CP4171 Chemistry I - Chemical Periodicity

Nuclear Charge, Shielding Effect & Effective Nuclear Charge

Comparing Valence and Core Electrons

Core electrons are more effective in shielding valence electrons from the nuclear charge due to their proximity to the nucleus.

Effective Nuclear Charge (ENC)

Effective nuclear charge (ENC) represents the net positive charge experienced by an electron in an atom. It considers both the attraction to the nucleus and the repulsion from other electrons.

Nuclear Charge

The attraction between protons in the nucleus and electrons is electrostatic, referred to as nuclear charge.

Shielding Effect

Shielding effect, also known as electron-electron repulsion or screening effect, occurs when an electron is partially shielded from the positive charge of the nucleus by other electrons. Core electrons are the most effective at shielding.

Effective Nuclear Charge (ENC) Explained

Electrons are simultaneously attracted to the nucleus and repelled by each other. Core electrons shield valence electrons from the full nuclear charge. The net nuclear charge acting on a valence electron is the Effective Nuclear Charge (ENC).

Trends in Effective Nuclear Charge (ENC)

  • Across a Period: ENC increases across a period because electrons are added to the valence shell while the number of core electrons remains the same, but the number of protons increases.

  • Down a Group: ENC remains approximately the same down a group because the number of valence electrons remains constant.

Formula for Effective Nuclear Charge

ENC=Atomic No.−No. of core e−sENC = Atomic \, No. - No. \, of \, core \, e^{-s}

Atomic Radii Trends

Atomic Radii Trend Across a Period

Going across a period (left to right), the atomic radius decreases; atoms get smaller.

  • The nuclear charge increases due to the increasing number of protons, but the shielding effect remains similar.

  • The effective nuclear charge (ENC) increases.

  • The number of shells remains the same.

  • As ENC increases, valence electrons are more strongly attracted to the nucleus, drawing them closer and reducing atomic size.

Atomic Radii Trend Down a Group

Going down a group (top to bottom), the atomic radius increases; atoms get bigger.

  • The nuclear charge increases due to the increasing number of protons.

  • The shielding effect increases due to the increase in core electrons.

  • The effective nuclear charge (ENC) remains similar.

  • The number of shells increases.

  • The distance from the nucleus to the valence electrons increases, leading to a larger atomic size.

First Ionization Energy Trends

General Trend

Generally, the first ionization energies increase from left to right across a period, with exceptions.

Factors Influencing Ionization Energy

  • Atomic Size: As atomic size decreases and effective nuclear charge (ENC) increases across a period, outer electrons are held more tightly, increasing ionization energy.

  • Electron Shells: The number of electron shells increases down a group. Therefore, the distance between the electrons and the nuecleus reduces the attraction between them which results to a lower ionization energy.

Equations for Ionization Energy

  • 1st ionization energy of Q: Q(g)+energy→Q+(g)+e−Q(g) + energy \rightarrow Q^+(g) + e^-

  • 2nd ionization energy of Q: Q+(g)+energy→Q2+(g)+e−Q^+(g) + energy \rightarrow Q^{2+}(g) + e^-

Exceptions to the Trend

Magnesium (Mg) and Aluminum (Al)

The decrease in ionization energy from Mg to Al is because the 3p subshell in Al is higher in energy and further from the nucleus than the 3s subshell in Mg. Thus, the outermost 3p electron of Al is more easily removed.

Phosphorus (P) and Sulfur (S)

In phosphorus, the electron to be removed is from a 3p orbital with 3 electrons. In sulfur, the 3p subshell contains 4 electrons, including one pair. Removing one of the paired electrons in sulfur is easier due to the reduction of electron-electron repulsion.

Difference Between Periods

There is a large difference in the first ionization energy between elements in different periods because of the increase in the number of electron shells, which increases the distance between the electrons and the nucleus, resulting in a weaker attraction and requiring less energy to remove an electron.

First Electron Affinity Trend

Definition

F(g)+e−→F−(g)+energyF(g) + e^- \rightarrow F^-(g) + energy

Electron Affinity

The atom with the highest electron affinity at the end of Period (eg Period 2-4) is the halogen atom. Noble gases have zero electron affinity.