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.
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
- 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: $$ P_{\text{total}} = P_1 + P_2 + P_3 + ... + P_n $$
- 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, $$ n_{\text{total}} = n_{\text{He}} + n_{\text{CH}_4} $$
- 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: $$ P_{\text{He}} = \frac{n}{2n} \cdot P_{\text{total}} = \frac{1}{2} P_{\text{total}} $$
- Calculating Partial Pressure
Using the established total pressure, $$ P_{\text{He}} = \frac{1}{2} \times 2 , \text{atm} = 1 , \text{atm} $$ Thus for CH₄: $$ P_{\text{CH}_4} = 1 , \text{atm} $$
- 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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