If we discuss the size of an atom, the distance we're referring to will depend on the element and how it exists naturally. For example, noble gases exist as individual atoms. If we cool a noble gas down to a low enough temperature that it forms a solid, the atomic radius would refer to the distance between the center of two adjacent noble gas atoms. This is called the non-bonding atomic radius or van der Waals radius. If we were to talk about the atomic size of bromine, however, we would refer to a covalent radius. Bromine exists as a diatomic molecule. So the atomic radius of bromine is the distance between the two nuclei in a bromine molecule. If we look at the trends in atomic radii for the elements, atomic radii tend to increase going down a column of the periodic table but decrease going to the right across a row of the periodic table. For example, as we move from lithium to neon, the atomic radius tends to decrease across the row. However, if we were to compare lithium and sodium, atomic radius increases going down a column. So atomic radius tends to increase going to the left and down in the periodic table. There are exceptions, but this is the general trend. Our atomic radius increases as we move down a column because our principal quantum number, n, increases. Remember that we said as n increases, electrons are held farther from the nucleus at a higher energy. So going down a column, n increases, and atomic radius increases. Atomic radius decreases going to the right across a period or row of the periodic table because effective nuclear charge increases. Effective nuclear charge is the net charge a valence electron "feels" from the nucleus, taking into account the repulsions from other electrons. For example, we can consider lithium, which has three electrons. The valence electron in the 2s subshell is more likely to be found farther from the nucleus than the two core electrons in the 1s subshell. The valence electron feels an attraction to the nucleus, which has a three plus charge, but it also experiences a repulsion from the two core electrons according to their combined 2charge. So the valence electron is shielded from the nucleus by these two core electrons and it only experiences a net pull toward the nucleus of about 1+ instead of 3+. We can compare this to beryllium. Beryllium has four electrons and four protons. Beryllium has two valence electrons in the 2s subshell, and these two valence electrons are shielded from the nucleus by the two core electrons in the 1s subshell. But the valence electrons do not effectively shield each other. So core electrons effectively shield valence electrons, but valence electrons do not effectively shield other valence electrons. The result is that the effective nuclear charge for beryllium is 2+. We'll find out as we move to the right across a row of the periodic table that effective nuclear charge generally increases. This increase in effective nuclear charge is what decreases atomic size moving to the right across a row of the periodic table because electrons feel a greater pull towards the nucleus.