Natural gas is almost entirely methane, CH4. What volume of natural gas at 20 C and 1.00 atm pressure is required to heat one quart of water from 20 C to 100 C? The density of methane at 20 C is 0.665 g/mL. The reaction for the combustion of methane is
CH4 (g) + 2 O2 (g) --> CO2 + 2 H2O (g)
3 answers
You have everything here except the heat of combustion of CH4.
I have no idea how to solve for the volume of methane. Any help?
Yes BUT you need the heat of combustion of CH4 to know howmuch heat it produces.
How much heat do you need? (1 qt H2O = 0.946 L and if the density of H2O is 1.00 g/mL that is 946 g)That's
q = 946 g H2O x specific heat H2O x (Tfinal-Tinitial). Let's suppose this is y joules.
g CH4 needed to produce y joules is
16 g CH4 x (y joules/heat combustion) = g CH4
Convert g CH4 to cc methane using the density. That is the volume.
You don't have heat combustion CH4, density of H2O (which I assumed as 1.00 g/mL), or specific heat H2O (which is 4.184 J/g or 1 cal/g)
How much heat do you need? (1 qt H2O = 0.946 L and if the density of H2O is 1.00 g/mL that is 946 g)That's
q = 946 g H2O x specific heat H2O x (Tfinal-Tinitial). Let's suppose this is y joules.
g CH4 needed to produce y joules is
16 g CH4 x (y joules/heat combustion) = g CH4
Convert g CH4 to cc methane using the density. That is the volume.
You don't have heat combustion CH4, density of H2O (which I assumed as 1.00 g/mL), or specific heat H2O (which is 4.184 J/g or 1 cal/g)
2 answers