The heat of formation of \(\mathrm{NH}_{3}(g)\) is \(-46 \mathrm{~kJ} \mathrm{~mol}^{-1}\). The \(\Delta H\) (in \(\mathrm{kJ} \mathrm{mol}^{-1}\) ) of the reaction, \(2 \mathrm{NH}_{3}(g) \longrightarrow \mathrm{N}_{2}(g)+3 \mathrm{H}_{2}(g)\) is : (a) 46 (b) \(-46\) (c) 92 (d) \(-92\)

Enthalpy change is a term used to describe the heat energy exchanged in a chemical process at constant pressure. It is denoted as \(\Delta H\), representing energy either released or absorbed by the system during a reaction. If \(\Delta H\) is negative, the process is exothermic, indicating that energy is released, often as heat, to the surroundings. Conversely, a positive \(\Delta H\) signals an endothermic process, where the system absorbs energy from the surroundings.

Understanding enthalpy changes is crucial when studying chemical reactions, as it helps predict how energy will flow. This knowledge, paired with the heat of formation values for substances involved, allows chemists to calculate the energy changes occurring in chemical processes.

A chemical reaction is the process where substances, known as reactants, transform into new substances, termed as products. These reactions involve the breaking of bonds in the reactants and the formation of new bonds in the products. This bond rearrangement is associated with energy changes; energy is required to break bonds and is released when new bonds form.

The stoichiometry of a reaction – the quantitative relationship between reactants and products in a chemical equation – is essential for determining the enthalpy change of the reaction. In our exercise, the stoichiometry of the decomposition reaction of \(\mathrm{NH}_{3}(g)\) into \(\mathrm{N}_{2}(g)\) and \(\mathrm{H}_{2}(g)\) helps calculate the overall \(\Delta H\) of the process.

Standard states refer to the physical states of substances at defined conditions of 1 atmosphere of pressure and a specified temperature, typically 298K (25 degrees Celsius). For pure substances, the standard state is the most stable form of the substance under standard conditions. For example, for hydrogen and nitrogen, it is the diatomic gases \(\mathrm{H}_{2}(g)\) and \(\mathrm{N}_{2}(g)\) respectively.

Heat of formation values, which are a form of enthalpy change, are always given for substances in their standard states. In the context of the exercise, it is important to note that the \(\Delta H\) provided for the formation of \(\mathrm{NH}_{3}(g)\) would apply when its constituent elements are in their standard states.

Reversible reactions are processes that can proceed in both forward and backward directions. In the context of enthalpy, when a reversible reaction proceeds in one direction, it has a specific \(\Delta H\), and when the reaction is reversed, the sign of \(\Delta H\) changes. This is because the energy absorbed in the forward reaction is released in the reverse reaction, and vice versa.

In our given exercise, the synthesis of \(\mathrm{NH}_{3}(g)\) from its elements is exothermic, with a \(\Delta H\) of -46 \(\mathrm{kJ mol^{-1}}\). The reverse reaction, which is the decomposition of \(\mathrm{NH}_{3}(g)\) into \(\mathrm{N}_{2}(g)\) and \(\mathrm{H}_{2}(g)\), has a \(\Delta H\) of +46 \(\mathrm{kJ mol^{-1}}\) per mole of \(\mathrm{NH}_{3}(g)\) decomposed. Thus, for two moles of \(\mathrm{NH}_{3}(g)\), the total enthalpy change would be double this value.