At `47^(@)C` and 16.0 atm, the molar volume of NH3 gas is about 10% less than the molar volume of an ideal gas. This is due to :

Options

A. `NH_(3)` decomposes to `N_(2) and H_(2)` at
B. The force of attraction between `NH_(3)` molecules is significant at this temperature and pressure
C. The volume occupies by `NH_(3)` molecules themselves is a significant fraction of the volume of the container at this temperature and pressure
D. at 16 atm , `NH_(3)` molecules no longer move randomly

Explanation

To solve the question regarding the molar volume of NH3 gas being about 10% less than that of an ideal gas at 47°C and 16.0 atm, we can follow these steps:

### Step 1: Understand Ideal Gas Behavior
Ideal gases are defined by the ideal gas equation:
\[ PV = nRT \]
In this equation, \( P \) is the pressure, \( V \) is the volume, \( n \) is the number of moles, \( R \) is the ideal gas constant, and \( T \) is the temperature in Kelvin. Ideal gases do not exhibit intermolecular forces and occupy the entire volume of the container.

**Hint:** Remember that ideal gases do not experience attractions or repulsions between molecules.

### Step 2: Identify Real Gas Behavior
Real gases, like NH3, deviate from ideal behavior under certain conditions, particularly at high pressures and low temperatures. Under these conditions, the forces of attraction between molecules become significant, which affects the volume they occupy.

**Hint:** Consider how intermolecular forces affect the behavior of gases, especially under high pressure and low temperature.

### Step 3: Analyze the Conditions Given
The question states that at 47°C (which is approximately 320 K) and 16 atm, the molar volume of NH3 is about 10% less than that of an ideal gas. This indicates that the conditions are conducive to real gas behavior, where intermolecular forces are significant.

**Hint:** High pressure compresses gas molecules, leading to increased interactions between them.

### Step 4: Apply Van der Waals Equation
The Van der Waals equation accounts for the volume occupied by gas molecules and the intermolecular forces:
\[ \left(P + \frac{a n^2}{V^2}\right)(V - nb) = nRT \]
Here, \( a \) accounts for the attractive forces between molecules, and \( b \) accounts for the volume occupied by the gas molecules themselves.

**Hint:** Recognize that the Van der Waals equation modifies the ideal gas law to account for real gas behavior.

### Step 5: Conclude the Reason for Deviation
The 10% reduction in molar volume compared to an ideal gas can be attributed to the significant attractive forces between NH3 molecules due to its polar nature. At high pressures, these forces pull the molecules closer together, reducing the volume they occupy compared to an ideal gas.

**Hint:** Focus on how molecular interactions can lead to deviations from ideal gas behavior.

### Final Answer
The reason for the molar volume of NH3 being about 10% less than that of an ideal gas at the given conditions is due to the significant attractive forces between NH3 molecules.