Without doing any calculations, predict the sign of \(\Delta H\) for each of the following reactions: $$\begin{array}{l}{\text { (a) } 2 \mathrm{NO}_{2}(g) \longrightarrow \mathrm{N}_{2} \mathrm{O}_{4}(g)} \\ {\text { (b) } 2 \mathrm{F}(g) \longrightarrow \mathrm{F}_{2}(g)} \\ {\text { (c) } \mathrm{Mg}^{2+}(g)+2 \mathrm{Cl}^{-}(g) \longrightarrow \mathrm{MgCl}_{2}(s)} \\ {\text { (d) } \mathrm{HBr}(g) \longrightarrow \mathrm{H}(g)+\mathrm{Br}(g)}\end{array}$$

Enthalpy change, denoted as \( \Delta H \), is a measure of heat energy released or absorbed during a chemical reaction at constant pressure. Understanding enthalpy change is crucial because it helps us determine whether a reaction is exothermic or endothermic, which in turn affects the surrounding environment. For instance, when you feel warmth from a reaction, such as when burning wood, it's because the reaction is exothermic and releasing heat. Conversely, when you add salt to ice, and it feels cold, it's endothermic; heat is absorbed from the surroundings.

Generally, if the total energy required to break bonds in the reactants is greater than the energy released when new bonds are formed in the products, the enthalpy change is positive, indicating an endothermic reaction. If more energy is released during product bond formation than is consumed in breaking reactant bonds, the enthalpy change is negative, signifying an exothermic reaction. It's this balance of energy during molecular interactions that determines the sign of \( \Delta H \).

Exothermic reactions are characterized by the release of heat energy to the surroundings, typically resulting in an increase in temperature. These reactions occur when the energy needed to form the bonds in the products is less than the energy released when the bonds in the reactants are broken. The classic examples of exothermic processes include combustion of fuels, explosive reactions, and many everyday chemical reactions like the rusting of iron.

The enthalpy change in these reactions is always negative (\(\Delta H < 0\)), which can be thought of as the system giving energy away. This concept is crucial for understanding such reactions because it impacts the feasibility and sustainability of many chemical processes, such as in industrial synthesis and energy production.

Endothermic reactions, in contrast to exothermic, absorb energy from the surroundings, which often results in a temperature decrease. These reactions require an input of energy to proceed, as forming the products' bonds requires more energy than what is released from breaking the reactants' bonds. Examples of endothermic reactions include photosynthesis, cooking an egg, or even the aforementioned melting of ice when salt is applied.

In these reactions, the sign of the enthalpy change is positive (\(\Delta H > 0\)), indicative of the system gaining energy from the environment. Understanding endothermic reactions is essential not just for academic purposes, but also for practical applications such as refrigeration and understanding climate change implications.

Chemical bonding is the force that holds atoms together in molecules and compounds, determining their structure, properties, and reactivity. Bonds form when atoms share or transfer electrons, and the type of bond – covalent, ionic, or metallic – depends on the nature of the interaction between atoms. Chemical bonding is at the heart of all the chemical reactions discussed; whether a reaction is exothermic or endothermic hinges on the changes in bond energy.

For example, in the formation of a molecule like \(\mathrm{F}_2\), the strong chemical bond between the two fluorine atoms stores less energy than the individual reactive atoms, resulting in an exothermic reaction. Similarly, the process of breaking the ionic bond in HBr molecule requires energy, leading to an endothermic reaction. Grasping the concept of chemical bonding is thus fundamental to predicting the behavior of atoms and molecules during chemical reactions.