The atomic radius generally decreases from left to right across a period in the periodic table, despite the increasing atomic number. This trend is primarily due to the increasing effective nuclear charge or the attraction between the positively charged protons in the nucleus and the negatively charged electrons in the electron cloud.
As you move from left to right across a period, the number of protons in the nucleus increases, adding positive charge to the nucleus. However, the number of electrons in the inner energy levels or shells remains relatively constant within a period. This results in an increase in the effective nuclear charge, which pulls the electrons closer to the nucleus.
The increased positive charge from the nucleus exerts a stronger pull on the electrons in the same energy level, causing the electron cloud to shrink. The outermost electrons experience this stronger pull and are more strongly attracted to the nucleus, resulting in a decrease in atomic radius.
Additionally, as you move across a period, electrons are added to the same energy level, occupying the same electron shell. However, since electrons are negatively charged, they repel each other. This electron-electron repulsion further compresses the electron cloud, leading to a decrease in atomic radius.
It's important to note that there can be exceptions to this trend in certain cases due to the influence of other factors, such as electron-electron repulsion and electron shielding effects. However, the general trend of decreasing atomic radius from left to right across a period holds for most elements in the periodic table.