Question

Using orbital box diagrams, depict an electron configuration for each of the following ions: (a) \(\mathrm{Na}^{+}\) (b) \(\mathrm{Al}^{3+},\) (c) \(\mathrm{Ge}^{2+},\) and \((\mathrm{d}) \mathrm{F}^{-}\).

Short answer

All ions listed have fully filled electron shells except Ge^{2+}, which lacks electrons in the 4p orbital.

Step-by-step solution

Understanding the Neutral Atom Configuration

First, identify the electron configuration of the neutral atoms before you consider the ion states. This provides the basis for removing or adding electrons.- **Na:** Neutral sodium has an electron configuration of \(1s^2 2s^2 2p^6 3s^1\).- **Al:** Neutral aluminum has an electron configuration of \(1s^2 2s^2 2p^6 3s^2 3p^1\).- **Ge:** Neutral germanium has an electron configuration of \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^2\).- **F:** Neutral fluorine has an electron configuration of \(1s^2 2s^2 2p^5\).

Adjust Electron Count for Ion Formation

Now adjust for each ion by removing or adding electrons as needed:- **Na^+:** Remove one electron (from 3s), resulting in \(1s^2 2s^2 2p^6\).- **Al^{3+}:** Remove three electrons (from 3p and 3s), resulting in \(1s^2 2s^2 2p^6\).- **Ge^{2+}:** Remove two electrons (from 4p), resulting in \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10}\).- **F^-:** Add one electron (to 2p), resulting in \(1s^2 2s^2 2p^6\).

Draw Orbital Box Diagrams

Depict the electron configurations using orbital box diagrams for each ion. These diagrams use boxes to represent orbitals and arrows to indicate electrons. - **Na^+:** All electrons occupy filled orbitals (1s, 2s, 2p boxes are fully filled). --------------------------------- | ↑↓ | 1s | --------------------------------- | ↑↓ | 2s | --------------------------------- | ↑↓ ↑↓ ↑↓ | 2p | --------------------------------- - **Al^{3+}:** Similar to Na^+ with all filled orbitals. --------------------------------- | ↑↓ | 1s | --------------------------------- | ↑↓ | 2s | --------------------------------- | ↑↓ ↑↓ ↑↓ | 2p | --------------------------------- - **Ge^{2+}:** The 4p orbital is empty while earlier orbitals are entirely filled. ------------------------------------ | ↑↓ | 1s | ------------------------------------ | ↑↓ | 2s | ------------------------------------ | ↑↓ ↑↓ ↑↓ | 2p | ------------------------------------ | ↑↓ | 3s | ------------------------------------ | ↑↓ ↑↓ ↑↓ | 3p | ------------------------------------ | ↑↓ | 4s | ------------------------------------ | ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ | 3d | ------------------------------------ | | 4p | (no electrons here) ------------------------------------ - **F^-:** Adding one electron fills the 2p shell completely. --------------------------------- | ↑↓ | 1s | --------------------------------- | ↑↓ | 2s | --------------------------------- | ↑↓ ↑↓ ↑↓ | 2p | ---------------------------------

Key concepts

Orbital Box Diagrams

Orbital box diagrams are an excellent tool for visualizing the electron configuration of atoms and ions. They help us understand the arrangement of electrons in atomic orbits through a simple box and arrow method. Each box represents an orbital, while arrows indicate electrons within those orbitals. For example, two arrows in a single box represent a pair of electrons in an orbital that are paired with opposite spins. It's important to note that each orbital can hold up to two electrons with spin pairing.

• Boxes are usually grouped together by subshell types, such as 1s, 2s, 2p, etc.
• The arrows (electrons) are distributed according to the Aufbau principle, filling lower-energy orbitals first before moving to higher energy ones.
• In diagrams, 'up' and 'down' arrows are used to indicate electron spin, following Hund's Rule by filling unpaired at first.

Orbital box diagrams offer a clear, visual representation of electron distribution, useful especially in predicting chemical reactivity and ion formation.

Ion Electron Configurations

When atoms gain or lose electrons to become ions, their electron configurations change. This occurs as electrons are added to or removed from their outer shells. The resulting electron configuration of an ion is essential for predicting an ion's chemical properties.

• For example, a neutral sodium atom (\[\text{Na}\]) loses one electron to become a positively charged ion (\[{\text{Na}^+}\]), modifying its electron configuration from \[1s^2 2s^2 2p^6 3s^1\] to \[1s^2 2s^2 2p^6\]. This makes \[{\text{Na}^+}\] a stable ion, mimicking the electron configuration of neon.
• Conversely, fluorine (\[\text{F}\]) adds an electron to become a negatively charged ion (\[\text{F}^-\]), changing its configuration from \[1s^2 2s^2 2p^5\] to \[1s^2 2s^2 2p^6\].

The determination of ion electron configurations helps to understand ionic behaviors, especially their tendency to achieve noble gas configurations.

Electron Removal and Addition

The process of electron removal and addition is crucial in forming ions. Electrons are removed from or added to the outermost orbitals of an atom, which impacts the overall charge and stability of the ion.

• When forming positive ions, electrons are removed, typically starting with the highest orbital energy level. For example, in the aluminum ion (\[\text{Al}^{3+}\]), three electrons are removed from the \[3p\] and \[3s\] orbitals.
• For negative ions, electrons are added until the outer orbitals are filled, like how fluorine becomes \[\text{F}^-\] by adding one electron to complete the \[2p\] shell.

Understanding electron removal and addition is central to predicting how atoms will react chemically to either become positively or negatively charged ions.

Filled and Empty Orbitals

Filled and empty orbitals play a significant role in determining the stability and reactivity of atoms and ions. An orbital is a region around an atom's nucleus where electrons are likely to be found. When orbitals are filled, they generally make the atom or ion more stable.

• Filled orbitals provide atoms with a more stable electron configuration, often resembling that of the nearest noble gas. For instance, the electron configuration of \[\text{F}^-\] is \[1s^2 2s^2 2p^6\], which is the same as neon (\[\text{Ne}\]), contributing to its stability.
• Conversely, empty or partially filled orbitals can make an atom or ion more reactive, as they tend to gain or lose electrons to achieve a filled configuration.

By examining which orbitals are filled or empty, you can predict the chemical behavior of different ions and understand why certain ions are more stable than others.