In Chemistry, periodic trends are the tendencies of certain elemental characteristics to increase or decrease as one progresses along a row or column of the periodic table of elements.

Atomic radius

The atomic radius is the distance from the atomic nucleus to the outermost stable electron orbital in an atom that is at equilibrium. The atomic radius tends to decrease as one progresses across a period because the effective nuclear charge increases, thereby attracting the orbiting electrons and lessening the radius. The atomic radius usually increases while going down a group due to the addition of a new energy level (shell). However, diagonally, the number of protons has a larger effect than the sizeable radius. For example, lithium (145 pm) has a smaller atomic radius than magnesium (150 pm). Atomic radii decrease left to right across a period, and also increase top to bottom down a group.

Ionization potential

The ionization potential (or the ionization energy) is the minimum energy required to remove one electron from each atom in a mole of atoms in the gaseous state. The first ionization energy is the energy required to remove one, the nth ionization energy is the energy required to remove the atom's nth electron, after the (n−1) electrons before it have been removed. Trend-wise, ionization potentials tend to increase while one progresses across a period because the greater number of protons (higher nuclear charge) attract the orbiting electrons more strongly, thereby increasing the energy required to remove one of the electrons.

As one progresses down a group on the periodic table, the ionization energy will likely decrease, since the valence electrons are farther away from the nucleus and experience a weaker attraction to the nucleus' positive charge. There will be an increase of ionization energy from left to right of a given period and a decrease from top to bottom. As a rule, it requires far less energy to remove an outer-shell electron than an inner-shell electron. As a result the ionization energies for a given element will increase steadily within a given shell, and when starting on the next shell down will show a drastic jump in ionization energy. Simply put, the lower the principal quantum number, the higher the ionization energy for the electrons within that shell. The exceptions are the elements in the boron and oxygen family which require slightly less energy than the general trend.

Electron affinity

The electron affinity of an atom can be described either as the energy gained by an atom when an electron is added to it, or conversely as the energy required to detach an electron from a singly-charged anion. The sign of the electron affinity can be quite confusing, as atoms that become more stable with the addition of an electron (and so are considered to have a higher electron affinity) show a decrease in potential energy; i.e. the energy gained by the atom appears to be negative. For atoms that become less stable upon gaining an electron, potential energy increases, which implies that the atom gains energy. In such a case, the atom's electron affinity value is positive. Consequently, atoms with a more negative electron affinity value are considered to have a lower electron affinity (they are more receptive to gaining electrons), and vice versa. However in the reverse scenario where electron affinity is defined as the energy required to detach an electron from an anion, the energy value obtained will be of the same magnitude but have the opposite sign. This is because those atoms with a high electron affinity are less inclined to give up an electron, and so take more energy to remove the electron from the atom. In this case, the atom with the more positive energy value has the higher electron affinity. As one progresses from left to right across a period, the electron affinity will increase, due to the larger attraction from the nucleus, and the atom "wanting" the electron more as it reaches maximum stability. Down a group, the electron affinity decreases because of a large increase in the atomic radius, electron-electron repulsion and the shielding effect of inner electrons against the valence electrons of the atom.

Electronegativity

Electronegativity is a measure of the ability of an atom or molecule to attract pairs of electrons in the context of a chemical bond. The type of bond formed is largely determined by the difference in electronegativity between the atoms involved, using the Pauling scale. Trend-wise, as one moves from left to right across a period in the periodic table, the electronegativity increases due to the stronger attraction that the atoms obtain as the nuclear charge increases. Moving down a group, the electronegativity decreases due to the longer distance between the nucleus and the valence electron shell, thereby decreasing the attraction, making the atom have less of an attraction for electrons or protons.

Metallic character

Metallic character refers to the chemical properties associated with elements classified as metals. These properties, which arise from the element's ability to lose electrons, are: thedisplacement of hydrogen from dilute acids; the formation of basic oxides; the formation of ionic chlorides; and their reduction reaction, as in the thermite process. As one moves across a period from left to right in the periodic table, the metallic character decreases, as the atoms are more likely to gain electrons to fill their valence shell rather than to lose them to remove the shell. Down a group, the metallic character increases, due to the lesser attraction from the nucleus to the valence electrons (in turn due to the atomic radius), thereby allowing easier loss of the electrons or protons.

Atomic emissions

In a recent paper published in August 2009 by the Journal of Light & Visual Environment (vol 32, No 2), Kitsinelis, et al. describe trends found in the periodic table of the elements for the most persistent neutral atomic emission lines. According to the paper, the general trend was observed for the wavelengths of the strongest lines of the neutral elements, which usually happen to be the result of transitions terminating at ground state, with s-block elements emitting with highest intensity mainly in the infrared and visible part of the spectrum, the p-block elements mainly in the ultraviolet, the d-block in the visible and ultraviolet and the f-block in the visible. It can also be said in the present case that the wavelengths of the strongest lines decrease going up a column (with decreasing atomic radius) and towards the right of a row (with increasing number of charges in the nucleus) something that was clearly noticeable for the s- and p- block elements while across the d- and f- block the values did not show any trend but remained in between those of the other blocks. The general trend observed in the decrease of the wavelengths, mainly for the s- and p- block, was similar to the one found for the electronegativity values or the atomic radii reversed.