the other hand is defined as a substance capable of
providing a hydroxyl ion (HO) on dissociation in aqueous
solutions.
 Bases which ionise completely are strong base and which
ionises partially are weak base
NaOH(aq)  Na+(aq) + OH-(aq)
KOH(aq)  K+(aq) + OH-(aq)
NH4OH(aq)  NH4+(aq) + OH-(aq)
Cu(OH)2(aq)  Cu2+(aq) + OH-(aq)

Limitation of Arrhenius concept
 Arrhenius concept is quite useful and explains the acid- base
behaviour to a good extent. However it has certain drawbacks
like,
 It is limited to only aqueous solutions and require dissociation
of the substance.
 It does not explain the acidic behaviour of some substances
which do not contain hydrogen. for example,AlCl3 .
 It does not explain the basic character of substances like NH3
and Na2CO3 which do not have a hydroxide groups.
 It does not explain the acidic and basic character of oxide

Bronsted concept of acid and base
 According to them, an acid is defined as a proton (H+)
donor, and a base is defined as a proton acceptor.
 The definition is sufficiently broad and removes the first
limitation of Arrhenius concept.
 Any hydrogen-containing molecule or ion capable of
donating or transferring a proton is an acid, while any
molecule or ion that can accept a proton is a base.

Water can act as bronsted acid / base
 Water as bronsted base
Here, water accepts a proton (H+), hence is bronsted base
CH3COOH + H2O CH3COO- + H3O+
Bronsted acid Bronsted base
 Water as bronsted acid
Here, water donates a proton (H+), hence is bronsted acid
CO3-- + H2O HCO3- + OH-
Bronsted acid Bronsted acid

Conjugated acid-base pair
 Bronsted acid react with bronsted base to give the conjugated
acid of bronsted base and conjugate base of Bronsted acid
respectively
 An acid and a base which differ only by a proton are called
a conjugate acid-base pair.Thus NH3 is called the conjugate
base of NH4
+, and NH4
+ is the conjugate acid of NH3. Similarly,
base of NH4
+, and NH4
+ is the conjugate acid of NH3. Similarly,
HCl is the conjugate acid of Cl–, and Cl– the conjugate base of
HCl.
HCl + NH3  NH4
+ + Cl-
Acid 1 Base 2 Acid 2 Base 1
Conjugated pair
Conjugated pair

Conjugated acid-base pair
 Strong acid has weak conjugate base
 Weak acid has strong conjugate base
 Strong base has weak conjugate acid
 Weak base has strong conjugate base

Problems: Conjugated acid-base pair
 Write the conjugated acid and base of NH3
 Which of the following are not acid-base pair, explain briefly
A. H2SO4 and SO4
– –
B. HBr and Br –
C. Ca(OH)2 and Ca(OH)+
2
D. CH4 and CH5
+
 What is the conjugate acid or the conjugate base of
1. HCl
2. CH3NH2
3. OH–
4. HCO3
–

Limitation of Bronsted concept
 It does not explain the acidic behavior of some
substances which do not contain hydrogen. eg:AlCl3,
BF3, FeCl3
 It does not explain the reaction between acidic oxide
 It does not explain the reaction between acidic oxide
and basic oxide

Advantages of Bronsted concept over
Arrhenius concept
 Bronsted concept is not limited to molecules only
but it can explain acidic and basic nature of ions
 It can explain the basic characteristic of the substance
like Na2CO3, NH3 etc
like Na2CO3, NH3 etc
 It can explain the acid base reaction in non-aqueous
medium or in the absence of solvent

Lewis concept of acid and base
 Those substances which has a tendency to accept lone
pair of electron are acid
 Those substance which has tendency to donate lone pair
of electron are base
B: A B+A-
B: A B+A-
Electron donor
Lewis base
Electron acceptor
Lewis acid
H+
O
H
H
O
H
H
H
+

Advantage of Lewis concept
 It is most general out of all the concept and can
explain acidic and basic nature of all those substance
which could not explain by earlier concept
 It can explain the acid-base reaction which other
 It can explain the acid-base reaction which other
concept cant explain

Limitation of lewis concept
 Lewis concept consider the formation of coordinate bond
during acid base reaction however
 In well known acid base reactions, there is no coordinate
bond
 Acid-base reaction is fast but formation of coordinate bond
 Acid-base reaction is fast but formation of coordinate bond
is slow
 Metals like Ni, Fe forms coordinatebond but are not
considered as acid-base reaction
 It can not explain the relative strength of acid and base

All Lewis bases are bronsed bases.
 According to the Bronsted concept bases are the substance
which can accept proton in aqueous solution
eg: NH3 + H2O  NH4
+ + OH-
In the relation ammonia accept proton and become NH4
+
In the relation ammonia accept proton and become NH4
+
 According to the Lewis concept, bases are the substances
which has tendency to donate lone pair of electron
eg: NH3 + H+  NH4
+
In this reaction, ammonia donates the lone pair electron to H+

pKa value of acid
 Negative log with base 10 of Ka of acid
pKa = - log (Ka)
pKa = - log (Ka)
 Higher Ka value lower is pKa value and
stronger is the acid
 It is a single number with no exponent.
Hence easy to compare the value

pKb value of base
 Negative log with base 10 of Kb of base
pKb = - log (Kb)
pKb = - log (Kb)
 Higher Kb value lower is pKb value and stronger is
the base
 It is single number with no exponent. Hence easy
to compare values

Ionization of water
 Water can act both as a weak acid as well a weak base. In a
sample of water a small number of water molecules undergo
autoionisation, in which half the ions act as an acid while the
other half acts as a base.
H2O + H2O H3O+ + OH-
 We can simplify this equation by writing hydronium ions,
H3O+, as simple hydrogen ions, H+
 The equilibrium expression for this reaction is
H2O H + + OH-
Kc = [H+][OH-]
[H2O]

Ionization of water
 The extend of ionization of pure water is very low
 So [H+] and [OH-] are very low
 Due to this, concentration of water, [H2O], is regarded as constant
Now, Kc. [H2O] = [H+] [OH-]
K = [H+] [OH-]
Kw = [H+] [OH-]
Kw is called the ionic product of water.At 298K it is1.0 x 10-14 M2
For ionization of H2O, equal H+ and OH- ions are produced
i.e. [H+] = [OH-]
so, Kw = [H+]2
 Hence, [H+] = [OH-] = 1.00 x10-7 M

pH and pH-scale
 It is convenient to express the concentration of H+ ions on a
logarithmic scale.This is the pH scale.
 The pH of a solution is defines as the negative logarithm to
base 10 of the hydrogen ion concentration in moles per liter.
i.e pH = -log[H+]
 i.e pH = -log[H+]
 pH < 7 is acidic, pH = 7 is netural and pH > is basicWhen it is
below 7, the solution is acidic
 Similarly pOH = -log[OH-]
 At 298K, pH + pOH = 14

pH Numerical
 If Normality (say X) of strong acid is given then use,
 Concentration of hydrogen = [H+] = Normality of acid
 pH = -log [H+] = -log(X)
If Molarity of strong acid is give then
 If Molarity of strong acid is give then
 Convert molarity to normaily ( N = M * Basicity) and use
above mentioned formula
 Or, convert the concentration (in molarity) of acid to the
concentration of hydrogen ion and put it into the formula
pH = -log[H+]

Numerical
 Calculate pH of 0.04N HNO3 solution, assuming HNO3 is
completely ionized.
 Calculate the strength in gram/litre of NaOH whose pH value is
11.
 Calculate the pH of 0.1M acetic acid solution having Ka = 1.8 *
 Calculate the pH of 0.1M acetic acid solution having Ka = 1.8 *
10-5
 What is the pH of solution of NaOH whose concentration is 0.4
gram/litre
 Calculate the pH of an aqueous solution containing 10-7 moles
of NaOH per liter

Numerical
 What mass of KOH should be dissolved in one liter of solution
to prepare a solution having pH 12 at 25 C?
 0.41 gm of NaOH is place in 100 ml of 0.1N H2SO4. Find the
pH of the resulting solution
 Calculate the hydrogen ion concentration of a solution whose
 Calculate the hydrogen ion concentration of a solution whose
pH is 9.5
 Calculate the pH of 0.1M H2SO4
 Calculate the pH of a 0.1M of a solution of acetic acid. At
25oC, the ionization constant of acetic acid is 1.7 x 10-5.

Numerical
 10cc N/2 HCl, 30 cc N/10 HNO3 and 60cc N/5 H2SO4 are
mixed together. Find pH of mixture.
 Calculate the hydroxide ion concentration of a solution having
pH 10.5
 The pH of 0.1M HCN solution is 5.2. What is the value of
 The pH of 0.1M HCN solution is 5.2. What is the value of
ionization constant for the acid.
 The pH of KOH solution is 10. Calculate the hydroxyl ion
concentration.
 At 298 K, 0.1 M solution of CH3COOH acid is 1.3% ionizes.
What is the ionization constant (Ka) for the acid.

Ionic equilibrium part 1

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Ionic equilibrium part 1