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Molar Enthalpy Combustion
Name: Jaye
Status: student
Grade: 9-12
Country: Canada
Date: April 2009
Question:
How can you find the molar enthalpy of combustion of
magnesium directly? (without using Hess's Law)?
Replies:
Burn a mole of magnesium.
Richard Barrans, Ph.D., M.Ed.
Department of Physics and Astronomy
University of Wyoming
Jaye,
You have identified the beauty of Hess' Law. It is the safe way -- paper, pencil,
and calculator.
No high school lab, (and probably very few university labs) should undertake the
experimental combustion of magnesium. Your teacher may have demonstrated the
combustion of magnesium in air -- lots of bright light, heat and smoke. Not the
kind of reaction students should do without more training.
Several years ago a truck containing metallic magnesium travelled through Chicago
on one of the expressways. Somehow, the magnesium ignited. The truck driver stopped
the truck under a bridge. After the fire department extinguished the flames,
observers noted that (1) much of the bridge had melted, and (2) the concrete highway
under the bridge had disintegrated.
Warren Young
The combustion of magnesium metal is the reaction: Mg + 1/2 O2 = MgO. Hess's Law
states that the enthalpy of a reaction is the sum of the enthalpies of formation
of the products minus the sum of the enthalpies of formation of the reactants,
using a properly balanced chemical equation. The enthalpy of formation of ELEMENTS
in their respective standard states (i.e. the form of the element at 1 atm pressure
at the temperature of interest -- usually 298.15 kelvins) is assigned a value of
zero . In the case of the combustion of Mg (the reaction above), the enthalpy of
combustion the enthalpy of formation of MgO (-601.6 k-J/mol MgO) are identical.
It's not clear to me why you want to shy away from Hess's Law, but it is possible
to measure the enthalpy of combustion of Mg using a reaction calorimeter (also
referred to as a "bomb calorimeter", since the reaction in the calorimeter is run
at constant volume rather than constant pressure, which is contained in the
definition of the enthalpy of reaction. The conversion from constant volume to
constant pressure is straightforward. See the web site for details about Hess's
Law, and bomb calorimetry:
http://www.mikeblaber.org/oldwine/chm1045/notes/Energy/HessLaw/
Energy04.htm
http://www.chem.hope.edu/~polik/Chem345-2000/bombcalorimetry.htm
Vince Calder
Jaye,
In essence this should be straightforward, although, if you actually tried to do
it, it could prove problematic.
You would need to burn a known mass of Mg in O2, you would need to use pure O2 as
Mg reacts with N2 as well as O2 when burnt in air, forming MgO and Mg3N2. You would
need to set up this reaction so that the energy released heated up a known mass of a
suitable substance (water is generally used) and then you can use Q = m.Cp.dT to
calculate the heat released. Then divide the heat by the moles of Mg you burnt. But
here are the problems:
How will you make sure all the heat released goes into the water?
How will you make sure all the Mg reacts?
How will you make sure no heat escapes to the environment, or could you allow for
heat loss in some way?
Mg + O2 generally takes a lot of energy to start the reaction, how will you put in
this energy and not use it to heat the water?
To help you answer these problems you could try looking up a device called a "bomb
calorimeter". Good luck!!
Tom Collins
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Update: June 2012
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