A mixture of 0.154 moles of C is reacted with 0.117 moles of O2 in a sealed, 10.0 L vessel at 500 K, producing a mixure of CO and CO2. The total pressure is 0.640 atm. What is the partial pressure of CO?
To find the partial pressure of CO, we need to use the ideal gas law equation, which states that PV = nRT.
First, we need to determine the number of moles of CO and CO2 produced in the reaction. From the balanced chemical equation, we can see that one mole of C reacts with one mole of O2 to form one mole of CO and one mole of CO2.
Given that we have 0.154 moles of C and 0.117 moles of O2, we can conclude that both reactants will be completely consumed, resulting in the formation of 0.117 moles of CO and 0.117 moles of CO2.
Next, we can use the ideal gas law to calculate the partial pressure of CO. Rearranging the equation to solve for pressure (P), we have:
P = (n * R * T) / V
Where:
P = pressure of the gas
n = number of moles of the gas
R = ideal gas constant (0.0821 L·atm/mol·K)
T = temperature in Kelvin
V = volume of the gas
Substituting the values into the equation:
n = 0.117 moles
R = 0.0821 L·atm/mol·K
T = 500 K
V = 10.0 L
P = (0.117 * 0.0821 * 500) / 10.0
P = 0.300 atm
Therefore, the partial pressure of CO is 0.300 atm.