Question

Draw Lewis symbols showing the valence electrons of the following ions. (a) \(\mathrm{P}^{3-}\) (c) \(\mathrm{Se}^{2-}\) (e) \(\mathrm{C}^{4-}\) (b) \(\mathrm{In}^{3+}\) (d) \(\mathrm{Be}^{2+}\)

Short answer

The Lewis symbols for the given ions are as follows: For \( \mathrm{P}^{3-}\), \( \mathrm{Se}^{2-}\), and \( \mathrm{C}^{4-}\), there should be 8 dots around the atom, representing 8 valence electrons. For \( \mathrm{In}^{3+}\) and \( \mathrm{Be}^{2+}\), there should be zero dots, representing no valence electrons after the atoms have lost their electrons to become ions.

Step-by-step solution

Determine the number of valence electrons for each ion

To determine the number of valence electrons in an ion, first find the number of valence electrons in the neutral atom of the same element, then adjust for any lost or gained electrons to reach the ionic state. For example, \( \mathrm{P}^{3-}\) means that phosphorus has gained 3 electrons, becoming an ion with 5 + 3 = 8 electrons in its outer shell.

Draw the Lewis symbols

The Lewis symbol is drawn by writing the symbol of the atom and surrounding it with dots representing each valence electron. Ions are drawn in brackets with their charge written outside the brackets. Following the method explained in step 1, draw the Lewis symbols for all the ions: - \( \mathrm{P}^{3-}\) will have eight dots surrounding it, meaning 8 valence electrons.- \( \mathrm{Se}^{2-}\) will have eight dots, meaning 8 valence electrons.- \( \mathrm{C}^{4-}\) will have eight dots, meaning 8 valence electrons.- \( \mathrm{In}^{3+}\) will have no dots since it has lost all its valence electrons.- \( \mathrm{Be}^{2+}\) will also have no dots, same reason as indium.

Key concepts

Valence Electrons

Valence electrons are the outermost electrons of an atom that participate in chemical bonding. Understanding how to count and illustrate valence electrons is essential in determining how atoms will interact with each other. For instance, a neutral phosphorus atom has five valence electrons, selenium has six, carbon has four, indium has three, and beryllium has two.

When atoms become ions, they lose or gain electrons to achieve a stable electron configuration, often resembling the nearest noble gas. In the exercise example, we notice that when phosphorus becomes a \(\mathrm{P}^{3-}\) ion, it gains three electrons. This results in eight valence electrons, which fills its outer shell. Similarly, \(\mathrm{Se}^{2-}\) and \(\mathrm{C}^{4-}\) gain two and four electrons respectively, each ending up with eight valence electrons. In contrast, \(\mathrm{In}^{3+}\) and \(\mathrm{Be}^{2+}\) lose electrons, ending up with zero valence electrons because they have lost their entire outer shell.

Visualizing this concept through Lewis symbols helps clarify the concept of valence electrons and predict ions' behavior in chemical reactions.

Ionic Charge

The ionic charge is the electrical charge an atom carries when it loses or gains electrons to form an ion. In a neutral atom, the number of protons (positive charges) equals the number of electrons (negative charges). However, when an atom forms an ion, this balance is disrupted.

For instance, the \(\mathrm{P}^{3-}\) ion carries a 3- charge because it has gained three additional electrons beyond its stable neutral state. This negative charge indicates an excess of electrons. Conversely, the \(\mathrm{In}^{3+}\) ion has a 3+ charge, implying it has lost three electrons and therefore has three more protons than electrons, giving it a positive charge.

In Lewis symbols, the ionic charge is typically indicated just outside the brackets that contain the element's symbol and its valence electrons. This portrays the atom's ionized state and is crucial in understanding how the ion will interact with others to form ionic bonds.

Electron Configuration

Electron configuration describes the distribution of electrons in the orbitals of an atom. For students new to chemistry, thinking of it as the atom's 'address book,' where electrons 'reside' in different 'homes' (orbitals), can be a helpful analogy. Neutral atoms fill their orbitals in a specific order, following rules such as the Aufbau principle, Hund's rule, and the Pauli exclusion principle.

When atoms become ions, their electron configuration changes as electrons are either added to or removed from their valence shell. This adaptation seeks to reach the most energetically favorable state, often modeled after the noble gases, known for their stability due to having complete valence shells.

In the context of our exercise, \(\mathrm{P}^{3-}\), \(\mathrm{Se}^{2-}\), and \(\mathrm{C}^{4-}\) ions have gained electrons to complete their valence shell similar to the noble gas configuration, whereas \(\mathrm{In}^{3+}\) and \(\mathrm{Be}^{2+}\) have lost all their valence electrons and have a configuration resembling the noble gas in the period before them. Understanding electron configurations is fundamental to predicting how atoms will bond and the properties of the resulting substances.