For the reversible reaction, `N_(2)(g)+3H_(2) (g) Leftrightarrow 2NH_(3) (g)+" heat"`. The equilibrium shifts in forward direction
(1) (1) By increasing the concentration of `NH_(3)(g)`
(2) By decreasing the pressure
(3) By decreasing concentration of `N_2(g) and H_2(g)`
(4) By increasing the pressure and decreasing temperature

Answer

To determine which condition will shift the equilibrium of the reaction \( N_2(g) + 3H_2(g) \leftrightarrow 2NH_3(g) + \text{heat} \) in the forward direction, we can apply Le Chatelier's Principle. This principle states that if an external change is applied to a system at equilibrium, the system will adjust itself to counteract that change and restore a new equilibrium.

Let's analyze each option step by step:

### Step 1: Analyze the first option - Increasing the concentration of \( NH_3(g) \)
- When we increase the concentration of \( NH_3(g) \), the system will respond by shifting the equilibrium to the left (backward direction) to reduce the concentration of \( NH_3 \).
- **Conclusion**: This option does not shift the equilibrium in the forward direction.

### Step 2: Analyze the second option - Decreasing the pressure
- Decreasing the pressure favors the side of the reaction with more moles of gas. In this reaction, we have 4 moles of reactants (1 mole of \( N_2 \) and 3 moles of \( H_2 \)) and 2 moles of products (2 moles of \( NH_3 \)).
- Since there are more moles on the reactant side, decreasing the pressure will shift the equilibrium to the left (backward direction).
- **Conclusion**: This option does not shift the equilibrium in the forward direction.

### Step 3: Analyze the third option - Decreasing the concentration of \( N_2(g) \) and \( H_2(g) \)
- Decreasing the concentration of reactants \( N_2 \) and \( H_2 \) will cause the system to shift in the direction that increases their concentration, which is the forward direction (to produce more \( NH_3 \)).
- However, since we are decreasing the concentration of the reactants, it will not favor the forward reaction effectively.
- **Conclusion**: This option does not effectively shift the equilibrium in the forward direction.

### Step 4: Analyze the fourth option - Increasing the pressure and decreasing temperature
- Increasing the pressure favors the side with fewer moles of gas. In this reaction, the product side has 2 moles (2 \( NH_3 \)) compared to 4 moles on the reactant side. Thus, increasing the pressure will shift the equilibrium to the right (forward direction).
- Decreasing the temperature favors exothermic reactions. Since the formation of \( NH_3 \) releases heat (exothermic), lowering the temperature will also shift the equilibrium to the right (forward direction).
- **Conclusion**: This option shifts the equilibrium in the forward direction.

### Final Answer:
Based on the analysis, the correct condition that shifts the equilibrium in the forward direction is:
**Option 4: By increasing the pressure and decreasing temperature.**

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