A gaseous mixture containing equal masses of He and CH4 is at a total pressure of 2 atm. Calculate partial pressure of He. What will be the pressure exerted by a mixture of 3.2 g o... A gaseous mixture containing equal masses of He and CH4 is at a total pressure of 2 atm. Calculate partial pressure of He. What will be the pressure exerted by a mixture of 3.2 g of methane and 4.4 g of carbon dioxide contained in a 9 dm³ flask at 27°C? What will be the pressure of the gaseous mixture when 0.5 L of H at 0.8 bar and 2.0 L of dioxigen at 0.7 bar are introduced in a 1L vessel at 27°C? Calculate the total pressure in a mixture of 8 g of dioxigen and 4 g of dihydrogen confined in a vessel of 1 dm³ at 27°C? Calculate the total pressure in a mixture of 3 g of dinitrogen, 3.0 g of dythydrogen and 8.0 g of dioxigen confined in vessel of 5 dm³ at 27°C? A mixture of dihydrogen and dioxigen are one bar pressure contains 20% by weight of dihydrogen. A neon-dioxigen mixture contains 70.6g dioxigen 167.5g neon. If prepared, what will be the total pressure of dioxigen and neon in the system? 1 mole of a gas 'AB' dissociates to an extent of 10% at 127°C according to AB ↔ A + B. Occupies a volume of 4 × 10⁴ ml. Find total pressure at this temp, assuming ideal gas behaviour. A closed bulb of capacity 200ml containing CH4, H2 and He at 300K. The ratio of partial pressures of CH4, H2 and He, respectively is 2 : 3 : 5. Calculate the ratio of their weights present in the container.

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Understand the Problem

The question primarily involves the principles of Dalton's Law of Partial Pressures and requires calculations related to gas mixtures and their individual pressures. Each sub-question deals with different aspects of gases, such as mixtures, total pressure, reactions, and behaviors under specific conditions.

Answer

The partial pressure of He is $1 \, \text{atm}$.
Answer for screen readers

The partial pressure of He is $1 , \text{atm}$.

Steps to Solve

  1. Understanding Dalton's Law of Partial Pressures

Utilize Dalton's Law, which states that the total pressure of a gas mixture is equal to the sum of the partial pressures of each individual gas. The formula is given as: $$ P_{\text{total}} = P_1 + P_2 + P_3 + ... + P_n $$

  1. Calculating the Partial Pressure of He and CH₄

Given that the total pressure (P_{\text{total}} = 2 , \text{atm}) and both gases have equal molar amounts, we denote:

  • Molar mass of He = 4 g/mol
  • Molar mass of CH₄ = 16 g/mol

To calculate the partial pressure of each gas: $$ P_{\text{He}} = \frac{n_{\text{He}}}{n_{\text{total}}} \cdot P_{\text{total}} $$ $$ P_{\text{CH}4} = \frac{n{\text{CH}4}}{n{\text{total}}} \cdot P_{\text{total}} $$

Since both gases are in equal amounts, $$ n_{\text{total}} = n_{\text{He}} + n_{\text{CH}_4} $$

  1. Finding the Moles of Each Gas

Let (n) be the number of moles for He and CH₄. Therefore, $$ n_{\text{total}} = n + n = 2n $$ Now, calculate the mole fraction: $$ P_{\text{He}} = \frac{n}{2n} \cdot P_{\text{total}} = \frac{1}{2} P_{\text{total}} $$

  1. Calculating Partial Pressure

Using the established total pressure, $$ P_{\text{He}} = \frac{1}{2} \times 2 , \text{atm} = 1 , \text{atm} $$ Thus for CH₄: $$ P_{\text{CH}_4} = 1 , \text{atm} $$

  1. Final Calculation for He's Partial Pressure

Since the question specifically requires you to find the partial pressure for He: $$ P_{\text{He}} \text{ (final answer) } = 1 , \text{atm} $$

The partial pressure of He is $1 , \text{atm}$.

More Information

The problem illustrates Dalton’s Law in gas mixtures and its application. This law is essential in understanding behaviors of gases in different conditions.

Tips

  • Forgetting to account for the proportions of gases when they are equal in mass but different in molecular weight.
  • Miscalculating total pressure or mole fractions. Always ensure to check calculations step-by-step.

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