Question

Give the formula of a negative ion that would have the same number of electrons as each of the following positive ions. a. \(\mathrm{Na}^{+}\) c. \(\mathrm{Al}^{3+}\) b. \(\mathrm{Ca}^{2+}\) d. \(\mathrm{Rb}^{+}\)

Short answer

The negative ions with the same number of electrons as the given positive ions are: a. \(\mathrm{O}^{2-}\) b. \(\mathrm{S}^{2-}\) c. \(\mathrm{O}^{2-}\) d. \(\mathrm{Se}^{2-}\)

Step-by-step solution

Determine the atomic number and number of electrons for each positive ion

First, we need to recall that the atomic number gives the number of protons in the nucleus. In a neutral atom, the number of protons equals the number of electrons. For a positive ion, the number of electrons is reduced by the number of positive charges. So, we will first find the atomic numbers of the given elements and then calculate the number of electrons in their ions. a. \(\mathrm{Na}^{+}\): Sodium (Na) has an atomic number of 11. In the neutral atom, there are 11 electrons. In the \(\mathrm{Na}^{+}\) ion, there are 11 - 1 = 10 electrons. b. \(\mathrm{Ca}^{2+}\): Calcium (Ca) has an atomic number of 20. In the neutral atom, there are 20 electrons. In the \(\mathrm{Ca}^{2+}\) ion, there are 20 - 2 = 18 electrons. c. \(\mathrm{Al}^{3+}\): Aluminum (Al) has an atomic number of 13. In the neutral atom, there are 13 electrons. In the \(\mathrm{Al}^{3+}\) ion, there are 13 - 3 = 10 electrons. d. \(\mathrm{Rb}^{+}\): Rubidium (Rb) has an atomic number of 37. In the neutral atom, there are 37 electrons. In the \(\mathrm{Rb}^{+}\) ion, there are 37 - 1 = 36 electrons.

Find the negative ions with the same number of electrons

Now, we will find a negative ion that has the same number of electrons as each given positive ion by adding the atomic number and the number of charges. a. For 10 electrons, we look for an element with atomic number 8 and a -2 charge: This is oxygen, and the ion we are looking for is \(\mathrm{O}^{2-}\). b. For 18 electrons, we look for an element with atomic number 16 and a -2 charge: This is sulfur, and the ion we are looking for is \(\mathrm{S}^{2-}\). c. For 10 electrons, we already found that the ion with the same number of electrons is \(\mathrm{O}^{2-}\). d. For 36 electrons, we look for an element with atomic number 34 and a -2 charge: This is selenium, and the ion we are looking for is \(\mathrm{Se}^{2-}\). So the respective negative ions with the same number of electrons as the given positive ions are: a. \(\mathrm{O}^{2-}\) b. \(\mathrm{S}^{2-}\) c. \(\mathrm{O}^{2-}\) d. \(\mathrm{Se}^{2-}\)

Key concepts

Isoelectronic Species

Isoelectronic species are interesting because they have the same number of electrons, resulting in similar electron configurations, yet they are different elements or ions. This means they can behave in similar ways in terms of chemical properties.

For instance, the ion \(\mathrm{Na}^+\) has 10 electrons just like \(\mathrm{O}^{2-}\). These are examples of isoelectronic species, meaning they have identical electron shells, notwithstanding their different charges and atomic compositions.

This concept is essential in chemistry because it helps us predict and understand the behavior of various ions and atoms based on their electron configuration rather than their atomic number or nuclear composition. Identifying isoelectronic species helps in understanding and predicting bonding and reactivity.

Atomic Number

The atomic number is fundamental in determining the identity of an element. It represents the number of protons found in the nucleus of an atom and is unique to each element.

In neutral atoms, this number also indicates the number of electrons, providing balance between positive and negative charges. However, when atoms form ions, the number of electrons either increases or decreases based on the charge. For example:

• Sodium (\(\mathrm{Na}\)), with an atomic number of 11, loses one electron to form \(\mathrm{Na}^+\)
• Sulfur (\(\mathrm{S}\)), with an atomic number of 16, gains two electrons to form \(\mathrm{S}^{2-}\)

Understanding atomic numbers is crucial for determining the electron configuration of ions and predicting their chemical properties.

Electron Configuration

Electron configuration is a representation of the distribution of electrons among the atomic orbitals in an atom or ion. These configurations are vital in understanding chemical bonding, ion formation, and reactivity.

For an ion, knowing the electron configuration allows us to predict how it will bind with other atoms. For example, both \(\mathrm{Na}^+\) and \(\mathrm{O}^{2-}\) have an electron configuration of \([\mathrm{Ne}]\) because both have 10 electrons. This shared electron configuration ties back to our earlier discussion of isoelectronic species.

Writing electron configurations involves understanding the periodic table's structure, including the filling order of different orbitals, which follows the "Aufbau principle," "Pauli exclusion principle," and "Hund's rule." Becoming adept at electron configurations provides insight into the chemical nature of an element or ion.

Anion

An anion is a negatively charged ion, typically formed when an atom gains one or more electrons. This increase in electrons leads to an overall negative charge as the number of electrons now outweighs the number of protons.

Common examples include:

• \(\mathrm{O}^{2-}\), formed when oxygen gains two electrons
• \(\mathrm{S}^{2-}\), formed when sulfur gains two electrons
• \(\mathrm{Se}^{2-}\), formed when selenium gains two electrons

Understanding how atoms become anions helps predict their behavior in chemical reactions. Anions often participate in forming ionic compounds, where they pair with positively charged cations. This is because opposite charges attract in an ionic bond, stabilizing the overall structure of a compound. Recognizing anions is a foundational skill in predicting and explaining phenomena in chemistry.