Are you struggling to understand Chemical Equilibrium in your Class 11 Chemistry syllabus? Don’t worry — this guide will make it crystal clear! 🌟
Here, you’ll find easy explanations, important formulas, derivations, and MCQs with answers to help you prepare for board exams and MDCAT 2025. Whether you’re learning about reversible reactions, Le Chatelier’s Principle, or the Law of Mass Action, everything is explained in a simple, student-friendly way. Let’s dive into equilibrium and make chemistry your strongest subject!
Chemical equilibrium class 11, XI chemistry notes, chemical equilibrium formulas, class 11 chemistry chapter 5, MDCAT chemistry, chemistry short questions, chemical equilibrium examples.
Chemical Equilibrium
Equilibrium means no change in state of body which may be
in rest or motion
If a body is
in rest and in rest = static
equilibrium (may be stable or unstable)
If a body is in motion and in motion =
dynamic equilibrium
Quasi-static
Despite the
underlying motion, the macroscopic
properties (like concentration, pressure, or color) remain
unchanged. The system behaves as if it’s in a static state, but this is just an illusion due
to the balance of opposing processes. “Quasi-static” captures this
beautifully—it looks still, but isn't truly at rest. Like a perfectly
choreographed dance.
Under given set of conditions if a reversible process or chemical
reaction is carried out in a closed container, a constancy in some observable
properties like colour intensity, pressure, density, is observed. Such a state
is referred to as an equilibrium state.
Definition of
Chemical Equilibrium
Chemical Equilibrium is the state of a reversible reaction in a closed vessel at which there is no nET observable change in the concentrations of reactants and products with time and rate of forward reaction is exactly equal to the rate of reverse reaction i.e. it an apparent state of rest in a reversible chemical reaction where the rate of forward reaction becomes equal to the rate of reverse reaction. Thus at equilibrium state:
Chemical equilibrium is a state of a reversible reaction in which all
reacting species are present with no net change in their concentration and this
occurs only if the two opposing reactions are going on with the same rate. It
shows that reactant and product are continuously interconvert to each other
with the same rate.
Other Ways of defining equilibrium
1. Equilibrium represents the state of a process in which the
measurable
properties like T, P, colour,
concentration of the system do not show any change with the passage of
time.
2. a reversible
reaction is said to be in equilibrium when the rate of transformation of reactants into products is just equal to the rate of transformation
of products into reactants (i.e. two opposing reactions
occur at the same rate) and the concentrations of reactants and products do not change with the passage of time and becomes constant.
3. An
equilibrium is said to have been established when velocities of opposing
reactions become equal
Graphical Representation of attainment of dynamic equilibrium
In a reversible reaction, dynamic
equilibrium is established before the completion of reaction. The rate of both
forward and reverse reaction becomes equal upon reaching the equilibrium point.
The following graph which is of concentration vs time, shows that the
concentrations of both reactants and products becomes constant at equilibrium.
Equilibrium can be attained in homogenous and heterogeneous system.
Effect of Catalyst on equilibrium
The state of equilibrium is not affected by the presence of catalyst. It only helps to attain the equilibrium state in less or
more time.
Necessary Conditions for Equilibrium
1. Must have a closed system
2. The rates of opposing changes are equal
3. constant observable (macroscopic) properties
of a system
4. Must have a constant temperature
5. Activation energy is low enough to
allow a reaction
(i) Chemical equilibrium is only established in a
closed vessel where no
substances (either reactant or product) can
enter or leave the system. (A close
system is a system that may exchange energy but
not matter with its surroundings).
(ii)
The forward and reverse rates of
reactions must be equal. Equilibrium is achieved in reversible process when opposing changes to a closed system occur
simultaneously at the same rate i.e. the rates
of opposing changes are equal. Equilibrium is a dynamic process
(iii)
The concentration of both reactants and products
should remain constant. The addition
or removal of any one of them
causes the equilibrium to be disturbed.
The observable (macroscopic) properties of a system at
equilibrium are constant (e.g. temperature, pressure, colour, mass, density, pressure,
pH, concentration etc.) i.e. When a chemical system is at equilibrium, there
are no visible changes in the system. The concentrations of reactants and
products are constant (Not equal!)
(iv)
Temperature, pressure and volume should
be constant at equilibrium. If any one of these variable is changed, the system will
not remain in equilibrium.
Characteristics of
Equilibrium
1.Approaching Equilibrium from either direction
2. change in free energy i.e. ∆G = 0.
3. change
in microscopic properties
1.Approaching Equilibrium from either direction
Equilibrium can be established from either direction i.e. chemical
equilibrium can be approached from both sides.
2. change in free
energy i.e. ∆G = 0.
Thermodynamically, at equilibrium the
Gibb’s free energy (G) is minimum and any change occurring at equilibrium
proceeds without change in free energy i.e. ∆G = 0.
3. change in microscopic properties
It is a microscopic property. When chemical equilibrium is established
even then minute changes continuously take place. These changes are called
microscopic properties.
Explanation
Equilibrium can
be established for both physical processes and chemical reactions. The reaction
may be fast or slow depending on the experimental conditions and the nature of
the reactants. When the reactants in a closed vessel at
a particular temperature react to give products, the concentrations of the
reactants keep on decreasing, while those of products keep on increasing for
some time after which there is no change in the concentrations of either of the
reactants or products. This stage of the system is the dynamic equilibrium and the rates of the forward and reverse reactions become equal. It is
due to this dynamic equilibrium stage that there is no change in the
concentrations of various species in the reaction mixture. This constancy in composition indicates
that the reaction has reached equilibrium.
The characteristics of system at equilibrium are better understood if we examine some physical processes. The most familiar examples are phase transformation processes, e.g.
With 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 reverse reaction,
Chemical systems at equilibrium have constant observable
properties. Nothing appears to be happening because the internal movement
involves entities that are too small to see. A critical task of chemical
engineers is to disturb (unbalance) chemical equilibria in industrial
reactions. Production of specific desired products is controlled by
manipulating the conditions under which reactions occur.
Dynamic Equilibrium
It is an equilibrium involving constant interchange of activated particles in motion. The chemical equilibrium is said to be in Dynamic State
because it involves constant
and continuous interchange (exchange) of activated (dynamic) molecules of reacting
substances (reactants and products) in motion i.e. forward and reverse reactions occur at equal rates in opposite directions (reaction is
continuously going on in the forward and reverse directions with equal rates).
A system at equilibrium is dynamic on the molecular level; no further net change is observed because changes in one
direction are balanced by changes in the other.
Apparently, it
appears that the equilibrium is dead or static and the reaction seems to be
cease but the equilibrium is dynamic and the molecules are still changing from
reactants to products and from products to reactants but with no net change in
their concentrations.
Although the
concentrations of the substances remain unchanged (as indicated by the term
“equilibrium”), there is still activity going on; both forward and backward
reactions are continually occurring (as indicated by the term “dynamic”) but
since they proceed at the same rate, each species is formed as fast as it is
consumed, resulting in a constant concentration term.
Activated or Dynamic Molecules
The small fractions of reacting molecules that successfully collide
to form products are called Activated or dynamic Molecules.
Equilibrium Mixture
A mixture of various substances at equilibrium in a closed vessel
is called equilibrium mixture. It is a mixture of various species in which the
chemical equilibrium exists. It is a mixture of reactants and products in the equilibrium state.
Equilibrium
Concentration
The concentrations of reactants and products at equilibrium state
are called Equilibrium Concentration.
equilibrium position
Each set of equilibrium concentrations is called an equilibrium
position. It is a particular set of equilibrium concentrations of reactant
and product species. the equilibrium position refers to the
relative amounts of reactants and products in the system at the point of
equilibrium
It is essential to distinguish between
the equilibrium constant and the equilibrium positions for a given reaction
system. There is only one
equilibrium constant for a particular system at a particular temperature, but
there are an infinite number of
equilibrium positions. The specific
equilibrium position adopted by a system depends on the initial concentrations, but the equilibrium constant does not depend on the initial concentrations.
Example: for the system A ⇌ B, the mass action expression is [B]/[A]. Let's
say K = 200. This means, at equilibrium, [B] will be 200 times
greater than [A]. That's what the equilibrium constant tells you – equilibrium
for this system lies to the right, and the K value is greater than 1.
There are infinitely many different equilibrium positions that
satisfy K. We could have [B] = 20M and [A] = 0.1M. Or we could
have [B] = 1M and [A] = 0.005M. Those are two different
equilibrium positions that are both at equilibrium. In both cases, Q = K and
the system is at equilibrium.
(i)
a reaction with an equilibrium position that favours the products:
[product] > [reactant] at equilibrium
equilibrium lies to the right
(ii) a
reaction with an equilibrium position that favours the reactants:
[reactant] > [product]
equilibrium lies to the left
Example of attainment of equilibrium for hydrogen iodide
formation from hydrogen & iodine
An example of reaction at equilibrium is a
reaction of hydrogen and iodine vapours at a high temperature of 500oC
in a closed container to produce hydrogen iodide. When certain amount of
hydrogen and iodine are mixed in a sealed container at 500oC, some
of their molecules react with each other to give hydrogen iodide. At the same
time, some of the hydrogen iodide molecules decompose back to hydrogen and
iodine.
At the start of reaction, there is a higher
concentration of hydrogen and iodine and after that the concentration of
decreases as hydrogen iodide is formed. The concentration of hydrogen iodide
increases as the forward reaction proceeds.
As hydrogen iodide is formed, the reverse
reaction is then able to occur. Although the rate of reverse reaction is quite
slow in the beginning due to low concentration of HI but as the time goes on,
the rate of the forward reaction will go on decreasing and the reverse reaction
will go on increasing and ultimately the two rates will become equal to each
other. Ultimately, the rate at which H2 and I2 react to
form HI (rate of forward reaction) becomes equal to the rate at which HI breaks
down back into H2 and I2 (rate of reverse reaction).
Thus, the equilibrium will set up and concentration of various species (H2,
I2, HI) becomes constant. The attainment of equilibrium can be seen
by the intensity of purple colour of iodine which decreases gradually until a
constant light purple colours is settled. It is represented as
Ways to recognize Chemical Equilibrium
The formation of a chemical equilibrium can be recognized by
following two ways:
(i) Physical method
In this method, specific radiations (UV, IR or visible) pass
through reaction mixture. Both reactants and products absorb radiations with
respect to their equilibrium concentration noted by spectrometer. The % absorbance of these radiations determines the equilibrium concentration
of reaction mixture.
Concentration of a chemical solution is directly proportional to
its absorption of light. There is a linear relationship between concentration and
absorbance of the solution, which enables the concentration to be calculated by
measuring its absorbance.
(ii) Chemical Method
for Determination of Kc for Ethyl Acetate Equilibrium by Experiment
For determining equilibrium constant by using chemical method let
us consider the esterification of ethyl alcohol and acetic to form ester and
water
CH3COOH(l)+C2H5OH(l)⇌CH3COOC2H5(l) + H2O(l
Acetic acid Ethyl alcohol Ethyl acetate (ester)
Since the esterification equation shows that 1 mole of acetic acid
and 1 mole of ethyl alcohol reacts to form 1 mole of ester and 1 mole of water
therefore we conclude that the amount of acid used up in the reaction is equal
to the amount of alcohol consumed. Thus at equilibrium we have (a–x) mole
acetic acid, (b–x) moles of alcohol and x moles of ester and x moles of water.
We make ICE (Initial, change and equilibrium concentration) table for the
reaction as
Now applying law of Mass Action to calculate Kc
If we repeat the same
experiment by taking different amount of CH3COOH and C2H5OH,
the value of Kc will be the same at constant temperature.
 |
Summary of Equilibrium
▶ state in a
reversible reaction
▶ At equilibrium, all measurable properties of reaction like mol,
concentration, pH etc. do not change
▶ Rate of forward reaction becomes equal to rate of backward
reaction
▶ established only in
closed vessel
▶ Constant measurable properties (moles, conc., mole fraction,
colour etc.) does not mean equal at all.
▶ Equilibrium is stable in
nature (reaction always tends to stay at equilibrium)
▶ Equilibrium is dynamic (particles always colliding,
reacting, and re-forming) but
Quasi-static in nature
▶ Molecules try to maximize entropy
▶ Catalyst helps achieves equilibrium sooner than expected.
▶ Molecules try to minimize energy
▶ Equilibrium can be achieved from any direction
Or
Rate of change of Reactants to products = Rate of change of
products to reactants
And
Constant concentration of reactants and products
Phase equilibrium
Phase
equilibrium involves a single chemical substance existing in more than one
phase in a closed system. Water placed in a sealed container evaporates until
the water vapour pressure (concentration of water in the gas phase) rises to a
maximum value, and then remains constant.
Solubility equilibrium
It involves a
single chemical solute interacting with a solvent substance, where excess
solute is in contact with the saturated solution.
chemical equilibrium
A chemical
equilibrium involves several substances: the reactants and products of a chemical
reaction.
Homogeneous Equilibrium
A chemical equilibrium in which the reactants and products are in
the same phase is called homogeneous equilibrium
e.g.
2SO2(g)
+ O2(g) ⇌ 2SO3(g)
Heterogeneous equilibrium
A chemical equilibrium in which the reactants and products are
present in different phases is called Heterogeneous equilibrium
e.g.
CaCO3(s)
⇌ CaO(s) + CO2(g)
Law of Mass Action/Law of Equilibrium (Guldberg-Waage Law)
Statement of LMA
In 1864, two Norwegian chemists Cato Maximillian (C.M) Guldberg (1836–1902) and Peter (P) Waage (1833–1900) studied experimentally a number of equilibrium reactions and put
forward their results as a generalization known as law of mass action.
The two Norwegian scientist C.M Guldberg (1836–1902) and Peter Waage (1833–1900) observed that reversible reaction reaches a state where
the ratio of its product concentration to that of reactant concentration
becomes constant. On the basic of their research
conclusions, they derived a quantitative relationship between the
rate of reaction and active masses (molar
concentration) of reacting substances in the
form of Law of Mass action or equilibrium law.
The rate of any reaction is directly proportional to its
active mass and the rate of a chemical reaction is directly proportional to the
product of the active masses or molar concentration (in mol/dm3) of
reacting substances raised to the power of their stoichiometric coefficients in the balanced chemical equation at constant temperature.
OR
At a given
temperature, the product of concentrations of the reaction products raised to
the respective stoichiometric coefficient in the balanced chemical equation
divided by the product of concentrations of the reactants raised to their
individual stoichiometric coefficients has a constant value. This is known as
the Equilibrium Law or Law of Chemical Equilibrium.
The Molar Concentration of substances in mol/dm3
(mol/litre) is termed as active mass which
represented by square brackets; [ ]. Molar concentration of different species
is indicated by enclosing these in square bracket and, it is implied that these
are equilibrium concentrations.
For a general reaction: A + B ⇌ Product, representing Molar concentration of A and B as [A]
and [B] respectively, then according to Law:
Rate
of Reaction a [A] [B]
Rate of Reaction a Molar Concentration
of reacting substances
Rate
of Reaction a Active mass of 1st
substance × Active mass of 2nd substance
Mathematical
Expression of LMA
For a general reaction: A +
B ⇌ Product; Law of Mass Action can be written expressed as
Rate
of Reaction a Molar Concentration of
reacting substances
Rate of
Reaction a Active mass of 1st
substance x Active mass of 2nd substance
Rate
of Reaction a [A] [B]
Here,
[A] = Molar concentration of A in mol/dm3
[B] = Molar concentration of B in mol/dm3
Importantly, the Equilibrium Law expresses Kc as a relationship between the concentrations of products and reactants in a system at equilibrium and it provides us with a quantifying means to determine the position of the equilibrium.
The magnitude of the equilibrium constant informs us of the relative proportion of products and reactants, providing us information on the extent of the reaction (but not reaction rate).
Generally the
subscript ‘eq’ (used for equilibrium) is omitted from the concentration terms.
It is taken for granted that the concentrations in the expression for Kc
are equilibrium values. While writing
expression for equilibrium constant, symbol for phases (s, l, g) are generally
ignored.
Active Mass
Active mass represent amount of substance. (Active mass of pure liquid and pure solid and solvent is 1. It may be expressed in two ways:
1. Active mass in terms of
concentration (mol L−1 or moldm−3)
2. Active mass in terms of partial pressure in atm only for gases.
1. Active mass in terms
of concentration
It is represented by [ ] = M = C = no of moles of substance per dm3
[ ] = mol/Vdm3 = W/M × V dm3
[ ] = W/Vdm3 × 1/M = ρ/M moldm−3 (where moldm−3)
If solid or pure liquid is taken then active mass = 1 because throughout reaction density of solid and pure liquid does not change.
[Solid] or [Pure Liquid] = 1
The active mass of gas does not remain constant in the reaction
because the density of the gas changes as the volume of container change.
Volume of gas = volume of container
2. Active mass in terms of Partial pressure only for gaseous reactions
In case of gaseous equilibrium, where reactants and products are in gaseous state, the concentrations of gaseous reactants and products can be expressed in terms of their partial pressures as at constant temperature the partial pressure of a gas is directly proportional to its molar concentration. i.e.
PV = nRT ⇒ P = (n/V) RT ⇒ P =
Molar concentration (C) RT [(n/V =C] OR P a Molar concentration (C)
Answer
Active mass of 2g NaOH(s) is 1.
Q2. Calculate active mass 2g NaOH dissolved in 2dm3
water.
Answer
Active mass of NaOH = [NaOH] = mol/molar mass × volume (dm3)
= 20/40×2 = ¼ or 0.25 moldm−3.
Derivation of Equilibrium Constant (Kc)
Expression/LMA Expression
Let us apply the law of Mass
Action to derive equilibrium constant (Kc), for a
general hypothetical reversible reaction in which reactants A and B combine
to form products C and D where a, b, c and d are numbers of moles needed to
balance a chemical equation at a certain temperature. At equilibrium state, the
concentrations of A, B, C and D become constant. Let [A], [B], [C] and [D] are
the active masses or molar concentrations in mole/dm3 of A, B, C and
D at equilibrium state respectively).
To illustrate law of mass action
mathematically, consider a general reversible reaction in which reacting
species A,B, C, and D exist in equilibrium state at a certain temperature:
According to the law of mass action, the rate of forward reaction
(Rf) is directly proportional to the product of active masses of
reacting substance A and B. Similarly the rate of backward reaction (Rb)
is directly proportional to the active masses of reacting substances C and D.
Then, according to Law of Mass Action, rate of forward reaction and reverse
reactions are given as:
Rate of forward reaction a [A]a [B]b OR Rf
= Kf [A]a[B]b (Kf = specific rate
constants for forward reaction)
Rate of reverse reaction a [C]c[D]d OR Rr =
Kr [C]c[D]d (Kr = specific rate constants for
reverse reaction)
Where Kf and Kr are the proportionality
constant and are known as the specific rate constant for forward reaction and
Specific Rate Constant for Reverse Reaction respectively. Their values depend
upon the nature of reactants and products.
Since, chemical equilibrium is dynamic, so at the equilibrium state Rate of Forward and reverse reaction becomes equal:
At any given temperature, both Kf and Kr are
constant, the ratio Kf/Kr of will also be constant and
collectively termed as equilibrium constant donated as Kc (or simply
K) where subscript ‘c’ indicates concentrations in mole/dm3. The
above expression is known as equilibrium law or Equilibrium Constant Expression or Kc–Expression or Law of Mass Action Expression or LMA-Expression. It shows that in a reversible reaction at equilibrium state at a
certain temperature, the ratio of active masses of products to that of
reactants becomes constant. The Kc–Expression is written by placing the active masses of products in the
numerator and active masses of reactants in the denominator with each
concentration term raised to a power equal to the coefficient of the substance
in the balanced equation.
Derivation of Different Types of Equilibrium Constant (Kc)
Expression/LMA Expression
A ⇌ B
Rf α
[A] OR Rf = Kf [A] (Kf
= forward rate constant)
Rb α [B] OR Rb = Kb [B] (Kb = backward rate constant)
Since, at the equilibrium state:
Equilibrium Constant (Kc)
The equilibrium constant (Kc) of a reversible
reaction is a constant ratio of Kf/Kr (specific
rate constant for forward reaction/specific rate constant for reverse reaction)
at constant temperature. Kc
is the ratio of the product
of the molar equilibrium concentrations (active masses) of the products and the product of molar
equilibrium concentrations (active masses) of reactants with each
concentration term raised to a power equal to
its numerical coefficient given in the balanced equation at constant temperature.
Thus Kc is directly proportional to molar equilibrium
concentrations (active masses) of products and inversely proportional to molar
equilibrium concentrations (active masses) of reactants.
For a general reversible reaction; aA + bB ⇌ cC + dD, Kc is given by
Characteristics of Equilibrium Constant
1. Expression for equilibrium constant is applicable only
when concentrations of the
reactants and products have attained
constant value at
equilibrium state.
2. Kc for any given reaction at a particular
temperature always has the same
value.
3. The value Kc is determined by experiment.
Factors that does not affect Equilibrium Constant
1. The value of
equilibrium constant is independent
of initial concentrations of the reacting species (reactants and
products) i.e. it does not depend on the initial concentrations of the
reactants.
2. Kc is independent of the number of intermediate steps in the
reaction mechanism.
3. The equilibrium
constant is independent of the presence of a
catalyst.
4. The equilibrium
constant for a reverse reaction is equal to the inverse of the equilibrium constant for the forward reaction
5. The equilibrium
constant is independent of the pressure and
volume
6. The equilibrium constant is independent of the of inert material.
Factors affecting Equilibrium Constant/ Effect of Change in
temperature on the Value of Kc
Equilibrium constant is temperature-dependent i.e. Kc changes with change in temperature having one
unique value for a particular reaction. The effect of temperature change on Kc
depends on the sign of DH for the reaction. The Kc for an exothermic reaction (negative DH) decreases as the temperature increases while Kc
for an endothermic reaction (positive DH) increases with the rise in temperature.
For example
(i) Kc for the
synthesis of ammonia by Haber’s process is 4.1 × 108 at 25oC but 0.5 at
400oC.
(ii) Kc for the decomposition of N2O4
at 25oC is 4.64 × 10−3 but at 127oC it is 1.53.
Importance of Kc
Kc
determines which in greater concentration at equilibrium – the products or the
reactants. In general:
1. Kc
> 102 (1); equilibrium lies to the right and favours the product.
2. Kc
< 10-2 (1); equilibrium
lies to the left and favours the reactants.
If Kc
is very large, the equilibrium mixture will contain mostly products while if Kc
is very small, the equilibrium mixture will contain mostly reactants.
Unit of Kc
1. The unit of Kc (and Kp) depend on
the specific reaction and the unit of Kc (and Kp) varies
depending on the terms in
its expression. The unit of Kc depends on the form of equilibrium
expression.
2. Kc has dimensions equal to
(concentration)∆n
or (concentration)np − nr or
(concentration)(c+d)‒(a+b)
or (concentration)c+d‒a‒b or
(mol/dm3)∆n
or (mol/L)∆n or (mol dm−3)∆n or “mol(c+d)‒(a+b) dm‒3[(c+d)‒(a+b)]”
Unit of KX is always unitless
as mole fraction is unitless.
where ∆n is equal to the total number of
moles of products minus the total number of moles of reactants.
3. If the number of
moles of reactants is equal to the number of moles of products, Kc
has no unit since concentration units
(mol/dm3) of all species are cancelled out by each other. Kc
will be dimensionless (has no
unit) only for those reactions for which a + b = c + d signifying that the total number of moles of reactants and products are equal.
4. If the number of
moles of reactants are different from the number of moles of products, then the
unit of Kc is
determined by using the formula (mol/dm3)∆n
4. In thermodynamics, Kc (and Kp) is defined to have no unit. In general practice the unit of Kc is not written.
Example
Units of Kc depend upon the
number of moles of reactants and products involved in the reaction.
(i) Reaction
without change in number of moles (Dn = 0) has Kc with no unit.
(ii) Reaction with
change in number of moles (Dn ≠ 0) has Kc with variable units depending upon change in moles.
For the homogeneous reaction 4NH3(g)+5O2(g) ⇌ 4NO(g) + 6H2O(g); the unit
of equilibrium constant Kc is calculated as:
The unit of the equilibrium constant Kc is (mol/litre)Δn.
Here, Δn is the difference in the total number of gaseous
products and the total number of gaseous reactants. For the reaction; Δn = 4 +
6 − (4 + 5) = 1.
Hence, the unit of the equilibrium
constant Kc will be (mol/litre)Δn =
(mol/litre)1 = mol/litre = conc.
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