Reaction: Mg(s)+2HCl(aq)→MgCl2(aq)+H2(g)?
I was given the mass of Mg: 0.25g
The molar mass of Mg: 24.31g/mol
The concentration of HCl: 3.0 mol/L
The volume of HCl: 0.0512 L
The specific heat capacity of water: 4.19 J/g°C
Temperature change: +23.0°C
Calculate the enthalpy change, in kJ, released per mole of Mg reacted in reaction 2. Show your work.
My answer was -480 kJ released per mole of Mg reacted in reaction two.
Using Hess’ Law, calculate the theoretical enthalpy change for reaction 2.
My answer was -456.7 kJ for reaction two.
Are these answers correct and if they are, how can the experimental value account for more heat lost than the theoretical value.
Please help, I have no idea what I did wrong for these questions!
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Verified answer
(0.25 g Mg) / (24.31 g Mg/mol) = 0.01028 mol Mg
(3.0 mol HCl/L) x (0.0512 L) = 0.1536 mol HCl
The density of 3.0 mol/L HCl is about 1.0495 g/mL.
So 0.0512 L of such a solution would weigh about 53.73 grams.
0.01028 mole of Mg would react completely with 0.01028 x (2/1) = 0.02056 mole of HCl, but there is more HCl present than that, so HCl is in excess and Mg is the limiting reactant.
(0.01028 mol Mg) x (1 mol H2 / 1 mol Mg) x (2.01588 g H2/mol) = 0.0207 g H2 lost
(4.19 J/g°C) x (0.25 g + 53.73 g - 0.02 g) x (23.0°C) = 5200 J
(5200 J) / (0.01028 mol Mg) = 505836 J/mol = -506 kJ/mol Mg
[The minus sign did not come from the calculations. It is the convention for exothermic processes.]
[I don't have enough data to use Hess' Law.]
One reason the observed enthalpy change might not match a Hess' Law calculation is that the usual tables used with Hess' Law are just averages taken over many kinds of reactions, so they are approximations at best.
Another source of error is the assumption that the resulting solution has the same heat capacity as pure water. It probably doesn't exactly.
Another is that accurately measuring the temperature change of a calorimeter can be tricky business.