Trends in the Periodic Table

• Trends are defined as patterns observed within the periodic table. Patterns may occur:

  • Top to bottom within a group.

  • Left to right within a period.

Key Properties to Investigate

1. Atomic Radius

2. Ionization Energy

3. Electronegativity

Atomic Radius

• Definition: The atomic radius refers to the size of an atom, measured as the distance from the nucleus to the furthest electron in the electron cloud.

Group 1 Example

• Hydrogen: Only 1 electron shell occupied (1 electron).

• Lithium: 3 electrons, requiring:

  • 2 electrons in the 1st shell.

  • 1 electron in the 2nd shell (total of 2 shells).

• Trend: As you move down the group (e.g., from hydrogen to cesium), there is an increase in the number of electron shells, leading to a larger atomic radius.

• Memory Aid: Think of a snowman - larger at the bottom, representing increased atomic radius.

Period Trend

• Moving from left to right in a period (e.g., period 2 from lithium to neon):

  • Atomic radius decreases.

  • Despite adding electrons, the increased nuclear charge (more protons) pulls electrons closer to the nucleus, resulting in a smaller radius.

  • Example: Neon is smaller than lithium in period 2.

• Overall Trends:

  • Top to Bottom: Atomic radius increases.

  • Left to Right: Atomic radius decreases.

  • Comparison: Helium (smallest) to francium (largest).

Ionization Energy

• Definition: Ionization energy is the energy required to remove an electron from an atom. It can be thought of as the "cost" of removing an electron.

Atom Size and Ionization Energy

• Small Atom: Electrons are close to the nucleus, generally requiring higher energy to remove.

• Large Atom: Valence electrons are farther away, making them easier to remove, hence lower ionization energy is needed.

• Shielding Effect: Interior electrons reduce the nuclear charge effect felt by valence electrons, making it easier for larger atoms to lose electrons.

Period and Group Trends

• Across a Period: Ionization energy increases (for example, from lithium to neon) because smaller atoms hold their electrons more tightly.

• Down a Group: Ionization energy decreases due to larger atomic size and increased shielding.

Electronegativity

• Definition: Electronegativity measures an atom's attraction to electrons. It is measured on a scale from 0 to 4.

  • 4 indicates a strong attraction for electrons (e.g., fluorine).

  • 0 indicates little to no attraction for electrons (e.g., elements on the left side).

Periodic Table Example

• Fluorine vs. Lithium:

  • Fluorine, with 7 valence electrons, seeks to gain one more, thus possesses a high electronegativity.

  • Lithium, with only 1 valence electron, prefers to lose it, leading to low electronegativity.

Trends in Electronegativity

• Across a Period: Electronegativity increases (from lithium to fluorine).

• Down a Group: Electronegativity decreases because the atoms become larger, making it harder for the nucleus to attract nearby electrons.

Summary of Trends

• Key Takeaway: Understanding these trends is crucial, but understanding why they exist (primarily related to atomic size and the distance of valence electrons from the nucleus) is even more important.