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Ionization enthalpy is defined as the smallest amount of energy required to remove the most loosely bonded electron from an isolated gaseous atom in order to transform it into a gaseous cation. Ionization enthalpy is measured in joules per kilogram of mass of a gaseous atom.
It is measured in electron volts (eV) per atom, kilocalorie per mole, or kilojoules per mole, among other units of measurement.
The force of attraction between electrons and the nucleus, as well as the force of repulsion between electrons, are all factors that influence the I.E of an atom.
The measurement principle is:
Typically, the ionization energy of a chemical element is measured in an electric discharge tube, in which a fast-moving electron generated by an electric current collides with a gaseous atom of the element, causing it to expel one of its electrons. The ionization energy is expressed in joules or electron volts.
Trends in ionization energy
The ionization energy is proportional to the radius of the atom. Since moving from right to left on the periodic table increases the atomic radius, and going from left to right in the periods and up the groups raises the ionization energy, traveling from right to left on the periodic table makes sense. The alkaline earth metals (group 2) and the nitrogen group elements are two examples of exceptions to this general trend (group 15). Typically, the ionization energy of group 2 elements is more than that of group 13 elements, while the ionization energy of group 15 elements is greater than that of group 16 elements. Groups 2 and 15 feature electrical configurations that are entirely and half-filled, respectively, and as a result, it takes more energy to remove an electron from completely filled orbitals than it does from incompletely filled orbitals.
Small ionization energy is observed in alkali metals (IA group), compared to halogens or the VII A group, which is observed in a large number of other elements. Beyond the radius (the distance between the nucleus and the electrons in the outermost orbital), the number of electrons between the nucleus and any electrons in the outermost orbital you’re examining has an effect on the ionization energy as well as the number of electrons between you and the electron(s) you’re examining in the outermost shell. Shielding is the term used to describe the phenomenon in which the entire positive charge of the nucleus is not felt by outer electrons as a result of the negative charges of inner electrons partially canceling out the positive charge. More electrons sheltering the outer electron shell from the nucleus means less energy is required to expel an electron from an atom with a large number of electrons. The lower the ionization energy is, the more the shielding effect exists. It is because of the shielding effect that the ionization energy drops from the top to the bottom of a group as it moves downward. According to this pattern, Cesium has the lowest ionization energy, while Fluorine has the highest ionization energy, and the opposite is true for oxygen (with the exception of Helium and Neon).
Factors on which ionization enthalpy depends
The Ionization Energy is determined by two factors, which are as follows:
- The force of attraction between electron and nucleus
- The force of repulsion between electrons
There will be a difference between the effective nuclear charge felt by the outermost electrons and the actual nuclear charge felt by them. Due to the fact that the inner electrons will shelter the outermost electrons by interfering with the passage of nuclear charge, this will occur. The shielding effect is the name given to this phenomenon. For example, the 3s1 electrons in Na will be sheltered by the core electrons of the atom (1s2, 2s2 and 2p6). It is generally accepted that a greater amount of shielding is produced when the inner orbitals are entirely occupied.
Period trends
Ionization enthalpy decreases when moving from top to bottom.
Ionization enthalpy increases when moving left to right across the period.
Trends in ionization enthalpy in a group
As we travel down in a set of elements, the initial ionization enthalpy of the elements drops. While advancing down in a group, the atomic number grows, as does the number of shells, which is a positive feedback loop.
Because the outermost electrons are located far away from the nucleus, they can be easily removed. The second component that contributes to a reduction in ionization energy is the shielding effect caused by an increasing number of shells as we proceed down the group hierarchy.
Trends in ionization enthalpy across a period
Increasing ionization energy of elements is observed as we progress from left to right across a period. As a result of the decrease in the size of atoms over time, this is the case.
As we proceed from left to right, the valence electrons of an atom get closer to the nucleus as a result of the higher nuclear charge on the nucleus. In order to remove one electron from the valence shell, more energy must be used in order to enhance the force of attraction between the nucleus and electrons.
Conclusion
Ionization enthalpy is defined as the smallest amount of energy required to remove the most loosely bonded electron from an isolated atom in order to transform it into a gaseous cation. The ionization energy is expressed in joules or electron volts. Ionization Energy is determined by two factors, which are a force of attraction and repulsion. Cesium has the lowest ionization energy, while Fluorine has the highest. The opposite is true for oxygen (with the exception of Helium and Neon), and for many other atoms.
The ionization energy of atoms changes as we move down the group hierarchy. As we advance down a group, the atomic number and number of shells grow, which is a positive feedback loop. This means that more energy must be used in order to enhance the force of attraction between the nucleus and electrons.
