Before delving into the definition of ionic radius, it might be necessary to brush up the basics and discuss the formation of an ion. An ion is formed when an atom loses or gains electrons from its valence electronic orbital to attain a stable electronic configuration. The loss of electrons results in a positively charged cation, whereas gain forms a negatively charged species called an anion. Hence, the ionic radius can be defined as the radial distance measured between the centre of the nucleus of an ion to the outermost electronic orbital where the electron cloud is still under the influence of the positive electric field of the nucleus.
There will be an obvious change in atomic properties due to loss or gain of electrons. Let us consider atomic size at first. The atomic size of a cation is smaller than the parent atom as the attractive force exerted by the positively charged nucleus on the electrons in the outer electron shell is unbalanced and greater than that of the electrons (They are less in number and the atom is not electrically neutral anymore). Similarly, for an anion, the repulsive force existent among the electrons is dominant over the nuclear attractive force (As the electrons are more in number), and as a result, anions are larger in size compared to parent atoms. An evident conclusion that can be drawn from here is that anionic radius> cationic radius.
Radius of Sodium atom = 227 pm.
Radius of Sodium cation= 186 pm.
Isoelectronic species are those having the same number of electrons in total. For instance, F- and Na+, both have 10 electrons. However, their atomic sizes differ due to the difference in effective nuclear charge. It follows the above trend, and hence, F- has a larger ionic radius compared to Na+.
Ionic radius trends refer to a predictable pattern change in the ionic radius of elements on moving down or across in the modern periodic table. Important analytical conclusions about chemical reactivity of elements can be drawn from this. In this context, it may be interesting to note an alternative definition for ionic radius, which states that it is half the distance between two ions hardly in contact with one another, placed in a crystal lattice.
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If we move down along a group in the periodic table, then the number of electronic shells keeps increasing by one with every group. As a result, the ionic radius of elements also increases on moving down a group.
While moving across any period in the periodic table, the nature of elements gradually changes from metallic to non-metallic. The number of electrons in the outermost shell keeps increasing until the last element (noble gas) is reached. Hence the effective nuclear charge increases on moving across a period and the ionization potential (The energy required to remove the loosely bound outermost electron from an isolated gaseous atom) increases. This implies that it becomes difficult to lose electrons as we move across a period. So, the elements in the left typically form cations while those in the later periods towards the right form anions. Thus the ionic radius initially decreases and later increases, followed by another decrease. A maximum radius is obtained in between for the anion with the maximum negative charge. The following table may be used for reference and better understanding of the trend.
The ionic radius trend can be observed to decrease, with increasing positive charge and, to increase with increasing negative charge.
For ions having the same or closely similar charges, the ionic radii decrease slowly with an increase in atomic number across the period for transition elements positioned in Groups 3-12 of the modern periodic table. The reason behind this behavioral trend of ionic radius can be attributed to the increase in effective nuclear charge on moving across the period.
Q: Arrange the following in order of increasing ionic radius:
N3-, Li+, Be2+, O2-
Following the periodic trend, cations have smaller radii than anions.
Hence, we may conclude that ionic radius is an important periodic property whose trends can be monitored and usefully put to application in predicting properties of elements. Interestingly all periodic properties are interlinked and, to some extent, interdependent.
Chemistry is all about exceptions, and the trends as mentioned earlier, although mostly generalized in the application, also show some exceptions.
The ionic radius of Oxide (O2-) is larger than Nitride (N3-).
The periodic table that we use in chemistry was first invented by Dmitri Mendeleev in the year 1869.
It was in 1886 when Antoine Becquerel first discovered the concept of radioactivity.
Q: What are the Periodic Trends in Chemistry?
Ans: Periodic trends are patterns of change in properties (both chemical and physical) observed when we move across or down in the modern periodic table. The most common trends include atomic radius, ionic radius, electron affinity, electronegativity, ionization potential, etc. The root cause behind these trends is the change in the atomic structure of elements within their periods. The observable pattern helps us predict properties and derive conclusive ideas about chemical reactivity shown by different elements. These laws enable the chemical elements to be organized in the periodic table based on their atomic structures and properties. The unknown properties of newly discovered elements can also be predicted using these trends.
Q: Is ionic Radius the Same as Atomic Radius?
Ans: No, the atoms are electrically neutral substances, while ions have electrons less or more than the parent atom in order to attain stability. The size of ions is either greater or lesser than the parent atom. The trend in atomic radius across the periodic table also varies from the trend in ionic radius. The atomic radius decreases from left to right across a period because the number of shells remains constant, but protonic charge keeps increasing, causing greater attractive pull on the outermost shell electrons. The ionic radius trend across a period varies accordingly, as mentioned above in the article.