How do I determine the standard enthalpy of combustion for magnesium in hydrochloric acid?

Reaction 1
mass= 0.13g
Starting temp= 23.0 C
Highest temp= 34.8 C
50mL of HCl
Reaction 2
Mass= 0.97
Starting temp= 23.0 C
Highest temp= 38.0 C
50mL of HCl

1 Answer
May 10, 2018

There isn't a way, as far as I can see, of finding the standard enthalpy of combustion from here. I have found the standard enthalpy of reaction below, although the numbers only seem to fit for the first experiment. Maybe I've misunderstood though;

regardless, the enthalpy of reaction is #-461" kJ mol"^-1#

Explanation:

Consider reaction 1

We will make a fairly reasonable assumption here, that the HCl is basically water. This will allow us to use water's density, and specific heat capacity.

Since #1ml=1cm^3#

#"volume of HCl"=50cm^3#

The density of water is #1" g cm"^-3#

#:."mass HCl"=1xx50#
#=50g#

Now we use #Q=mcDeltaT#
#Q=50xx-4.18xx(34.8-23)#
#=-2466.2J# (write all your sig figs at this point).
#=-2.4662kJ#

To find the enthalpy of reaction, we need to divide by the moles of magnesium (the unit is kJ per mole).

#"mol Mg"=0.13/24.3#
(I'm taking the #A_r# of magnesium to be 24.3)

#"mol Mg"=5.350xx10^-3" mol"#
I've only written 4sf, so keep this number in your calculator; use the ANS button.

So the enthalpy of reaction:

#:. Delta_rH=-2.4662/(5.350xx10^-3)#
#=-461" kJ mol"^-1 " " (3"sf")#

Here we round our answer.

Consider reaction 2
You should do the same thing again for this question. I have included my calculation, but the answer is off by a long shot. I don't know if this is my calculation or an error with the question - either way, the first reaction enthalpy we calculated seems more accurate.

#Q=mcDeltaT#
#=50xx-4.18xx(38-23)#
#=-3135J#
#=-3.135kJ#

#"mol Mg"=0.97/24.3#
#=0.0399" mol"#

#Delta_rH=-3.135/0.0399#
#=-78.57" kJ mol"^-1#
(see what I mean?)