Law Of Chemical Equilibrium And Equilibrium Constant

Step 1: Understand the equilibrium concept. In this reaction, as the system approaches equilibrium, the volume decreases because the number of moles of gas decreases. This leads to a decrease in pressure.
Step 2: Conclusion. Thus, the pressure decreases as the reaction progresses.

Step 1: Le Chatelier's Principle. According to Le Chatelier's principle, when the pressure is increased on ice at constant temperature, the equilibrium will shift to the side with fewer molecules of gas, which is the solid phase.
Step 2: Conclusion. Thus, the water will freeze.

Step 1: Effect of temperature on equilibrium.
For an endothermic reaction (positive ), increasing the temperature shifts the equilibrium to the right, favoring the forward reaction.

Step 2: Conclusion.
Thus, the correct answer is option (B).

Final Answer:

Step 1: Understanding the equilibrium.
The reaction is: If 80% of hydrogen reacts, this gives us the concentration of reactants and products, which can be used to calculate .

Step 2: Conclusion.
Thus, the correct value of is 64. Hence, the correct answer is option (A).

Final Answer:

Step 1: Temperature dependence of equilibrium constant.
The equilibrium constant depends on temperature, as shown by the van't Hoff equation. A change in temperature alters the position of equilibrium.

Step 2: Conclusion.
Thus, the correct answer is option (D).

Final Answer:

Step 1: Use relation between and equilibrium constant.

Step 2: Convert into J.

Step 3: Substitute values.


Step 4: Convert to .

Final Answer:

Step 1: Given equilibrium constant.
For reaction:

Step 2: Reverse the reaction.
Reversing gives:

So new constant becomes:

Step 3: Divide the reaction by 2.
Required reaction is:

So equilibrium constant becomes:

Final Answer:

Step 1: Use Henderson-Hasselbalch equation.

Step 2: Find .

Step 3: Substitute given pH.


Step 4: Calculate moles of acid.

Step 5: Moles of salt needed.

Step 6: Find salt concentration.
Salt volume = .

Thus as per answer key:

Final Answer:

Step 1: Use definition of buffer capacity.
Buffer capacity is:

where = moles of acid/base added, = volume in litres.
Step 2: Calculate moles of acetic acid added.
Molar mass of .

Step 3: Given pH change.

Volume:

Step 4: Compute buffer capacity.

Final Answer:

Step 1: Given equilibrium.
Step 2: Required reaction is reverse and half.
Target:
This is exactly of the reverse of the given reaction.
Step 3: Reverse reaction constant.
Reverse of given reaction:
Step 4: Take half reaction.
When coefficients are divided by 2, equilibrium constant becomes square root:
Final Answer:

Step 1: Fundamental relation.
Thermodynamics gives:
Step 2: Rearranging for .
Taking exponential:
Final Answer:

Step 1: Analyze effect of temperature (Le Chatelier principle).
Reaction gives on product side, so reaction is exothermic.
For exothermic reactions, lowering temperature shifts equilibrium towards products.
Step 2: Analyze effect of pressure.
Reactant moles of gas:
Product moles of gas:
Since products have fewer moles, increasing pressure shifts equilibrium towards products.
Step 3: Combine both effects.
To maximize product yield:
- low temperature (favours exothermic direction)
- high pressure (favours fewer moles side)
Final Answer:

Step 1: Meaning of equilibrium constant.
If is very large ( ), equilibrium lies far to the right, meaning products dominate.
If , reactants dominate.
Step 2: Compare values.
Largest value is .
Step 3: Conclusion.
So reaction goes farthest to completion when .
Final Answer:

Step 1: Understand Le-Chatelier's Principle.
According to Le-Chatelier's principle, if a reaction is exothermic ( is negative), lowering the temperature will favour the product formation. For reactions involving gases, high pressure favours the side with fewer moles of gas.
Step 2: Conclusion.
Thus, high pressure and low temperature would favour the formation of C, but this is not exactly one of the provided options. The most fitting option is (A).
Final Answer:

Step 1: Understand the reaction.
The reaction is: At equilibrium, the total pressure is the sum of the pressures of CO and CO.
Step 2: Calculation.
Given the initial pressure and the equilibrium pressures, calculate the equilibrium constant for this reaction.
Final Answer:

By examining the equations and their corresponding types of processes, we match them based on the definitions of spontaneity, equilibrium, and other thermodynamic principles.

Step 2: Conclusion.
The correct matching of equations to processes is option (c).

This reaction is reversible at high temperatures, where magnesium oxide reacts with chlorine gas to produce magnesium chloride and oxygen. The reaction is temperature dependent, and high temperatures favor the forward direction.

Step 2: Conclusion.
The reversible reaction that exhibits high-temperature behaviour is , corresponding to option (a).

Step 1: Analyzing the reaction.
When is heated, it dissociates into ammonia and HCl. The ammonia diffuses faster than HCl, which changes the litmus paper.

Step 2: Conclusion.
The correct answer is that the ammonia diffuses faster than HCl, corresponding to option (3).

The equilibrium constant for a reversible reaction is the ratio of the rate constants for the forward and reverse reactions: where is the rate constant for the forward reaction and is the rate constant for the backward reaction.

Increasing the amount of carbon shifts the equilibrium towards the products, increasing the concentrations of CO and according to Le Chatelier's principle.

Step 1: Equilibrium Expression.
The equilibrium expression for this reaction is given by: Using the mole fractions and pressures at equilibrium, we can derive the expression for .
Step 2: Conclusion.
The correct answer is (A), .

Step 1: Le Chatelier’s Principle.
According to Le Chatelier’s principle, increasing the temperature of an exothermic reaction (like the formation of ammonia) will shift the equilibrium to the left, decreasing the concentration of ammonia at higher temperatures.
Step 2: Conclusion.
The correct answer is (B), vs (200°C to 400°C) is decreasing.

Step 1: Apply Le Chatelier's principle.
To increase the quantity of , we can shift the equilibrium towards the product side. Adding more will push the reaction towards the formation of .
Step 2: Conclusion.
Thus, adding more will increase the quantity of .
Final Answer:

Step 1: Understand the relationship between and .
The relation between and is given by: where is the change in the number of moles of gas, and is the gas constant.
Step 2: Conclusion.
Thus, the value of depends on the total gas pressure and the change in the number of moles.
Final Answer:

Step 1: In reactions with a positive , the equilibrium position shifts with temperature. The pressure of is influenced by the temperature at equilibrium.

Final Answer:

Step 1: The equilibrium constant for the new reaction can be derived by taking the square root of the equilibrium constant for the original reaction.
Step 2: Since the stoichiometry is halved, the equilibrium constant for the new reaction is:
Final Answer:

Step 1: According to Le Chatelier’s principle, increasing the volume of a system at equilibrium will shift the equilibrium to the side with more moles of gas.
Step 2: In the case of , the left side has 4 moles of gas, and the right side has 2 moles. Therefore, increasing the volume will shift the equilibrium to the right.

Final Answer:

We can use the relationship between and for ideal gases: where is the change in the number of moles of gas, , and . After calculating , we find that the value of is .
Final Answer:

For a reaction involving gases, according to Le Chatelier’s principle, reducing the volume shifts the equilibrium towards the side with fewer moles of gas. In this case, the right side has fewer moles of gas, so the equilibrium shifts to the right.
Final Answer:

- The reaction of ethylbenzene with results in oxidation, forming a carboxylic acid group. - The subsequent reaction with leads to substitution at the benzylic position. - Finally, the hydration of the product gives the final product as , i.e., a bromo-substituted product.

Step 2: Conclusion.
Thus, the correct product is (4) Br.

The reaction is exothermic. Lowering temperature favors product formation, increasing .

1 mole of methane needs 2 moles of oxygen. So, 2 moles need 4 moles of oxygen.

Strong acid + its salt does not form a buffer.

Increasing pressure shifts equilibrium towards the side with fewer moles of gas to reduce pressure.