Question

Explain briefly in relates to the atomic properties and trends of atomic size

Answer 1

Atomic size, or atomic radius, refers to the distance from the nucleus of an atom to the boundary of its surrounding cloud of electrons. Various trends influence atomic size across the periodic table, primarily determined by the number of electron shells and the effective nuclear charge.

1. **Trends in Periods (Left to Right)**:
- As you move across a period from left to right, atomic size generally decreases. This is because additional protons are added to the nucleus, increasing the nuclear charge. The increased positive charge attracts the electrons more strongly, pulling them closer to the nucleus and resulting in a smaller atomic radius.

2. **Trends in Groups (Top to Bottom)**:
- As you move down a group in the periodic table, atomic size increases. This is due to the addition of electron shells (energy levels). Each new electron shell is located farther from the nucleus, leading to an overall increase in atomic size, despite the increase in nuclear charge.

3. **Effective Nuclear Charge (Z_eff)**:
- The effective nuclear charge, which considers both the total positive charge of the nucleus and the shielding effect of inner-shell electrons, plays a crucial role in determining atomic size. A higher effective nuclear charge increases the attraction between the nucleus and the outer electrons, leading to smaller atomic radii.

4. **Shielding Effect**:
- Inner electrons can shield outer electrons from the full effect of the nuclear charge, which is why atomic size increases down a group. The more electron shells present, the greater the shielding effect, which reduces the effective nuclear charge felt by the outermost electrons.

In summary, atomic size decreases across periods due to increased nuclear charge without significant additional shielding while increases down groups due to the addition of electron shells, which outweighs the effects of nuclear charge.