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Where does ΔH fit into this? is it interchangeable with Q? 
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#1
Jun2613, 10:41 PM

P: 91

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. 


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