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rewona [7]
3 years ago
7

One mole of an ideal gas is sealed in a 22.4-L container at a pressure of 1 atm and a temperature of 273 K. The temperature is t

hen increased to 304 K , but the container does not expand. What will the new pressure be?
Part A

The most appropriate formula for solving this problem includes only which variables?

Enter the required variables, separated by commas (e.g., P,V,T).

Q2)A sample of nitrogen gas in a 1.69-L container exerts a pressure of 1.37 atm at 17 ∘C.

-What is the pressure if the volume of the container is maintained constant and the temperature is raised to 327 ∘C?

Q3)A gas mixture with a total pressure of 770 mmHgcontains each of the following gases at the indicated partial pressures: 120 mmHg CO2, 227mmHg Ar, and 190 mmHg O2. The mixture also contains helium gas

-What mass of helium gas is present in a 14.0-L sample of this mixture at 282 K ?

Q4)

A

Calculate the density of oxygen, O2, under each of the following conditions:

STP

1.00 atm and 35.0 ∘C

Express your answers numerically in grams per liter. Enter the density at STP first and separate your answers by a comma.

B

To identify a diatomic gas (X2), a researcher carried out the following experiment: She weighed an empty 4.1-L bulb, then filled it with the gas at 2.00 atm and 24.0 ∘C and weighed it again. The difference in mass was 9.5 g . Identify the gas.

Express your answer as a chemical formula.
Chemistry
1 answer:
Ket [755]3 years ago
5 0

Answer:

1.11 atm

(P, T)

2.83 atm

0.740 g

0.179 g/L, 0.158 g/L

N₂

Explanation:

<em>One mole of an ideal gas is sealed in a 22.4-L container at a pressure of 1 atm and a temperature of 273 K. The temperature is then increased to 304 K, but the container does not expand. What will the new pressure be?</em>

Assuming ideal behavior, we can calculate the new pressure (P₂) using Gay-Lussac's law.

\frac{P_{1}}{T_{1}} =\frac{P_{2}}{T_{2}} \\P_{2}=\frac{P_{1}}{T_{1}}.T_{2}=\frac{1atm}{273K} .304K=1.11atm

<em>The most appropriate formula for solving this problem includes only which variables?</em>

Gay-Lussac's law includes pressure (P) and absolute temperature (T).

<em>Q2) A sample of nitrogen gas in a 1.69-L container exerts a pressure of 1.37 atm at 17 °C.  What is the pressure if the volume of the container is maintained constant and the temperature is raised to 327 °C? </em>

Initially the system is at 17°C (290 K) and the temperature is raised to 327°C (600 K). We can calculate the new pressure using Gay-Lussac's law.

\frac{P_{1}}{T_{1}} =\frac{P_{2}}{T_{2}} \\P_{2}=\frac{P_{1}}{T_{1}}.T_{2}=\frac{1.37atm}{290K} .600K=2.83atm

<em>Q3) A gas mixture with a total pressure of 770 mmHg contains each of the following gases at the indicated partial pressures: 120 mmHg CO₂, 227mmHg Ar, and 190 mmHg O₂. The mixture also contains helium gas .</em>

What mass of helium gas is present in a 14.0-L sample of this mixture at 282 K?

First, we have to calculate the pressure of Helium. We know that the total pressure is the sum of partial pressures.

Ptotal = pCO₂ + pAr + pO₂ + pHe

pHe = Ptotal - pCO₂ - pAr - pO₂

pHe = 770mmHg - 120mmHg - 227mmHg - 190mmHg=233mmHg

We can calculate the moles of Helium using the ideal gas equation.

P.V=n.R.T\\n=\frac{P.V}{R.T} =\frac{233mmHg.14.0L}{(0.08206atm.L/mol.K).282K} .\frac{1atm}{760mmHg} =0.185mol

The molar mass of He is 4.00g/mol.

0.185mol.\frac{4.00g}{mol} =0.740g

<em>Calculate the density of oxygen, O₂, under each of the following conditions: </em>

  • <em> STP </em>
  • <em>1.00 atm and 35.0 ∘C </em>

<em> Express your answers numerically in grams per liter. Enter the density at STP first and separate your answers by a comma.</em>

<em />

STP stands for Standard Temperature and Pressure. The standard temperature is 273 K and the standard pressure is 1 atm.

We can calculate the density using the following expression:

\rho=\frac{P.M}{R.T} =\frac{1.00atm.4.00g/mol}{(0.08206atm.L/mol.K).273K} =0.179 g/L

<em>At 1.00 atm and 35.0 °C (308 K)</em>

\rho=\frac{P.M}{R.T} =\frac{1.00atm.4.00g/mol}{(0.08206atm.L/mol.K).308K} =0.158 g/L

<em>To identify a diatomic gas (X₂), a researcher carried out the following experiment: She weighed an empty 4.1-L bulb, then filled it with the gas at 2.00 atm and 24.0 ∘C and weighed it again. The difference in mass was 9.5 g . Identify the gas.  Express your answer as a chemical formula.</em>

We will look for the molar mass of the compound using the ideal gas equation.

P.V=n.R.T=\frac{m}{M} .R.T\\M=\frac{m.R.T}{P.V} =\frac{9.5g \times (0.08206atm.L/mol.K)\times 297K }{2.00atm \times 4.1L} =28g/mol

If the molar mass of X₂ is 28 g/mol, the molar mass of X is 14 g/mol. Then, X is nitrogen and X₂ is N₂.

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