1. Periodic Trends

2. Valence Electron Configurations
They can be determined from the group number of the element.
▪ Example: In Group 1, all the elements have an ns1 valence configuration, where n is the period number.
They are slightly less systematic for d-block metals. Here, the filling of (n-1)d orbitals occurs. In Period 4, they are as follows:

3. Valence Electron Configurations

4. Atomic Radii
Trend: Across a period, atomic radius decreases.
▪ Explanation: Across a row, the principal quantum number (n) remains essentially
the same, whereas the effective nuclear charge (Zeff) is increasing as more and more protons are added to the nucleus. This larger value of Zeff exerts a stronger pull on
the electrons, shrinking the atomic radius.

5. ▪ Trend: Down a group, atomic radius increases.
Atomic Radii
▪ Trend: Down a group, atomic radius increases.
▪ Explanation: Down a group, Zeff varies only slightly, while n is increasing. As n increases, the outer-most electron will be added to an orbital having a larger average radial probability. Thus, the atomic radius increases down a column.
Lanthanide Contraction
▪ Atomic radii in the 5d series of the d block are very similar to their congeners in the 4d series even though the atoms have a greater number of electrons.
▪ For example, the radii of Mo and W in Group 6 are 140 and 141 pm, respectively.
This reduction of radius below that expected is called the lanthanide contraction.
▪ Explanation: It occurs due to the presence of 4f electrons in the lanthanoids. The poor shielding properties of f electrons results in a higher effective nuclear charge than expected on the basis of a simple extrapolation from other atoms.

6. Radius Related Trends
Melting Point of Some Ionic Solids
▪ The melting point of each crystalline solid depends on the strength of the electrostatic forces holding the ions together.
▪ The strength of the cation–anion attractive force is expected to have an inverse
dependence on the ionic radius of the halide, according to Coulomb’s law of electrostatic attraction.
▪ Thus, the melting points increase as the ionic radius of the anion decreases.

7. Radius Related Trends
Lattice Energy of Some Ionic Solids
Ionic Solid
Lattice Energy (kJ/mol)
Halide Radius (pm)
MgF2
-2957
119
MgCl2
-2526
167
MgBr2
-2440
182
MgI2
-2327
206
▪ The lattice energy of an ionic solid is defined as the amount of energy gained when the gaseous ions are brought together from infinity to form the ionic solid.
▪ Because the electrostatic attraction between gaseous ions follows a similar relationship to Coulomb’s law, the magnitude of the lattice energy should show an inverse dependence on the distance between the ions in the crystalline solid.
▪ Thus, the lattice energy increases as the ionic radius of the anion decreases.

8. Radius Related Trends
Hydration Enthalpies of Selected Ions
▪ The hydration energy is the change in Gibbs energy when an ion or molecule is transferred from a vacuum (or the gas phase) to a solvent.
▪ The smaller the ionic radius, the higher the charge density (Z/r) of the ion will be. Ions with large charge densities can form stronger ion-dipole forces with the polar water molecule.
▪ Thus, the enthalpy of hydration is expected to become more negative (more favorable from a thermodynamic standpoint) as the ionic radius decreases.