Periodicity is the repeating pattern of element properties as atomic number increases. It explains why elements in the same group often react in similar ways and why properties change predictably across a row. Trends such as atomic radius, ionization energy, electronegativity, and metallic character help chemists predict bonding, reactivity, and compound formation.
These patterns make the periodic table more than a list of elements, because it becomes a map of chemical behavior.
The main cause of periodicity is the arrangement of electrons in shells and subshells. Across a period, protons are added while electrons enter the same main energy level, so the effective nuclear charge generally increases and pulls electrons closer. Down a group, new electron shells are added, which increases distance and shielding from the nucleus.
The balance between nuclear attraction, shielding, and electron shell structure produces the major periodic trends.
Key Facts
- Atomic radius generally decreases from left to right across a period and increases from top to bottom down a group.
- Ionization energy generally increases from left to right across a period and decreases from top to bottom down a group.
- Electronegativity generally increases from left to right across a period and decreases from top to bottom down a group.
- Metallic character generally decreases from left to right across a period and increases from top to bottom down a group.
- Effective nuclear charge can be estimated as Z_eff = Z - S, where Z is nuclear charge and S is shielding.
- First ionization energy is the energy for M(g) -> M+(g) + e-.
Vocabulary
- Atomic radius
- Atomic radius is a measure of the size of an atom, usually based on the distance between nuclei of bonded atoms.
- Ionization energy
- Ionization energy is the energy required to remove an electron from a gaseous atom or ion.
- Electronegativity
- Electronegativity is the ability of an atom in a chemical bond to attract shared electrons.
- Effective nuclear charge
- Effective nuclear charge is the net positive attraction felt by a valence electron after shielding by inner electrons is considered.
- Shielding
- Shielding is the reduction of nuclear attraction on outer electrons caused by repulsion from inner electrons.
Common Mistakes to Avoid
- Saying atomic radius increases across a period, because added protons actually pull electrons in the same shell closer when shielding changes only slightly.
- Treating all periodic trends as perfectly smooth, because electron subshell stability and electron pairing can create small exceptions.
- Confusing ionization energy with electronegativity, because ionization energy removes an electron from an isolated atom while electronegativity describes attraction for shared electrons in a bond.
- Ignoring shielding down a group, because added electron shells reduce the nucleus's pull on valence electrons even though nuclear charge increases.
Practice Questions
- 1 Arrange Li, Be, B, and C in order of increasing atomic radius, and briefly state the trend used.
- 2 Using Z_eff = Z - S, estimate the effective nuclear charge for a sodium valence electron if Z = 11 and S = 10. Then estimate it for a chlorine valence electron if Z = 17 and S = 10.
- 3 Magnesium is more metallic than sulfur even though both are in period 3. Explain this using valence electrons, effective nuclear charge, and the direction of metallic character across a period.