For the equilibrium `2NO_(2)(g)hArrN_(2)O_(4)(g) + 14.6` Kcal the increase in temperature would

Options

A. Favour the formation of `N_(2)O_(4)`
B. Favour the decomposition of `N_(2)O_(4)`
C. Not alter the equilibrium
D. Stop the reaction

Answer

To solve the question regarding the equilibrium reaction:

\[ 2NO_2(g) \rightleftharpoons N_2O_4(g) + 14.6 \text{ kcal} \]

we need to analyze the effect of temperature on this exothermic reaction.

### Step-by-Step Solution:

**Step 1: Identify the nature of the reaction.**
- The given reaction is exothermic because it releases heat (14.6 kcal).

**Hint:** Recall that exothermic reactions release heat, which can be represented as a product in the equilibrium expression.

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**Step 2: Apply Le Chatelier's Principle.**
- Le Chatelier's Principle states that if a system at equilibrium is subjected to a change in conditions (such as temperature, pressure, or concentration), the system will adjust to counteract that change and restore a new equilibrium.

**Hint:** Remember that increasing temperature in an exothermic reaction shifts the equilibrium to favor the endothermic direction (which absorbs heat).

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**Step 3: Determine the effect of increasing temperature.**
- Since the reaction is exothermic, increasing the temperature will shift the equilibrium to the left (towards the reactants) to absorb the added heat.

**Hint:** Think about how the system responds to added heat; it will favor the direction that consumes heat.

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**Step 4: Conclude the effect on the products and reactants.**
- As a result of the temperature increase, the equilibrium will favor the decomposition of \( N_2O_4 \) back into \( NO_2 \).

**Hint:** Visualize the reaction: if you add heat, the system will try to use that heat up by producing more reactants.

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**Final Answer:**
The increase in temperature would favor the decomposition of \( N_2O_4 \) back into \( NO_2 \).

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