Assuming the ideal gas behavior of CH4 and Cl2, the partial pressure "p" can be calculated using the ideal gas law:
"p=\\frac{nRT}{V}" ,
where "n" is the number of the moles of a gas, "R" is the ideal gas constant, 8.314 m3·Pa/(mol·K) , "T" is the temperature in kelvin, 35+273.15 = 308.15 K and "V" is the volume of the system, 5 L or 5·10-3 m3.
The number of the moles of a gas "n" is:
"n = \\frac{m}{M}",
where "m" is the mass and "M" is the molar mass of a gas. The molar masses of CH4 and Cl2 are 16.04 g/mol and 70.91 g/mol, respectively. Therefore:
"n(CH_4) = \\frac{100 \\text{ g}}{16.04 \\text{ g\/mol}} = 6.23" mol,
"n(Cl_2) = \\frac{100\\text{ g}}{70.91\\text{ g\/mol}} = 1.41" mol.
The partial pressures of CH4 and Cl2 are:
"p_{CH_4} = \\frac{6.23\u00b78.314\u00b7308.15}{5\u00b710^{-3}} = 3.19\u00b710^6" Pa
"p_{Cl_2} = \\frac{1.41\u00b78.314\u00b7308.15}{5\u00b710^{-3}} = 7.22\u00b710^5" Pa.
According to the Dalton's law of partial pressures, the total pressure is the sum of the partial pressures:
"p_{tot} = p_{CH_4} + p_{Cl_2} = (3.19 + 0.722)\u00b710^6 = 3.91\u00b710^6" Pa.
Answer: the partial pressures of CH4 and Cl2 are 3.19·106 Pa and 7.22·105 Pa, respectively. The total pressure of the system is 3.91·106 Pa.
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