Question

In Chapter \(13,\) you learned that the bonding forces in ionic solids such as \(\mathrm{NaCl}\) are very strong, yet many ionic solids dissolve readily in water. Explain.

Short answer

Ionic solids like NaCl have strong bonding forces due to the electrostatic attraction between their positively and negatively charged ions. However, they can still dissolve readily in water because water is a polar solvent. The polar water molecules surround the ions in the ionic compound through a process called hydration, where the negatively charged oxygen atoms attract the positive ions and the positively charged hydrogen atoms attract the negative ions. The energy released during this process, called hydration energy, helps to overcome the lattice energy holding the ions together in the ionic compound. Therefore, ionic compounds like NaCl dissolve in water when the hydration energy is sufficient to overcome the lattice energy.

Step-by-step solution

Understand Ionic Bonds

Ionic bonds are formed between elements with a large difference in electronegativity, usually a metal and a non-metal. In the case of NaCl (sodium chloride), the positively charged sodium cations (Na+) are attracted to the negatively charged chloride anions (Cl-) through electrostatic forces, forming a highly ordered, stable crystal lattice structure.

Water as a Polar Solvent

Water is a polar molecule, which means it has a partial positive charge on the hydrogen atoms and a partial negative charge on the oxygen atom, due to the difference in electronegativity between hydrogen and oxygen. This polarity makes water an excellent solvent for ionic compounds.

Process of Dissolution

When ionic compounds like NaCl come into contact with water, the polar water molecules surround the ions in the crystal lattice. The negatively charged oxygen atoms of water molecules are attracted to the positively charged sodium ions (Na+), while the positively charged hydrogen atoms are attracted to the negatively charged chloride ions (Cl-). This process is called hydration.

Hydration Energy

The energy released during the hydration process, called hydration energy, helps to overcome the lattice energy of the ionic compound (the energy that holds the ions together in the crystal lattice). When the hydration energy is greater than or comparable to the lattice energy, the ionic compound will dissolve in water.

Ionic Solids Dissolving in Water

In the case of NaCl and many other ionic compounds, the hydration energy is sufficient to overcome the lattice energy, allowing the ions to separate from the crystal lattice and become surrounded by water molecules. As a result, the ionic compound dissolves readily in water, despite the strong bonding forces holding the ions together within the solid.

Key concepts

Polar Solvents

Polar solvents, such as water, are crucial in the process of dissolving ionic compounds like sodium chloride (NaCl). These solvents have molecules with an uneven distribution of electrical charge, meaning one part of the molecule is more negative, and another part is more positive. This characteristic stems from the difference in electronegativity between the atoms that make up the molecule.

For instance, water's oxygen atom has a higher electronegativity compared to its hydrogen atoms. As a consequence, the electrons spend more time near the oxygen atom than the hydrogen atoms, creating a partial negative charge near the oxygen and a partial positive charge near the hydrogens. This separation of charge allows the water molecules to interact with and attract the oppositely charged ions in an ionic compound. These interactions are essential for the dissolution process as they enable the solvent to surround and isolate the ions from the solid structure.

Hydration Energy

When discussing solubility, hydration energy plays a pivotal role. It is the energy released when water molecules bond to the ions of a solute. This occurs because the partial positive charge of the water's hydrogen atoms attracts the negatively charged anions, while the partial negative charge on the oxygen atom attracts positively charged cations.

The strength of this attraction and the resultant energy release can be powerful enough to break the ionic bonds within a solid crystal, facilitating dissolution. As a benchmark, if the hydration energy exceeds or is on par with the lattice energy of the ionic compound, the compound is likely to dissolve in water. The greater the hydration energy in comparison to the lattice energy, the more soluble the compound will be.

Lattice Energy

Lattice energy is the amount of energy that holds an ionic solid's structure together. In an ionic crystal, each ion is bonded to its neighbors by electrostatic forces—the attraction between positively and negatively charged ions. This ordered three-dimensional arrangement of ions forms a lattice structure that is very stable and requires a significant amount of energy to break apart.

In the context of solubility, the lattice energy must be compared to hydration energy. If the hydration energy provided by a polar solvent like water is high enough, it can break apart the lattice, dissolving the ionic compound. Essentially, the competition between the lattice energy trying to hold the ions together and the hydration energy trying to pull them apart determines an ionic compound's solubility.

Electronegativity

Electronegativity is a measure of an atom's ability to attract and hold onto electrons within a chemical bond. Atoms with high electronegativity, such as oxygen, tend to pull electrons closer, while those with low electronegativity, such as sodium, are less adept at attracting electrons.

This difference in electronegativity not only determines the polarity of a molecule (like water) but also the bond type that forms between atoms—in this case, the ionic bond in NaCl. Sodium (Na) has a low electronegativity and readily loses an electron, becoming a cation (Na+), whereas chlorine (Cl) has a high electronegativity and gains an electron, becoming an anion (Cl-). It is the stark difference in electronegativity between the metal (Na) and the non-metal (Cl) that results in the ionic bond and the formation of the NaCl crystal lattice.