Periodic trends are changes in the properties of chemical elements along the periodic table of elements. Major periodic trends include electronegativity, ionization energy, electron affinity, atomic radius, ionic radius, metallic character, and chemical reactivity.
These trends reflect the changes in atomic properties along the periods (horizontal rows) and groups (vertical columns) in the periodic table. They give a qualitative assessment of the properties of each element. However, not all properties strictly follow these trends, for example, the ionization energies in group 3, the electron affinities in group 17 or the densities in group 1 elements (alkali metals).
History
The periodic trends are based on the Periodic Law, which states that if the chemical elements are listed in order of increasing atomic number, many of their properties go through cyclical changes, with elements of similar properties recurring at intervals.[1] This principle was discovered by Russian chemist Dmitri Mendeleev in 1871 after a number of investigations by scientists in the 19th century. Mendeleev also proposed a periodic system of elements that was based not only on atomic weights but also on the chemical and physical properties of the elements and their compounds.[2] In 1913, Henry Moseley determined that periodicity depends on the atomic number rather than atomic weight. Lothar Meyer presented his table several months after Mendeleev, but opposed Mendeleev's Periodic Law. This empirical periodicity was later explained by the periodic recurrence of similar electronic configurations in the outer shells of the elements.
Atomic radius
The atomic radius is the distance from the atomic nucleus to the outermost electron orbital in an atom. Moving from left to right on the periodic table, one proton is added to the nucleus and one electron is added to the same shell as the former. The addition of the extra proton to the nucleus results in a stronger attraction force of the nucleus on the electrons. The additional electron was added to the same shell and therefore does not contribute largely to an increasing atomic radius. The atomic radius usually increases while going down a group due to the addition of a new electron shell.
Atomic radius is also occasionally defined in other ways; the definition used depends on the application. Below is a list of common alternative definitions.
- 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 of elements of an ionic compound.
- Atomic radius: the distance between the clear air and radiation poisoning.
Ionization energy
The ionization energy is the minimum amount of energy required to remove an electron from an isolated atom. The first ionization energy is the energy required to remove the first electron, the energy required to remove the second electron is the second ionization energy, etc. Ionization energy tends to increase while one progresses across a period, because the higher atomic number causes a stronger nuclear attraction force, thereby increasing the energy required to remove an electron.
As one progresses down a group on the periodic table, the ionization energy decreases since the valence electrons are farther away from the nucleus and experience a weaker attraction to the nucleus's positive charge. This effect is due to the core-electrons shielding the valence electrons from the positive charge of the nucleus. As a result, the ionization energies for a given element will increase steadily within a given shell, and increase sharply when starting on the next shell down. 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 released by an atom when an electron is added to it, conversely as the energy required to detach an electron from a singly charged anion.[3] 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. In such a case, the atom’s electron affinity is positive. 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 is negative.[4] 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 a 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 (Cl) has the greatest electron affinity.
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.[5] 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 an increase in the distance between the nucleus and the valence electron shell, thereby decreasing the atom's attraction to electrons.
However, in the group (iii) elements, electronegativity increases from aluminium to thallium.
The element having highest electronegativity is Fluorine.
Valence electrons
Valence electrons are the electrons in the outermost electron shell of an element. In a period, the number of valence electrons increases as we move from left to right. However, in a group this periodic trend is constant, that is the number of valence electrons remains the same.
Valence
Valence, or valency, is a measure of an element's capacity to combine with others when forming molecules or chemical compounds.
Traveling across a period, valency will first increase and then decrease. There is no change going down a group.
However, this periodic trend is sparsely followed for heavier elements (elements with atomic numbers greater than 20), especially for the lanthanide and actinide series.
The greater the number of core electrons, the greater the shielding of electrons from the core charge of the nucleus. For this reason, ionization energy is lower for elements lower down in a group, and polarizability is higher for elements lower down in a group. The valency does not change going down a group, since the bonding behavior is not affected by the core electrons. However, non-bonding interactions such as those just cited are affected by core electrons.
Metallic and non-metallic properties
Metallic properties generally increase down groups, as decreasing attraction between the nuclei and outermost electrons cause these electrons to be more loosely bound and thus able to conduct heat and electricity. Across each period, from left to right, the increasing attraction between the nuclei and the outermost electrons causes the metallic character to decrease.
Conversely, non-metallic character generally decreases down groups and increases across a period.
Most metals are lustrous (when freshly fractured, polished, or prepared), ductile, malleable, and sonorous, while most nonmetals are not.
See also
References
- ^ Sister, Harry H. (1963). Electronic structure, properties, and the periodic law. New York: Reinhold publishing corporation.
The physical and chemical properties of elements are periodic functions of the charges on their atomic nuclei i.e. their atomic numbers.
- ^ Sauders, Nigel (2015). Who Invented The Periodic Table?. Encyclopedia Britannica. pp. 26–29. ISBN 9781625133168.
- ^ Rennie, Richard; Law, Jonathan (2019). A Dictionary of Physics. Oxford University Press. ISBN 9780198821472.
- ^ "Atomic Structure". SparkNotes.com. 27 November 2015. Retrieved 2021-06-07.
- ^ Allred, A. Louis (2014). Electronegativity. McGraw-Hill Education. ISBN 9780071422895.