Detailed Notes on Electron Shells, Ionization Energy, and Core Charge

Electrons reside in shells around the nucleus.
Each shell can hold a limited number of electrons.
Electrons fill the innermost shells first before occupying outer shells.
Electrons in outer shells are at higher energy levels and are less strongly attracted to the nucleus.

The outermost shell of an atom is called the valence shell.
Electrons in the valence shell are called valence electrons.
Valence electrons determine the chemical properties of an atom.
Atoms in the same group of the periodic table have the same number of valence electrons, leading to similar chemical properties.
- Example: Fluorine (F) and Chlorine (Cl) have seven valence electrons and similar ionization energies.
Electrons not in the valence shell are called core electrons.

Ionization energies increase as we move across the periodic table.
This is due to increasing nuclear charge ( $+Z$ ) while electrons are in the same shell (approximately the same distance from the nucleus).
Increased nuclear charge leads to a stronger attraction and higher ionization energies.

Ionization energies decrease significantly as we move from the rightmost element (Neon) to the next element (Sodium) on the left of the next period.
The added electron in the next element occupies a new shell farther from the nucleus.
Increased distance reduces the attraction to the nucleus and lowers the ionization energy.

Neon (Ne) has a higher ionization energy than Fluorine (F).
Fluorine (F):
- Nuclear charge: +9
- Electron configuration: 2 electrons in the first shell, 7 in the valence shell.
Neon (Ne):
- Nuclear charge: +10
- Electron configuration: 2 electrons in the first shell, 8 in the valence shell.
The model suggests that the increased nuclear charge in Neon results in greater attraction for electrons.

Electrons are repelled by other electrons (negative charges).
In Neon, an electron is attracted to a +10 nucleus but repelled by 9 other electrons.
In Fluorine, an electron is attracted to a +9 nucleus but repelled by 8 other electrons.
The additional nuclear charge is accompanied by an additional electron, leading to increased electron-electron repulsion.
It raises the question: Why does the increase in nuclear attraction outweigh the increase in electron repulsion as atomic number increases?

Successive ionization energies involve removing multiple electrons from an atom or ion.
First ionization energy ( $IE_1$ ): Energy to remove the first electron.
- $X(g) \rightarrow X^+(g) + e^-$
Second ionization energy ( $IE_2$ ): Energy to remove the second electron from the +1 ion.
- $X^+(g) \rightarrow X^{2+}(g) + e^-$
And so on for subsequent ionizations.

A significant jump occurs between the first and second ionization energies (a factor of ~9).
Smaller increases are observed between subsequent ionization energies.
Sodium has a +11 nuclear charge and 11 electrons (2 in the inner core, 8 in the next shell, 1 valence electron).
Removing the first valence electron requires a certain amount of energy based on attraction to the nucleus and repulsion from other electrons.
Removing a second electron (from the core) requires significantly more energy because it is closer to the nucleus.
The large jump between $IE1$ and $IE2$ indicates that Sodium has one valence electron.

The number of valence electrons in each element can be determined by looking for large jumps in successive ionization energies.
Magnesium (Mg): Large jump between $IE2$ and $IE3$ , indicating 2 valence electrons.
Sulfur (S): Large jump between $IE6$ and $IE7$ , indicating 6 valence electrons.

Removing an electron closer to the nucleus wouldn't cause a factor of 9 increase, so there must be more contributing to the increase in ionization energy.
This is because any given valence electron is repulsed by the core electrons. The increased ionization energy is a direct result of removing the valence electron, leaving only strongly bound core electrons.
The definition of core charge is the nuclear charge ( $+Z$ ) minus the number of core electrons.
The core charge attempts to quantify the effective nuclear attraction experienced by valence electrons, considering the shielding effect of core electrons.
Sodium Core Charge: +11 (nuclear charge) - 10 (core electrons) = +1
The valence electron is effectively feeling an attraction to a net +1 charge.

Removing the valence electron creates a sodium ion ( $Na^+$ ) with 10 core electrons.
In the sodium ion, the outermost shell is closer to the nucleus; the electron you are trying to remove is significantly closer to the nucleus, and therefore much harder to remove than the original valence electron.
The core charge for electrons in the valence shell of the sodium ion is +11 (nuclear charge) - 2 (core electrons) = +9 (enormous increase).
This significant increase in core charge accounts for the large increase in ionization energy.

The number of total electron electron repulsions also plays a role.
In a sodium atom, the electron trying to be removed is repulsed by 10 core electrons.
Once that electron is removed, any remaining electrons are only repulsed by 8 electrons.
Even after all valence electrons are removed, successive ionization energies continue to increase due to valence electron - valence electron repulsion.

Removing an electron from a positively charged ion is more difficult than from a neutral atom due to increased attraction.

Electron-electron repulsion has a twofold effect:
- Repulsion of valence electrons by core electrons (accounted for by core charge).
- Repulsion amongst valence electrons.
The model accounts for variations in ionization energies by considering valence shells and core charges.
Much of chemistry can be understood by understanding the ionization energies of atoms because they are related to how strongly attracted the electrons are to each atom.