Explanation: Chemical equilibrium is dynamic in nature. At equilibrium, the reaction does not stop; instead, the rate of the forward reaction becomes exactly equal to the rate of the backward reaction. As a result, the concentrations of reactants and products remain constant over time.

Explanation: A catalyst increases the rates of both forward and backward reactions equally by lowering the activation energy barrier. It does not affect the position of equilibrium or the value of the equilibrium constant (K_c). It only helps in **attaining the equilibrium faster** (decreases the time required).

Explanation: For K_p = K_c, the change in the number of gaseous moles (Δn_g) must be zero.
– In reaction C: H₂(g) + I₂(g) &rightleftharpoons 2HI(g)
Δn_g = 2 – (1 + 1) = 0.
Therefore, K_p = K_c(RT)⁰ = K_c.

Explanation: The unit of K_p is given by (atm)^Δn_g or (bar)^Δn_g.
For the given reaction: Δn_g = 2 – (1 + 3) = -2.
Thus, the unit of K_p is atm⁻² (or bar⁻²).

Explanation: According to the properties of the equilibrium constant, when a chemical equation is reversed, the equilibrium constant of the reverse reaction becomes the reciprocal of the forward reaction’s constant.
Thus, K_reverse = 1 / K.

Explanation: When a balanced chemical equation is multiplied by a factor ‘n‘, the new equilibrium constant is raised to the power of ‘n‘.
Here, the equation is multiplied by 2, so the new constant K_2 = K_1².

Explanation: When two or more chemical equations are added to yield a net chemical equation, the equilibrium constant of the net equation is the product of the equilibrium constants of the individual steps.
Hence, K = K₁ × K₂.

Explanation: According to Le Chatelier’s Principle:
– **Pressure effect:** Increasing pressure shifts equilibrium to the side with fewer gas moles (from 4 moles of reactants to 2 moles of product). Thus, **high pressure** is favorable.
– **Temperature effect:** Since the reaction is exothermic, decreasing the temperature shifts the equilibrium in the forward direction. Thus, **low temperature** is favorable.

Explanation: Adding an inert gas at constant volume increases the total pressure, but it does not alter the partial pressures or molar concentrations of the reacting species. Since the active concentrations remain unchanged, there is **no effect** on the state of equilibrium. (Inert gas addition at constant pressure shifts the equilibrium to the side with a greater number of moles).

Explanation: Predicting the direction of a chemical reaction:
– **Q_c < K_c:** The reaction proceeds in the forward direction.
– **Q_c = K_c:** The system is at equilibrium.
– **Q_c > K_c:** The concentration of products is higher than at equilibrium. To attain equilibrium, the reaction proceeds in the **backward/reverse direction**.

Explanation: In heterogeneous equilibrium, the active mass (molar concentration) of pure solids and pure liquids is constant and taken as **unity (1)**.
Since CaCO₃(s) and CaO(s) are pure solids, [CaCO₃] = 1 and [CaO] = 1.
Thus, the equilibrium constant simplifies to: K_c = [CO₂] (and K_p = p_CO₂).

Explanation: For a given chemical reaction, the value of the equilibrium constant is independent of initial concentrations, volume, pressure, or the presence of a catalyst. Its value changes only with a change in **temperature** (as described by the van ‘t Hoff equation).

Explanation: The general relation is α = (D – d) / ((n – 1)d), where ‘n‘ is the number of product molecules formed from one reactant molecule.
For PCl₅(g) &rightleftharpoons PCl₃(g) + Cl₂(g), n = 2.
Substituting n = 2 into the formula: α = (D – d) / ((2 – 1)d) = (D – d) / d.

Explanation: According to Le Chatelier’s Principle, increasing pressure shifts the equilibrium to the side with lower volume (higher density).
Since liquid water has a higher density than ice, it occupies a smaller volume. Increasing pressure shifts the equilibrium towards the side of lower volume, i.e., water. Consequently, **more ice melts** into water, decreasing the melting point of ice.

Explanation: Vapour pressure at a given equilibrium is an intensive property and depends solely on the **nature of the liquid and the temperature**. It remains independent of the volume of the vessel, surface area, or the amount of liquid present.

Explanation: Utilizing the relation ΔG° = -2.303 RT log K_c:
– If K_c > 1, then log K_c is positive.
– The negative sign in the equation ensures that ΔG° becomes **negative** (ΔG° < 0), meaning the reaction is thermodynamically spontaneous under standard conditions.

Explanation:
– Formula: K_p / K_c = (RT)^Δn_g
– For the reaction: Δn_g = 1 – (1 + ½) = -½
– Therefore, K_p / K_c = (RT)^-1/2 or 1 / √(RT).

Explanation: Pure solids like NH₄HS(s) have no partial pressure in the gas phase (their active mass is constant and taken as 1). Hence, the solid reactant is omitted from the K_p expression, which depends only on the partial pressures of the gaseous products: K_p = p_NH₃ × p_H₂S.

Explanation: In the heterogeneous physical equilibrium H₂O(l) &rightleftharpoons H₂O(g), the concentration of the pure liquid water is constant and taken as 1.
Thus, K_p = p_H₂O(g), which is the vapor pressure of water at that specific temperature.

Explanation: The magnitude of the equilibrium constant indicates the extent of a reaction.
– If K_c > 10³: Products predominate over reactants. A very large value (like 10¹²) indicates that the reaction goes nearly to completion.
– If K_c < 10⁻³: Reactants predominate, meaning the reaction barely proceeds.

Explanation: Sodium oxalate reacts with Fe³⁺ to form a stable complex, thereby decreasing the concentration of free Fe³⁺ ions in the equilibrium mixture. According to Le Chatelier’s Principle, decreasing reactant concentration shifts the equilibrium in the **backward direction**, decreasing the concentration of the red-colored [Fe(SCN)]²⁺ complex.

Explanation:
– Relationship: K_p / K_c = (RT)^Δn_g
– In this reaction, Δn_g = (1 + 1) – (1 + 1) = 0.
– Therefore, K_p / K_c = (RT)⁰ = 1.

Explanation: According to Le Chatelier’s Principle, increasing the pressure shifts the equilibrium towards the side with fewer moles of gas.
In A(g) &rightleftharpoons 2B(g):
– Reactant gas moles = 1
– Product gas moles = 2
Since there are fewer gas moles on the reactant side, increasing pressure will shift the equilibrium to the **backward direction**.

Explanation: In a homogeneous chemical equilibrium, all reactants and products are present in the same phase (either all gases or all in a single solution).
– In reaction B, N₂, H₂, and NH₃ are all in the gaseous phase. Hence, it is homogeneous.
– The other reactions (A, C, D) contain different phases (solids/liquids alongside gases) and are classified as heterogeneous.

Explanation: Once chemical equilibrium is achieved, the rates of forward and backward reactions are equal. Consequently, the concentrations of both reactants and products **become constant** (do not change with time). They do not need to be equal to each other, just constant.