Periodic Trend: Ionic Radius (Simplified): Videos & Practice Problems

Ionic radius refers to the distance between an ion's nucleus and its outer shell. Cations, or positive ions, are smaller than their neutral forms due to the loss of electrons, resulting in a decrease in ionic radius. Conversely, anions, or negative ions, are larger than their neutral counterparts because gaining electrons increases the ionic radius. Thus, the trend shows that more electrons lead to a larger ionic radius, emphasizing the importance of electron count over periodic table position in determining ionic size.

Periodic Trend: Ionic Radius (Simplified) Concept 1

Understanding ionic radius is crucial in the study of chemistry, particularly when examining the behavior of ions. The ionic radius is defined as the distance from the nucleus of an ion to its outermost electron shell. Unlike many periodic trends that can be inferred from the periodic table, the ionic radius trend is determined by the number of electrons present in the ion.

As a general rule, the ionic radius increases with the addition of electrons. This means that anions, which are negatively charged ions formed by gaining electrons, tend to have a larger radius than their neutral parent atoms. For example, when oxygen (which has the electron configuration of 1s² 2s² 2p⁴) gains two electrons to become an oxide ion, its outer shell now contains eight electrons. This increase in electron count results in a larger ionic radius compared to the neutral oxygen atom.

Conversely, cations, which are positively charged ions formed by losing electrons, exhibit a smaller ionic radius than their neutral counterparts. Taking lithium as an example, the lithium atom has the electron configuration of 1s² 2s¹. When it loses one electron to form a lithium ion, it only retains one electron shell, leading to a decrease in size. Thus, the trend indicates that losing electrons results in a smaller ionic radius.

In summary, the ionic radius is influenced by the charge of the ion: cations are smaller than their neutral forms due to the loss of electrons, while anions are larger due to the gain of electrons. Therefore, when assessing ionic radius, focus on the total number of electrons in the ion rather than relying on the periodic table.

Periodic Trend: Ionic Radius (Simplified) Example 1

In understanding ionic radii, it's essential to recognize the differences between cations and anions compared to their neutral atoms. Cations, which are positively charged ions formed by the loss of electrons, exhibit a smaller ionic radius than their neutral counterparts. For example, calcium (Ca), which is in group 2A and has a charge of +2, will have a reduced ionic radius compared to the neutral calcium atom due to the loss of electrons.

Conversely, anions are negatively charged ions that result from the gain of electrons, leading to an increase in ionic radius. For instance, the sulfide ion (S²⁻) is larger than the neutral sulfur atom because the addition of electrons expands the electron cloud, making the ionic radius larger.

When examining specific ions, the cadmium ion (Cd²⁺) is also a cation and, like other cations, it has a smaller ionic radius than the neutral cadmium atom. Therefore, the statement claiming that the cadmium ion has a larger ionic radius than the cadmium atom is false.

In summary, the key takeaways are that cations are smaller than their neutral atoms due to electron loss, while anions are larger than their neutral atoms due to electron gain. This fundamental concept is crucial for understanding the behavior of ions in various chemical contexts.

Which of the following statements is false?
a) The phosphide ion has a smaller ionic radius than the phosphorus atom.
b) The phosphide ion has a larger ionic radius than the phosphorus atom.
c) The phosphide ion's ionic radius cannot be compared to the phosphorus atom.
d) The phosphide ion has an equal ionic radius than the phosphorus atom.

• A. A and C
• B. B and C
• C. A, C, and D
• D. B, C, and D
• E. All of the statements are false.

Answer: (C)

Here’s what students ask on this topic:

What is the ionic radius and how does it differ between cations and anions?

The ionic radius is the distance between an ion's nucleus and its outer shell of electrons. For cations, which are positive ions, the ionic radius is smaller than their neutral forms because they lose electrons, resulting in fewer electron shells or a reduced electron cloud. Conversely, anions, which are negative ions, have a larger ionic radius than their neutral forms because they gain electrons, increasing the electron cloud and repulsion among electrons. Thus, cations are smaller, and anions are larger compared to their neutral counterparts.

Why do cations have a smaller ionic radius compared to their neutral atoms?

Cations have a smaller ionic radius compared to their neutral atoms because they lose one or more electrons. This loss reduces the number of electron shells or decreases the electron cloud's size. For example, a neutral lithium atom (Li) has the electron configuration 1s² 2s¹. When it loses one electron to become a Li⁺ ion, it loses its entire second shell, resulting in a smaller ionic radius. The reduced electron-electron repulsion also contributes to the smaller size of cations.

How does the ionic radius change when an atom becomes an anion?

When an atom becomes an anion, its ionic radius increases. This is because the atom gains one or more electrons, which adds to the electron cloud and increases electron-electron repulsion. For instance, a neutral oxygen atom (O) has the electron configuration 1s² 2s² 2p⁴. When it gains two electrons to become an O^2- ion, the added electrons increase the size of the electron cloud, resulting in a larger ionic radius. Thus, gaining electrons leads to a larger ionic radius for anions.

Why is the periodic table not used to determine the trend in ionic radius?

The periodic table is not used to determine the trend in ionic radius because the ionic radius depends on the number of electrons in the ion rather than its position on the periodic table. The trend shows that more electrons lead to a larger ionic radius, and fewer electrons result in a smaller ionic radius. Therefore, to determine the ionic radius, one must consider whether the ion is a cation or anion and the total number of electrons it has, rather than its placement on the periodic table.

How does the number of electrons affect the ionic radius?

The number of electrons directly affects the ionic radius. When an ion gains electrons (becoming an anion), the increased electron-electron repulsion causes the electron cloud to expand, resulting in a larger ionic radius. Conversely, when an ion loses electrons (becoming a cation), the reduced electron-electron repulsion and fewer electron shells cause the electron cloud to contract, resulting in a smaller ionic radius. Thus, more electrons lead to a larger ionic radius, and fewer electrons lead to a smaller ionic radius.