Summary of "Atomic Radius - Basic Introduction - Periodic Table Trends, Chemistry"

Summary of “Atomic Radius - Basic Introduction - Periodic Table Trends, Chemistry”

Main Ideas and Concepts

Definition of Atomic Radius

Atomic radius is the distance from the nucleus (center) of an atom to the outer edge of its electron cloud. It is often visualized as the radius of a circle representing the atom.

Calculation of Atomic Radius

The atomic radius of an element can be calculated by taking half the distance between the nuclei of two bonded atoms of the same element. Example: In a bromine molecule, the nuclei are 228 pm apart, so the atomic radius is 114 pm.

Periodic Table Trends in Atomic Radius

Across a Period (Left to Right):
- Atomic radius decreases.
- Reason: Increasing nuclear charge (more protons) pulls electrons closer, reducing size despite electrons being added to the same energy level.
- Effective nuclear charge (Z_eff) increases because electrons added to the same shell do not shield each other well.
- Example values from Period 2:
  - Lithium (150 pm) > Beryllium (113 pm) > Boron (88 pm) > Carbon (77 pm) > Nitrogen (70 pm) > Oxygen (66 pm) > Fluorine (64 pm)
  - Neon is a slight exception, slightly larger than fluorine (69 pm).
Down a Group (Top to Bottom):
- Atomic radius increases.
- Reason: Addition of energy levels (shells) increases orbital size, outweighing increased nuclear charge.
- Effective nuclear charge on valence electrons remains roughly constant within a group because increased protons are offset by increased inner shell electrons (shielding).
- Example: Sodium (186 pm) > Lithium (150 pm); Potassium (227 pm) > Sodium (186 pm).

Effective Nuclear Charge (Z_eff)

Calculated as: Number of protons (Z) minus number of inner (shielding) electrons.
Determines the net positive charge felt by valence electrons.
Higher Z_eff pulls electrons closer, reducing atomic radius.

Shielding Effect

Inner electrons repel outer electrons, reducing the effective nuclear charge felt by valence electrons.

Methodology / Instructions for Understanding Atomic Radius Trends

Calculating Atomic Radius from Diatomic Molecule:
1. Measure the distance between nuclei of two identical atoms in a molecule.
2. Divide this distance by two to get the atomic radius.
Determining Trends Across Periods and Groups:
- Across a period: Identify increasing proton number and constant energy level; expect decreasing radius.
- Down a group: Identify increasing energy levels; expect increasing radius.
Comparing Atomic Sizes Between Elements:
- Use position on periodic table:
  - Left to right → size decreases
  - Top to bottom → size increases
- Use known atomic radius values for confirmation.
Ranking Elements by Atomic Radius:
- Place elements on the periodic table.
- Use trends (down and left = bigger; up and right = smaller) to order them.
- Example ranking (smallest to largest): Neon < Chlorine < Selenium < Tin < Cesium.

Example Problem Solutions

Calcium vs. Magnesium: Calcium is larger (197 pm vs. 116 pm) because it is below magnesium in the same group.
Silicon vs. Phosphorus: Silicon is larger (117 pm vs. 110 pm) because it is to the left of phosphorus in the same period.
Strontium vs. Sulfur: Strontium is larger (215 pm vs. 104 pm) because it is below and to the left of sulfur.
Magnesium vs. Iodine: Magnesium is slightly larger (160 pm vs. 153 pm) because it is further left and above iodine.

Speakers / Sources Featured

Primary Speaker: Unnamed instructor or chemistry educator explaining atomic radius concepts and periodic trends.
No other speakers or external sources are explicitly mentioned.