Chemistry_UG

1. An amazing fact
Feeding the earth’s enormous (and growing)
population
Simple chemical reaction involving
the reaction of nitrogen with hydrogen to form
ammonia
Chemistry
The Molecular Nature of Matter
Neil D. Jespersen
St. John’s University, New York
James E. Brady
St. John’s University, New York
In collaboration with Alison Hyslop
St. John’s University, New York
100 million tons of ammonia annually, with most of it going to agricultural uses as a
fertilizer

2. THEORETICAL PRINCIPLES OF REACTIONS
IN
SOLUTION EQUILIBRIA
Dr. B. Santosh Kumar
&
Dr.K.V.Nagalakshmi
PHYSICAL CHEMISTRY

3. TOPICS
Law of mass action
Factors effecting chemical reactions in solution
Electrolytic dissociation
Solubility product
Effects of acids and Temperature on precipitates

4. TOPICS
Acid-Base equilibria in water
Oswald dilution law
Strengths of acids and bases
Dissociation of poly proticacid
Titration curve of weak polyprotic acid
Common ion effect
Ionic product of water
Hydrogen ion exponent
Hydrolysis of salts
Degree of hydrolysis
Buffer solution

5. Chemical Equilibrium
Every reversible chemical reaction can reached the
state where the concentrations of products and
reactants remain constant over time and there are
no visible changes in the system.
That particular state is called equilibrium state and
there are no observable changes as times goes by.
Reactants Products

6. Characteristics of chemical equilibrium
• Dynamic in nature
• Constancy in some observable properties
like color, pressure, concentration, density
and temperature etc..
• Equilibrium state can be affected by
altering some factors like pressure,
volume, concentration and temperature.

7. Law of mass action
“The rate of any chemical reaction is
proportional to the product of the active
masses of the reacting substances, with each
mass raised to a power equal to the
coefficient that occurs in the chemical
equation”
aA + bB cC + dD
Cato M.Guldberg and Peter Waage

8. Chemical reactions in solution
Electrolyte (AB) is dissolved in water and
a small
fraction of it dissociates to form ions (A+
and B–).
When the equilibrium has been reached
between the undissociated and the free
ions
AB A+ + B-
The fraction of the amount of the electrolyte in
solution present as free ions is called the Degree
of dissociation.

9. The degree of dissociation (x) =
amount dissociated (mol/L)
initial concentration (mol/L)
On applying the Law of Mass Action
[A+][B+]
[AB]
K =
x =

10. Factors affecting the reactions in solutions
• Nature of Solute
• Nature of the solvent
• Concentration
• Temperature
Strong acids and strong bases, and the salts obtained
by their interaction are almost completely dissociated
in solution. On the other hand, weak acids and weak
bases and their salts are feebly dissociated.
This effect of the solvent is measured by its ‘dielectric
constant’. The dielectric constant of a solvent may be
defined as its capacity to weaken the force of
attraction between the electrical charges
The extent of dissociation of an electrolyte is inversely
proportional to the concentration of its solution.
The higher the temperature greater is the
dissociation. At high temperature the increased
molecular velocities overcome the forces of attraction
between the ions and consequently the dissociation is
great.

11. Solubility product
An ionic solid substance (AgCl) dissolves in water, it
dissociates to give separate cations (Ag+) and anions (Cl-)
As the concentration of the ions in solution increases, they
collide and reform the solid phase.
Dynamic equilibrium is established between the solid phase
and the cations and anions in solution
AgCl Ag+ + Cl-

12. A Saturated solution is a solution in which the dissolved
and undissolved solute are in equilibrium.
A saturated solution represents the limit of a solute’s
ability to dissolve in a given solvent.
This is a measure of the “solubility” of the solute.
The Solubility (S) of a substance in a solvent is the
concentration in the saturated solution.
Solubility of a solute may be represented in grams per
100 ml of solution. It can also be expressed in
moles per litre.
The value of solubility of a substance depends on the
solvent and the temperature.

13. Applying the Law of Mass Action
[Ag+][Cl-]
[AgCl]
K =
The amount of AgCl in contact with saturated
solution does not change with time and the
factor [AgCl] remains the same.
Thus the equilibrium expression becomes
Ksp = [Ag+] [Cl–]
The product [Ag+] [Cl–] in the Ksp expression above is
also called the Ionic Product or Ion Product.

14. The Ksp expression may be stated as the product
of the concentration of ions (mol/l) in the
saturated solution at a given temperature is
constant.
This is called the Solubility product principle.

15. ACID-BASE EQUILIBRIUM IN WATER
Autoionization of Water
Sensitive instruments, pure water is observed to
weakly conduct electricity, indicating the presence of
very small concentrations of ions.
They arise from the very slight self-ionization, or
autoionization

16. On applying Law of Mass Action
water is a pure liquid, with a constant 55.6 molar
concentration, it does not appear in this equilibrium law
Because of the importance of the autoionization
equilibrium,
its equilibrium constant is given the special symbol, Kw,
that is called the ion product constant of water.
According to equilibrium law

17. Omit the water molecule that carries the hydrogen ion
and write H+ in place of H3O+
The equilibrium equation for the autoionization of water
then simplifies to
In pure water, the concentrations of H+ and OH- produced by
the autoionization are equal.
The concentrations have the following values at 25 °C.
[H+] = [OH-] = 1.0 × 10-7 mol L-1
Kw = (1.0 × 10-7) × (1.0 × 10-7) = 1.0 × 10-14
Therefore, at 25 °C,

18. Criteria for Acidic, Basic, and Neutral Solutions
Neutral solution [H3O+] = [OH-]
Acidic solution [H3O+] > [OH-]
Basic solution [H3O+] < [OH-]
consequences of the autoionization of water
In any aqueous solution, there are always both H3O+
and OH- ions, regardless of what solutes are present.
This means in a solution of the acid HCl there is some
OH-, and in a solution of the base NaOH, there is
some H3O+.
We call a solution acidic or basic depending on which
ion has the largest concentration.

19. Ostwald’s Dilution Law
According to the Arrhenius theory of dissociation,
an electrolyte dissociates into ions in water solutions.
These ions are in a state of equilibrium with the
undissociated molecules.
This equilibrium is called the Ionic equilibrium.
Ostwald noted that the Law of Mass Action can be
applied to the ionic equilibrium as in the case of
chemical equilibria.

20. Let us consider a binary electrolyte AB which dissociates
in solution to form the ions A+ and B–.
Let C moles per litre be the concentration of the electrolyte and
α (alpha) its degree of dissociation.
The concentration terms at equilibrium may be written as
[AB] = C (1 – α) mol litre– 1
[A+] = C α mol litre– 1
[B–] = C α mol litre– 1
Applying the Law of Mass Action :

21. The equilibrium constant “Kc” is called the Dissociation constant
or Ionization constant. Constant at a constant temperature
If one mole of an electrolyte be dissolved in V litre of the solution,

22. V is known as the Dilution or the solution.
This expression which correlates the variation of the degree of dissociation
of an electrolyte with dilution, is known as Ostwald’s Dilution Law.
For Weak Electrolytes
For weak electrolytes, the value of α is very small as compared to 1,
so that in most of the calculation we can take (1 – α) = 1.
Thus the Ostwald’s Dilution Law expression becomes
It implies that the degree of dissociation of a weak electrolyte is proportional
to the square root of the dilution

23. For Strong Electrolytes
For strong electrolytes, the value of α is large and it cannot be neglected
in comparison with 1
which gives a quadratic equation
α2 + α KcV – KcV = 0
from this equation the value of α can be evaluated.

24. Limitation of Ostwald’s Law
Ostwald’s Dilution law holds good only for weak
electrolytes and fails completely when applied to
strong electrolytes.
For strong electrolytes, which are highly ionised in
solution, the value of the dissociation constant K, far
from remaining constant, rapidly falls with dilution.

25. Factors that explain the failure of Ostwald’s law in case
of strong electrolytes
(1) The law is based on Arrhenius theory which assumes that only
a fraction of the electrolyte is dissociated at ordinary dilutions
and complete dissociation occurs only at infinite dilution.
(2) The Ostwald’s law is derived on the assumption that the Law
of Mass Action holds for the ionic equilibria as well. But when
the concentration of ions is very high, the presence of charges
affects the equilibrium. Thus the Law of Mass Action in its
simple form cannot be applied.
(3) The ions obtained by dissociation may get hydrated and may
affect the concentration terms. Better results are obtained by
using activities instead of concentrations.

26. For weak acids and bases, the molar concentrations of
H+ and OH- are very small
Writing and comparing two exponential values is not
always easy
A Danish chemist, S. P. L. Sørenson suggested an easier
approach to make comparisons of small values of [H+]
easier, called the pH of the solution
pH = -log [H+]
[H+] = 10-pH
Hydrogen ion exponent

27. pH is that it is a measure of the acidity of a solution. Hence, we may
define acidic, basic, and neutral in terms of pH values.
At 25 °C, in pure water, or in any solution that is neutral,
[H+] = [OH-] = 1.0 × 10-7 M
Therefore, the pH of a neutral solution at 25 °C is 7.00.
Acidic solution [H+] is larger than 10-7 M and so has a pH less than
7.00.
Thus, as a solution’s acidity increases, its pH decreases.
A basic solution is one in which the value of [H+] is less than 10-7 M
and so has a pH
that is greater than 7.00. As a solution’s acidity decreases, its pH
increases. At 25 °C:
pH = 7.00 Neutral solution
pH < 7.00 Acidic solution
pH > 7.00 Basic solution

28. Add a concentrated solution of a soluble lead compound,
such as Pb(NO3)2
increased concentration of Pb2+ in the PbCl2 solution
will drive the position of equilibrium to the left, causing
some PbCl2 to precipitate.
Le Châtelier’s principle the net result being that PbCl2 is
less soluble in a solution that contains Pb2+ from another
source than it is in pure water.
If a concentrated solution of a soluble chloride salt such
as NaCl is added to the saturated PbCl2 solution. The
added Cl- will drive the equilibrium to the left, reducing
the amount of dissolved PbCl2.
Common ion effect

29. Addition of a few drops of concentrated HCl, forced the
equilibrium to shift to the left. This caused some white
crystals of solid NaCl to precipitate. (Michael Watson)
Common ion effect

30. “The reduction of the degree of dissociation of a
salt by the addition of a common-ion is called
the Common-ion effect”
The common ion effect can dramatically lower
the solubility of a salt, as Example 18.5
demonstrates.
Common ion effect

31. Hydrolysis of salts
The salt of a weak acid, HA and a strong base dissolves in
water to form the anion A–. The A– anion tends to react
with water by drawing a proton (H+) from its molecule to
form the unionised molecule.
The reaction of an anion or cation with water
accompanied by cleavage of O–H bond is called
Hydrolysis.
the salt of a weak base, BOH, and a strong acid dissolves in
water to form the cation B+. The cation B+ reacts with water
by accepting OH– ions from its molecule.

32. The term hydrolysis is derived from hydro, meaning
water, and lysis, meaning breaking.
It may be noted that in anionic hydrolysis In solution
becomes slightly basic (pH > 7) due to the generation
of excess OH– ions.
In cationic hydrolysis, there is excess of H+ ions which
makes the solution slightly acidic (pH < 7).
Classification of different salts yypes according to their hydrolytic
behaviour:
(1) Salts of Weak acids and Strong bases
(2) Salts of Weak bases and Strong acids
(3) Salts of Weak acids and Weak bases

33. Why NaCl solution is neutral ?
NaCl dissociates in water to give the anion Cl–. HCl and
Cl– constitute an acid-base conjugate pair.
Since HCl is a strong acid, Cl– is very weak base. Cl– is unable
to accept a proton (H+) from an acid, particularly water. That is
why Cl– does not hydrolyse. It cannot generate OH– ions as
follows:

34. Salts of Weak acids and Strong bases
Eg: Sodium acetate, CH3COONa, and sodium cyanide, NaCN
CH3COO- + H2O ⎯⎯→ CH3COOH + -OH
It ionises in aqueous solution to form the anion CH3COO–. Being
the conjugate base of a weak acid, CH3COOH, it is a relatively
strong base. Thus CH3COO– accepts H+ ion from water and
undergoes
hydrolysis.
CH3COONa. This is a salt of the weak acid, CH3COOH, and strong
base, NaOH.
The resulting solution is slightly basic due to excess OH– ions
present.

35. Salts of Weak bases and Strong acids
Eg: Ammonium chloride, ferric chloride, aluminium
chloride, and copper sulphate.
NH4+ is a Bronsted conjugate acid of the weak base
NH4OH. Therefore, it is a relatively strong acid.
It accepts OH– ion from water (H2O) and forms the
unionised NH4OH and H+ ion.
The accumulation of H+ ions in solution makes it
acidic.

36. Salts of Weak acids and weak bases
Eg: Ammonium acetate, ammonium cyanide and
ammonium fluoride
Both the anion and the cation produced by ionisation
of the salt undergo hydrolysis.
The resulting solution is neutral, basic or acidic
depending on the relative hydrolysis of the anions and
the cations.

37. Many chemical and biological systems are quite sensitive to
pH
pH of your blood should be, within the range of 7.35 to
7.42, either to 7.00 or to 8.00 causes death.
Lakes and streams with a pH less than 5 often cannot
support fish life.
Thus, a change in pH can produce unwanted effects, and
systems that are sensitive to pH must be protected from the
H+ or OH- that might be formed or consumed by some
reaction.
Buffer Solutions
“Buffer solution is one which maintains its pH
fairly constant even upon the addition of small
amounts of acid or base”

38. A buffer contains solutes that enable it to resist large changes in pH
when small amounts of either strong acid or strong base are added
to it.
Ordinarily, the buffer consists of two solutes, one a weak acid and
the other its conjugate base.
If the acid is molecular, then the conjugate base is supplied by a
soluble salt of the acid.
Types of buffer solutions:
(1) a weak acid together with a salt of the same acid with a strong
base. These are called Acid buffers
e.g., CH3COOH + CH3COONa.
(2) a weak base and its salt with a strong acid. These are called
Basic buffers. e.g., NH4OH + NH4Cl.

39. How a Buffer Works ??
Buffer must be able to neutralize either a strong acid or strong
base that is added. This is precisely what the conjugate acid and
base components of the buffer do.
Lets consider a buffer composed of acetic acid, CH3COOH, and
acetate ion, CH3COO-, supplied by a salt such as CH3COONa.
If we add extra H+ to the buffer (from a strong acid) the acetate
ion (the weak conjugate base) can react with it as follows.
H+(aq) + CH3COO-(aq) ----→ CH3COOH(aq)
The added H+ changes some of the buffer’s CH3COO-, to its
conjugate (weak) acid, CH3COOH.
This reaction prevents a large buildup of H+ and a corresponding
decrease in pH.

40. A similarly, the addition of strong base to the buffer. The OH-
from the strong base will react with some CH3COOH.
Thus, one member of a buffer team neutralizes H+ that might
get into the solution, and the other member neutralizes OH-.
CH3COOH(aq) + OH-(aq) ----→ CH3COO-(aq) + H2O
Understanding buffers is an important tool for chemists to
use in applications ranging from the protocols of a research
project to designing a consumer product.
Here the added OH- changes some of the buffer’s acid, CH3COOH, into
its conjugate base, CH3COO-. This prevents a buildup of OH-, which
would otherwise cause large increase in the pH.

41. Calculation of pH of buffer solution
The pH of an acid buffer can be calculated from the dissociation
constant, Ka, of the weak acid and the concentrations of the acid
and the salt used.
This relationship is called the Henderson-Hasselbalch equation or
simply Henderson equation.

42. Loss of the first proton to yield the HSO4- ion is complete, but
the loss of the second proton is incomplete and involves an
equilibrium.
Dissociation of polyprotic acids
Many acids capable of supplying more than one H+ per molecule.
Recall that these are called polyprotic acids.
Eg: sulfuric acid, H2SO4, carbonic acid, H2CO3, and phosphoric
acid, H3PO4.
These acids undergo ionization in a series of steps, each of
which releases one proton.
For weak polyprotic acids, such as H2CO3 and H3PO4, each
step is an equilibrium. Even sulfuric acid, which we consider a
strong acid, is not completely ionized.

43. Let’s consider the weak diprotic acid, H2CO3. In water, the acid
ionizes in two steps, each of which is an equilibrium that
transfers an H+ ion to a water molecule.
As usual, we can use H+ in place of H3O+ and simplify these
equations to give
Each step has its own ionization constant, Ka,
which we identify as Ka1 for the first step and
Ka2 for the second. For carbonic acid,
Each step has its own ionization constant, Ka, which we identify as
Ka1 for the first step and Ka2 for the second.

44. For carbonic acid,

45. Triprotic Acids:
A triprotic acid (H3A) can undergo three dissociations and will
therefore have three dissociation constants: Ka1 > Ka2 > Ka3. Take,
for example the three dissociation steps of the common triprotic
acid phosphoric acid

46. Fractional Concentration of Conjugate Base Species
For example, a generic diprotic acid will generate three species
in solution: H2A, HA–, and A2-, and the fractional concentration
of HA–, which is given by:

47. Thank you…