Standard Enthalpy of Formation of H2(g) vs H(g) My question is somewhat related to this quesion. When I look at a table of standard values there are two entries for Hydrogen:


$\ce {H2(g)} $  - 0 kJ/mol
$\ce {H(g)} $  -  218 kJ/mol


I understand the $\ce{H2(g)} $ is 0 because it is used as a reference, why is monatomic $\ce{ H(g)}$ given in the table and what does it mean? When would one use it in a calculation of enthalpy change?

From what I understand Hydrogen only exists as a monatomic gas at very high temperatures, the Standard enthalpies of formation are given at 1 bar, 298.15 K.
The value may be used whenever atomic hydrogen is present as a (probably intermediate) species. E.g. take the reaction

$$\ce{2H &lt;=&gt; H2}$$

which is the basic reaction in atomic hydrogen welding.