Routing every reaction by hand would be exhausting, so chemistry tabulated the routes once. A standard enthalpy of formation, written $\Delta H^{\circ}_{f}$, is the enthalpy change when one mole of a compound forms from its elements in their standard states, at 25 °C and 100 kPa. The degree symbol is said standard, the subscript f is formation, and the unit is kJ/mol. It follows immediately that an element in its standard state has $\Delta H^{\circ}_{f} = 0$ — forming oxygen from oxygen costs nothing. That zero is not an approximation; it is the floor the whole table is measured from, and tables usually do not bother printing it.

With one floor for everything, the routing collapses to a subtraction: $\Delta H^{\circ}_{\text{rxn}} = \sum \Delta H^{\circ}_{f}(\text{prod}) - \sum \Delta H^{\circ}_{f}(\text{react})$, where each sum runs over that side of the equation with every value multiplied by its coefficient from the balanced equation. Products minus reactants, in that order, because that is the direction the arrow points — reverse it and you have computed the reverse reaction, a perfectly good number attached to the wrong question. Two habits keep it honest: balance the equation before you touch the table, and multiply before you add.