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The question in the image asks to calculate the volume of dinitrogen monoxide gas collected, given that 62.6 mmol of gas is evolved at 16.0 °C and 1 atm pressure.

We can solve this using the **Ideal Gas Law**:

\[
PV = nRT
\]

Where:
- \(P\) = pressure (1 atm)
- \(V\) = volume (what we are solving for)
- \(n\) = moles of gas (62.6 mmol = 0.0626 mol)
- \(R\) = gas constant (\(0.0821 \, \text{L·atm/mol·K}\))
- \(T\) = temperature in Kelvin (16.0 °C = 16.0 + 273.15 = 289.15 K)

Now, let's solve for \(V\):

\[
V = \frac{nRT}{P}
\]

Substituting the known values:

\[
V = \frac{(0.0626 \, \text{mol}) \times (0.0821 \, \text{L·atm/mol·K}) \times (289.15 \, \text{K})}{1 \, \text{atm}}
\]

Let's calculate:

\[
V = 1.482 \, \text{L}
\]

Thus, the volume of dinitrogen monoxide gas collected is **1.48 L** (rounded to three significant figures).

Would you like a more detailed explanation or any other clarifications?

Here are 5 related questions to think about:
1. How does the ideal gas law change if the pressure is not constant?
2. What happens to the volume if the temperature is doubled while keeping the amount of gas constant?
3. How does the volume change if we increase the number of moles of gas in the same conditions?
4. How would the answer differ if the gas constant \(R\) were given in different units (e.g., J/mol·K)?
5. What are the limitations of the ideal gas law in real-world conditions?

**Tip**: Always convert temperature to Kelvin when using the ideal gas law to avoid errors in your calculations.

A reaction at 16.0 °C evolves 62.6 mmol of dinitrogen monoxide gas. Calculate the volume of dinitrogen monoxide gas that is collected. You can assume the pressure in the room is exactly 1 atm.

Learn how to calculate the volume of dinitrogen monoxide gas collected from a reaction at 16.0 °C and 1 atm pressure. This problem uses the ideal gas law (PV = nRT), and involves converting mmol to moles and Celsius to Kelvin. Suitable for students in Grades 10-12.

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