Atomic radius decreases
Moving from left to right across the periodic table, the atomic radius generally decreases. This trend is observed due to the following factors:

Effective nuclear charge: The effective nuclear charge experienced by the valence electrons increases as we move from left to right across a period. This is because the number of protons in the nucleus increases, while the shielding effect remains relatively constant. The increased positive charge in the nucleus attracts the electrons more strongly, resulting in a decrease in atomic radius.

Electron-electron repulsion: As we move across a period, the number of electrons in the valence shell increases. These electrons repel each other, causing the electron cloud to spread out. This repulsion counteracts the attractive force of the nucleus, leading to a larger atomic radius.

Trends in atomic radius:
- Atomic radius generally decreases from left to right across a period.
- Atomic radius tends to increase from top to bottom within a group.

Periodic table visualization:
To visualize this trend, let's consider the periodic table. Starting from the left side, we observe that the elements in the first period (hydrogen and helium) have the smallest atomic radii. Moving to the second period, the atomic radii gradually increase from left to right. This trend continues as we move across subsequent periods.

Exceptions:
There are a few exceptions to the general trend of decreasing atomic radius. For example, the atomic radius of helium is smaller than that of hydrogen, despite being in the same period. This is due to the presence of a full 1s orbital in helium, which causes greater electron-electron repulsion and a smaller atomic radius.

In conclusion, as you move from left to right across the periodic table, the atomic radius generally decreases. This is primarily due to the increasing effective nuclear charge and the electron-electron repulsion experienced by the valence electrons.