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A sample of nitrogen gas is stored in a 500.0ml flask at 108kpa and 10.0c the gas is transferred to a 750.0ml flask at 21.0c what is the pressure of nitrogen in the second flask?
To find the pressure of the nitrogen gas in the second flask, you can use the combined gas law equation: P1V1/T1 = P2V2/T2, where P1, V1, and T1 are the initial pressure, volume, and temperature, and P2, V2, and T2 are the final pressure, volume, and temperature. Plug in the given values to find the final pressure of nitrogen in the second flask.
Why should the flask be covered not stoppered?
Covering the flask instead of stoppering it allows for any gas produced during the reaction to escape, preventing pressure build-up that could lead to the flask exploding. Stoppering the flask could trap the gas and result in a dangerous situation.
Why did the water rush into the erlenmeyer flask when it was submerged in the cold water?
1. When the flask was placed into the cold water, the colder air molecules in the flask move slower, putting out less pressure. With the decrease in air pressure inside the flask, the now greater pressure outside pushes water into the flask until the pressure inside equals the pressure outside.
A flask with a capacity of 1.00dm contains 5.00g of ethane The flask will burst if the pressure exceeds 1.00x10 Pa at what temperature will the pressure of the gas exceed the bursting temperature?
To find the temperature at which the pressure of the gas will exceed 1.00x10 Pa, you can use the Ideal Gas Law: PV = nRT. First, calculate the number of moles of ethane using its molar mass. Next, rearrange the formula to solve for temperature (T = PV / nR), where P is the bursting pressure, V is the volume, n is the number of moles, and R is the ideal gas constant. Plug in the values and solve for T.
A sample of HCl gas in placed in a 326-mL flask where it exerts a pressure of 67.5 mm Hg What is a pressure of this gas sample if it is transferred to a 135-mL flask at the same temperature?
Use Boyle's Law: p*V=constant at same temperatureSo let p2 be the wanted pressure in the 2nd flask (135 mL)then: p2 * 135 = 67.5 * 326 , so p2 = 67.5 * 326/135 = 163 mmHg
When the water level is higher inside than outside the flask is the gas pressure in the flask higher or lower or the same as the atmospheric pressure?
When the water level is higher inside the flask than outside, the gas pressure in the flask would be lower than the atmospheric pressure. This is because the water exerts a partial vacuum on the gas in the flask, reducing its pressure compared to the external atmospheric pressure.
When the water level is higher inside than outside the flask is the gas pressure in the flask higher than lower than or the same as the atmospheric pressure?
When the water level is higher inside than outside the flask, the gas pressure in the flask is lower than the atmospheric pressure. This is because the weight of the column of water inside the flask creates an additional pressure on the gas inside, reducing its pressure relative to the atmospheric pressure outside.
A sample of nitrogen gas is stored in a 500.0ml flask at 108kpa and 10.0c the gas is transferred to a 750.0ml flask at 21.0c what is the pressure of nitrogen in the second flask?
To find the pressure of the nitrogen gas in the second flask, you can use the combined gas law equation: P1V1/T1 = P2V2/T2, where P1, V1, and T1 are the initial pressure, volume, and temperature, and P2, V2, and T2 are the final pressure, volume, and temperature. Plug in the given values to find the final pressure of nitrogen in the second flask.
What is the total pressure in the flask?
The total pressure in a flask is the sum of the partial pressures of all the gases present in the flask. It can be calculated using the ideal gas law equation, PV = nRT, where P is the total pressure, V is the volume, n is the number of moles of gas, R is the gas constant, and T is the temperature in Kelvin.
What is the initial pressure of H2S gas in the flask with a Kp value of 0.120 at 25 deg C?
To determine the initial pressure of H2S gas in the flask, we need the total pressure and the partial pressure of another gas in equilibrium with H2S. Without the partial pressure of the other gas, we can't determine the initial pressure of H2S with just the Kp value and temperature provided.
Why should the flask be covered not stoppered?
Covering the flask instead of stoppering it allows for any gas produced during the reaction to escape, preventing pressure build-up that could lead to the flask exploding. Stoppering the flask could trap the gas and result in a dangerous situation.
What If the manometer reading is 305 mm and the atmospheric pressure is 1.03 ATM what is the pressure of the gas in the flask?
478 mm hg
Assume you seal 1.0 g of diethyl ether in an evacuated 100 mL flask If the flask is held at 30 degrees Celsius what is the approx gas pressure in the flask?
At 30 degrees C, the vapor pressure of ethe is about 590 mm Hg. (The pressure requires 0.23 g of ether in the vapor phase at the fiven conditions, so there is sufficient ether in the flask.) At 0 degrees C, the vapo pressure is about 160 mm Hg, so some ether condenses when the tempeature declines.
Is the volume occupied by the gas in the flask approxmately the same greater or less than before it was heated?
Assuming the flask is sealed - the volume remains the same but the pressure increases
Why did the water rush into the erlenmeyer flask when it was submerged in the cold water?
1. When the flask was placed into the cold water, the colder air molecules in the flask move slower, putting out less pressure. With the decrease in air pressure inside the flask, the now greater pressure outside pushes water into the flask until the pressure inside equals the pressure outside.
When all the liquid has evaporated the vapor totally fills the florence flask and exerts a pressure equal to the atmospheric pressure why?
I'm guessing you are analyzing an experiment where you are determining the molecular mass of an organic liquid. You heated the flask and the liquid evaporated filling the flask, but escaping through a small hole in the covering. 1. Gases always fill the container. So, if the liquid evaporated and formed a gas (vapor), it filled the flask, 2. The pressure on the outside the flask is air pressure. since the vapor isn't pushing off the cover, the pressure is not higher than the air pressure. But since the extra escaped, it cannot be less than the air pressure. Therefore, it is the same.
What is the pressure exerted by 1.55 g Xe gas at 20 C in a sealed 560 ml flask?
Using the ideal gas law, PV = nRT, where P is pressure, V is volume, n is number of moles, R is the gas constant, and T is temperature in Kelvin, we can calculate the pressure. First convert grams of Xe to moles using the molar mass of Xe. Then rearrange the ideal gas law to solve for pressure P. Plug in the values for volume, temperature, moles, and gas constant to find the pressure in the flask.
