CHEM 1411 Concept Reviews: Periodic Properties of the Elements Part II

Valence Orbitals and Shielding Effects

Valence Orbitals: The outermost occupied orbitals of an atom. The electrons in these orbitals are called valence electrons and are the electrons involved in bonding and chemical reactions. Since the d orbitals and f orbitals always fill after the outer shell (higher energy level) s and p orbitals, ONLY s and p orbitals can ever be properly considered “valence”.

In an atom, the nucleus contains all the positively charged protons. Thus, the nucleus has a strong attractive force that it exerts on the negatively-charged electrons orbiting around it. However, outer shell electrons can be shielded from the nucleus’ attractive positive charge by inner shell electrons. Electrons in the same shell (energy level) do not shield each other very well. Therefore, to keep things simple, we usually only consider inner shell electrons when determining the effect of shielding. In general, the effective nuclear charge (Z_eff), the amount of positive nuclear charge “felt” by a given electron, can be determined by the following equation:

$LaTeX: Z_{eff\:}=Z-S$ $Z_{eff\:}=Z-S$

where Z is the nuclear charge (# of protons), S is the shielding constant (# of inner shell electrons), and Z_eff is the effective nuclear charge.

Example: What is the effective nuclear charge experienced by the valence electrons of a Sulfur atom?

The electron configuration for Sulfur: 1s²2s²2p⁶3s²3p⁴ The bolded orbitals are the valence orbitals.

This gives us 6 valence or outer shell electrons and 10 inner shell electrons. We can then estimate our screening constant S to be “10” since we have 10 inner shell electrons. Sulfur has a total of 16 protons, therefore Z is 16.

Using our equation above, $LaTeX: Z_{eff\:}=16-10=$ $Z_{eff\:}=16-10=$ 6

Going across the period, one will notice that Z increases while S remains the same. As such, effective nuclear charge increases across a period. This trend is the reason behind many other trends we see in the periodic table.

Trends in the Periodic Table

There are two primary factors that are responsible for the various trends seen in the periodic table. They are effective nuclear charge and the number of filled shells (energy levels) of electrons.

Atomic Radii: usually refers to the bonding radius of an atom which is one half the distance between the two bonded nuclei of that atom.

Trend: In general, since effective nuclear charge increases going left to right across a period, the atomic radii decreases going left to right across a period. This is because a larger effective nuclear charge will cause electrons in an atom to be pulled in closer to the nucleus due to increased electrostatic force, thereby decreasing the size of the radii. Going down a group, atoms have more filled shells of electrons. Therefore, atomic radii increase going down a group.

Ionization Energy: The minimum energy required to remove an electron from an isolated gaseous atom or ion. The first ionization energy of an atom refers to the energy needed to remove a first electron from that gaseous atom. The second ionization energy refers to the energy required to remove a second electron after a first electron has already been removed. Shown below are the equations that represent the first and second ionization energies of sodium.

First ionization energy: $LaTeX: Na\left(g\right)\longrightarrow Na^+\left(g\right)+e^-$ $Na\left(g\right)\longrightarrow Na^+\left(g\right)+e^-$ I₁ = 496 kJ/mol

Second ionization energy: $LaTeX: Na^+\left(g\right)\longrightarrow Na^{2+}\left(g\right)+e^-$ $Na^+\left(g\right)\longrightarrow Na^{2+}\left(g\right)+e^-$ I₂ = 4562 kJ/mol

Trend: In general, as effective nuclear charge increases, removing an electron from an atom becomes more difficult, thus requiring more energy. Since effective nuclear charge increases going left to right across a period, ionization energy increases going left to right across a period. However, electrons that are farther away from the nucleus are held less strongly by the nucleus and are therefore easier to remove, thus requiring less energy. Since atomic size increases going down a group, ionization energy decreases going down a group.

Electron Affinity: The energy change that occurs when an electron is added to an isolated gaseous atom. As an odd twist, since energy is released when an atom with a greater attraction to electrons gains an electron, a more negative electron affinity is considered to be a “greater electron affinity”. An example is shown below:

Electron Affinity for Oxygen: $LaTeX: O\left(g\right)+e^-\longrightarrow O^-\left(g\right)$ $O\left(g\right)+e^-\longrightarrow O^-\left(g\right)$ $LaTeX: \Delta E=-141\:kJ/mol$ $\Delta E=-141\:kJ/mol$

Trend: In general, as effective nuclear charge increases, the more an atom is attracted to electrons. As such, since effective nuclear charge increases going left to right across a period, electron affinity increases (becomes more negative) going left to right across a period. In contrast, the more filled shells of electrons an atom has, the more difficult it becomes to attract electrons. As such, electron affinity decreases (becomes less negative or more positive) going down a group.

CHEM 1411 Ch 8-2 img 1.png

Important Exceptions: Since half-filled and completely filled orbitals have an enhanced stability. Therefore, breaking such a configuration will require much more energy. Conversely, forming such a configuration will require much less energy. This leads to disturbances in the trends near groups 2 and 5 (group 2 elements have a completely filled s-orbital and group 5 elements have a half-filled p-orbitals).