Smallest and Largest Atomic Radiusįrancium has the largest atomic size on the periodic table, and helium has the smallest atomic size. The Trend on a GraphĪs shown in the graph below, the atomic radius is largest at the first element in each period, and it decreases down each period. As electron cloud sizes increase, so do atomic radii. This is because between each group, electrons occupy successively higher energy levels. Group Trendĭown a group, atomic radii increase. This is why the difference in atomic radii decreases down each period. One thing to note is that the effect of the attraction between the positively charged nucleus and the electrons is slightly countered by the repulsion of electrons as they are successively added. This increased positive charge attracts or pulls, the electrons in closer to the nucleus, decreasing the atomic radius. Down the period, however, the number of protons also increases. This is because while the number of electrons increases down the period, they only add to the same main energy level, and therefore do not expand the electron cloud. For example, ionization energy, electronegativity, and of course atomic radius which we will discuss now. There are many trends on the periodic table. Let’s break down the trend into its period and group trends. Atoms decrease in size across the period and increase in size down the group. Atomic Radius Trend on the Periodic TableĪtomic radii increase toward the bottom left corner of the periodic table, with Francium having the largest atomic radius. This is because, down a group, the principal quantum number (n) increases which results in an increase in the distance between the nucleus and valence electrons. On the other hand, the atomic radius generally increases down a group. Thus the atomic radius is measured as shown in the diagram below. This decreases the radius of the elements that go from left to right. This is because the borders of orbitals are quite fuzzy, and they also change under different conditions. While your initial thought may have been to measure the distance from the center of an atom’s nucleus to the edge of its electron cloud, this is inaccurate and not feasible. The atomic radius is measured as half the distance between two nuclei of the same atoms that are bonded together. Simply put, the factor of increasing n down a group is greater than the effect caused by the increasing Z eff, thus causing the atomic radius of atoms to increase down a group in the periodic table.Let’s discuss the definition of the atomic radius, also called atomic size, and the atomic radius trend on the periodic table. The principal quantum number, n, of electron orbitals that increases down a group and due to the quantum mechanical nature of electrons, the radius of these electron orbitals increases with increasing n, thus increasing in size as you go down the group. The effective nuclear charge, Z eff, increases down a group which draws electrons closer towards the nucleus, decreasing atomic radius. Thus we have a quantum mechanical picture whereby the principle quantum number of electron orbitals affects their radius.Įssentially we have two competing factors that affect the trend of atomic radii in question: This comparison can be further extended where the 3s orbitals are larger in radius than 2s and these are larger than 1s. The quantum mechanical description shows that the 1s orbital of a hydrogen atom is a lot smaller in radius than the 2s orbital of a lithium atom for example. The principal quantum number, n, of electron orbitals affects their size. ![]() However, what we have failed to consider is the size of the electron orbitals in question. This, as you quite rightly mention, should suggest that because Z eff increases down a group, a greater force of attraction would be experienced between the outer shell valence electrons and the nucleus, thus decreasing the atomic radius. You are right in thinking that the effective nuclear charge increases down as a group as the increasing nuclear charge has a greater effect on Z eff than the shielding effect of more inner electron shells. The overall effective nuclear charge experienced by the valence electron in question is Z eff. Just so that we are on the same page before I attempt to answer this question, the definition of effective nuclear charge we are considering Z eff=Z−σ, where Z is the nuclear charge (number of protons) and σ is the shielding factor of the inner shell electrons.
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