When hydrochloric aicd is addded to cobalt and nitrate solution at room temperautre, the following reaction takes place and the reaction mixture becomes blue. On cooling the mixture it becomes pink. On the basis of this information mark the corect ansewer.
`underset("pink")[[Co(H_(2)O)_(6)]]^(3+)(aq)+4CI^(-)hArr`
`underset("blue")[CoCI_(4)]^(2-)(aq)+6H_(2)O(l)` (A)Δ H > 0 for the reaction (B)Δ H < 0 for the reaction (C)Δ H = 0 for the reaction (D)The sign of Δ H cannot be predicted on the basis of this information

Options

A. `DeltaHgt0` for the reaction
B. `DeltaHlt0` for the reaction
C. `DeltaH=0` for the reaction
D. The sign of `DeltaH` cannot be predicted on the basis of this information.

Explanation

To solve the problem step by step, we will analyze the information provided and apply the principles of chemical equilibrium, particularly Le Chatelier's principle.

### Step 1: Understand the Reaction
The reaction provided is:
\[ \text{[Co(H}_2\text{O)}_6]^{3+}(aq) + 4 \text{Cl}^-(aq) \rightleftharpoons \text{[CoCl}_4]^{2-}(aq) + 6 \text{H}_2\text{O}(l) \]

When hydrochloric acid (HCl) is added to a cobalt nitrate solution, the solution initially turns blue due to the formation of the complex ion \([CoCl_4]^{2-}\).

**Hint:** Identify the color changes associated with the reactants and products to understand the direction of the reaction.

### Step 2: Analyze the Color Change
- The blue color indicates the presence of \([CoCl_4]^{2-}\).
- Upon cooling, the mixture turns pink, indicating the formation of \([Co(H_2O)_6]^{3+}\), which is the reactant.

**Hint:** Color changes can indicate shifts in equilibrium; cooling favors the formation of reactants in this case.

### Step 3: Apply Le Chatelier's Principle
According to Le Chatelier's principle, if a system at equilibrium is subjected to a change in temperature, pressure, or concentration, the system will adjust to counteract that change.

- Since cooling the mixture causes the reaction to shift towards the reactants (pink color), this suggests that the forward reaction (formation of products) is endothermic.

**Hint:** Remember that endothermic reactions absorb heat, so lowering the temperature shifts the equilibrium towards the reactants.

### Step 4: Determine the Sign of ΔH
For an endothermic reaction:
- When the temperature decreases, the equilibrium shifts to the left (towards the reactants).
- This means that the enthalpy change (\(ΔH\)) for the reaction is greater than zero.

Thus, we conclude that:
\[ ΔH > 0 \]

**Hint:** Endothermic reactions have a positive \(ΔH\) because they require heat to proceed.

### Step 5: Choose the Correct Answer
Based on the analysis:
- (A) \(ΔH > 0\) for the reaction is correct.
- (B) \(ΔH < 0\) is incorrect (exothermic).
- (C) \(ΔH = 0\) is incorrect (no heat change).
- (D) The sign of \(ΔH\) cannot be predicted is incorrect (we have determined it).

**Final Answer:** The correct option is (A) \(ΔH > 0\).