Periodic trends

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This article refers to periodic trends in chemistry, for other periodic trends see trend (disambiguation).
The Periodic Trends

When comparing the properties of the chemical elements, recurring ('periodic') trends are apparent. This led to the creation of the periodic table as a useful way to display the elements and rationalize their behavior. When laid out in tabular form, many trends in properties can be observed to increase or decrease as one progresses along a row or column.

These period trends can be explained by theories of atomic structure. The elements are laid out in order of increasing atomic number, which represents increasing positive charge in the atomic nucleus. Negative electrons are arranged in orbitals around the nucleus; recurring properties are due to recurring configurations of these electrons.

These periodic trends are distributed among 3 different properties, namely physical properties, chemical properties and on the basis of chemical reactivity. In chemical properties, it is classified on the basis of two; i.e. periodicity of valence or oxidation states, anomalous properties of second period elements. In this article the periodic trends with respect to their physical properties are briefly discussed.

Atomic radius[edit]

Main article: 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 radii tend to decrease across a period from left to right. The atomic radius usually increases while going down a group due to the addition of a new energy level (shell). However, atomic radii tend to increase diagonally, since the number of electrons has a larger effect than the sizeable nucleus. For example, lithium (145 picometer) has a smaller atomic radius than magnesium (150 picometer).

Atomic radius can be further specified as:

  • Covalent radius: half the distance between two atoms of a diatomic compound, singly bonded.
  • Van der Waals radius: half the distance between the nuclei of atoms of different molecules in a lattice of covalent molecules.
  • Metallic radius: half the distance between two adjacent nuclei of atoms in a metallic lattice.
  • Ionic radius: half the distance between two nuclei

Ionization energy[edit]

Main article: Ionization energy

The ionization potential is the minimum amount of 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 two, the 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 energy tends 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. Ionization energy and ionization potentials are completely different. The potential is an intensive property and it is measured by "volt"; whereas the energy is an extensive property expressed by "eV" or "kJ/mole".

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's 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[edit]

Main article: 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.[1] Consequently, atoms with a more negative electron affinity value are considered to have a higher 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.

Although it may seem that Fluorine should have the greatest electron affinity, the small size of fluorine generates enough repulsion that Chlorine has the greatest electron affinity.


Main article: 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 in 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.

However, in the group 13 elements electronegativity increases from aluminium to thallium, and in group 14 electronegativity of lead is lower than that of tin.

Valence electrons[edit]

Main article: Valence electrons

Valence electrons are the electrons in the outermost electron shell of an isolated atom of an element. Sometimes, it is also regarded as the basis of Modern Periodic Table. In a period, the number of valence electrons increases (mostly for light metal/elements) as we move from left to right side. However, in a group this periodic trend is constant, that is the number of valence electrons remains the same.

However, this periodic trend is sparsely followed for heavier elements (elements with atomic number greater than 20), especially for lanthanide and actinide series.

It is also important to consider the core electrons when speaking about the valence electrons.[why?]

Metallic and non-metallic properties[edit]

Metallic properties increase down groups as decreasing attraction between the nuclei and the outermost electrons causes the outermost electrons to be loosely bound and thus able to conduct heat and electricity. Across the period, increasing attraction between the nuclei and the outermost electrons causes metallic character to decrease.

Non-metallic property increases across a period and decreases down the group due to the same reason.

See also[edit]

Further reading[edit]


  1. ^ SparkNotes Editors (27 November 2015). "SparkNote on Atomic Structure". Retrieved 29 November 2015.