# Where does ΔH fit into this? is it interchangeable with Q?

## Main Question or Discussion Point

hi,

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?

thanks.