Bond dissociation energies of HF, HCl, HBr follows the order \(\ldots \ldots\) (a) \(\mathrm{HBr}>\mathrm{HCl}>\mathrm{HF}\) (b) \(\mathrm{HCl}>\mathrm{HBr}>\mathrm{HF}\) (c) \(\mathrm{HF}>\mathrm{HBr}>\mathrm{HCl}\) (d) \(\mathrm{HF}>\mathrm{HCl}>\mathrm{HBr}\)

Bond strength is a key factor in determining a molecule's properties, including its reactivity and stability. In the context of hydrogen halides like HF, HCl, and HBr, the bond strength refers to the energy required to separate the hydrogen atom from its halogen partner. Stronger bonds have higher dissociation energies, meaning they require more energy to break apart.

One might ask, what influences bond strength? There are several factors at play, including the elements involved in the bond, the distance between the nuclei (bond length), and the match of orbitals between the bonding atoms. Generally, when the bonded atoms have a small size and the bond length is short, the bond is strong due to the increased overlap of the electron orbitals. This principle leads to the observation that HF possesses the strongest bond among the given molecules, as fluorine is the smallest and most electronegative of the halogens.

Studying halogen bond trends is a fascinating facet of inorganic chemistry. Within the group of halogens on the periodic table, there is a clear trend in the size and reactivity of atoms. As we move down the group from fluorine to chlorine to bromine, and further to iodine, the size of the atoms increases. This is due to the addition of electron shells.

As a consequence of increasing atomic size, bond lengths to hydrogen also increase down the group, correlating inversely with bond strength. This relationship illustrates why HF has the highest bond dissociation energy and why the energy decreases through HCl to HBr. It's essential to understand these periodic trends, as they are consistent and predictable elements of bonding behavior in halogens, significantly affecting their chemical reactivity and interactions with other elements.

Periodic trends are tendencies or patterns that can be observed in the elements on the periodic table, affecting their chemical and physical properties. These trends arise due to the periodic nature of the atomic structure of elements, particularly the configuration of electrons.

Key periodic trends include atomic size, ionization energy, electronegativity, and of course, bond dissociation energy. Understanding these trends allows chemists to predict how different elements will behave and interact. For instance, it is known that across a period from left to right, the atomic size decreases while electronegativity and ionization energy typically increase. Conversely, down a group, atomic size increases, and electronegativity and ionization energy usually decrease. These trends are crucial when exploring why a bond like HF is stronger than HBr, despite bromine being a larger atom than fluorine. It is the nuanced understanding of these patterns that enhances the comprehension of chemical behaviors and assists in predicting the nature of bonds within different molecules.