Ionic Equilibrium
Equilibrium established
between the ions and the undissociated electrolyte in a solution is called the
ionic equilibrium. In other words, Ionic equilibrium is the study of solutions
of electrolyte.
Electrolyte: Substances which furnish
ions in aqueous solution or in fused state is called the electrolytes. These
substances conduct electricity in their aqueous solutions or in their fused
state.
Degree of dissociation (α): The number of molecules
dissociated out of total number of molecules originally taken is called the
degree of dissociation.
Mathematically,
It is expressed in
percentage or as fraction or decimals. If 40 molecules are dissociated out of
100 molecules than the degree of dissociation is expressed as 40% or 0.4.
Strong
Electrolytes: These substances dissociate almost completely into ions in aqueous
solutions and hence are good conductor of electricity. Ex: KCl, NaCl, KOH, NaOH
etc.
Weak Electrolytes: These
substances dissociate to a small extent in aqueous solutions and accordingly
conduct electricity to a small extent. Ex: NH4OH, CH3COOH
etc.
A
strong electrolyte dissociates completely in aqueous solution, therefore their
ionization is represented by single right handed arrow which suggests that the
solution contains almost ions.
HCl +H2O → H3O+ +
Cl —
Here the H+ ion from HCl
is received by water to form H3O+. Thus here we may consider
H+ and H3O+
the same.
Similarly, NaOH
(aq.) → Na+ (aq) + OH— (aq.)
But
a weak electrolyte dissociates to a small extent. Some of the ions formed react
back to give undissociated electrolyte and hence equilibrium is established in
this process. Such an equilibrium is called the ionic equilibrium. The
equilibrium constant in this case is called the ionisation constant. Thus the
ionisation of a weak electrolyte AB can be represented as:
AB ⇌ A+ + B—
The
equilibrium or ionisation constant for this equilibrium is represented as:
Ki = [A+] . [ B—] / [AB], where [AB] represents concentration of AB in mol/ltr and so on
Ionisation
of weak Acid - Ostwald’s Dilution Law:
Consider the dissociation of a weak
acid such as CH3COOH. Let its degree of dissociation is α. It means α mole of the acid will
dissociate out of each 1 mole initially taken. Thus if C mole of
the acid is initially taken then Cα moles of the acid will be dissociated. In the following dissociation, the concentration of water is in excess and hence can be considered as constatnt. Accordingly,
Applying law of chemical equilibrium and considering H+ and H3O+ the same,
K = {[CH3COO—] . [H+]} / {[CH3COOH] . [H2O]}
But K . [H2O] = Constant = Ka
Where Ka = Acid constant or dissociation constant of acid
=> Ka = [CH3COO—] . [H+] / [CH3COOH]
=> Ka =
(Cα . Cα) /
{C (1 — α)}
= Cα2 / {1 — α}
In case of weak
electrolyte, the value of α is very small and can be neglected as compared to 1.
i.e., 1 — α =
1
=> Ka =
Cα2 ............... eq. 1
=> α = (Ka /C)1/2
Hence,
Ostwald’s dilution Law states that the degree of dissociation of weak
electrolyte is inversely proportional to the square root of its concentration.
We know, C = n/V.
If n = 1 mole then, C = 1/V
=> α = (Ka .V)1/2
The [H+] =
C α = (KaC)1/2
In a similar way the
degree of dissociation of a weak base, NH4OH (a weak electrolyte)
can be written as,
=> α = (Kb / C)1/2
= (Kb .V)1/2,
where Kb =
Disociation constant of base
=> [OH—] =
C α = (KbC)1/2
Relative strength of weak acids:
Note: Form eq. 1, it is clear that, higher
is the value of acid constant of an acid higher is the value of α (degree of dissociation) and thus
its relative strength to furnish H+ ion is stronger. Thus an acid with higher
acid constant is relatively a stronger acid.
Similar conclusion can be drawn between two bases.
Concept of Acids and Bases:
Arrhenius Concept:
This concept proposed by Swedish
chemist Svante Arrhenius in 1884, is one of the earliest theories explaining
the behavior of acids and bases in aqueous solutions. According to this
concept:
Arrhenius Acid:
An acid is defined as a substance
that, when dissolved in water, increases the concentration of hydrogen or
hydronium ions (H₃O⁺).
The typical chemical equation for the dissociation of an Arrhenius acid in water is:
Acid → H⁺ + Anion
For example, hydrochloric acid (HCl)
dissociates in water as follows:
HCl → H⁺ + Cl⁻
Arrhenius Base:
A base is defined as a substance
that, when dissolved in water, increases the concentration of hydroxide ions
(OH⁻).
The typical chemical equation for the
dissociation of an Arrhenius base in water is:
Base → OH⁻ + Cation
For example, sodium hydroxide (NaOH)
dissociates in water as follows:
NaOH → Na⁺ + OH⁻
Neutralization Reaction:
When an Arrhenius acid reacts with an
Arrhenius base, a neutralization reaction occurs, giving rise to the formation
of water and a salt.
The general form of the
neutralization reaction is:
Acid + Base → Salt + Water
For example, the reaction between
hydrochloric acid (HCl) and sodium hydroxide (NaOH) is a neutralization
reaction:
HCl + NaOH → NaCl + H₂O
While the Arrhenius concept is useful
for understanding acid-base behavior in aqueous solutions, it has limitations.
It is primarily applicable to substances that ionize or dissociate in water,
and it does not account for the behavior of acids and bases in non-aqueous
solvents. Later theories, such as the Brønsted-Lowry and Lewis concepts,
provide broader and more versatile definitions of acids and bases.
Bronsted – Lowry Concept:
This concept was formulated
independently by Danish chemist Johannes Brønsted and English chemist Thomas
Lowry in 1923, is a more general and versatile theory compared to the Arrhenius
concept. According to the Brønsted-Lowry definition:
Brønsted-Lowry Acid:
An acid is a substance that can
donate a proton (H⁺ ion) to another substance.
The acid is often referred to as a
proton donor.
Brønsted-Lowry Base:
A base is a substance that can accept
a proton (H⁺ ion) from another substance.
The base is often referred to as a
proton acceptor.
In the Brønsted-Lowry theory, the
acid-base reaction involves the transfer of protons from the acid to the base.
This concept extends the understanding of acids and bases beyond aqueous
solutions and allows for the characterization of acid-base reactions in various
solvents.
General Reactions:
An acid (HA) donates a proton to a
base (B):
HA→A⁻+H⁺
A base (B) accepts a proton from an
acid (HA):
HA+B→A⁻+HB⁺
Examples:
Acid-Base Reaction in Water:
HCl+H₂O→Cl⁻+H₃O⁺
In this reaction, HCl (hydrochloric
acid) donates a proton to water, forming chloride ion (Cl⁻) and hydronium ion
(H₃O⁺).
Ammonia Acting as a Base:
NH₃+H₂O→NH₄⁺+OH⁻
In this reaction, ammonia (NH₃)
accepts a proton from water, forming ammonium ion (NH₄⁺) and hydroxide ion
(OH⁻).
Conjugate Acid – Base Pair:
In this concept, a conjugate
acid-base pair consists of two substances related to each other by the transfer
of a proton (H⁺ ion) during an acid-base reaction. Specifically:
Conjugate Acid
is the species formed when a base accepts a proton (H⁺).
Example: In
the reaction NH₃+H₂O→NH₄⁺+OH⁻, NH₄⁺ is the conjugate acid of NH₃.
Conjugate Base
is the species formed when an acid donates a proton (H⁺).
Example: In
the reaction HCl+H₂O→Cl⁻+H₃O⁺, Cl⁻ is the conjugate base of HCl.
The relationship between a conjugate
acid-base pair can be represented using the following general reaction:
Acid+Base⇌Conjugate Base+Conjugate Acid
For example, considering the reaction
of acetic acid (CH₃COOH) with water:
CH₃COOH+H₂O⇌CH₃COO⁻+H₃O⁺
Here, CH₃COOH is the acid, H₂O is the base, CH₃COO⁻ is the conjugate base, and H₃O⁺ is the conjugate acid.
Note: If an acid is strong its
conjugate base must be weak so that it does not react back with the H+
ion to give rise to the same acid again. Similarly if a base is strong then its
conjugate acid should be weak and vice versa.
The Brønsted-Lowry concept is more
inclusive than the Arrhenius theory because it applies to a wider range of
solvents and allows for a more comprehensive understanding of acid-base
reactions.
Lewis Concept:
This concept was proposed by American
chemist Gilbert N. Lewis in 1923. Unlike the Brønsted-Lowry and Arrhenius
theories, which focus on proton transfer, the Lewis theory defines acids and
bases based on electron pair interactions. According to the Lewis definition:
Lewis Acid:
A Lewis acid is a substance that can
accept an electron pair.
Lewis acids are electron pair
acceptors.
All cations (such as Na+) and
electron deficient species (molecules having incomplete octate such as BF3)
are Lewis acids.
Lewis Base:
A Lewis base is a substance that can
donate an electron pair.
Lewis bases are electron pair donors.
All anions (such as Cl-) and molecules having lone pairs of electrons (such as NH3, H2O) are Lewis bases.
In this theory, the acid-base
interaction is characterized by the sharing or transfer of electron pairs.
Lewis acids and bases can form coordinate covalent bonds, where both electrons
in the bond come from the Lewis base.
General Reactions:
Formation of a Coordinate Covalent
Bond: The Lewis acid accepts an electron pair from the Lewis base to form a
complex.
Examples:
Reaction between Boron Trifluoride
(BF₃) and Ammonia (NH₃):
BF₃+NH₃→BF₃NH₃ (i.e., H3N:
--->BF3)
In this reaction, BF₃ acts as a Lewis
acid by accepting a lone pair of electrons from NH₃, which acts as a Lewis
base.
Formation of Hydronium Ion (H₃O⁺):
H₂O+H⁺→H₃O⁺
In this reaction, H⁺ acts as a Lewis
acid by accepting a lone pair of electrons from H₂O, which acts as a Lewis
base.
The Lewis acid-base theory is more
inclusive than the Brønsted-Lowry and Arrhenius theories because it doesn't
rely on the presence of protons or the requirement of aqueous solutions.
Ionisation
constant or Ionic Product (KW) of water:
Water undergoes a process of self-ionization,
where a small fraction of water molecules dissociate into hydronium ions (H₃O⁺)
and hydroxide ions (OH⁻). The ionization constant of water is expressed by the
following equilibrium reaction:
H2O⇌H++OH−
Applying law of chemical equilibrium,
K={[H+].[OH−]}/[
H2O]
Since water is a very weak electrolyte, the
concentration of undissociated water [H2O]
is almost constant.
Thus K[
H2O] = Constant = KW = [H+].[OH−]
KW is called the ionic product of
water or Ionisation constant of water.
The equilibrium constant (Kw) for this
reaction is defined as the product of the concentrations of hydronium ions ([H3O+] or [H+])
and hydroxide ions in water at a given temperature:
At 25 degrees Celsius, the ionization constant
of water (Kw) is approximately 1.0×10−14 mol2/L2
The expression for Kw at any temperature is
given by the product of the concentrations of H⁺ and OH⁻ ions, and it can be
used to calculate one of the concentrations if the other is known.
The value of Kw changes with change in temperature.
pH :
The term "pH" stands for ‘’power
of hydrogen’’ or "potential of hydrogen" which is a measure of the
acidity or basicity of a solution or to quantify the concentration of hydrogen
ions (H⁺) in a solution.
. Mathematically, it is defined as
the negative logarithm of hydrogen ion concentration in moles per litre.
pH = - log [H+]
pH scale:
The pH scale is a measure used to
specify the acidity or basicity (alkalinity) of a solution. It is a logarithmic
scale that ranges from 0 to 14, with 7 considered neutral. The pH scale is
based on the concentration of hydrogen ions (H⁺) in a solution. Here's an
overview of the pH scale:
For a neutral solution, [H+]
= [OH-]
and we know [H+] X [OH-]
= 10 -14
Thus [H+] = [OH-]
= 10 -7
Thus pH = - log [H+] = - log 10 -7 = 7
for Neutral solution
For Acidic solution, [H+]
> 10 -7
That means, [H+] = 10 -6,
10 -5, 10 -4 and so on
Thus for acidic solution, pH is less than 7
Similarly, for basic solution, pH is greater than 7
A solution with a pH of 1 is highly
acidic.
A solution with a pH of 7 is neutral.
A solution with a pH of 14 is highly
basic.
Accordingly, we can write pOH = -log [OH-]
Also pH + pOH = 14
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