C6H5OH(s) + 7O2(g) --> 6CO2(g) + 3H2O(g)
A 2.000 gram sample of pure phenol, C6H5OH(s), is completely burned according to the equation above. If the volume of the combustion chamber is 10.0L, calculate the final pessure in the container when the temperature is 110 degrees C. (Assume no oxygen remains unreacted and that all products are gaseous)
Convertr 2.000 g phenol to mols. mols = grams/molar mass.
Convert mols phenol to mols CO2 and mols H2O using the coefficients in the balanced equation.
Use PV = nRT to calculate P (in atmospheres). n is the total mols calculated above, V is known, R is known (0.08205 L*atm/mol*K) and T is known. Don't forget to convert 110 degrees C to Kelvin. Post your work if you get stuck.
1.2
To calculate the final pressure in the container, we can use the ideal gas law equation: PV = nRT, where P is the pressure, V is the volume, n is the number of moles, R is the ideal gas constant, and T is the temperature in Kelvin.
Step 1: Convert the temperature from Celsius to Kelvin.
Given temperature = 110 degrees C
Convert to Kelvin: T = 110 + 273.15 = 383.15 K
Step 2: Calculate the number of moles of CO2 produced.
From the balanced equation, we can see that 1 mole of C6H5OH produces 6 moles of CO2.
Molar mass of C6H5OH = 94.11 g/mol
Molar mass of CO2 = 44.01 g/mol
Number of moles of CO2 = (2.000 g C6H5OH) * (6 mol CO2 / 94.11 g C6H5OH) = 0.1272 mol CO2
Step 3: Calculate the number of moles of H2O produced.
From the balanced equation, we can see that 1 mole of C6H5OH produces 3 moles of H2O.
Number of moles of H2O = (2.000 g C6H5OH) * (3 mol H2O / 94.11 g C6H5OH) = 0.06072 mol H2O
Step 4: Calculate the total number of moles of gas produced.
Total number of moles of gas = moles of CO2 + moles of H2O
Total number of moles of gas = 0.1272 mol CO2 + 0.06072 mol H2O = 0.1879 mol
Step 5: Substitute the values into the ideal gas law equation.
PV = nRT
P * 10.0 L = (0.1879 mol) * (0.0821 L/mol·K) * (383.15 K)
P * 10.0 L = 5.8567 L·atm
P = 5.8567 L·atm / 10.0 L
P = 0.586 atm
Step 6: Convert the pressure to the appropriate units.
0.586 atm is the final pressure in the container when the temperature is 110 degrees C.
To calculate the final pressure in the container, we can use the ideal gas law equation:
PV = nRT
where P is the pressure, V is the volume, n is the number of moles, R is the ideal gas constant, and T is the temperature in Kelvin.
First, we need to calculate the number of moles of the gaseous products produced in the reaction.
We start with 2.000 grams of phenol (C6H5OH) and convert it to moles using the molar mass of phenol.
Molar mass of phenol (C6H5OH):
Carbon (C): 12.01 g/mol
Hydrogen (H): 1.008 g/mol
Oxygen (O): 16.00 g/mol
Molar mass of phenol:
C6H5OH = 6(12.01 g/mol) + 6(1.008 g/mol) + 16.00 g/mol = 94.11 g/mol
Number of moles of phenol:
n = mass / molar mass
n = 2.000 g / 94.11 g/mol = 0.02125 mol
According to the balanced chemical equation, 6 moles of CO2 and 3 moles of H2O are produced per mole of phenol combustion.
Number of moles of CO2 produced:
n_CO2 = 6 * 0.02125 mol = 0.1275 mol
Number of moles of H2O produced:
n_H2O = 3 * 0.02125 mol = 0.06375 mol
Now, we need to convert the temperature from Celsius to Kelvin:
T(K) = T(C) + 273.15
T(K) = 110 °C + 273.15 = 383.15 K
Finally, we can substitute the known values into the ideal gas law equation to calculate the final pressure.
P * 10.0 L = (n_CO2 + n_H2O) * 0.0821 L·atm/mol·K * 383.15 K
Rearranging the equation to solve for P:
P = [(n_CO2 + n_H2O) * 0.0821 L·atm/mol·K * 383.15 K] / 10.0 L
P = [(0.1275 mol + 0.06375 mol) * 0.0821 L·atm/mol·K * 383.15 K] / 10.0 L
P = (0.19125 mol * 0.0821 L·atm/mol·K * 383.15 K) / 10.0 L
P ≈ 6.2919 atm
Therefore, the final pressure in the container would be approximately 6.3 atm when the temperature is 110 degrees C.