Ionization Enthalpy and factors affecting ionization enthalpy

Ionization energy, often known as ionisation enthalpy, is defined as the amount of energy that an isolated gaseous atom needs to lose an electron in its ground state.

The energy required by an atom to produce A+ ions is known as the first ionisation energy of element A. The ionisation energy is measured in kilo Jule. 

A+ (g) + e– A (g) A+ (g) A+ (g) A+ (g) A+ (g) A+ (g

Second ionisation energy, on the other hand, is defined as the energy required to remove the second electron from its valence shell. The following given  equation can be used to describe it: 

A+ (g) + A2+ (g) + e– e– e– e– e– e– e– e– e–

Because it takes a specified amount of energy to remove an electron from an atom, the ionisation enthalpies of chemical elements are always positive. The nucleus will attract the second outer electron more than the first outer electron. As a result the second ionisation energy will be higher than the first one. Similarly. the third ionisation enthalpy will be higher than the second.

Ionization Energy Influencing Factors

The energy of ionisation is determined by two factors:

1.Electrons and the nucleus are attracted to each other.

2.The repelling force between electrons.

The outermost electrons will feel a lower effective nuclear charge than the real nuclear charge. Because the inner electrons will obstruct the nuclear charge path, the outermost electrons will be shielded. The shielding effect is the name for this phenomenon. In Na, for example, the 3s1 electrons are sheltered by the core electrons (1s2, 2s2 and 2p6). When the inner orbitals are totally occupied, the shielding effect is more noticeable.

Ionization Energy Trend in Periodic Table

Periodic tendencies in general:

1. Ionization enthalpy trends in a group:

As we proceed down in a group, the initial ionisation enthalpy of elements drops. The atomic number and the number of shells both grow as you move down in a group. Because the outermost electrons are far from the nucleus, they can be easily removed. The shielding effect caused by an increasing number of shells as we progress down a group is the second component that reduces ionisation energy.

2. Changes in ionisation enthalpy over time:

The ionisation energy of elements increases as we move from left to right across a period. This is related to the reduction in the size of atoms throughout time. As we proceed from left to right, the valence electrons approach closer to the nucleus of an atom due to rising nuclear charge. The nucleus’s force of attraction with the electrons increases, requiring more energy to remove an electron from the valence shell.

Ionisation enthalpy

The first ionisation enthalpy is the enthalpy change associated with the loss of the first electron from an isolated gaseous atom in its ground state. The amount of energy required to remove an electron from an isolated gaseous atom in its gaseous form is known as the ionisation enthalpy of an element.

The following factors influence ionisation enthalpy:

• Effect of Penetration
• Effect of a shield

1. Effect of Penetration

The proximity of an electron in an orbital to the nucleus is known as  penetration. It shows the relative density of electrons near the nucleus of an atom for each shell and subshell. The radial probability distribution functions can help us to  observe that the electron density of s orbitals is closer than the electron density of p and d orbitals.

2s > 2p > 3s > 3p > 4s > 3d will be the order of penetration power.

2. The Protective Effect

The shielding effect is defined as an effect in which inner electrons form a shield for outer shell electrons, preventing the proper nuclear charge from reaching the outermost electrons. The outermost electrons  have a low effective nuclear charge rather than the actual nuclear charge. 

3. Configuration of electronic devices

Stable elements have orbitals that are half-filled or fully filled. When we  try to take an electron out of these orbitals, it will make them less stable. As a result  removing an electron from these orbitals requires more energy. Due to this process the ionisation energy is increased.

Trend in the Periodic Table

As we move  from left to right the  atomic radius decreases.The size of an atom shrinks, the attractive attraction between the nucleus and the outermost electrons rises, resulting in an increase in ionisation energy throughout a period in the periodic table.

There is a difference in the trend  of ionisation enthalpy from boron to beryllium in the second period. The ionisation enthalpy of boron should, in theory, be greater than that of beryllium, but this is not the case. The reason for this is that beryllium contains fully filled subshells, and it also has a penetrating effect. Boron has both 2s and 2p orbitals, whereas beryllium only contains 2s orbitals. The 2s orbital has a higher penetrating power than the 2p orbital. As a result, removing an electron from the 2p subshell in beryllium will be easier than from the 2s subshell. As a result of these two considerations, beryllium’s ionisation enthalpy will be higher than that of boron.

As the number of shells in a group grows, the ionisation energy of elements drops. The outermost electrons will be furthest from the nucleus, resulting in a lower effective nuclear charge. Second, as the number of shells in the group grows, the shielding effect increases, resulting in a decrease in ionisation energy.

Conclusion 

Therefore, Ionization Enthalpy is considered a measure of the tendency of an atom or ion to surrender an electron or the strength of the electron binding. The greater the ionization energy, the more difficult it is to remove an electron. The ionization energy may be an indicator of the reactivity of an element.