NT

Topic 3 Periodicity: Everything You Need to Know for the IB Exam

IB Chemistry Topic 3: Periodicity

1. Arrangement of the Periodic Table

  • Groups and Periods:

    • Groups (columns) have the same number of valence electrons.

      • Group 1: 1 valence electron

      • Group 2: 2 valence electrons

    • Periods (rows) indicate the energy level of valence electrons.

      • Period 1: 1st energy level

      • Period 2: 2nd energy level

2. Blocks of the Periodic Table

  • S, P, D, and F blocks:

    • Group 1 & 2: S Block

    • Groups 3-12: D Block

    • Groups 13-18: P Block

    • Lanthanides and Actinides: F Block

  • Electron Configurations:

    • Elements in S block end in 'S'

    • Elements in D block end in 'D'

    • Elements in P block end in 'P'

    • Elements in F block end in 'F'

3. Trends in Periodicity

3.1 Effective Nuclear Charge

  • Definition: The strength of the nucleus in attracting valence electrons.

  • Trend:

    • Decreases down a group

    • Increases across a period

  • Example:

    • Lithium (Li) vs. Francium (Fr):

      • Li has high effective nuclear charge due to fewer inner electrons.

      • Fr has lower effective nuclear charge because of increased shielding from inner electrons.

3.2 Atomic Radius

  • Definition: The size of the atom.

  • Trend:

    • Increases down a group (adding electrons to higher energy levels increases size)

    • Decreases across a period (increased nuclear charge pulls electrons closer)

  • Example:

    • Hydrogen vs. Francium:

      • Hydrogen is much smaller due to fewer total electrons.

3.3 Ionic Radius

  • Definition: The size of ions formed when atoms lose or gain electrons.

  • Trends:

    • Positive ions (cations) are smaller than neutral atoms.

    • Negative ions (anions) are larger than neutral atoms.

  • Trends in Groups:

    • Ionic radius decreases from groups 1 to 14 for cations

    • Increases for anions up to group 17 then decreases.

4. Ionization Energy

  • Definition: The energy required to remove an electron from an atom.

  • Trend:

    • Increases across a period (metals to non-metals)

    • Decreases down a group (increased shielding leads to easier electron removal)

  • Example:

    • Lithium (Li) requires less energy to ionize than Francium (Fr) due to fewer shielding electrons.