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Chapter 8: Problem 16

What do we mean when we say that two ions or an atom and an ion are isoelectronic?

Short Answer

Expert verified
Two ions or an atom and an ion are isoelectronic when they possess the same number of electrons or identical electronic structure.

Step by step solution

01

Definition of Isoelectronic Species

Isoelectronic refers to two different elements (atoms or ions) possessing the same number of electrons or same electronic structure. For example, N3-, O2-, F-, Ne, and Na+ are all isoelectronic because they all have the same number of electrons (10).
02

Factors to Remember

1. Isoelectronic elements may have different atomic numbers. 2. It is important to remember that the number of protons may vary, leading to differences in atomic number. 3. In the case of ions, the element will have gained or lost electrons, to reach a stable configuration, similar to another element or ion.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Electrons
Electrons are subatomic particles with a negative charge, which orbit around the nucleus of an atom. They play a crucial role in chemistry, as they are involved in forming chemical bonds between atoms. The number of electrons in an atom determines its electronic structure, which directly influences the atom's chemical properties.

When atoms or ions become isoelectronic, they share the same number of electrons. This similarity can lead to comparable chemical behaviors, as it means having identical electronic arrangements. For instance, the ions N3-, O2-, F-, and Ne all have 10 electrons, hence they exhibit similar stability and electronic characteristics.
Ions
Ions are atoms or molecules that have gained or lost electrons, resulting in a net charge. The process of gaining electrons will give a negative charge, forming an anion. Conversely, losing electrons results in a positive charge, creating a cation.

For example, when a neutral sodium atom loses one electron, it becomes Na+, a cation. Isoelectronic ions may belong to different elements, but they share the same electronic configuration. Understanding ions is important as their properties, like charge and size, influence chemical reactions and the formation of compounds.
Electronic structure
The electronic structure of an atom or ion is essentially the arrangement of electrons around the nucleus. This arrangement is dictated by quantum mechanics and the nature of electron shells and subshells. An atom's or ion's specific electronic structure determines how it interacts in chemical reactions.

Isoelectronic species have identical electronic structures, meaning their electrons are arranged in the same pattern. For example, the ions N3-, O2-, F-, and Ne, although different in terms of protons, each possess an electronic configuration of 1s虏 2s虏 2p鈦�. This configuration leads to similar chemical properties and reactivity.
Atomic number
The atomic number of an element is defined as the number of protons in the nucleus of its atoms. It uniquely identifies a chemical element. A change in the atomic number means a different element altogether.

Interestingly, isoelectronic species might have different atomic numbers because they belong to different elements. While these species have the same number of electrons, the difference in their atomic numbers means their nuclei are different in terms of protons.
  • For example, N3- has an atomic number of 7.
  • On the other hand, Na+ has an atomic number of 11.
This diversity in atomic numbers doesn't affect the shared trait of having the same electronic structure.

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Most popular questions from this chapter

Use the third period of the periodic table as an example to illustrate the change in first ionization energies of the elements as we move from left to right. Explain the trend.

State whether each of the following properties of the representative elements generally increases or decreases (a) from left to right across a period and (b) from top to bottom within a group: metallic character, atomic size, ionization energy, acidity of oxides.

Arrange the following in order of increasing first ionization energy: \(\mathrm{Na}, \mathrm{Cl}, \mathrm{Al}, \mathrm{S},\) and \(\mathrm{Cs}\)

Two atoms have the electron configurations \(1 s^{2} 2 s^{2} 2 p^{6}\) and \(1 s^{2} 2 s^{2} 2 p^{6} 3 s^{1} .\) The first ionization energy of one is \(2080 \mathrm{~kJ} / \mathrm{mol}\), and that of the other is \(496 \mathrm{~kJ} / \mathrm{mol} .\) Match each ionization energy with one of the given electron configurations. Justify your choice.

The only confirmed compound of radon is radon difluoride, \(\operatorname{RnF}_{2}\). One reason that it is difficult to study the chemistry of radon is that all isotopes of radon are radioactive so it is dangerous to handle the substance. Can you suggest another reason why there are so few known radon compounds? (Hint: Radioactive decays are exothermic processes.)
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