[sammiegirl]

I actually have two questions....

1) A chemical engineer studying the properties of fuels placed 1.146g of a hydrocarbon in the bomb of a calorimeter and filled it with O2 gas.  The bomb was immersed in 2.982L water and the reacation initiated.  The water temperature rose from 21.67 degrees Celsius to 27.90 degrees Celsius.  If the calorimeter (excluding the water) had a heat capacity of 405 J/K, what was the heat of combustion of the fuel (in J/g)?


For this problem, I understand that the heat released must equal the heat absorbed.  So the q combustion will equal the opposite of the q calorimeter and q water.  Using the equation of q=m x c x delta t i solve the heat of combustion to be -80253J.  I then divide this number by 1.146 grams since that is the amount I have.  Is this right? Or where am I making my mistake?

2) Pure liquid octane (C8H18, d=.702 g/mL) is used as the fuel in a test of a new automobile drive train.  What is the magnitude of the energy released by the complete combustion of all the octane in a 20.9-gal tank (delta H = 5.45E3 kJ/mol)?

For this question, I convert gallons into grames using the density and then from grams into moles where I get 48.456 moles of octane.  Then I multiply the number of moles by the delta H to find the magnitude of the energy.  I feel like I am doing this right but I am obviously not.

[sdekivit]

1) use the fact that the amount of heat released by the fuel due to combustion is taken up by the calorimeter and the water:

Q_combustion = Q_calorimeter + Q_water

thus:

Q = C * delta T + c * m * delta T

then calculate the amount of joules released per g fuel.

2) Calculate mass of octane with the density: (d = m / V, with V in mL!!!). Then use the fact that per mol octane 5,45 x 10³ J is released.