Trends in the Periodic Table

The reason for the development of the periodic table was to arrange the elements in a way that had meaning. Thus, specific trends are apparent in the table. The major trends associated with the periodic table are atomic radius, nuclear charge, electronegativity, electron affinity, and ionization energy.

Atomic radius is half the distance between the nuclei of two identical atoms. The edge of the atom is difficult to define because the position of any electron at a given time cannot be precisely known. Atomic radius is therefore described in some different ways: Covalent radius is the half the distance between the nuclei of two identical atoms in a covalent bond. Van der Waals radius is half the distance between two identical nuclei that are not bonded but are as close as possible.

Atomic radius decreases from left to right across the periodic table and increases from top to bottom down the table. Atomic radius decreases from left to right because, within a period, the number of electrons increases within a shell while the number of protons increases in the nucleus. More protons and electrons mean the electromagnetic attraction between them also increases, thereby pulling the electron shell closer to the nucleus and decreasing the atomic radius. Within a group, however, electrons are added to higher valence shells. A higher valence shell is farther away from the nucleus than existing shells. These electrons at the higher valence shell are farther away from the nucleus, increasing atomic radius. Nuclear charge (Z) is the total charge of all protons within the nucleus. It is therefore equal to atomic number. Nuclear charge increases from left to right across the periodic table and from top to bottom down the periodic table. Effective nuclear charge (Z_eff) is the net positive charge of an atom when electron shielding is considered. Electron shielding is the decrease in attraction from the nucleus experienced by valence electrons due to the presence of inner electrons between the nucleus and the valence electrons. Effective nuclear charge is calculated by subtracting the number of nonvalence electrons (S) from the nuclear charge (Z). For example, neon has 10 protons and 10 electrons. Two of these electrons are nonvalence electrons. ${\rm {Z}}_{\rm{eff}}=10-2=8$ for neon. Like nuclear charge, effective nuclear charge increases from left to right across the periodic table and from top to bottom down the periodic table.

Electronegativity is the tendency of an atom to attract electrons toward itself when forming bonds. It is a qualitative rather than quantitative attribute; that is, it cannot be measured numerically. In the 1930s American chemist Linus Pauling devised a scale that ranks elements according to electronegativity, which is still in use today.

Most elements form bonds with other elements to satisfy the octet rule: they attempt to fill eight electrons in their valence shell. Elements on the left side of the periodic table have fewer than four valence electrons, so they tend to lose electrons when forming bonds, while those on the right side of the periodic table have more than four valence electrons, so they tend to gain electrons when forming bonds. Thus, electronegativity increases from left to right across the periodic table, with a few exceptions. The noble gases have full valence shells, so they do not commonly form bonds with other elements at all and thus have very low electronegativity. The transition metals form vertical groups 3 to 12. Between groups 3 and 11, electronegativity of transition metals generally increases. Group 12 transition metals have lower electronegativity compared to group 11. Elements of group 13, the first group of p-block elements, have electronegativities very similar to those of group 12 transition metals. Lanthanides and actinides are a general exception to electronegativity. They have much more complicated chemistry and do not show any trends. Lanthanides and actinides are considered to have no electronegativity.

Electronegativity decreases from top to bottom down the periodic table. This happens because the atoms of elements lower in a group have a larger atomic radius and thus an increased distance between the valence electrons and the nucleus. As a result, atoms with a larger atomic radius exert a weaker force on nearby atoms. They are therefore less likely to form bonds with other elements compared to atoms with a smaller atomic radius. Electron affinity is the change in energy that occurs when an atom of a neutral gas gains an electron. It is a quantitative attribute, and thus can be measured numerically. Electron affinity indicates the ability of an atom to accept an electron. It increases from left to right across the periodic table and decreases from top to bottom down the periodic table. These trends are due to atomic radius: a smaller atomic radius is associated with a higher electron affinity, because less energy is required to add electrons to a valence shell closer to the nucleus compared to one that is farther away. However, electron affinities are difficult to measure accurately, so the trends are a generalization.

Ionization energy is the amount of energy required to remove an electron from a gaseous atom or ion. It can be thought of as the opposite of electron affinity. It is a quantitative attribute, expressed either in joules (J) or electron volts (eV). Each electron in an atom has its own ionization energy; however, the energy needed to remove the first electron is usually referenced when discussing any element. The first ionization energy is the energy required to remove an electron from each atom in one mole of atoms in a gaseous state, resulting in one mole of ions with a 1+ charge. The second ionization energy is the energy required to remove a second electron from each ion in one mole of 1+ ions in a gaseous state, resulting in a mole of ions with a 2+ charge. Note that second ionization energy is not the total energy required to remove two electrons; it is the difference between the energy required to remove the first and second electrons. There are additional ionization energies for each electron in the atom.

Ionization energy depends on the charge of the nucleus, the atomic radius, and the electron configuration of the atom. Elements on the left side of the periodic table have their valence shells less than half full, so they readily give up an electron. In contrast, elements on the right side of the periodic table have their valence shells more than half full, so they do not give up electrons easily. This explains why metals are more likely to form cations and nonmetals are more likely to form anions. In addition, elements with valence shells farther away from the nucleus, that is, those lower on the periodic table, give up those valence electrons more easily than elements with valence shells closer to the nucleus. This is due to not only the distance between the nucleus and these electrons, but also the effect of electron shielding of electrons in nonvalence shells. However, after the first electron is given up, the remaining electrons are held more firmly because the charge on the nucleus does not change. This is also true for each additional electron given up.

Thus, ionization energy increases from left to right across the periodic table and decreases from top to bottom down the periodic table. When all of the trends on the periodic table are viewed together, relationships among the properties of elements become clear. Electron affinity, effective nuclear charge, ionization energy, and electronegativity generally increase from left to right across the periodic table because of the way in which electrons fill the valence shell across a period. Effective nuclear charge and atomic radius generally increase down a group because of the increase in distance between valence electrons and the nucleus.