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Chapter 9: Problem 19

All the isoelectronic species illustrated in the text had the electron configurations of noble gases. Can two ions be isoelectronic without having noble-gas electron configurations? Explain.

Short Answer

Expert verified
Yes, two ions can be isoelectronic without having noble-gas electron configurations. This is possible when two ions have the same total number of electrons but these electrons do not occupy a completely filled electron shell configuration, as is the case with noble gasses.

Step by step solution

01

Understanding Isoelectronic Terms

Isoelectronic refers to different atoms, ions, or molecules that have the same total number of electrons. The challenge here is to think about whether these atoms or ions must necessarily have a noble-gas electron configuration.
02

Analyzing the Electron Configurations

The electron configurations of noble gases are characterized by completely filled electron shells. An atom becomes an ion by losing or gaining electrons to achieve a 'stable' configuration, usually similar to the nearest noble gas. However, this process is not limited to imitating noble gas configurations.
03

Conclusion

So, two ions can indeed be isoelectronic without having noble gas electron configurations. This becomes possible when two ions have the same total number of electrons, but those electrons do not fully fill their respective shells (which would be the case for a noble gas configuration). In other words, they may both lack electrons to become truly stable (resembling a noble-gas configuration), but they could still have the same total number of electrons.

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Key Concepts

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

Electron Configuration
Electron configuration is the distribution of electrons in an atom or ion. It tells us how electrons are arranged around the nucleus in different shells and subshells. The order in which electrons fill various orbitals follows the rules based on increasing energy levels and principles like the Aufbau principle, Hund's rule, and the Pauli exclusion principle.
For example, the electron configuration for hydrogen is written as \(1s^1\), indicating one electron in the first shell's s subshell. Understanding this concept is crucial as it helps us predict chemical properties and reactivity of elements.
For ions, electron configurations change as they gain or lose electrons. An ion's electron configuration can mirror that of a noble gas when it reaches a more stable form, reflecting the completion of an energy shell.
Noble Gases
Noble gases belong to Group 18 of the periodic table and are known for their characteristic full electron shells, making them chemically inert under most conditions. Their general electron configuration is \(ns^2np^6\), which indicates completely filled valence orbitals, contributing to their lack of reactivity.
Noble gases are the standard for electronic stability, which many other elements strive to achieve by forming ions. For example, sodium (Na) loses one electron to resemble neon (Ne), thus, achieving a noble gas electron configuration. This desired stable state makes noble gases essential in understanding the concepts of ion formation and chemical bonding.
  • Helium (He): \(1s^2\)
  • Neon (Ne): \(1s^2 2s^2 2p^6\)
  • Argon (Ar): \(1s^2 2s^2 2p^6 3s^2 3p^6\)
These configurations display full shell configurations, which are energetically favorable.
Ion Stability
Ion stability relates to how energetically favorable an ion's electron configuration is, often aiming for a state similar to that of a noble gas. Stability in ions generally happens when they have a complete set of electrons filling their outer shell, reducing the energy of the system.
However, ions can achieve stability even when their configurations do not match those of noble gases. For example, the sulfide ion (S虏鈦�) and chloride ion (Cl鈦�) are isoelectronic, each having 18 electrons, but only chloride matches a noble gas configuration (argon).
Ionic stability is also influenced by factors such as:
  • Charge Balance: The overall energy reduction in a system when oppositely charged ions balance each other.
  • Coordination Number: The number of ions bonded to a central ion.
  • Crystal Lattice: Arrangement in solid structures that provide stability.
Understanding ion stability helps us comprehend how atoms interact, form compounds, and stabilize in different environments.

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Most popular questions from this chapter

The most negative electron affinity of the following elements is that of (a) \(\mathrm{Br} ;\) (b) \(\mathrm{Sn} ;\) (c) \(\mathrm{Ba} ;\) (d) Li.

How much energy, in joules, must be absorbed to convert to \(\mathrm{Na}^{+}\) all the atoms present in \(1.00 \mathrm{mg}\) of gaseous Na? The first ionization energy of Na is \(495.8 \mathrm{kJ} / \mathrm{mol}.\)

In multielectron atoms many of the periodic trends can be explained in terms of \(Z_{\text {eff }}\) Consider the following statements and discuss whether or not the statement is true or false. (a) Electrons in a \(p\) orbital are more effective than electrons in the \(s\) orbitals in shielding other electrons from the nuclear charge. (b) \(\mathrm{Z}_{\text {eff }}\) for an electron in an \(s\) orbital is lower than that for an electron in a \(p\) orbital in the same shell. (c) \(Z_{\text {eff }}\) is usually less than \(Z.\) (d) Electrons in orbitals having \(\ell=1\) penetrate better than those with \(\ell=2.\) (e) \(\mathrm{Z}_{\text {eff }}\) for the orbitals of the elements \(\mathrm{Na}(3 s)\) \(\mathrm{Mg}(3 s), \mathrm{Al}(3 p), \mathrm{P}(3 p),\) and \(\mathrm{S}(3 p)\) are in the order \(Z_{\text {eff }}(\mathrm{Na})<\mathrm{Z}_{\text {eff }}(\mathrm{Mg})>\mathrm{Z}_{\text {eff }}(\mathrm{Al})<\mathrm{Z}_{\text {eff }}(\mathrm{P})>\mathrm{Z}_{\text {eff }}(\mathrm{S}).\)

Briefly describe each of the following ideas or phenomena: (a) the periodic law; (b) ionization energy; (c) electron affinity; (d) paramagnetism.

Explain the important distinctions between each pair of terms: (a) actinide and lanthanide element; (b) covalent and metallic radius; (c) atomic number and effective nuclear charge; (d) ionization energy and electron affinity; (e) paramagnetic and diamagnetic.
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