Trends Across a Period and Down a Group
All periodic trends can be explained by two key factors:
- Effective Nuclear Charge ($Z_\text{eff}$):
- Higher $Z_\text{eff}$ pulls electrons closer, decreasing atomic and ionic radii, and increasing ionization energy and electronegativity.
- Shielding Effect:
- Inner electrons shield outer electrons from the nucleus, reducing $Z_\text{eff}$ and making it easier to remove or add electrons.
Effective nuclear charge
Effective nuclear charge is the net positive charge experienced by valence electrons.
Atomic Radius: The Size of an Atom
Atomic radius
The atomic radius is the distance from the nucleus of an atom to the outermost electron.
While we cannot measure this directly (since electrons exist in a cloud), we approximate it based on the distances between two bonded atoms.
Trend Across a Period
- As you move across a period from left to right, the atomic radius decreases.
- Each successive element adds a proton to the nucleus and an electron to the same energy level.
- This increases the effective nuclear charge ($Z_\text{eff}$), pulling the electrons closer to the nucleus.
- Consider sodium (Na) and chlorine (Cl), both in Period 3.
- Sodium has an atomic radius of 186 pm, while chlorine’s radius is only 99 pm.
- The additional protons in chlorine's nucleus exert a stronger pull on its electrons, shrinking the atomic radius.
Trend Down a Group
- As you move down a group, the atomic radius increases.
- This is because each successive element adds a new energy level (or shell), increasing the distance between the nucleus and the outermost electrons.
- Although the nuclear charge also increases, the additional inner shells provide shielding, reducing the pull on the valence electrons.
- Think of the nucleus as a magnet and the electrons as metal balls.
- Adding more layers of insulation (energy levels) weakens the magnet’s pull on the outermost balls.
Ionic Radius: The Size of Ions
Ionic radius
Ionic radius is the measure of an ion's size, defined as the distance from its nucleus to the outermost electron, influenced by its charge and electron configuration.
Cations (positively charged ions) are smaller than their parent atoms, while anions (negatively charged ions) are larger.
Trend Across a Period
- For cations, the ionic radius decreases across a period due to increasing nuclear charge.
- For anions, the same trend holds, but anions are always larger than cations within the same period.
Consider $Na^{+}$ (cation) and $Cl^{-}$ (anion).
- $Na^{+}$ has lost its outermost electron, reducing electron-electron repulsion and shrinking its radius to 102 pm.
- $Cl^{-}$, however, has gained an electron, increasing repulsion and expanding its radius to 181 pm.
Trend Down a Group
Ionic radius increases down a group for both cations and anions, following the same reasoning as atomic radius (additional energy levels and shielding).
Ionic Radius for Isoelectronic Ions
What Are Isoelectronic Ions?
Isoelectronic ions
Isoelectronic ions are ions that have the same number of electrons but different nuclear charges.
Despite having the same electron configuration, their ionic radii vary due to differences in nuclear attraction on the electron cloud.
Trend Across Isoelectronic Series
- In a series of isoelectronic ions (e.g., $O^{2-}$, $F^-$, $Na^+$, $Mg^{2+}$), the ionic radius decreases as nuclear charge increases.
- This is because a greater positive charge pulls the electrons closer to the nucleus, reducing the radius.
- $O^{2-}$ has the largest radius due to its lower nuclear charge.
- $Mg^{2+}$ has the smallest radius as its higher nuclear charge pulls electrons more tightly.
