Based on the answer to Practice Problem \(12.4\), is the energy required for solute separation greater for \(\mathrm{NaCl}\) or \(\mathrm{MgCl}_{2}\) ? Explain.

Understanding the Lattice Energy Formula is key to making sense of the forces that stabilize ionic compounds. Simply put, the formula provides insight into the amount of energy that would be needed to break an ionic solid apart into isolated gaseous ions.

The general formula is given by:
LE = \( \frac{k_{c} \cdot q_{1} \cdot q_{2}}{r} \)

where \(LE\) stands for Lattice Energy, \(k_{c}\) is Coulomb's constant, \(q_{1}\) and \(q_{2}\) represent the charges of the ions involved, and \(r\) is the distance between the centers of the ions. A higher charge and closer proximity (a smaller distance) mean stronger ionic bonds and higher lattice energy. When applied to different compounds, this formula allows us to predict which compound would require more energy to separate its constituents; that's exactly what facilitates the comparison between the lattice energies of different ionic solids, such as NaCl and MgCl₂ in the given problem.

Ions are atoms or molecules with an electric charge due to the loss or gain of electrons, and the force that holds oppositely charged ions together in a compound is known as an ionic bond.

When exploring ionic bonds, elements from the periodic table often undergo ionic bonding to reach a stable electronic configuration, typically achieving a full outer electron shell. For example:

• In sodium chloride (NaCl), sodium (Na⁺) donates an electron and becomes a positively charged ion, while chlorine (Cl⁻) accepts an electron, becoming negatively charged. They come together to form NaCl, which has a neutral overall charge.
• Magnesium chloride (MgCl₂), however, involves a magnesium ion (Mg²⁺) bonding with two chloride ions to maintain charge neutrality. Each chloride ion contributes a -1 charge, making a total of a -2 charge to balance the +2 charge from magnesium.

Understanding ionic bonds is crucial for analyzing the strength of interactions between ions, which in turn influences properties like melting points, boiling points, and lattice energy of the compounds.

Solute separation energy essentially refers to the amount of energy needed to overcome the forces holding solute particles together in order to dissolve the solid into its constituent ions.

This concept is vitally important when studying how ionic compounds interact with solvents. For compounds like NaCl and MgCl₂, which exhibit ionic bonding, the solute separation energy is closely linked to the Lattice Energy. A higher lattice energy implies more energy is needed to separate the ions, suggesting that the compound is more stable.

In practice, comparing the solute separation energy between NaCl and MgCl₂ gives us insight into why MgCl₂'s higher charge product results in significantly more energy needed for dissolution. This characteristic also impacts the solubility and dissolution rate; MgCl₂ takes more energy to dissolve in solvents compared to NaCl, which relates to the stronger ionic bond in the former.