Question

How many unpaired electrons are in the following ions? (a) \(\mathrm{Hg}^{2+}\) (b) \(\mathrm{F}^{-}\) (c) \(\mathrm{Sb}^{3+}\) (d) \(\mathrm{Fe}^{3+}\)

Short answer

Answer: (a) Hg2+ ion has 0 unpaired electrons, (b) F- ion has 0 unpaired electrons, (c) Sb3+ ion has 0 unpaired electrons, and (d) Fe3+ ion has 5 unpaired electrons.

Step-by-step solution

(a) Hg2+ ion

Step 1: Find atomic number of Hg Hg (mercury) is in the periodic table with an atomic number of 80. Step 2: Determine electron configuration of the neutral atom The electron configuration of the neutral Hg atom is \(\mathrm{[Xe] 4f^{14} 5d^{10} 6s^2}\). Step 3: Electron configuration of the ion For the \(\mathrm{Hg}^{2+}\) ion, we need to remove 2 electrons. They are removed from the highest energy level, which is 6s. So, the electron configuration of the ion is \(\mathrm{[Xe] 4f^{14} 5d^{10}}\). Step 4: Count unpaired electrons There are no unpaired electrons in \(\mathrm{Hg}^{2+}\) ion as all the electrons are in fully-filled orbitals.

(b) F- ion

Step 1: Find atomic number of F F (fluorine) is in the periodic table with an atomic number of 9. Step 2: Determine electron configuration of the neutral atom The electron configuration of the neutral F atom is \(\mathrm{1s^2 2s^2 2p^5}\). Step 3: Electron configuration of the ion For the \(\mathrm{F}^{-}\) ion, we need to add 1 electron. It will be added to the 2p orbital. So, the electron configuration of the ion is \(\mathrm{1s^2 2s^2 2p^6}\). Step 4: Count unpaired electrons There are no unpaired electrons in \(\mathrm{F}^{-}\) ion as all the electrons are in fully-filled orbitals.

(c) Sb3+ ion

Step 1: Find atomic number of Sb Sb (antimony) is in the periodic table with an atomic number of 51. Step 2: Determine electron configuration of the neutral atom The electron configuration of the neutral Sb atom is \(\mathrm{[Kr] 4d^{10} 5s^2 5p^3}\). Step 3: Electron configuration of the ion For the \(\mathrm{Sb}^{3+}\) ion, we need to remove 3 electrons. They are removed first from the 5s orbital, and then from the 5p orbital. So, the electron configuration of the ion is \(\mathrm{[Kr] 4d^{10} 5p^0}\). Step 4: Count unpaired electrons There are no unpaired electrons in \(\mathrm{Sb}^{3+}\) ion as all the electrons are in fully-filled orbitals.

(d) Fe3+ ion

Step 1: Find atomic number of Fe Fe (iron) is in the periodic table with an atomic number of 26. Step 2: Determine electron configuration of the neutral atom The electron configuration of the neutral Fe atom is \(\mathrm{[Ar] 3d^6 4s^2}\). Step 3: Electron configuration of the ion For the \(\mathrm{Fe}^{3+}\) ion, we need to remove 3 electrons. They are removed first from the 4s orbital, and then from the 3d orbital since 4s orbital has slightly higher energy level. So, the electron configuration of the ion is \(\mathrm{[Ar] 3d^5}\). Step 4: Count unpaired electrons There are 5 unpaired electrons in \(\mathrm{Fe}^{3+}\) ion as all the electrons are in the 3d orbitals and none of them are fully-filled.

Key concepts

Unpaired Electrons

Unpaired electrons are electrons in an atom that are not paired with another electron in an orbital. Each orbital can hold two electrons, and if there's only one electron in an orbital, it is considered an unpaired electron. Unpaired electrons play a significant role in determining the magnetic properties of an atom or an ion.

Unpaired electrons occur in the valence shell, which is the outermost shell of an atom. Elements and ions have varying numbers of unpaired electrons depending on their electron configuration. For example:

• The \(\mathrm{Hg}^{2+}\) and \(\mathrm{F}^{-}\) ions studied have no unpaired electrons because their electron configurations are in filled orbitals.
• In contrast, the \(\mathrm{Fe}^{3+}\) ion has 5 unpaired electrons, represented by its \(3d^5\) configuration.

Understanding unpaired electrons helps in predicting the chemical bonding and magnetic properties of elements and ions.

Ions

Ions are atoms or molecules that have gained or lost electrons, resulting in a net positive or negative charge. This change occurs to satisfy the octet rule, achieving a stable electronic configuration much like the nearest noble gas.

There are two types of ions:

• Cations: Positively charged ions formed when an atom loses electrons. For instance, \(\mathrm{Hg}^{2+}\) and \(\mathrm{Fe}^{3+}\) are cations because they have lost electrons.
• Anions: Negatively charged ions formed when an atom gains electrons. \(\mathrm{F}^{-}\) is an example of an anion as it has gained an extra electron.

Ions are essential in various chemical processes and are found extensively in biochemical and industrial applications.

Periodic Table

The periodic table is a systematic arrangement of elements in order of increasing atomic number. This table is divided into periods (rows) and groups (columns), where elements in the same group often share similar chemical properties.

The periodic table is important for predicting the behavior of elements, their electron configurations, and understanding trends such as:

• Atomic size
• Electronegativity
• Ionization energy

For example, elements like mercury (Hg), fluorine (F), antimony (Sb), and iron (Fe) can be located on the periodic table to determine their atomic number and properties. This helps in identifying how many electrons they possess and how they will form ions.

Atomic Number

The atomic number of an element is the number of protons in the nucleus of an atom. It uniquely identifies a chemical element and is represented by the symbol \(Z\) on the periodic table.

The atomic number indicates the number of electrons in a neutral atom. For ions, the atomic number remains the same even if the number of electrons changes due to gain or loss. Here are a few examples:

• Mercury (\(\mathrm{Hg}\)), with atomic number 80, has 80 protons and typically 80 electrons.
• Fluorine (\(\mathrm{F}\)), with atomic number 9, contains 9 protons and usually 9 electrons in its neutral state.
Understanding atomic numbers is crucial for describing elements, understanding their placement on the periodic table, and predicting their behavior in chemical reactions.