Jun26-13, 11:41 PM
this might be a stupid question. if so, bear with me.
i've been using the following equation in my coursework (basic chemistry):
Q = m c ΔT
where, of course
m = mass
c = specific heat capacity
ΔT = change in temperature (°C)
Q, as i understand it, is simply the quantity of thermal energy in joules or kJ.
here's where i'm a little confused:
where does ΔH fit into this? is it interchangeable with Q?
here's some context (random solved question):
a sample of anthracene (c14h10) undergoes complete combustion in a calorimeter (which is made of aluminimum [c=0.900J/g°C] and has a mass of 948g).
the calorimeter contains 1.50 L of water (1500g) which had an initial temp of 20.5°C and ends up with a final temp of 34.3°C.
[ΔT = 13.8°C]
find the molar enthalpy (molar heat of combustion of anthracene)
using the thermal energy formula above (i don't know its proper name), i got these results:
water: Q = 86.5 kJ
calorimeter: Q = 11.8 kJ
the two of them together: Q = 98.3 kJ
molar mass of anthracene = 178g/mol
the given mass of the sample of anthracene is equal to 0.014 moles
so the heat released per mole of anthracene dissolved in water = 7020 kJ/mol
would this final value be the ΔH value? if so, why? is it simply the units?
or are delta H and Q interchangeable?
Jun27-13, 03:25 AM
Check the definition of enthalpy. Under constant pressure, and assuming there is only expansion work done (that is, the only work done on/by the system is the one related to the volume change) ΔH=ΔQ.
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