A. Why Is Liquid Ammonia a Poor Solvent for Ionic Compounds?
- Dielectric Constant and Ion Dissolution:
The dielectric constant of a solvent measures its ability to reduce the electrostatic forces between charged particles (ions). Water has a high dielectric constant (~80), which means it can effectively separate and stabilize ions, allowing many salts to dissolve easily. Liquid ammonia, however, has a much lower dielectric constant (~22). This means it cannot separate ions as effectively as water. When ionic compounds dissolve, their positive and negative ions tend to stay close together in ion pairs or clusters rather than separating freely in the solution. As a result, many ionic salts are poorly soluble in liquid ammonia. - Selective Solubility of Certain Salts:
Despite the low dielectric constant, some salts do dissolve well in liquid ammonia. For example:- Nitrates, thiocyanates, perchlorates, and cyanides dissolve because their ions interact better with ammonia molecules or have ionic lattices that are easier to break.
- Salts like oxides, hydroxides, carbonates, phosphates, sulfates, and sulfides remain insoluble because their strong ionic bonds and large lattice energies aren’t overcome by the solvent’s limited ability to separate ions.
- Halide Solubility Trends:
- Iodides (I⁻) dissolve well because larger iodide ions interact favorably with ammonia.
- Bromides (Br⁻) dissolve sparingly; partial interaction occurs.
- Chlorides (Cl⁻) and fluorides (F⁻) are mostly insoluble because these smaller, more tightly held ions don’t interact well with ammonia, except some exceptions like BaCl₂ and NaCl which dissolve to some extent.
- Solubility of Ammonium Salts and Organics:
- Ammonium salts (NH₄⁺ salts) dissolve well because ammonium ions form strong hydrogen bonding and favorable interactions with ammonia molecules.
- Organic compounds such as alcohols, ethers, ketones, halogens, and phenols dissolve readily due to similar polarity or hydrogen bonding capability.
- Aromatic hydrocarbons are sparingly soluble due to their nonpolar nature and poor interaction with ammonia.
- Non-metals like sulfur, phosphorus, and iodine dissolve but often chemically react with ammonia, not just physically dissolve.
- Increased Basicity of Liquid Ammonia:
Compared to water, ammonia is a stronger base because the nitrogen atom has a lone pair of electrons that more readily accepts protons. This makes liquid ammonia a more basic medium, affecting the chemical reactions and equilibria in the solvent. - Ion Pairing and Clustering:
Because ammonia can’t fully separate ions due to its low dielectric constant, ions tend to stay paired or form clusters rather than being freely mobile. This influences conductivity, reactivity, and the overall behavior of salts in liquid ammonia.
Summary:
- Liquid ammonia’s low dielectric constant makes it a less effective solvent for many salts compared to water.
- Some salts dissolve well due to specific ion interactions or lattice energies.
- Ammonia’s higher basicity alters the chemical environment compared to water.
- Ion pairing/clustering is common, even for strong acids, bases, and salts, changing the dynamics of solutions in liquid ammonia.
B. Solubility of Ionic Compounds (Inorganic Salts) in Liquid SO₂
1. Iodides and Thiocyanates — Most Soluble:
Iodide (I⁻) and thiocyanate (SCN⁻) ions are relatively large and polarizable. Their large ionic radii and diffuse charge distribution reduce the lattice energy of their salts, making the lattice easier to break apart. Moreover, liquid SO₂, being a polar but aprotic solvent with moderate dielectric constant (~17), can stabilize these large, polarizable anions through dipole-induced interactions. This enhances their solubility.
2. Sulfates, Sulfides, Oxides, and Hydroxides — Partially Insoluble:
These ions generally form salts with higher lattice energies due to stronger ionic bonds and smaller ionic sizes, leading to more rigid crystal lattices. The relatively low dielectric constant of SO₂ is insufficient to fully separate and stabilize these ions. Additionally, these ions often have higher charge density, so the solvent’s polarity is not enough to overcome the strong ionic attractions, causing poor solubility.
B. Solubility of Non-Ionic (Covalent) Compounds in Liquid SO₂
1. Covalent Halides (e.g., IBr, BCl₃, AlCl₃, AsCl₃, PBr₃, CCl₄, SiCl₄, SnCl₄) — Soluble:
These compounds are molecular and often polar or polarizable. SO₂, having a permanent dipole moment and the ability to act as a Lewis base, can interact with covalent halides via dipole-dipole interactions or coordinate bonding (especially with Lewis acidic compounds like AlCl₃). This stabilizes them in solution, increasing solubility.
2. Organic Compounds (amines, ethers, alcohols, benzene, alkenes, pyridine, quinoline, halogen derivatives, acid chlorides) — Soluble:
These molecules vary in polarity but many have lone pairs or π-electrons capable of interacting with SO₂’s dipole. Polar organics like amines and alcohols can engage in dipolar or weak hydrogen bonding with SO₂. Aromatic and alkene compounds dissolve due to van der Waals forces and some dipole-induced dipole interactions.
3. Alkanes — Insoluble:
Alkanes are nonpolar, saturated hydrocarbons with only weak London dispersion forces. SO₂ is a polar solvent, so it does not effectively solvate nonpolar alkanes, leading to poor solubility (like “like dissolves like” principle).
C. Solubility of Metals
Metals are insoluble in liquid SO₂ because metallic bonding involves a sea of delocalized electrons not easily disrupted by molecular solvents. SO₂ cannot sufficiently solvate metal atoms or ions without prior oxidation or chemical reaction.
D. Conductivity of Salt Solutions in Liquid SO₂
1. Effect of Cation Size on Conductivity:
Conductivity depends on ion mobility and degree of ion dissociation. Larger cations such as (CH₃)₄N⁺ (tetramethylammonium) have lower charge density and weaker solvation shells, allowing them to move more freely, increasing conductivity. Smaller ions like Na⁺ have stronger solvation and higher charge density, hindering mobility and lowering conductivity.
Increasing order of cation conductivity:
Na⁺ < NH₄⁺ < K⁺ < (CH₃)₃S⁺ < (CH₃)₄N⁺
2. Effect of Anion Size on Conductivity:
Larger anions (like I⁻) are more polarizable and less tightly solvated, increasing their mobility. Smaller anions (like SCN⁻) interact more strongly with solvent molecules or associate more, reducing mobility.
Increasing order of anion conductivity:
SCN⁻ < ClO₄⁻ < Cl⁻ < I⁻
Summary
- Liquid SO₂ is a moderately polar solvent with limited ability to solvate highly charged or small ions, so solubility depends heavily on ion size, charge, and polarizability.
- Covalent and organic compounds dissolve well due to dipole or coordinate interactions with SO₂.
- Metals do not dissolve due to the nature of metallic bonding.
- Conductivity in liquid SO₂ solutions depends on ion size, solvation, and mobility, with larger ions typically showing higher conductivity.