Chemistry - Periodicity

Atomic radius

• ← Increases

  • Outer electrons are in the same shell, have more protons in the nucleus, and experience the same amount of shielding. This results in a stronger attraction between the nucleus and the outer shell electrons, pulling them closer to the nucleus.

• ↓ increases 

• increased e- shells - further from nucleus 

• e- shielding - lessens nuclear attraction

Ionisation energy

• → increases

• ↓ increases

  • The atomic radius is larger, the smaller the nuclear charge experienced by outer electrons, and therefore the lower the ionisation energy (IE).

• Nuclear charge is higher, the higher the nuclear charge, and therefore the higher the IE.

Electron shielding is greater with more inner shells, resulting in a lower nuclear charge.

Groups 2 to 3:

• Electrons are lost from the p orbital in Group 3.

• Electrons are lost from the s orbital in Group 2.

The p orbital is higher in energy than the s orbital, so it is easier to lose an electron from the p orbital.

Groups 5 to 6:

• Group 6 elements lose electrons from the p orbital with two electrons (p4).

• Group 5 elements lose electrons from the p orbital with one electron (p3).

Extra electron-electron repulsions make it easier to lose an electron from the p4 orbital than from the p3 orbital.

Melting point across groups 2 + 3

• The trend increases from the metals to a peak at the group 14 elements (carbon and silicon), before sharply dropping for the non-metals.

E-/neg trend

• → increases

  • As the number of protons in the nucleus increases, the positive nuclear charge becomes stronger. This increased charge pulls the bonding electrons closer to the nucleus, resulting in a decrease in the atomic radius. 

• ↓ decreases

  • As you move down a group, atoms have more electron shells. This increases the distance between the nucleus and the outer electrons, providing more inner-shell electrons to “shield” them from the nucleus’s positive charge. This weakens the attraction for bonding electrons. 

Electron affinity meaning

the energy needed for an atom to form a 1- ion

1st Ionisation energy meaning

the energy needed for an atom to form a 1+ ion

E-/Neg meaning

the tendency of an atom to attract a shared pair of electrons in a chemical bond

Electron affinity trend

• → GENERALLY increases

  • As you move from left to right across a period, the number of protons (and the nuclear charge) increases. This stronger attraction pulls the electrons, including a new one, more strongly towards the nucleus. Consequently, more energy is released when an electron is added, resulting in a higher (more exothermic) electron affinity.

• ↓ GENERALLY decreases

  • • Trend in Group: Atomic radius increases down a group, weakening the nucleus’s attractive force and resulting in less energy released when an electron is added.

  • • Exception to the Trend: Second-period elements in a group can sometimes have lower electron affinity than third-period elements due to increased electron-electron repulsion.

  • • Factors Affecting Electron Affinity: Nuclear charge, atomic radius, and electron shielding all influence the strength of an element’s electron affinity.