English chapter we are going to discuss about the reduction in the oxidation their heat evolution changes occurrence and about their reducing agent and oxidization

1. Equilibriums
book 1

2. Introduction INTRODUCTION
 When a liquid evaporates in a closed container
molecules with relatively higher kinetic energy escape
the liquid surface into the vapour phase strike the
liquid surface and are retained in the liquid phase .it
gives rise to constant vapour pressure because of an
equilibrium in which the number of molecules leaving
the liquid equal the number returning to liquid from the
vapour. We say that the system has reached an
EQUILBIRUM STATE.

3.  The mixture of reactants and products is called
equilibrium mixture .
 The stage where there is no change in the concentrations is
called a dynamic equilibrium.
 Equilibriums involving ions are called ionic equilibrium.

4. Equilibriums in physical processes
 The most famous examples of physical equilibrium is the
phase transformation processes.
That is given below:-
solid liquid
⇋
liquid gas
⇋
solid gas
⇋

5. The Concept of Equilibrium
 As a system approaches
equilibrium, both the forward
and reverse reactions are
occurring.
 At equilibrium, the forward
and reverse reactions are
proceeding at the same
rate.

6. A System at Equilibrium
 Once equilibrium is
achieved, the amount
of each reactant and
product remains
constant.
 Here NO2 and N2O4
has reached
equilibrium.

7. Depicting Equilibrium
In a system at equilibrium, both the forward and reverse
reactions are being carried out; as a result, we write its equation
with a double arrow
Here the partial pressures of the above gases are also equal.
N2O4 (g) 2 NO2 (g)
⇋

8. Equilibrium

9. Equilibriums involving dissolution of solid or gases
in liquids
 Solids in liquids:-
When no more solute can be dissolved in the solution it is called a
saturated solution.
That is,
sugar(solution) sugar(solid)
⇋
And also the rate of dissolution of sugar = rate of crystallization of
sugar.
The ratio of radioactive molecules to the nonradioactive molecules
in the solution increases till attains a constant value.

10. Gases in liquids and henry’s law
 There is a certain equilibrium between the
molecules of the gaseous state and molecules
of the liquid state.
Carbon dioxide(gas) carbon dioxide(in solution)
⇋
Henry’s law:-states that the mass of a gas dissolved
in a given mass of a solvent at any temperature is
proportional to the pressure of the gas above the
solvent.
The amount decreases with the increase in
temperature.

11.  (1) for solid ⇋liquid equilibrium , there is only one
temperature(melting point) at 1 atm (1.013 bar) at which the two
phases can coexist. If there is no exchange of heat with the
surroundings , the mass of the two phases remains constant.
 (2) For liquid vapour equilibirum the vapour pressure is constant
⇋
at a given temperature.
 (3) for dissolution of solids in liquids, the solubility is constant at a
given temperature.
 (4) For dissolution of gases in liquids, the concentration of a gas in
liquid is proportional to the pressure of the gas over the liquid.

12. General characteristics of equilibria involving physical
processes
 Equilibrium is only possible only in a closed system at a given
temperature.
 Both the opposing processes occur at the same rate and there
is a dynamic but a stable condition. When equilibrium is attained
for a physical process, it is characterized by constant value of
one of its parameters at a given temperature The magnitude of
such quantities at any stage indicates the extent to which the
physical process has proceeded before reaching equilibrium.

13. Dynamic equilibrium
 General case of an equilibrium reaction:-
A+B C+D
⇋
With the passage of time, there is accumulation of the
products C and D and depletion of the reactants A and B.
This leads to a decrease in the rate of forward reaction
and an increase in the rate of the reaction.
Finally, the two reactions occur at the same rate and the
system reaches a state of equilibrium.

14. Law of chemical equilibrium and equilibrium
constant
 We know that for a reversible reaction:-
A+B C+D
⇋
These reactants and products are balanced and hence we
can relate them by a equilibrium equation.
Where Kc is called the equilibrium constant and the
expression on the right side is called equilibrium
constant expression

15.  At a given temperature the product of concentrations of the
reactions products raised to the respective stoichiometric
coefficient in the balanced chemical equation divided the by the
product of concentrations of the reactants raised to their
individual stoichiometric coefficients has a constant value . This is
known as law of chemical equilibirum.
 Equilibirum constant for a general reaction can be written
as:- a A+b B ⇋ c C+ d D
Equilibirum constant for the reverse reaction is the inverse of the
equilibrium constant in the forward direction

16. What Does the Value of K Mean?
 If K >> 1, the reaction is
product-favored; product
predominates at
equilibrium.
 If the reverse takes place
than the former
predominates.

17. What Does the Value of K Mean?
 If K >> 1, the reaction is
product-favored; product
predominates at
equilibrium.
 If K >> 1, the reaction is
product-favored; product
predominates at
equilibrium.
• If K << 1, the reaction is
reactant-favored; reactant
predominates at
equilibrium.

18. The range of equilibrium constants
small K
small K large K
intermediate K

19. Homogeneous equilibria
 In an homogeneous system the reactants and the products
are in the same phase(solution phase) .
and also we can say in the next reaction that the ions are
equal

20. Equilibrium constant in gaseous systems
 We know that
So,
Here p is the pressure in Pascal's ,”n”is the number of moles and
volume is in cubic metre. And T is the temperature in Kelvin.
If c is the concentration in mol/L and p is in bar. then
Here R=0.0831 bar litre/mol K

21. Heterogeneous equilibria
 Equilibrium in a system having more than one phase is
called heterogeneous equilibrium.
 Example :- the equilibrium between water vapour and liquid
water in a closed container is an example of heterogeneous
equilibrium.

22. Applications of equilibrium constants
 Expression for equilibrium constant is applicable only when
concentrations of the reactants and products have attained constant
value at equilibirum state.
 The value of equilibrium constant is independent of initial
concentrations of the reactants and products.
 Equilibrium constant is temperature dependent having one unique value
for a particular reaction is represented by a balanced equation at a given
temperature
 The equilibrium constant for a reverse reaction is equal to the inverse of
the equilibrium constant for the forward reaction

23. Predicting the extent of the reaction
 The magnitude of equilibirum constants (kc and kp) is
directly proportional to the concentrations of the products.
and inversely proportional to the concentrations of the
reactants.
 If Kc is very large then the reaction proceeds nearly to
completion.
 If Kc is very small then the reaction proceeds rarely.

24. Predicting the direction of the reaction
 The equilibrium constant helps in predicting the direction in
which a given reaction will proceed at any stage. For this
purpose we calculate the reaction quotient Q.
 It is denoted with Qc with molar concentrations.

If Qc >kc then the reaction will proceed in the direction of
reactants( reverse reaction)
 If Qc < kc the reaction will proceed in the direction of products
(forward reaction)

25.  If Qc >kc net reaction goes from left to right
 If Qc < kc net reaction goes from right to left
 If Qc = kc no net reaction occurs

26. Calculating equilibrium concentrations
 Write the balanced chemical equation for the reaction.
 Under the equation make the ICE table.
(a)The initial concentrations
(b)The change in concentrations on going to equilibrium
(c) Equilibrium concentration
Substitute the equilibrium concentrations into the equilibirum
equation and solve for x. if you are to solve a quadratic
equation choose the

27. mathematical solution that makes chemical sense.
• Calculate the equilibirum concentrations form the calculated
value of x.
• Check the results by substituting them into the equilibrium
equation.

28. Relationship between equilibrium constant (K),
reaction quotient (Q) and Gibbs energy (G)
 If, ∆G is negative then the reaction is spontaneous and proceeds in the
forward direction.
 If ,∆G is positive then the reaction is considered non-spontaneous ,
instead as reverse reaction would have negative ∆G , the products of
the forward reaction shall be converted to reactants.
 If , ∆G = 0 reaction has achieved equilibrium, and at this point, there is
no longer any free energy left to drive the reaction.

29.  A mathematical expression of this thermodynamic view can be
written as:-
 At equilibrium ∆G =0 and Q = Kc
Then the equation becomes:-
Which is equal to :
Therefore:-
Taking antilog on both sides we get :-

31. Factors affecting equilibria
 Le Chateliers’s principle:- it
states that a change in any of the
factors that determine the
equilibrium conditions of a system
will cause the system to change in
such a manner so as to reduce or
to counteract the effect of the
change.

32. Effect of concentration change
The chatelier’s principle predicts that:-
 The concentration stress of an added reactant/product is relieved by net
reaction in the direction that consumes the added substance.
 The concentration stress of a removed reactant/product is relieved by net
reaction in the direction that replenishes the removed substance. Or in other
words
 “ When the concentration of any reaction or products in a reaction at
equilibrium is changed , the composition of the equilibrium mixture changes
so as to minimize the effect of concentration changes”.

33. Effect of temperature changes
 In general temperature depends on the ∆H for the reaction:-
 The equilibrium constant for an exothermic reaction decreases as the
temperature increases in this case ∆H is negative
 The equilibrium constant for an endothermic reaction increases as
temperature increases and in this case ∆H is positive.
 Raising the temperature shifts the equilibrium to left and decreases its
concentration

34. EFFECT OF INERT GASADDITION
 If the volume is kept constant and an inert gas such as
argon is added which does not take part in the
reaction, the equilibrium remains undisturbed. It is
because the addition of an inert gas at constant
volume does not change the partial pressures or the
molar concentrations of the substance involved in the
reaction.
 The reaction quotient changes only if the added gas is
a reactant or product involved in the reaction

35. EFFECT OF CATALYST
 A catalyst increases the rate of the chemical reaction by making available a
new low energy pathway for the conversion of reactants to products.
 It increases the rate of forward and reverse reactions that pass through the
same transition state and does not affect equilibrium.
 Catalyst lowers the activation energy for the forward and reverse reactions
by exactly the same amount.
 Catalyst does not affect the equilibrium composition of a reaction mixture. It
does not appear in the balanced chemical equation or in the equilibrium
constant expression.
 If a reaction has an exceedingly small K, a catalyst would be of little help.

36. ACIDS BASES AND SALTS
 Acid + base salt +water.
⇒
This reaction is called neutralization reaction.
Dielectric constant:-the ability of a polar solvent to dissociate in
water is called dielectric constant.
It is a measure relative permeability.
Separation of ions in water is called ionization or dissociation. When
ions of water are separated it is known as hydration.
Organic acids like acetic acid and formic acid cannot ionize . They
can only be partially ionized. That is (<5%)

37. Arrhenius concept of acids and bases
 According to Arrhenius theory:-
“acids are substances that dissociate in water
to give H+
(aq) and bases are substances that
produce hydroxyl ions OH-
(aq)”.
HX→ H+
+X-
or
HX+H2O→H3O+
+ X-
Water forms bond with hydrogen ion to form Hydronium ion. It has the shape of trigonal bipyramidal.

38. The Brönsted Lowry theory of acids and bases
 According to Brönsted Lowry theory, “acid is a
substance that is capable of donating a hydrogen ion and
bases are substances capable of accepting hydrogen
ions.”
 In short acids are proton donors and bases are
proton acceptors.

39. Lewis acids and bases
 According to Lewis:- “He defined acids
as a species which accepts electron
pair base which donates and electron
pair”.
 In Lewis concepts acids do not have a
proton at all.
 In bases a lone pair is provided in such
cases
 BF3 does not have a proton but still acts
as an acid and reacts with NH3 by

40. IONIZATION OF ACIDS AND BASES
 Always the stronger acid donates a proton to the stronger base.
 Strong acids dissociate very easily in water resulting in a
formation of a weak base. So always strong acids have weak
conjugate bases
 Examples are perchloric acid, hydroiodic acid, nitric acid etc:-
 Weak acids have very strong conjugate bases
 Examples are hydrofluoric acid, acetic acid.

41. Ionization constant and ionic product
 In pure water one water molecule donates proton and acts as
an acid and another water molecules accepts a proton and acts
as a base at the same time . The following equilibrium exists
as:-
The dissociation can be represented as:-

42.  The concentration of water is omitted from the denominator as water is a
pure liquid and its concentration remains constant. Water is incorporated
within the equilibrium constant to give a new constant Kw which is called the
ionic product of water.
 W can distinguish acidic, basic and neutral solutions by the relative
values of and concentrations:-
Acidic:-
Basic:-
Neutral:-

43. Ionization constants of weak acids
 General formula for dissociation of weak acids is given by the
formula:-
 Ka is called the dissociation or the ionization constant.
 At a given temperature T, Ka is a measure of the strength of an
acid HX that is larger the value of Ka the stronger is the acid. It is
a dimensionless quantity with all species of concentrations 1M.

44.  The pH scale for hydrogen ion concentration has been extended to other species and
quantities.
Steps to evaluate pH for weak electrolytes:-
A. The species present before dissociation are identified as Brönsted Lowry acids/bases.
B. Balanced equations for possible reactions i.e with a species acting both as acid as
well as a base are written.
C. The reaction with higher ionization constant is identified as the primary reaction
Whilst the other is a subsidiary reaction.
D. Enlist in a tabular form the following values of each of the species in the primary
reaction.
(i) Initial concentration
(ii) (ii)change in concentrations into equilibriums in terms of α, degree of ionization.
(iii)Equilibrium concentration.

45. E. Substitute equilibrium concentrations into equilibrium constant
equation for principal reaction and solve for α.
F. Calculate pH by the formula pH=-log(H+
)

46. Acid and Base Strength
Acid and Base Strength
 Strong acids are completely
dissociated in water.
Their conjugate bases are
quite weak.
 Weak acids only dissociate
partially in water.
Their conjugate bases are
weak bases.

47. Acid and Base Strength
In any acid-base reaction, the equilibrium will favor the reaction
that moves the proton to the stronger base.
HCl(aq) + H2O(l) → H3O+
(aq) + Cl−
(aq)
H2O is a much stronger base than Cl−
, so the
equilibrium lies so far to the right K is not
measured (K>>1).

48. Ionization of weak bases
 The equilibrium constant for base ionization is called
base ionization constant and is represented by Kb.
 Alternatively if c = initial concentration of base and α =
degree of ionization of base that is the extent to which
the base ionizes then equilibirum constant can be
written as:-

49. Relation between Ka and Kb
 The equilibrium constant for a net reaction obtained after adding two or
more reactions equal the product of the equilibrium constants for
individual reactions.
 In case of a conjugate acid base pair:-
Therefore :-

50. Di-and polybasic acids and di-and polyacidic
bases
 Some acids like oxalic acid, sulphuric acid , phosphoric acid
have more than one ionizable proton per molecule of the acid.
Such acids are called polybasic or polyprotic acids.
 Dibasic acids are those acids which have more than 2
ionizable protons and a tribasic acid has more than 3
ionizable protons.
 Higher order ionization constants are smaller than the lower
order ionization constants.
 Hence it is difficult to remove a positively charged proton
from a negative ion due to electrostatic forces.

51. Polyprotic Acids
 often acid molecules have more than one ionizable H – these are
called polyprotic acids
 the ionizable H’s may have different acid strengths or be equal
 1 H = monoprotic, 2 H = diprotic, 3 H = triprotic
 HCl = monoprotic, H2SO4 = diprotic, H3PO4 = triprotic
 polyprotic acids ionize in steps
 each ionizable Hydrogen removed sequentially
 removing of the first Hydrogen automatically makes removal of the
second Hydrogen harder.
 H2SO4 is a stronger acid than HSO4
−

52.  Extent of dissociation of an acid depends on the strength and
polarity of the H-A bond.
 Strength of the H-A bond decreases as the energy required to
break the bond decreases. That is HA becomes a stronger
acid.
 As the size of A increases down a group in the periodic table
H-A bond strength decreases and so the acid strength
increases.
 As electronegativity of A increases, the strength of the acid
also increases.

54. General Trends in Acidity
 The stronger an acid is at donating H+
, the
weaker the conjugate base is at accepting H+
 Higher the oxidation number of an acid , stronger is
the oxyacid
H2SO4 > H2SO3 ; HNO3 > HNO2
 Cations are stronger than acids than neutral
molecule and neutral acids are stronger than
anions.
H3O+
> H2O > OH-
; NH4
+
> NH3 > NH2
-
Bases have an opposite trend.

55. Common ion effect
 Reducing the concentration of hydrogen ions is called
the common ion effect
 It can be defined as a shift of the equilibrium on adding
a substance that provides more of an ionic species
already present in the dissociation equilibrium.
 Common ion effect is based on the chatelier’s
principle.

56. Hydrolysis of salts and the pH of their solutions
 Salts are formed by the reactions between acids and
bases in definite proportions which afterwards undergo
ionization in water.
 The cations and anions formed on ionization of salts
either exist as hydrated ions in aqueous solution or
interact with water to reform corresponding
acids/bases. The later process of interaction between
water and cations or anions or both of the salts is
called hydrolysis.

57.  We can say that the degree of hydrolysis is independent of
concentration of solution and pH of such solutions is
determined by their pK values.
 The pH of solution can be greater to 7 , if the difference is
positive and it will be less than 7 if the difference is negative.
)

58. Buffer solutions
 The solutions which resist change in pH on dilution or with the
addition of small amount of acid or alkali are called buffer
solutions.
 Buffer solutions are prepared from the knowledge of pKa and
pKb of base and controlling the ratio of salt and acid or salt and
base.

59. Solubility equilibria of sparingly soluble salts.
 Solubility of salts depends on lattice enthalpy of the
salts and the solvation enthalpy of the ions in the
solution.
 For a salt to dissolve in a solvent the strong forces of
attraction between its ions must be overcome by the
ion-solvent interactions.
 The solvation enthalpy of ions is referred to as
solvation which is always negative that is energy
released in the process of solvation.

60.  The amount of solvation enthalpy depends on the
nature of the solvent.
 In case of non polar solvent, solvation enthalpy is
small and hence not sufficient to overcome the lattice
enthalpy of the salt.
 Salt does not dissolve in a non polar solvent.
 As a general rule:- For a salt to be able to dissolve in a
particular solvent its solvation enthalpy must be
greater than its lattice enthalpy so that the latter may
be overcome by former.
 Solubility of salts depend on its temperature.

61. Solubility product constant
 Solubility product constant is denoted by Ksp .
 A solid salt has the formula of :- with molar solubility
S in equilibrium with its saturated solution may be
represented by the equation:-
Where

62.  Solubility product constant is given by :-
 Therefore,
The term Ksp in the equation is given by Qsp when the
concentration of one or more species is not the
concentration under equilibrium. Obviously then Ksp = Qsp
but otherwise it gives the direction of the process of
precipitation or dissolution.

63. Common ion effect of solubility of ionic salts
 If we increase the concentration of any one of the ions , it should combine
with the ion of its opposite charge and some of the salt will be precipitated
once again.
 Similarly if the concentration of one of the ions is decreased more salt will
dissolve to increase the concentration of both the ions where it will be again
precipitated.
 The solubility of salts of weak acids like phosphates increases at lower pH.
 This is because at lower pH the concentration of the anion decreases due to
protonation. This inturn will increase the solubility of the salt.
 Protonation:- protonation is the process is the addition of a proton (H+
)
to an atom molecule , or ion, forming the conjugate acid