Ionization energy: types, factors, calculations, exceptions to trends (2023)

This is the energy required to remove one or more electrons from a neutral atom to produce a positively charged ion that changes the chemical behavior of the atom.

Ionization energy (IE) is necessary to remove one or more electrons from neutral valence to produce a positively charged ion and change the chemical properties of an atom.

It is often expressed in kJ/mol, and the measurement is based on a single gas atom. All elements in the periodic table have a characteristic known as ionization energy, sometimes called ionization potential.

i.e. can be represented as shown

• X + first ionization energy → X⁺+ e-

Where

• X is a neutral atom.
• X⁺is an ion of atom X with a single positive charge.
• is an electron with a single negative charge.

Types of ionization energy

First ionization energy

The "first ionization energy" of an element refers to the energy required to remove the outermost valence electron from a neutral atom in the gas phase.

In the equation, it refers to I.E. required to remove the first electron from a neutral atom, yielding an ion with a single positive charge. The process by which the first I.E. the measured sodium content will be represented by the following equation:

Na (g) + energy→Already⁺(g) + e^–

Second ionization energy

The second ionization energy needed to remove the next electron is always higher than the first ionization energy because it takes even more energy to remove an electron from a cation than from a neutral atom.

In the case of an alkali metal atom, since its loss gives the atom a stable electron shell, removal of the first electron is quite simple. The removal of the second electron creates a new electron shell that is tighter and more closely bound to the atomic nucleus.

The process by which the second I.E. the measured sodium content will be represented by the following equation:

Na (g) + energy→Already²⁺(g) + e^–

Third ionization energy

The "third ionization energy" of an element refers to the energy required to remove the third electron. It has to be formed³⁺cations.

The process by which the third I.E. the measured sodium content will be represented by the following equation:

Already²⁺(g) + energy→Already³⁺(g) + e^–

All elements have a first ionization energy, even atoms that do not form positive ions in test tubes.

Throughout the periodic table, the first ionization energy varies continuously. IE decreases from top to bottom in groups and increases from left to right in a period. Therefore, helium has the highest first ionization energy, while francium has one of the lowest.

Factors affecting the amount of ionization energy

Ionization energy is a measure of the force needed to eject a particular electron from the gravitational pull of the nucleus. The high price of I.E. shows a strong attraction between the electron and the nucleus. The size of this attraction will be determined as follows:

The distance of the electron from the nucleus:

The distance between the electrons and the nucleus affects the attraction. It decreases rapidly as space increases. An electron closer to the nucleus is attracted more strongly than an electron farther away.

Nuclear charge:

When the number of protons in a nucleus increases, its positive charge increases and vice versa. Electrons tend to be strongly attracted to positively charged nuclei.

The number of electrons between the outermost electrons and the nucleus:

Let's look at the sodium atoms with the electron structure 2,8,1.

There are two electron shells between the outer electron and the nucleus. The 11 protons in the sodium nucleus are reduced by the inner electrons. The outer electron only feels a net attraction of about 1+ from the center. This reduction in the strength of the nucleus by the inner electrons is known as shielding or shielding.

Penetration Electron effect
A single atom can contain various subshells, including the s, p, d, and f subshells. Unlike the p and other subshells, the s subshell is known to be closer to the nucleus. Unlike the other subcortex, the s subcortex has a stronger affinity for the nucleus, making it difficult to remove electrons from it. Therefore, additional ionization energy is required to remove electrons from the s subshell.

Fully and half-filled electrons

I.E is greater in fully filled orbitals than in partially filled or partially filled orbitals. Partially filled orbitals are less stable than fully filled orbitals.

Calculation of ionization energy

The following equation can be used to calculate the ionization potential of hydrogen:

E =doc R_H( 1/n²), Where

• Mis the energy of the electron (or the amount of energy needed to remove an electron, ionization energy)
• theis Planck's constant = 6.626 * 10^-34Js (seconds)
• Dois the speed of light = 3.00 * 10⁸m/s (meters/second)
• R_His the Rydberg constant = 1.097 * 10⁷M^-1(1/meter)
• Nis the principal quantum number (or energy level) of the electron

After applying constant values, the equation becomes:

E = (2,18 * 10^-18J)(1/n²)

From here, you can enter a value for the electron's energy level to calculate how much power is needed to remove it.

A unit of ionization energy

Ionization energy is usually expressed in kcal/mol, kJ/mol, or electron volts (eV) per atom. Mathematics,

A single electron volt (eV) is 3.827*10^-20per person.

= 3827 x 10^-20x 4.184 calories per person (∵cal=4.184 J)

= 1602 x 10^-19J per person

= 1602 x 10^-19x 6022 *10²³J/mol

= 96472 J mol^-1

= 96 472 kJ times^-1

Energy potential of ionizationPeriodic table

• Moving from left to right over the period of an element (series), the ionization energy tends to increase. The reason for this is that the radius of an atom decreases as we go through a period. This occurs when more protons are supplied, increasing the nucleus-electron attraction and pushing the electron shells closer together.
• The ionization energy generally decreases moving from top to bottom of a group of elements (column) as the principal quantum number of the outermost (valence) electron increases moving downwards. Atoms have more protons moving down the group which attract electron shells. The outer electrons are further away from the nucleus because each row adds a new shell.
• Helium, one of the noble gases, has the highest ionization energy and is found in the upper right corner of the periodic table. Fransium, the alkali metal on the left side of the map, has one of the lowest ionization energies.
• Typically, group 2 elements have a higher ionization energy than group 13 elements, and group 15 elements have a higher ionization energy than group 16 elements.
• Groups 2 and 15 have fully and half-filled electron configurations, respectively, so more energy is required to remove an electron from fully filled orbitals than from incompletely filled orbitals.
• The alkali metals (group IA) have low ionization energies, especially compared to halogens or group VIIA.

Turn off the ionization energy trend in the periodic table

• Boron has a lower first I.E. instead of beryllium. By comparing their electron configuration, we can analyze differences in their ionization energies. Boron with outer electrons in the "s" orbital is very close to the nucleus, while beryllium releases its first electron from the p orbital, which is relatively far from the nucleus.
• Nitrogen has three p-orbital electrons while oxygen has four electrons. Nitrogen should have a lower IE, but instead has a lower first ionization energy than oxygen due to the symmetrical arrangement of the three electrons in the p orbital of the nitrogen atom, according to Hund's rule. Another reason for this exception is electron-electron repulsion.
• Lanthanides and actinides have lower ionization energies due to the presence of an inner "f" electron shell.

1st, 2nd and 3rd ionization energies

• I₁ = first ionization energy (energy needed to remove an electron from a neutral atom)
• I₂is the second ionization energy (the energy needed to remove an electron from an atom with a +1 charge).

Each subsequent ionization energy is greater than the previous one.

That impliesI₁23<...N will always be true.

Here is an example showing how the ionization energy increases as electrons are ejected sequentially.

Mg (g) → Mg⁺(g) + e^–I₁= 738 kJ/mol

Mg⁺(g) → Mg²⁺(g) + e⁻I₂= 1451 kJ/mol

See the first, second and third ionization energies of elements/ions

bibliographic references

• F. Albert Cotton i Geoffrey Wilkinson, Advanced Inorganic Chemistry (wyd. 5, John Wiley 1988) s. 1381.
• Lang, Peter F.; Smith, Barry C. "Ionization energies of atoms and atomic ions." Journal of Chemical Education. 80 (8).
• Cotton, F. Albert; Wilkinson, Geoffrey (1988). Advanced Inorganic Chemistry (5η έκδ.). John Wiley. ISBN0-471-84997-9.
• Housecroft, CE; Sharpe, AG (1 listopada 1993).Inorganic Chemistry(eBook). Inorganic Chemistry. Volume 3 (15th). Switzerland: Pearson Prentice-Hall. pp. 536, 649, 743.
• https://brilliant.org/wiki/ionization-energy/
• https://chemistrytalk.org/ionization-energy-trend/