Le-Chatelier’s Principle for MDCAT/ECAT:| Definition, Explanation, Examples & Applications

Welcome to Learn Chemistry by Dr. Inam Jazbi! In this article, we’ll explore one of the most important concepts in Chemical Equilibrium — Le-Chatelier’s Principle. This principle explains how a system at equilibrium responds to external changes such as temperature, pressure, or concentration. Understanding this concept helps predict the direction of chemical reactions and is essential for both exam preparation and industrial chemistry applications like the Haber Process and Contact Process.

Learn the definition, rule, examples, and real-life uses of Le-Chatelier’s Principle in a clear and simple way below 👇

Le-Chatelier’s Principle/ Le- Chatelier’-Brawn’s Principle/Law of mobile equilibrium (Factors Affecting Balance of Chemical Equilibrium)

Statement of the Principle

An equilibrium system becomes temporarily unbalance by the change in experimental conditions such as concentration, temperature and pressure of reacting species and to restore its balance again the system shifts to either right or left by speeding up either forward or backward directions. These equilibrium constraints or reaction parameters may affect the position of equilibrium.

The quantitative influence of the change of conditions like concentration, temperature and pressure on a equilibrium was generalized by a French chemistry Le-Chatelier in 1884 in a general rule or  principle called Le-Chatelier’s Principle or Le- Chatelier’-Braun’s Principle put forward by a French chemist, Henri Louis Le-Chatelier’s and Karl Ferdinand Braun in 1884.

If an external stress or constraint such as concentration, temperature or pressure is applied to a system at equilibrium, the equilibrium is disturbed and tends to shift in a direction to offset the effects of imposed stress.

OR

If a system in equilibrium is altered by changing concentration, pressure or temperature or the introduction of a catalyst called constraint or stress, the equilibrium position changes or shifts in the direction that will minimize or remove (undo) the effects of that constraint or stress.

It states that changes in the temperature, pressure, volume, or concentration of a system will result in predictable and opposing changes in the system in order to achieve a new equilibrium state.

 

Explanation

There are two parts in the statement of Le-Chatelier’s principle, which need further elaboration.

First part which refers to ‘disturbance on equilibrium (equilibrium constraints or reaction parameters or applied stress)’ implies that an altering in the experimental conditions push the system temporary out of equilibrium so that Q_c will no more equal to K_c. The term constraint we mean the changes in concentration, pressure or temperature.

The second part which specifies the ‘shifting of equilibrium position to left or right’ means system tends to repair its equilibrium by reducing the stress and tends to attain a new equilibrium position where Q_c=K_c again. Shifting means “net reaction occurs” either in the forward (shift right) or reverse (shift left).

According to Le-Chatelier’s principle, if one of the factors involved in affecting chemical equilibrium is altered, the equilibrium shifts towards right (forward reaction) or left (reverse reaction) in order to restore the balance of equilibrium. The change in concentration and change in pressure cause shift in equilibrium without affecting the equilibrium constant. The change in temperature alters the position of equilibrium by altering the equilibrium constant. Catalyst neither shifts the equilibrium nor does it alter the equilibrium constant but it merely enables a quicker attainment of equilibrium.

Prediction of equilibrium change based on pushes and pulls

The simplest way of predicting the equilibrium change is based on pushes and pulls.

Imagine that you have an office chair sitting in front of you. If you push the chair, it moves away from you – the act of pushing moves the chair away from the disturbance the pushing force, you .If you pull the chair, it moves towards you – the act of pulling moves the chair towards the disturbance the pulling force. This is a common, everyday situation we have all experienced at one time or another. We can apply the ideas of pushes and pulls to equilibrium by making the following associations:

if I increase the amount of chemical substances or heat temperature then I am pushing the reaction. The reaction responds to this push by moving away from the push.

An analogy of Le-Chatelier’s principle in Physics

An analogy of Le-Chatelier’s principle in Physics is Newton’s Third law of Motion. 

Detailed Explanation of the Principle

When a chemical system at dynamic equilibrium is disturbed by changing the concentration of either reactants or products; or by changing the partial pressures of any of gaseous reactants or of gaseous products; or temperature, the position of equilibrium is changed in that direction so as to establish a new equilibrium state i.e., either forward reaction or backward reaction is favoured.

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.

If a system at equilibrium is disturbed or changed by a change in temperature, pressure or concentration, the system will shift to a new equilibrium position to counteract the disturbance (or change).

Thus, according to Le Chatelier’s principle, reversible reactions are self-correcting; when they are thrown out of balance by a change in concentration, temperature, or pressure, the system will naturally shift in such a way as to “re-balance” itself after the change.

Applications of the Principle

Le Chatelier’s principle can be used in practice to understand reaction conditions that will favour increased product formation. The influence of various factors such as changes in concentration, pressure or temperature or catalyst on a system in equilibrium is governed by Le-Chatelier’s Principle. When using Le-Chatelier’s Principle, the applied stress is identified first. Each applied stress has a predictable effect on an equilibrium position.

1.   Effect of Concentration Change

When  a system at equilibrium is disturbed by increasing or decreasing the concentrations of one or more species involved in the reaction, the equilibrium tends to shift towards left or right in order to reduce the effect of this stress and readjust itself until Q_c = K_c. Thus equilibrium is self-correcting. In general, whenever we add or remove some of the reacting species from a system at equilibrium, the system reacts in a particular direction to reduce the amount of added substance or to produce certain amount of removed substance so that the stress imposed on the system will be offset.

In general:

(i) The equilibrium position shifts towards right if the amount of reactant is added or product is removed.

(ii) The equilibrium position shift to the left if the amount of reactant is removed or product is added.

Generally, rate of forward reaction will increase if an additional amount of any reactant is added at equilibrium to a system. Generally, rate of reverse reaction will increase if an additional amount of any product is added to a system at equilibrium.

Explanation

Consider a general reversible reaction:

At equilibrium, there is a constant ratio between concentrations of products and reactants and it is called Equilibrium Constant represented by K_c. If the concentration of either reactants or products is altered, then equilibrium will disturb and equilibrium will shift in forward or reverse direction in order to restore the constant value of K_c and to re-establish the equilibrium.

if we increase the concentration of reactants i.e., A or B or both, the value of concentration quotient or reactant quotient no longer remains equal to K and the system no longer remains in a state of equilibrium. In order to attain the state of equilibrium again, the concentrations of C and D will increase and the equilibrium shifts in the forward direction. Similarly, if we increase the concentration of the products, either C or D or both, the equilibrium shifts in the backward direction.

According to Le-Chatelier’s principle, when the concentration of one of the substances in, a system at equilibrium is increased, then the equilibrium will shift so as to use up the substance added. Suppose at equilibrium one of the reactants is added, the equilibrium will shift in the direction that consumes reactants, i.e., the forward direction. The more of the reactants would be converted into products. On the other hand, if one of the products is added the equilibrium will shift in the backward direction because it consumes the products.

Illustration

To understand the guideline provided by Le-Chatelier’s principle for the stress caused by the change in the concentration, consider the following system at equilibrium

CO_(g) + 3H_2(g) ⇌ CH_4(g) +H₂O_(g)

When certain amount of CO or H₂ is added to the system, the value of reacting quotient (Q_c) is lowered than its K_c and reaction is no longer at equilibrium. The stress caused by addition of CO or H₂ is offset by shifting the equilibrium to the right side. In doing so some amount of CO or H₂ has consumed turning the value of Q_c back to K_c.

In kinetic aspect, the addition of CO or H₂ increases the rate of forward reaction by consuming added CO or H₂ thereby producing more CH₄ and H₂O until at a certain point a new equilibrium will be established. 

1.    Increase in the concentration of the reactant or decrease in the concentration of the product will shift the equilibrium towards right (product side).

2.    Decrease in the concentration of the reactants or increase in the concentration of the product will shift the equilibrium towards left (reactant side).

By using le Chatelier's principle, the effect of change in concentration on systems at equilibrium can be explained as follows:

1)   When the concentration of reactant(s) is increased, the system tries to reduce their concentration by favouring the forward reaction.

2)  When the concentration of product(s) is increased, the system tries to reduce their concentration by favouring the backward reaction.

3)   When the concentration of reactant(s) is decreased, the system tries to increase their concentration by favouring the backward reaction.

4)   When the concentration of product(s) is decreased, the system tries to increase their concentration by favouring the forward reaction.

 

 2.   Effect of Temperature Change

Effect of change in temperature is related to the nature of reaction whether it is an endothermic reaction or exothermic reaction. In term of enthalpy change (∆H^o), reactions are classified into exothermic and endothermic. Reversible reactions are exothermic in one direction and endothermic in other direction. However, when the equilibrium is shown in an equation form, the sign of enthalpy change (∆H^o) refers to the forward direction. Temperature changes affect the equilibrium constant and rates of reactions.

According to Le-Chatelier’s Principle, when temperature of a system is increased at equilibrium, the equilibrium will shift to that direction in which heat is absorbed i.e. the increase in temperature shifts the equilibrium in the direction of the endothermic reaction while decrease in temperature shifts the equilibrium in the direction of exothermic reaction.

Raising The Temperature Increases K_eq For An Endothermic Reaction And Lowers K_eq For An Exothermic Reaction (Lowering The Temperature Has The Opposite Effect.)

 

Explanation

Consider the following reaction

PCl_3(g) + Cl_2(g) ⇌ PCl_5(g)  (∆H^o = −ve)    K_c = [PCl₅]/[PCl₃][Cl₂]

The sign of ∆H^o shows that the reaction is exothermic in the forward direction. According to Le-Chatelier’s principle, increase in the temperature on this equilibrium system in the form of some heat will speed up the backward reaction where added heat is absorbed and the stress of increased temperature is reduced. Thus the rate of decomposition of PCl₅ becomes faster than its formation thus numerator value in the equilibrium constant expression becomes smaller than denominator (K_c decreases). Finally, the system reaches to a new equilibrium state with lower value of K_c.

Contrarily, if we lower the temperature by removing heat from the system at constant pressure, then according to Le-Chatelier’s principle, the equilibrium shifts to the right and rate of forward reaction becomes faster producing more PCl₅. In this way, the value K_c to attain new equilibrium will be enhanced. 

 

 

Effect of Temperature on Exothermic Reactions

In Exothermic reactions, K_c decreases with the rise of temperature i.e. concentration of products decrease and the reaction shifts towards reverse direction in order to nullify the change. Thus decrease in temperature always favours an exothermic reaction.

Effect of Temperature on Endothermic Reactions

In endothermic reactions, K_c increases with the rise of temperature i.e. concentration of products increases and the reaction shifts towards forward direction in order to nullify the change. Thus increase in temperature always favours an endothermic reactions.

Summary of Effect of temperature

(i) If ∆H = +ive (endothermic), an increase in temperature shifts the reaction in forward direction.

(ii) If ∆H = -ive (exothermic), an increase in temperature shifts the reaction in backward direction.

Remember that one direction of a reaction is always exothermic and the other direction is endothermic i.e. the reverse of exothermic is endothermic and the reverse of endothermic is exothermic. The endothermic direction has the larger activation energy.

When temperature increases, both rates (forward and reverse) increase but the rate of the endothermic reaction increases more! Equilibrium shifts in the endothermic direction. (Le-Chatelier’s says that this consumes the added heat!)

When temperature decreases, both rates (forward and reverse) decrease but the rate of the endothermic reaction decreases more! Equilibrium shifts in the exothermic direction.  (Le-Chatelier’s says that this partially replenishes the removed heat!)

 

Explanation

The effect of temperature on equilibrium has to do with the heat of reaction. Recall that for an endothermic reaction, heat is absorbed in the reaction, and the value of ΔH is positive. Thus, for an endothermic reaction, we can picture heat as being a reactant:

                          Heat + A ⇌  B                ΔH = +

For an exothermic reaction, the situation is just the opposite. Heat is released in the reaction, so heat is a product, and the value of ΔH is negative:

                                  A ⇌  B + heat                 ΔH = −

If we picture heat as a reactant or a product, we can apply Le-Chatelier’s principle just like we did in our discussion on raising or lowering concentrations.

For instance, if we raise the temperature on an endothermic reaction, it is essentially like adding more reactant to the system, and therefore, by Le-Chatelier’s principle, the equilibrium will shift the right.  Conversely, lowering the temperature on an endothermic reaction will shift the equilibrium to the left, since lowering the temperature in this case is equivalent to removing a reactant.

For an exothermic reaction, heat is a product. Therefore, increasing the temperature will shift the equilibrium to the left, while decreasing the temperature will shift the equilibrium to the right.

3.  Effect of Pressure Change

Change of pressure is related to the change of volume. In applying Le Chatelier’s principle to a heterogeneous equilibrium the effect of pressure changes on solids and liquids can be ignored because the volume (and concentration) of a solution/liquid is nearly independent of pressure.

The change of pressure will have no effect on the state of equilibrium, if:

(a) The system does not involve gaseous components.

(b) The number of moles of gaseous reactants is equal to the number of moles of gaseous products.

Pressure changes only affect homogenous gaseous equilibria in which there is a volume (or number of moles) change. The reason is that gases are highly compressible therefore the pressure applied to the system is directly proportional to the concentration of reacting species but inversely proportional to the volume at given constant temperature (P = (n/V) RT and P = CRT). The effect of changing the pressure on a gas-phase reaction depends on the stoichiometry of the reaction. A change in pressure or volume will result in an attempt to restore equilibrium by creating more or less moles of gas.

In gaseous equilibria reactions involving volume change, an increase in pressure will shift the equilibrium toward the side of decreasing volumes or fewer moles of gas. In such reactions, an increase in pressure shifts the equilibrium in the direction of decreasing volume or fewer moles of gas where as a decrease pressure shifts the equilibrium in the direction of increasing volume or greater moles of gas.

 

Explanation

To understand the effect of pressure on an equilibrium let us consider the following gaseous system in a cylinder fitted with a movable position: 

When the external pressure increases at constant temperature, the piston move downward causing a decrease in the volume which in general increases the concentration of all components of reacting mixture but since the number of moles of products are lesser than reactant, the denominator value exceeds the numerator. Thus the system is no longer in equilibrium and to reduce this stress the reaction tends to shift on right side.

Summary of Effect of Pressure

(i) When ∆n = 0;  no effect

(ii) When ∆n ≠ 0; Increase in pressure shifts the reaction in the direction of lesser number of moles

Decrease in pressure shifts the reaction in the direction of greater number of moles

 

Ways of introducing a pressure change

A pressure change can be introduced through the following ways:

1. Addition or removal of a gaseous component of the equilibrium mixture at constant volume

2. Expansion or compression of the reaction vessel

3. Addition of an inert gas at constant volume

4. addition of an inert gas at constant pressure

 

4. Effect of Catalyst

1. A substance, which increases the rate of a chemical reaction without itself being consumed in the reaction, is called a catalyst.

2. A catalyst speeds up the rates of both the forward and reverse reactions to the same extent. It simply enables equilibrium to be reached more quickly by decreasing the energy of activation of both reactions thereby hastening the attainment of the equilibrium.

3. A catalyst speeds up the rate at which equilibrium is achieved. i.e. it enables a quicker attainment of equilibrium. A catalyst only aids a reaction in attaining equilibrium in a shorter time.

4. A catalyst does not change composition of the equilibrium mixture, the position of equilibrium and the value of K_c.  Thus a catalyst does not change the yield of a reaction.

5. A catalyst lowers the activation energies of both the forward and backward reactions to the same extent. This, in turn, leads to the rates of both the forward and backward reaction being increases to the same extent. 

 

Industrial Application of Le-Chatelier’s Principle in Haber’s Process For the Manufacture of Ammonia

 

Principle

Ammonia is manufactured on large scale by Haber’s Process by the direct union of nitrogen and hydrogen. The reaction is reversible, exothermic involves gaseous equilibrium occurring with decrease in volume of products. Le-Chatelier’s Principle suggests the following conditions for the maximum yield of ammonia:

(i) High concentrations of N₂ and H₂ in the ratio of 1:3 by volume

(ii) Removal of ammonia after regular intervals.

(iii) Low temperature is maintained; optimum temperature is 400°C–500°C

(iv)High pressure is maintained; optimum pressure is 400-1000 atm (200-300 atm).  

(v)  Suitable catalyst in the form of finely divided iron.

(vi) Promoter in the form of MgO, Al₂O₃ and SiO₂

(vii) Yield of ammonia is 35%.

Hence, high reactant concentration, high pressure, low temperature and continuous removal of NH₃ will give best yield of NH₃.

 

Factors enhancing the yield of NH₃

For maximum yield of ammonia Le-Chatelier’s Principle suggests the following conditions:

 

1.   Effect of Concentration

When concentration of N₂ or H₂ is increased at equilibrium, then according to Le-Chatelier’s Principle, the equilibrium is forced toward right in order to reduce the concentration of added N₂ or H₂, thereby increasing the yield of NH₃. In actual practice, however, there is no practical advantage in using excess of N₂ or H₂.  The two gases are mixed in the theoretical ratio of 1:3 by volume, which considerably increases the yield of NH₃. By continuous removal of NH3 from the reaction mixture, reaction moves in forward direction.

 

2.   Effect of Temperature

The formation of NH₃ is exothermic. According to Le-Chatelier’s Principle, at low temperature more NH₃ will be formed while at high temperature NH₃ will decompose into N₂ and H₂. Thus the low temperature helps in stabilization of NH₃ i.e. favours the formation of NH₃. [However, at low temperature, equilibrium will be established in longer time and rate of formation of NH₃ will be low]. In actual practice, the reaction is carried out at high temperature of 450°–500°C that produces reduced yield at much higher rate.

3.   Effect of Pressure

The synthesis of NH₃ occurs with decrease in volume of products (4 volumes of N₂ and H₂ give 2 volumes of NH₃). According to Le-Chatelier’s Principle, the equilibrium state will shift towards right on increasing the pressure, which lowers the volume (concentration) of N₂ and H₂. For industrial synthesis of NH₃, the pressure on the equilibrium is maintained from 400–1000 atmosphere.

 

4.   Effect of Catalyst

Finely divided iron or platinum is used in Haber’s Process to speed up the reaction.

 

Industrial Application of Le-Chatelier’s Principle in Manufacture of Sulphuric Acid (H₂SO₄) by Contact Process

Principle

Sulphuric acid is manufactured on large scale by the Contact Process which is multi step process, the most important step of which is the oxidation of sulphur dioxide to sulphur trioxide by atmospheric oxygen. The reaction is reversible, exothermic involves gaseous equilibrium occurring with decrease in volume of products. Le-Chatelier’s Principle suggests the following conditions for the maximum yield of sulphur trioxide:

 

(i) High concentrations of oxygen (11 volumes of O₂ are used).  

(ii) Removal of SO₃ after regular intervals.

(iii) Low temperature of 400°C –500°C is maintained; optimum temperature is 673 K (400°C)

(iv) High pressure of 1-1.5 atm is maintained.

(v) Suitable catalyst in the form of vanadium pentaoxide (V₂O₅)

(vi) Promoter in the form of K₂SO₄

 

The above reaction is reversible, exothermic all the components of reaction are in gaseous state and reaction takes place with decrease in volume of products.

Factors increasing the yield of SO₃

For maximum yield of SO₃, Le-Chatelier’s Principle suggests the following conditions:

1.   Effect of Concentration

When concentration of O₂ is increased at equilibrium, then according to Le-Chatelier’s Principle, the equilibrium is forced towards right in order to reduce the concentration of added O₂, thereby increasing the yield of SO₃. In actual practice oxygen is taken in excess. (Usually 11 volumes of O₂ is used, one volume is consumed and 10 volumes are left).

2.   Effect of Temperature

The formation of SO₃ is exothermic. According to Le-Chatelier’s Principle, at low temperature more SO₃ will be formed while at high temperature SO₃ will decompose into SO₂ and O₂ again. Thus low temperature helps in stabilization of SO₃ i.e. favours the formation of SO₃. [However, at low temperature, equilibrium will be established in longer time and rate of formation of SO₃ will be low]. In actual practice, the reaction is carried out at high temperature of 400–450°C which produces reduced yield at much higher rate (since a fast rate is more important than a high yield).

3.   Effect of Pressure

The synthesis of SO₃ takes place with decrease in volume of products (3 volumes of SO₂ and O₂ yield 2 volumes of SO₃). According to Le-Chatelier’s Principle, the equilibrium state will shift towards right on increasing the pressure which lowers the volume (concentration) of SO₃ and O₂. For industrial point of view, the pressure on the equilibrium mixture is maintained at 1-1.5 atmosphere.

4.   Effect of Catalyst

Finely divided Pt or Vanadium Pentaoxide (V₂O₅) is used in Contact Process to speed up the reaction.

Q1. Write down the effect of change in temperature on equilibrium.

Answer

Effect of Temperature Change

Effect of change in temperature is related to the nature of reaction whether it is an endothermic reaction or exothermic reaction. Temperature changes affect the equilibrium constant and rates of reactions.

According to Le-Chatelier’s Principle, when temperature of a system is increased at equilibrium, the equilibrium will shift to that direction in which heat is absorbed i.e. the increase in temperature shifts the equilibrium in the direction of the endothermic reaction while decrease in temperature shifts the equilibrium in the direction of exothermic reaction.

Raising The Temperature Increases K_eq For An Endothermic Reaction And Lowers K_eq For An Exothermic Reaction (Lowering The Temperature Has The Opposite Effect). 

Effect of Temperature on Exothermic Reactions

In Exothermic reactions, K_c decreases with the rise of temperature i.e. concentration of products decrease and the reaction shifts towards reverse direction in order to nullify the change. Thus decrease in temperature always favours an exothermic reaction.

T ↑ K_c ↓ Backward reaction ↑

Effect of Temperature on Endothermic Reactions

In endothermic reactions, K_c increases with the rise of temperature i.e. concentration of products increases and the reaction shifts towards forward direction in order to nullify the change. Thus increase in temperature always favours an endothermic reactions.

T↑ K_c↑ Forward reaction↑

 

Q2. State Le-Chatelier’s Principle. Describe the effect of change in concentration and change in pressure on Equilibrium. Explain the industrial application of Le-Chatelier’s principle using Haber’s process.

Answer

Statement of the Principle

If an external stress or constraint such as change in concentration, temperature or pressure is applied to a system at equilibrium, the equilibrium is disturbed and tends to shift in a direction to offset the effects of imposed stress in order to achieve a new equilibrium state.

Effect of Concentration Change

When  a system at equilibrium is disturbed by increasing or decreasing the concentrations of one or more species involved in the reaction, the equilibrium tends to shift towards left or right in order to reduced the effect of this stress and readjust itself until Q_c = K_c. Thus equilibrium is self-correcting. In general:

(i) The equilibrium position shifts towards the right if the amount of reactant is added or product is removed. Generally, rate of forward reaction will increase if an additional amount of any reactant is added at equilibrium to a system.

(ii) The equilibrium position shift to the left if the amount of reactant is removed or product is added. Generally, rate of reverse reaction will increase if an additional amount of any product is added to a system at equilibrium.

 

Illustration

consider the following system at equilibrium

CO_(g) + 3H_2(g) ⇌ CH_4(g) +H₂O_(g)

 

When certain amount of CO or H₂ is added to the system, the value of reacting quotient (Q_c) is lowered than its K_c and reaction is no longer at equilibrium. The stress caused by addition of CO or H₂ is offset by shifting the equilibrium to the right side. In doing so some amount of CO or H₂ has consumed turning the value of Q_c back to K_c. In kinetic aspect, the addition of CO or H₂ increases the rate of forward reaction by consuming added CO or H₂ thereby producing more CH₄ and H₂O until at a certain point a new equilibrium will be established. 

Effect of Pressure Change

The change of pressure will have no effect on the state of equilibrium, if:

(a) The system does not involve gaseous components.

(b) The number of moles of gaseous reactants is equal to the number of moles of gaseous products.

Pressure changes only affect homogenous gaseous equilibria in which there is a volume (or number of moles) change (∆n≠0). The effect of changing the pressure on a gas-phase reaction depends on the stoichiometry of the reaction.

In gaseous equilibria reactions involving volume change, an increase in pressure will shift the equilibrium toward the side of decreasing volumes or fewer moles of gas.

 

Manufacture of Ammonia by Haber’s Process

Principle

Ammonia is manufactured on large scale by Haber’s Process by the direct union of nitrogen and hydrogen. The reaction is reversible, exothermic involves gaseous equilibrium occurring with decrease in volume of products. Le-Chatelier’s Principle suggests the following conditions for the maximum yield of ammonia:

(i)  High concentrations of N₂ and H₂ in the ratio of 1:3 by volume

(ii)    Low temperature of 400°C –500°C is maintained.

(iii)  High pressure of 400-1000 atm is maintained.

(iv)   Suitable catalyst in the form of finely divided iron.

 

Q4. Using Le-Chatelier’s principle, explain three ways in which yield of SO₃ can be increased in  Contact process.                       2SO₂ + O₂ ⇌ 2SO₃ (∆H^o = ive)

Answer

Manufacture of Sulphuric Acid (H₂SO₄) by Contact Process

Sulphuric acid is manufactured on large scale by the Contact Process which is multi step process, the most important step of which is the oxidation of sulphur dioxide to sulphur trioxide by atmospheric oxygen. The reaction is reversible, exothermic involves gaseous equilibrium occurring with decrease in volume of products. Le-Chatelier’s Principle suggests the following conditions for the maximum yield of sulphur trioxide:

(i)     High concentrations of oxygen (11 volumes of O₂ are used). 

(ii)    Low temperature of 400°C –500°C is maintained.

(iii) High pressure of 1-1.5 atm is maintained.

(iv)   Suitable catalyst in the form of vanadium pentaoxide (V₂O₅) promoted by K₂SO₄. 

Q13. In the synthesis of nitric acid by Ostwald process, one of the important reactions is the oxidation of nitric oxide to nitrogen dioxide.  2NO_(g)  + O_2(g)  ⇌ 2NO₂   ∆H^o = − 114 kJ/mol

Use Le-Chatelier’s principle to predict the direction of reaction when the equilibrium is disturbed by

(a) Increasing the pressure                            

(b) Increasing the temperature

(c) Adding O₂                                                           

(d) Removing NO (Example 7.6; Page # 155)

Answer

Predicting effect of Increasing Pressure using Le-Chatelier’s Principle

Since reaction occurs by decrease in moles or volume from 3 moles of gaseous reactants to 2 moles of gaseous products, an increasing pressure shifts the equilibrium to the side where there is less no. of moles i.e. to the right thus more product (NO₂) will be formed. 

Predicting effect of increasing temperature using Le-Chatelier’s Principle

Since the forward reaction is exothermic as shown by the negative sign of ∆H, so increasing temperature shifts the equilibrium to the side where heat is absorbed i.e. to the left thus more reactants (NO and O₂) will be formed.

Predicting effect of adding more O₂ (reactant)

Adding more reactant here O₂ gas in the reaction mixture shifts the reaction to the right giving more products (NO₂).

Predicting effect of removing NO gas (product)

Removing product here NO gas from the reaction mixture shifts the system to the left producing more reactants. 

MDCAT-style MCQs on Le-Chatelier’s Principle with clear explanations 👇

MCQ 1

When the pressure on a gaseous system at equilibrium is increased, the equilibrium will shift:
A. Towards the side with more moles of gas
B. Towards the side with fewer moles of gas
C. Towards the side with equal moles of gas
D. Does not change

✅ Answer: B
🧠 Explanation:
According to Le-Chatelier’s Principle, increasing pressure favors the side with fewer gas molecules, to reduce pressure.

Example:

$$N_2+3H_2\rightleftharpoons 2N{H}_{\mathrm{3}}$$

(Left: 4 moles → Right: 2 moles)
If pressure increases → equilibrium shifts right, forming more NH₃.

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MCQ 2

In the equilibrium

$$2S{O}_2+O_2\rightleftharpoons 2S{O}_3+\text{heat}$$

What happens if temperature is increased?
A. More SO₃ is formed
B. More SO₂ and O₂ are formed
C. No change
D. Reaction becomes faster

✅ Answer: B
🧠 Explanation:
This is an exothermic reaction (heat on right).
If temperature increases → system opposes by shifting left (endothermic side) → more SO₂ & O₂ form.

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MCQ 3

If a catalyst is added to a system in equilibrium, the equilibrium:
A. Shifts to the right
B. Shifts to the left
C. Does not shift
D. Depends on reaction

✅ Answer: C
🧠 Explanation:
Catalyst does not change the equilibrium position; it only helps reach equilibrium faster by lowering activation energy.

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MCQ 4

For the equilibrium

$$N_2{O}_4(g)\rightleftharpoons 2N{O}_2(g)$$

Colorless → Brown
When pressure is increased, color of the mixture:
A. Becomes darker
B. Becomes lighter
C. No change
D. Turns colorless instantly

✅ Answer: B
🧠 Explanation:
Left side: 1 mole of gas, Right side: 2 moles.
↑ Pressure → shift to left → more colorless N₂O₄ → mixture becomes lighter.

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MCQ 5

In the reaction

$$H_2+I_2\rightleftharpoons 2HI$$

If some HI is removed from the system at equilibrium, what happens?
A. Reaction shifts to right
B. Reaction shifts to left
C. No change
D. Reaction stops

✅ Answer: A
🧠 Explanation:
Removing product HI disturbs equilibrium.
System shifts right to produce more HI — restoring balance.

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MCQ 6

In the equilibrium

$$CaC{O}_3(s)\rightleftharpoons CaO(s)+C{O}_2(g)$$

If CO₂ gas is removed continuously, the decomposition of CaCO₃:
A. Stops
B. Increases
C. Decreases
D. Remains same

✅ Answer: B
🧠 Explanation:
Removing product CO₂ shifts equilibrium right, causing more CaCO₃ to decompose.

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MCQ 7

At constant temperature, increasing volume of a gas system causes equilibrium to:
A. Shift to more moles
B. Shift to fewer moles
C. No effect
D. Shifts randomly

✅ Answer: A
🧠 Explanation:
Increasing volume = decreasing pressure → system shifts toward more gas moles to increase pressure again.

MCQ 8

For the reaction

$$2NOCl(g)\rightleftharpoons 2NO(g)+C{l}_2(g)$$

What happens when the pressure is increased?

A. Shifts to right

B. Shifts to left

C. No effect

D. Depends on catalyst

✅ Answer: B

🧠 Explanation:

Left = 2 moles, Right = 3 moles of gas.

Increasing pressure favors the side with fewer moles → equilibrium shifts left.

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MCQ 9

In the reaction

$$N_2+O_2\rightleftharpoons 2NO(\text{endothermic})$$

Increasing temperature will:

A. Increase yield of NO

B. Decrease yield of NO

C. No effect

D. Stop reaction

✅ Answer: A

🧠 Explanation:

Reaction is endothermic, so heat is absorbed on the forward side.

Increasing temperature favors the endothermic (forward) direction → more NO is produced.

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MCQ 10

Which of the following changes will not affect equilibrium position?

A. Pressure

B. Temperature

C. Catalyst

D. Concentration

✅ Answer: C

🧠 Explanation:

Catalyst changes rate, not position of equilibrium.

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MCQ 11

In a closed container:

$$CO(g)+H_{2}O(g)\rightleftharpoons C{O}_2(g)+H_2(g)$$

If steam (H₂O) is added, equilibrium will:

A. Shift right

B. Shift left

C. No effect

D. Stop

✅ Answer: A

🧠 Explanation:

Adding reactant (H₂O) shifts equilibrium right to use up extra steam, forming more CO₂ and H₂.

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MCQ 12

In the equilibrium mixture of

$$2N{O}_2(g)\rightleftharpoons N_2{O}_4(g)$$

Colorless ⇌ Brown

If temperature decreases, what happens to color?

A. Becomes darker

B. Becomes lighter

C. No effect

D. Turns brown

✅ Answer: B

🧠 Explanation:

Formation of N₂O₄ is exothermic.

Lowering temperature → shifts right → more colorless N₂O₄ → lighter color.

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MCQ 13

When equilibrium is disturbed by adding more product, the system will:

A. Shift right

B. Shift left

C. No change

D. Stop reaction

✅ Answer: B

🧠 Explanation:

Adding product increases its concentration → equilibrium shifts left to reduce it (reverse reaction).

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MCQ 14

For the reaction

$$H_2(g)+C{l}_2(g)\rightleftharpoons 2HCl(g)$$

If volume is decreased, what happens?

A. Shifts to right

B. Shifts to left

C. No effect

D. Reaction stops

✅ Answer: C

🧠 Explanation:

Both sides have 2 moles of gas → changing pressure/volume has no effect.

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MCQ 15

In which of the following cases does temperature not affect equilibrium position?

A. Exothermic reaction

B. Endothermic reaction

C. Reaction with no heat change

D. All of these

✅ Answer: C

🧠 Explanation:

If no heat is absorbed or released (ΔH = 0), temperature has no effect on equilibrium.

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MCQ 16

At equilibrium, the rate of forward reaction compared to reverse reaction is:

A. Greater

B. Smaller

C. Equal

D. Zero

✅ Answer: C

🧠 Explanation:

At equilibrium, both reactions continue but at equal rates, keeping concentrations constant.

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MCQ 17

For reaction

$$2HI(g)\rightleftharpoons H_2(g)+I_2(g)$$

If temperature is increased, equilibrium shifts:

A. To right

B. To left

C. No change

D. Stops

✅ Answer: B

🧠 Explanation:

Decomposition of HI is endothermic in reverse, so increasing temperature shifts left, forming more H₂ and I₂.

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MCQ 18

Which condition favors maximum yield of ammonia in the Haber process?

$$N_2+3H_2\rightleftharpoons 2N{H}_3+\text{heat}$$

A. High temperature, high pressure

B. Low temperature, low pressure

C. High pressure, low temperature

D. Low pressure, high temperature

✅ Answer: C

🧠 Explanation:

Reaction is exothermic and produces fewer gas moles.

→ High pressure favors right side.

→ Low temperature favors exothermic side (NH₃ formation).