Part 8(C): Ionic Radii & Lattice Energy of Transition Metal Ions скачать в хорошем качестве

Part 8(C): Ionic Radii & Lattice Energy of Transition Metal Ions

The ionic radius is half the distance between atomic ions in a crystal lattice. To find the value, ions are treated as if they were hard spheres. The size of an element's ionic radius follows a predictable trend on the periodic table. As you move down a column or group, ionic radius increases. Lattice Energy Lattice Energy is a type of potential energy that may be defined in two ways. In one definition, the lattice energy is the energy required to break apart an ionic solid and convert its component atoms into gaseous ions. This definition causes the value for the lattice energy to always be positive, since this will always be an endothermic reaction. The other definition says that lattice energy is the reverse process, meaning it is the energy released when gaseous ions bind to form an ionic solid. As implied in the definition, this process will always be exothermic, and thus the value for lattice energy will be negative. Its values are usually expressed with the units kJ/mol. Lattice Energy is used to explain the stability of ionic solids. Some might expect such an ordered structure to be less stable because the entropy of the system would be low. However, the crystalline structure allows each ion to interact with multiple oppositely charge ions, which causes a highly favorable change in the enthalpy of the system. A lot of energy is released as the oppositely charged ions interact. It is this that causes ionic solids to have such high melting and boiling points. Some require such high temperatures that they decompose before they can reach a melting and/or boiling point. The Limitations Of Crystal Field Theory The crystal field theory is highly useful and more significant as compared to the valence bond theory. Even after such useful properties, it has many limitations. The following points will clearly state the limitations of crystal field theory: The assumption that the interaction between metal-ligand is purely electrostatic cannot be said to be very realistic. This theory takes only d-orbitals of a central atom into account. The s and p orbits are not considered for the study. The theory fails to explain the behaviour of certain metals which cause large splitting while others show small splitting. For example, the theory has no explanation as to why H2O is a stronger ligand as compared to OH–. The theory rules out the possibility of having p bonding. This is a serious drawback because is found in many complexes. The theory gives no significance to the orbits of the ligands. Therefore it cannot explain any properties related to ligand orbitals and their interaction with metal orbitals.