Atomic Structure and Electron Behavior: Key Properties

Electron Motion: Bohr Model vs. Quantum Theory

The description of electron motion around the nucleus of an atom differs significantly between the Bohr model and modern quantum theory.

Bohr's Planetary Model

In Bohr's planetary model, the electron's position and velocity can be precisely determined at any given time, allowing for the prediction of its linear motion in a fixed orbit around the nucleus of the atom.

Modern Quantum Theory and Orbitals

Modern quantum theory introduces the concept of an orbital, where the electron in its motion around the nucleus can take any path randomly and therefore does not follow a predetermined trajectory as in a classical orbit. An orbital provides the probability of finding the electron at a certain distance from the nucleus. This information is obtained from the magnitude of the wave function squared (Ψ²), where Ψ is the wave function derived from the Schrödinger equation.

The orbital concept is perfectly in accord with the two great pillars of modern quantum theory:

• De Broglie Hypothesis: Describes the dual corpuscular and wave-like behavior of the electron.
• Heisenberg Uncertainty Principle: States the impossibility of simultaneously determining an electron's exact position and momentum.

Key Atomic Properties

Ionization Energy (IE)

Ionization energy (IE) is the minimum energy required to remove an electron from a gaseous atom or ion in its ground state.

• Across a Group: In a family group (vertical column) of the periodic table, ionization energy generally decreases as the atomic number (Z) increases. This is because as the atomic radius increases, the outermost electrons are further from the nucleus and experience weaker attraction, making them easier to remove.
• Across a Period: Within the same period (horizontal row), there is a general tendency for ionization energy to increase with increasing atomic number (Z). However, there are some small decreases for certain atoms, particularly those with full or half-full electron sublevels, which enjoy greater stability. For these more stable configurations, more energy is required to remove an electron from their valence shell because the nucleus exerts a stronger force of attraction on electrons closer to it, making the atom more stable.

Electron Affinity (EA)

Electron affinity (EA) is the energy change that occurs when an electron is added to a gaseous atom to form a negative ion (anion) in the gaseous state.

• Across a Period: Generally, electron affinity increases (becomes more negative, indicating a greater release of energy) across a period as the atomic number (Z) increases. This is because a higher nuclear charge leads to a stronger attraction for an additional electron, making the resulting anion more stable.
• Across a Group: Electron affinity tends to decrease down a group.

Electronegativity

Electronegativity is a measure of an atom's tendency to attract a shared pair of electrons towards itself in a chemical bond.

• Across a Period: Electronegativity generally increases across a period as the atomic number (Z) increases. This is due to the increasing nuclear charge and decreasing atomic radius, which results in a stronger attraction for bonding electrons.
• Across a Group: Electronegativity generally decreases down a group as the atomic radius increases and the outermost electrons are further from the nucleus.

Atomic Radius

Atomic radius is a measure of the size of an atom, typically defined as half the distance between the nuclei of two identical atoms bonded together.

• Across a Period: Atomic radius generally decreases across a period as the atomic number (Z) increases, due to the increasing nuclear charge pulling the electron shells closer.
• Across a Group: Atomic radius generally increases down a group as new electron shells are added.

Ionic Radius Considerations

• Cations (Positive Ions): For positive ions, such as Na⁺, the ionic radius is smaller than the neutral atom (Na). This occurs because the atom loses an electron from its outermost shell, and the remaining electrons are more strongly attracted by the nucleus, pulling them closer.
• Anions (Negative Ions): For negative ions, such as Cl^-, the ionic radius is larger than the neutral atom (Cl). This is due to the addition of an electron, which increases electron-electron repulsion within the electron cloud, causing it to expand.