Dec3-09, 07:23 PM
1. The problem statement, all variables and given/known data
Calculate the heat absorbed by the water in the calorimeter and the heat released by 1,0g of burning ethanol. Then, calculate the molar heat of combustion (kJ/mol) of the ethanol.
The calorimeter used was very makeshift: ethanol in an alcohol burner was lit under a can containing water, both of which were placed inside a hollowed out larger can.
Here is what is known:
mass of ethanol burned = 20g
mass of heated water = 280g
change in temperature = 30"degrees"C
2. Relevant equations
Q(water) = -Q(ethanol)
mc"delta"T = - (mc"delta"T)
280g*4.184J/g"degree"C*30 = -Q(ethanol)
3. The attempt at a solution
Q(water) = mc"delta"T = 280g*4.184J/g"degree"C*30"degrees"C = 35145.6J = 35.1kJ
Q(ethanol) = -Q(water) = -35.1kJ (for 20g)
Q(per gram of ethanol) = -35.1kJ / 20g = -1.755kJ/g = -1.76kJ/g
Q or "delta"H? (per mol of ethanol) = -1.76kJ/g * 46.07g/mol = 81.08kJ/mol
--> Does this make sense? Is it right? And what is the difference between Q and "delta"H?
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