could not figure out how to add the 4 subscript for CH

What is the volume (in mL) of 0.0100 moles of CH_{4} gas at STP?

Here, we know to use the ideal gas law equation, which relates the volume, temperature, and pressure of an ideal gas:
**PV=nRT** ,

where **P** = the pressure of the gas in atmospheres (atm)

**V** = the volume of the space the gas is contained in in liters (L)

**n** = the number of moles of the gas (mol) = **
0.0100 mol**

**R** = the universal gas constant, **0.0821 **^{L·atm}**/**_{K·mol }

**T** = the absolute temperature of the gas in Kelvins (K)

A gas at **standard temperature and pressure** (STP) is defined at a temperature of
**0°C** and at a pressure of **1 atm**.

Since the the ideal gas law requires the absolute temperature of a gas, we have to convert the temperature of the gas at STP from °C to Kelvins:

K = °C + 273.15 ==> K = 0°C + 273.15 = **
273.15 K**

Sine we are solving for the volume of the gas, we take the equation PV=nRT and divide both sides by P to solve for V:

**V = nRT/P**

**V** = [(0.0100 mol)*(0.0821 ^{L·atm}/_{K·mol})*(273.15 K)] / (1 atm) = 0.22425615 L ≈ 0.22426 L

(0.22426 L) * (1000 mL / 1 L) = **224.26 mL**

## Comments

I love using the STP trick! It's important to understand how and when you use PV=nRT, but if they give you an unknown mass at STP, why go through all of the extra steps (which gives the opportunity for mistakes)? Chances are using the 1 mol @ STP method is what they want you to do for this question, and all they'll give you time to do for a question like this on the exam.