Boundless Chemistry

STP most commonly is used when performing calculations on gases, such as gas density is calculated using stp = Volume of Gas*(273/Temperature of Gas)*(Pressure of Gas/100). To calculate STP, you need Volume of Gas (V), Temperature of Gas (T) and Pressure of Gas (P).

Answer: Explain:- 22.4L Imol Page 10 Many gas law problems involve calculating the volume of a gas produced by the reaction of volumes of other gases.

1. The ideal gas equation is given by PV=nRT P V = n R T .
2. PV=nRT.
3. 8.3145L⋅kPaK⋅mol=0.0821L⋅atmK⋅mol=62.4L⋅mm HgK⋅mol.

It can be written as: V = nRT/P. “P” is pressure, “V” is volume, n is the number of moles of a gas, “R” is the molar gas constant and “T” is temperature.

22.4L

Gas density is defined as the mass of the gas occupying a certain volume at specified pressure and temperature. The density is usually represented in units of lbm/ft3. Another common density representation is the “gas gradient” that is expressed in units of psi/ft.

Multiply the volume and pressure and divide the product by the temperature and the molar gas constant to calculate moles of the hydrogen gas. In the example, the amount of hydrogen is 202,650 x 0.025 / 293.15 x 8.314472 = 2.078 moles.

The equations describing these laws are special cases of the ideal gas law, PV = nRT, where P is the pressure of the gas, V is its volume, n is the number of moles of the gas, T is its kelvin temperature, and R is the ideal (universal) gas constant.

11.2 L

22.41 L

Argon

Answer. Explanation: 22.4 liters pero bat wala sa mga choices..

Answer. Answer: The volume occupied is 33.6 Liters. Explanation: STP conditions are known as standard temperature and pressure.

The molar volume of a gas is the volume of one mole of a gas at STP. Avogadro’s hypothesis states that equal volumes of any gas at the same temperature and pressure contain the same number of particles. At standard temperature and pressure, 1 mole of any gas occupies 22.4 L.

Standard Pressure is 1 Atm, 101.3kPa or 760 mmHg or torr. At STP, 1 mole of any gas occupies 22.4L.

A gas industry reference base, and may vary from country to country. Normal, or ambient temperature is 70 °F (21° C). Normal pressure is one atmosphere (1013 hPa) or 14,696 psia.

Standard temperature and pressure (STP) is defined as exactly 100 kPa of pressure (0.986 atm) and 273 K (0°C). This makes for a very useful approximation: any gas at STP has a volume of 22.4 L per mole of gas; that is, the molar volume at STP is 22.4 L/mol (Figure 6.3 “Molar Volume”).

Hence, at 750 K and 20 kPa conditions of temperature and pressure does a sample of helium behave most like an ideal gas.

The ideal gas works properly when the inter-molecular interactions between the gas molecules and volume of gas molecule will be negligible. This is possible when pressure is low and temperature is high. Therefore, the correct option is (3) 500 K and 0.1 atm.

In summary, a real gas deviates most from an ideal gas at low temperatures and high pressures. Gases are most ideal at high temperature and low pressure.

Real gas behaves like ideal gas at high temperature and low pressure.

helium

At relatively low pressures, gas molecules have practically no attraction for one another because they are (on average) so far apart, and they behave almost like particles of an ideal gas. At higher pressures, however, the force of attraction is also no longer insignificant.

Why do real gases behave so differently from ideal gases at high pressures and low temperatures? Because the molecules of an ideal gas are assumed to have zero volume, the volume available to them for motion is always the same as the volume of the container.

They never show ideal behavior because of inter-molecular forces (but less stronger than the forces at fluid and at solid states). These forces increase at low temperature because the movement of molecules decrease at these conditions, and the distances between the molecules decrease under high pressure.

Gases are unlike other states of matter in that a gas expands to fill the shape and volume of its container. For this reason, gases can also be compressed so that a relatively large amount of gas can be forced into a small container.