In chemistry, periodicity refers to repeating trends in element properties on the periodic table. Basically, what this means is if you drop down a row (period) on the table and move across it, elements follow the same trend as other periods. Periodicity reflects Periodic Law. Periodic Law states that chemical and physical properties of elements repeat in a predictable way when elements are arranged by increasing atomic number.
Why Periodicity Is Important
In essence, periodicity is the guiding principle behind the organization of the modern periodic table. Elements within a group (column) display similar characteristics. The rows in the periodic table (the periods) reflect the filling of electrons shells around the nucleus, so when a new row begins, the elements stack on top of each other with similar properties.
Because of recurring trends, you can predict the properties and behavior of an element, even if it’s new. Chemists can use periodicity to determine the likelihood of a chemical reaction occurring or chemical bonds forming. Early on, scientists used gaps in the periodic table to know where elements should be and what their properties would be.
Simple Periodicity Example
Because of periodicity, you can tell from the periodic table that both sodium and lithium are highly reactive metals, with an oxidation state of +1. Similarly, you know beryllium is less reactive than lithium, but still a metal.
Periodicity allows predictions for the behavior of elements that haven’t been synthesized in large enough amounts to study directly. Chemists can tell oganesson (element 118) will have some properties of the elements above it on the table (the noble gases). It probably won’t be as reactive as, for example, tennessine (element 117), which is a halogen.
What Are the Periodic Properties?
Several element properties display periodicity. The key recurring trends are:
- Electronegativity – Electronegativity is a measure of how readily an atom forms a chemical bond. Electronegativity increases moving left to right across a period and decreases moving down a group. Or, you could say electropositivity decreases moving left to right and increases moving down the periodic table.
- Atomic Radius – This is half the distance between the middle of two atoms just touching each other. Atomic radius decreases moving left to right across a period and increases moving down a group. Even though you’re adding more electrons moving across a period, atoms don’t get bigger because they don’t get additional electron shells. The increasing number of protons draws the electrons closer, shrinking the atom size. Moving down a group, new electron shells get added and atom size increases.
- Ionic Radius – Ionic radius is the distance between ions of the atoms. It follows the same trend as atomic radius. Although it might seem like increasing the number of protons and electrons in an atom would always increase its size, the atom size doesn’t increase until a new electron shell is added. Atom and ion sizes shrink moving across a period because the increasing positive charge of the nucleus pulls in the electron shell.
- Ionization Energy – Ionization energy is the energy required to remove one electron from an atom or ion. It is a predictor of reactivity and the ability to form chemical bonds. Ionization energy increases moving across a period and decreases moving down a group. There are some exceptions, mainly due to Hund’s rule and electron configuration.
- Electron Affinity – This is a measure of readily an atom accepts an electron. Electron affinity increases moving across a period and (mostly) decreases moving down a group. Nonmetals usually have higher electron affinities than metals. The noble gases are an exception to the trend since these elements have filled electron valence shells and electron affinity values approaching zero. However, the behavior of the noble gases is periodic. In other words, even though an element group might break a trend, the elements within the group display periodic properties.
- Metallic Character – Metallic character or metallicity describes properties of metals, such as luster, conductivity, and high melting/boiling points. Also, metals readily accept electrons from nonmetals to form ionic compounds. The most metallic element is francium (lower left-hand side of the periodic table), while the least metallic element is fluorine (upper righthand side of the table).
- Group Properties – Elements in a column belong to the same element group. Each group displays characteristic properties. For example, halogens tend to be highly reactive nonmetals with a -1 oxidation state (valency), while noble gases are nearly inert and exist as gases under standard conditions.
Summary of Periodicity Trends
The periodicity of these properties follows trends as you move across a row or period of the periodic table or down a column or group:
Moving Left → Right
- Ionization Energy Increases
- Electronegativity Increases
- Atomic Radius Decreases
- Metallic character decreases
Moving Top → Bottom
- Ionization Energy Decreases
- Electronegativity Decreases
- Atomic Radius Increases
- Metallic character increases
Discovery of Periodic Law
Scientists discovered periodicity in the 19th century. Lothar Meyer and Dmitri Mendeleev independently formulated Periodic Law in 1869. Chemists of this era arranged elements by increasing atomic weight, because the proton and atomic number had not yet been discovered. Even so, periodic tables of the day displayed periodicity. The reason for recurring trends wasn’t understood until the 20th century, which brought the description of electron shells.
References
- Allred, A. Louis (2014). Electronegativity. McGraw-Hill Education. ISBN 9780071422895.
- Mendeleev, D. I. (1958). Kedrov, K. M. (ed.). Периодический закон [The Periodic Law] (in Russian). Academy of Sciences of the USSR.
- Rennie, Richard; Law, Jonathan (2019). A Dictionary of Physics. Oxford University Press. ISBN 9780198821472.
- Sauders, Nigel (2015). “Who Invented The Periodic Table?”. Encyclopedia Britannica. ISBN 9781625133168.