Answer: The atomic radius generally decreases across a period from left to right on the periodic table. This means that as you move across a period, the atoms become smaller. There are a few key factors that contribute to this trend: 1. Increasing nuclear charge: As you move across a period, the number of protons in the nucleus increases. This leads to a greater attractive force on the electrons, pulling them closer to the nucleus. As a result, the atomic radius decreases. 2. Constant shielding effect: Although the number of electrons also increases across a period, the shielding effect remains relatively constant. Shielding refers to the ability of inner electrons to shield the outer electrons from the attractive force of the nucleus. Since the shielding effect doesn't increase significantly, the electrons are not effectively shielded from the increasing nuclear charge. This further contributes to the decrease in atomic radius. 3. Energy level difference: As you move across a period, the electrons are added to the same energy level, or shell. However, the number of protons in the nucleus is increasing, which results in a stronger attractive force on the electrons. This causes the electrons to be pulled closer to the nucleus, resulting in a smaller atomic radius. To illustrate this trend, let's consider the example of the second period (from lithium to neon). As you move from left to right, the atomic number (number of protons) increases from 3 to 10. This increase in nuclear charge causes the atomic radius to decrease. For example, the atomic radius of lithium is larger than that of beryllium, and the atomic radius of beryllium is larger than that of boron. Overall, the atomic radius decreases across a period due to the increasing nuclear charge and constant shielding effect. It's important to note that there are exceptions to this trend in some cases, such as the noble gases, which have relatively large atomic radii compared to other elements in the same period.