ionisation energy, which is also called ionisation enthalpy of elements, can be defined as the amount of energy needed by an isolated gaseous atom to lose the electron in its ground state. Electron loss results in cation formation. The initial/first ionisation energy or otherwise called the Ei of a molecule/atom is the energy used to detach one mole of electrons from one mole of isolated gaseous ions or atoms.
Ionisation Energies of First, Second, and Subsequent Ones
The energy needed to remove the outermost valence electron from the neutral atom is defined as the first ionisation energy. The second ionisation energy is the one needed to remove the next electron, and so on. At the same time, the second ionisation energy is always higher compared to the first ionisation energy. For example, an alkali metal atom can be taken. Removing the first electron is relatively easy due to the reason its loss gives the atom, which is a stable electron shell. And, removing the second electron includes a new electron shell, which is closer and more tightly bound to the atomic nucleus.
The first ionisation energy of the hydrogen can be represented by the equation as given below:
H(g) → H+(g) + e-
ΔH° = -1312.0 kJ/mol
Factors Affecting the Ionisation Energy
ionisation energy primarily depends on two factors:
The effective nuclear charge, which is felt by the outermost electrons, will be less compared to the actual nuclear charge. This is due to the inner electrons will shield the outermost electrons by hindering the nuclear charge path. This effect is referred to as the shielding effect. For suppose, in the Na compound, the 3s1 electrons will be shielded with its core electrons (1s2, 2s2, 2p6). In general, the shielding effect is very prominent when the inner orbitals are filled completely.
Ionisation Energy Trend in Periodic Table
General Periodic Trends
While moving from top to bottom in a group, it decreases.
It increases starting from left to right across a period.
Trends in Ionisation Enthalpy in a Group
In a group, the first ionisation enthalpy of the elements decreases while we move down. Whereas, while moving down in a group, the atomic number will increase, and the number of shells will also increase. From the nucleus, the outermost electrons are far away and therefore can be removed easily. The second or another factor that decreases the ionisation energy is the shielding effect as we move down a group because of an increasing number of shells.
Trends in Ionisation Enthalpy Across a Period
While we move from the left to right across a period, the element's ionisation energy increases. This happens because of the decrease in the size of atoms across a period. And, the valence electrons get closer to the atom's nucleus while we move from the left to right because of the increased nuclear charge. The attraction force between the electrons and the nucleus increases, and thus more energy is needed to remove an electron from the valence shell.
Exceptions to the Ionisation Energy Trend
If we look at the first ionisation energies chart, two exceptions to the trend seem to be readily apparent. The first ionisation energy of the boron is less compared to the beryllium, and the first ionisation energy of the oxygen is less compared to nitrogen.
The reason for the discrepancy is because of the Hund's rule and electron configuration of these elements. The first ionisation potential electron for beryllium comes from the 2s orbital, although the boron ionisation involves a 2p electron. For both oxygen and nitrogen, the electron comes from the 2p orbital, whereas the spin is similar for all 2p nitrogen electrons, while there is a paired electron set in one of the 2p oxygen orbitals.
Let us look at the major key points that are related to the periodic trends in ionisation enthalpy as follows:
The most common units of the ionisation energy are given as either electron volts (eV) or kilojoules per mole (kJ/M).
ionisation energy is explained as the minimum energy needed to remove an electron from either an ion or atom in the gas phase.
The general trend for the ionisation energy is to increase moving from left to right direction across an element period. In contrast, the atomic radius decreases while moving left to the right direction across a period. Thus, the electrons are more attracted (closer) to the nucleus.
On the periodic table, ionisation energy shows periodicity.
The general trend is for ionisation energy to decrease when moving from the top to bottom direction down a periodic table group. Whereas, moving down a group, a valence shell can be added. Further, the outermost electrons are from the positive-charged nucleus; hence they are easier to remove.