The combustion of propane is given by the following reaction:

C3H8 (g) + 5 O2 (g) ---> 3 CO2 (g) + 4 H2O (g)
If 5.00 L of propane are burned in in the presence of excess O2, how many liters of carbon dioxide will be formed?

5.00L propane x (3 mols CO2/1 mol C3H8) = 5.00 x 3/1 = ?

To determine the number of liters of carbon dioxide formed, we need to use the coefficients from the balanced chemical equation.

The balanced equation shows that for every 1 mole of C3H8 burned, we produce 3 moles of CO2.

First, we need to convert the volume of propane (C3H8) to moles using the ideal gas law equation:

PV = nRT

Since we want to convert from volume to moles, we rearrange the equation to:

n = PV / RT

Where:
P is the pressure (in atm),
V is the volume (in L),
n is the number of moles,
R is the ideal gas constant (0.0821 L*atm/(mol*K)),
T is the temperature (in Kelvin).

Next, we calculate the number of moles of propane:
n(C3H8) = (5.00 L) / (0.0821 L*atm/(mol*K) * T)

Assuming the temperature is constant, we can substitute the values into the equation. Let's assume a temperature of 298 K:
n(C3H8) = (5.00 L) / (0.0821 L*atm/(mol*298 K))

Calculate the value of n(C3H8) using the above equation.

Once we have the number of moles of propane, we can use stoichiometry to find the number of moles of carbon dioxide produced. Since the molar ratio between C3H8 and CO2 is 1:3, we can determine the number of moles of carbon dioxide:

n(CO2) = n(C3H8) * (3 moles CO2 / 1 mole C3H8)

Finally, we can convert the number of moles of carbon dioxide to volume using the ideal gas law equation again. Assuming constant temperature and pressure, we can rearrange the equation to find the volume:

V = (n * R * T) / P

Where:
V is the volume (in L),
n is the number of moles,
R is the ideal gas constant (0.0821 L*atm/(mol*K)),
T is the temperature (in Kelvin),
P is the pressure (in atm).

Calculate the volume of carbon dioxide in liters using the above equation.

This will give you the volume of carbon dioxide produced when 5.00 L of propane is burned in the presence of excess O2.