From the following data, calculate the average bond enthalpy for the \(\mathrm{N}-\mathrm{H}\) bond: \(\begin{aligned} \mathrm{NH}_{3}(g) & \longrightarrow \mathrm{NH}_{2}(g)+\mathrm{H}(g) & \Delta H^{\circ} &=435 \mathrm{~kJ} / \mathrm{mol} \\ \mathrm{NH}_{2}(g) & \longrightarrow \mathrm{NH}(g)+\mathrm{H}(g) & \Delta H^{\circ} &=381 \mathrm{~kJ} / \mathrm{mol} \\ \mathrm{NH}(g) & \longrightarrow \mathrm{N}(g)+\mathrm{H}(g) & \Delta H^{\circ} &=360 \mathrm{~kJ} / \mathrm{mol} \end{aligned}\)

Enthalpy change is a measure of heat energy released or absorbed when a chemical reaction occurs at constant pressure. It's an essential concept in thermodynamics and chemistry, providing insight into the energy changes that accompany chemical processes.

Imagine a reaction where a substance transforms into another; energy may be required to break certain bonds in the reactants, or energy may be released when new bonds form in the products. If the reaction absorbs energy from the surroundings, it's endothermic, and the enthalpy change (Delta H) is positive. Conversely, if it releases energy, it's exothermic, with a negative enthalpy change.

To calculate the enthalpy change for breaking bonds—like the N-H bond in our exercise—you must consider the energy absorbed to break the bonds. For the given reactions, the increase in enthalpy reflects the energy input needed to cleave the N-H bonds and form new species. The enthalpy change values (435 kJ/mol, 381 kJ/mol, and 360 kJ/mol) are indicative of the bond strength and the energy required to break each bond in the sequence of reactions.

Chemical bonds are the forces holding atoms together in molecules and compounds. They are formed by the interactions between the electrons of the bonding atoms, resulting in a stable electronic configuration. Bonds can be covalent, where atoms share electron pairs, ionic, arising from electrostatic attractions, or metallic, featuring a 'sea' of delocalized electrons.

In our NH3 breaking down exercise, we're looking at covalent N-H bonds. The strength of a chemical bond is linked to how much energy is needed to break it—higher energy means a stronger bond. The average bond enthalpy is an average measure of the bond strength based on various molecules' data, providing a useful estimate for reactions involving similar bonds.

Understanding the intricacies of bond enthalpy can also illuminate reaction mechanics. For instance, in the three reactions provided, we're breaking the N-H bonds step by step. Each of these represents a distinct chemical change, involving the making and breaking of bonds, and thus, separate enthalpy changes.

Enthalpy calculation is a fundamental aspect of thermodynamics used in chemistry to predict the energy changes occurring during a reaction. To calculate the average bond enthalpy, one should consider the individual steps of a reaction and the total energy change associated with breaking the bonds.

For our exercise involving NH3, NH2, and NH, the calculation begins by summing the enthalpy changes as we break each N-H bond one at a time. With the total energy required to break all three bonds, we can find the average by dividing this total by the number of bonds broken, providing a simplified assessment of the bond's energetic characteristics.

Practical Application

Let's apply this to the exercise: the sum of the enthalpy changes is 1176 kJ/mol for three N-H bonds. Dividing this by the number of bonds (3), we find an average bond enthalpy of 392 kJ/mol. This figure represents the theoretical amount of energy typically needed to break a single N-H bond, giving us insight into the bond's strength and stability.