What is the molar enthalpy of fusion of magnesium?

I don't know who did this experiment, but someone died that day. Adding molten magnesium to water would certainly make someone run for their lives as the #"H"_2# gas gets produced.

Well, let's pretend we have a FANTASTIC blast shield, and that we're standing 10 feet away. THEN, let's pretend that we remotely (ala MythBusters) drop the #"52.0 g"# of #"Mg"(l)# into the #"185 mL"# of water.

The heat involved assumes the density of water is #"0.998 g/mL"# so that #"185 mL water"# #=# #"184.6 g" -= m_w# and:

#q_w = m_wc_wDeltaT#

#= "185 g" cdot "4.184 J/g"^@ "C" cdot 22.4^@ "C"#

#=# #"17303.8 J"#

This heat apparently came out of the liquid magnesium, so the rapid fusion process sucked #"17.3 kJ"# out of #"52.0 g"# of #"Mg"(l)#. We expect conservation of energy...

Therefore, #DeltaH_(fus) -= |q_(Mg)| stackrel(?" ")(=) |q_w|#, and so:

#color(blue)(DeltaH_(fus)) = "17.3 kJ"/"52.0 g" = "0.333 kJ/g"#

#= "0.333 kJ"/cancel"g Mg" xx (24.305 cancel"g Mg")/"mol"#

#=# #color(blue)("8.09 kJ/mol")#

The actual value is #"4.79 kJ/mol"#, so it looks like the remaining #"3.30 kJ/mol"# became the explosion.