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

  1. Jun 26, 2013 #1

    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?

  2. jcsd
  3. Jun 27, 2013 #2


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    Staff: Mentor

    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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