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

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In summary, the conversation discusses the use of the thermal energy formula Q = mcΔT in basic chemistry coursework and how it relates to enthalpy. The user is confused about where ΔH fits into the equation and if it is interchangeable with Q. The conversation then provides a solved question as context for finding the molar enthalpy of combustion of anthracene. The final value is calculated to be 7020 kJ/mol, and the question is raised if this is the value for ΔH and if it is simply a matter of units or if ΔH and Q are interchangeable.
  • #1
quicksilver123
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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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  • #2
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.
 

What is ΔH and how does it relate to Q?

ΔH is the symbol for enthalpy, which is a measure of the total energy of a thermodynamic system. Q, on the other hand, represents heat, which is a form of energy transfer. ΔH and Q are related through the equation Q = ΔH + ΔnRT, where Δn is the change in the number of moles of gas and R is the gas constant.

Can ΔH be used interchangeably with Q?

No, ΔH and Q cannot be used interchangeably. While they are related, they represent different quantities. ΔH is a measure of the total energy of a system, while Q represents the transfer of thermal energy.

In what situations would ΔH be more useful than Q?

ΔH is more useful than Q when studying chemical reactions or phase transitions, as it takes into account the total energy of the system, including any changes in the number of moles of gas. Q is more useful when studying heat transfer in a system.

How is ΔH measured in a reaction?

ΔH is typically measured through calorimetry, which involves measuring the change in temperature of a reaction vessel and using that information to calculate the heat released or absorbed by the reaction. It can also be calculated using standard enthalpies of formation.

Is ΔH affected by changes in pressure and volume?

Yes, ΔH is affected by changes in pressure and volume. This is because it takes into account the total energy of a system, including any changes in the number of moles of gas, which can be affected by changes in pressure and volume.

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