However, there are several practical ways to define the radius of atoms and, thus, to determine their relative sizes that give roughly similar values. The quantum mechanical picture makes it difficult to establish a definite size of an atom. These properties vary periodically as the electronic structure of the elements changes. An understanding of the electronic structure of the elements allows us to examine some of the properties that govern their chemical behavior. Going down the elements in a group, the number of electrons in the valence shell remains constant, but the principal quantum number increases by one each time. Going across a period from left to right, a proton is added to the nucleus and an electron to the valence shell with each successive element. This similarity occurs because the members of a group have the same number and distribution of electrons in their valence shells. The elements in groups of the periodic table exhibit similar chemical behavior. This is because the number of electrons in the outermost principal energy level is nearly constant. The radii of most transition elements, however, stay roughly constant across each row. Thus, as the effective nuclear charge increases steadily, the shielding of outer electrons becomes less, and this leads to a decrease in atomic radii. Across the period, the nuclear charge increases while the number of inner shell electrons remains constant. Thus, the effective nuclear charge, the charge felt by an outer electron is lesser than the actual nuclear charge.Įlectrons in the same valence shell do not shield one another very effectively. In any multi-electron atom, the inner shell electrons partially shield the outer shell electrons from the pull of the nucleus. Recall the concept of an effective nuclear charge. The decreasing atomic radii across a period can be explained by the effective nuclear charge. This trend is demonstrated by the entire periodic table.įurther, the plot reveals that the atomic radius is maximum for each alkali metal and falls to a minimum with each noble gas across the period. Thus, as outer electrons get farther from the nucleus, the atomic radius increases down the group.įor example, moving down group 1, the atomic radius increases from lithium to cesium. Moving down a group, the principal quantum number, n, increases by one for each element. The trend in atomic radii for main group elements down the columns is depicted here. The periodic table depicts variations in covalent radii that are often called atomic radii, which are influenced by two factors the number of principal energy levels of valence electrons, and the effective nuclear charge. In nonmetal, diatomic molecules, the radius is described as one-half of the distance between the centers of bonded atoms. In metals, the radius is described for atoms in their crystal structure as one-half of the distance between the centers of two neighboring atoms. Nonbonding atomic radius, or van der Waals radius of an atom, is one-half of the distance between adjacent nuclei in the atomic solid.Ĭonversely, a bonding atomic radius, or covalent radius, distinguishes between metals and nonmetals. So how is the atomic size defined, and what influences it?Īn atomic radius can be described in two ways. However, orbitals do not describe a confined space, but rather the statistical probability of where an electron can be found. An atom’s size is dictated by the electrons or their orbitals.
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