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Periodic trends refers to the specific patterns that are present in the periodic table and it describes different properties of an element, including their size and their electronic properties. Some very common periodic trends include: atomic radius, electronegativity, ionization energy, electron affinity. Periodic trends arise from the arrangement of the elements in the periodic table; this provides chemists with an invaluable tool to quickly understand an element’s properties. These trends arise due to the same atomic structure of the elements within their group families or periods, and also due to the periodic nature of the elements.
Electronegativity Trends
Electronegativity can be well understood as a chemical property that describes an atom’s ability to attract and bind with electrons. Electronegativity represents a qualitative property, there is not any standard method for calculating the electronegativity. Moreover, the common scale for calculating electronegativity is the Pauling scale; it was named after the chemist Linus Pauling. The Pauling scale assigned numbers that were basically dimensionless due to the qualitative nature of electronegativity. Electronegativity values of each element can be easily calculated on certain periodic tables. An example is provided below.
Electronegativity is used to measure an atom’s tendency to attract and form bonds with electrons. This property exists mainly due to the electronic configuration of atoms. Majority of atoms follow the octet rule (they usually have the valence, or outer, shell composed of 8 electrons). As the elements present on the left side of the periodic table possess less than a half-full valence shell, the energy needed to gain electrons is generally higher as compared to the energy needed to lose electrons. Because of which those elements present on the left side of the periodic table generally lose electrons while forming bonds. Conversely, elements present on the right side of the periodic table are more energy-efficient in gaining electrons to form a complete valence shell of 8 electrons. The nature of electronegativity is discussed below:
As one moves from left to right within a period of elements, electronegativity increases.
Moving from top to bottom down a group, electronegativity decreases.
An important exception to the above rules includes the noble gases, lanthanides, and actinides. Noble gases, lanthanides, and actinides do not possess electronegativity values.
For the transition metals, however they have electronegativity values, there is still a little variation among them across the period and up and down a group.
Ionization Energy Trends
Ionization energy refers to the energy needed to remove an electron from a neutral atom in its gaseous phase. Theoretically, ionization energy is the opposite of electronegativity. The lower the ionization energy the more readily the atom becomes a cation. Thus, the higher this energy is, the more unlikely the atom becomes a cation. Mainly elements present on the right side of the periodic table possess a higher ionization energy as their valence shell is nearly filled. Elements present on the left side of the periodic table possess low ionization energies because of their willingness to lose electrons and become cations. Therefore, ionization energy keeps on increasing from left to right on the periodic table.
Some elements possess various different ionization energies; these varying energies are known as the first ionization energy, the second ionization energy, third ionization energy and so on. The first ionization energy refers to the energy needed to remove the outermost, or highest, energy electron, the second ionization energy represents the energy required to remove any subsequent high-energy electron from a gaseous cation, etc.
Electron Affinity Trends
Electron affinity refers to the ability of an atom to accept an electron. Besides electronegativity, electron affinity represents a quantitative measurement of the energy change that occurs when an electron is added to a neutral gaseous atom. The more negative the electron affinity value, the higher an atom’s affinity for electrons. Electron affinity keeps on increasing from left to right within a period. This is resulted by a decrease in atomic radius. From top to bottom within a group the electron affinity decreases . This is resulted by an increase in the atomic radius.
Atomic Radius Trends
The atomic radius refers to one-half the distance between the nuclei of two atoms (similar to a radius is half the diameter of a circle). Although, this concept is complex by the fact that not all atoms are generally bound together in the same way. Some are bound by covalent bonds in molecules, few of them are attracted to each other in ionic crystals, and others are held in metallic crystals. Nonetheless, it is possible for a vast majority of elements to form covalent molecules in which two like atoms are bound together by a single covalent bond. The covalent radii of these molecules are generally referred to as atomic radii. This distance can be measured in picometers. Atomic radius patterns can be easily observed throughout the periodic table. Atomic radius decreases upon moving from left to right within a period. This is resulted by the increase in the number of protons and electrons across a period. On moving from top to bottom within a group the atomic radius increases. This is caused by electron shielding.
Conclusion
The arrangement of the periodic table helps us to understand certain trends among the atoms. The vertical columns (i.e. groups) of the periodic table are arranged so that all its elements possess the same number of valence electrons. All elements lying within a certain group therefore share similar properties. Here we come to an end of this topic. We hope that you were able to grasp a clear concept of the trends in the periodic table.
