Which Of The Following Compounds Is Least Soluble In Water

Which ofthe following compounds is least soluble in water? This question frequently appears in high‑school chemistry exams and serves as a gateway to understanding solubility rules, intermolecular forces, and the practical implications of chemical behavior. In this article we will explore the underlying principles that govern solubility, examine a representative set of common ionic compounds, and determine which one exhibits the lowest water solubility. By the end, readers will not only know the answer but also grasp the why behind it, equipping them to tackle similar problems with confidence.

Introduction to SolubilitySolubility describes the maximum amount of a solute that can dissolve in a given quantity of solvent at equilibrium. For aqueous solutions, solubility is usually expressed in grams of solute per 100 mL of water at a specific temperature, typically 25 °C. Several factors influence solubility:

• Nature of the solute – ionic compounds, molecular substances, and network solids each interact differently with water.
• Temperature – most solid solutes become more soluble as temperature rises, though exceptions exist.
• Pressure – primarily affects gases, not solids or liquids.
• Presence of other ions – common‑ion effect and complex ion formation can shift equilibrium.

Understanding these variables helps answer the core query: which of the following compounds is least soluble in water?

Solubility Rules: A Quick Reference

Before diving into specific compounds, it is useful to recall the general solubility rules that apply to ionic substances in water:

1. Nitrates (NO₃⁻), acetates (CH₃COO⁻), and most perchlorates (ClO₄⁻) are soluble.
2. Group 1 cations (Li⁺, Na⁺, K⁺, Rb⁺, Cs⁺) form soluble salts.
3. Ammonium (NH₄⁺) behaves like Group 1 cations and yields soluble salts.
4. Chlorides (Cl⁻), bromides (Br⁻), iodides (I⁻) are generally soluble, except when paired with Ag⁺, Pb²⁺, or Hg₂²⁺.
5. Sulfates (SO₄²⁻) are soluble except with Ba²⁺, Sr²⁺, Pb²⁺, and Ca²⁺ (the latter being only slightly soluble).
6. Carbonates (CO₃²⁻), phosphates (PO₄³⁻), chromates (CrO₄²⁻), and hydroxides (OH⁻) are mostly insoluble, forming precipitates with many metal cations.

These rules provide a framework for predicting which substances will dissolve readily and which will remain largely undissolved.

Representative Compounds and Their Solubilities

Consider the following set of compounds often used in classroom demonstrations:

• Sodium chloride (NaCl)
• Potassium nitrate (KNO₃)
• Calcium carbonate (CaCO₃)
• Silver chloride (AgCl)
• Barium sulfate (BaSO₄)

Each of these substances belongs to a different category of solubility behavior, making them ideal for illustrating the concept of least solubility Simple, but easy to overlook..

Group 1 Salts: Highly Soluble

Both NaCl and KNO₃ belong to the Group 1 cation family. Experimental data show that NaCl dissolves at approximately 35.Even so, according to the solubility rules, all salts of Li⁺, Na⁺, K⁺, Rb⁺, and Cs⁺ are soluble. Think about it: 9 g per 100 mL of water at 25 °C, while KNO₃ reaches about 31. Consider this: 6 g per 100 mL. These values place them firmly in the highly soluble category But it adds up..

Carbonates: Sparingly Soluble

Calcium carbonate (CaCO₃) is a classic example of a compound that is sparingly soluble. Day to day, its solubility at room temperature is roughly 0. 013 g per 100 mL, a value that reflects the strong lattice energy of the Ca²⁺ and CO₃²⁻ ions and the relatively weak hydration energy of the carbonate anion. Because of this, CaCO₃ forms the familiar chalky precipitate observed in limestone caves and marine sediments.

Halides with Heavy Metals: Low Solubility

Silver chloride (AgCl) and barium sulfate (BaSO₄) are notorious for their extremely low solubility. AgCl’s solubility is about 0.0019 g per 100 mL, while BaSO₄’s is even lower at approximately 0.0002 g per 100 mL. Both compounds precipitate instantly when mixed in aqueous solution, a property exploited in qualitative analysis and photographic processes Not complicated — just consistent..

Identifying the Least Soluble Compound

When the question asks which of the following compounds is least soluble in water, the answer hinges on comparing the quantitative solubility values. The compound with the smallest amount of solute that can dissolve is the least soluble It's one of those things that adds up..

From the table, barium sulfate (BaSO₄) exhibits the lowest solubility, with only 0.On the flip side, 0002 g dissolving in 100 mL of water. This minuscule value places BaSO₄ at the extreme end of the solubility spectrum, making it the correct answer to the query *which of the following compounds is least soluble in water?

Why Is BaSO₄ So Insoluble?

The remarkable insolubility of BaSO₄ can be explained by two inter

Thefirst factor is the exceptionally strong lattice energy of BaSO₄. Barium carries a +2 charge and the sulfate anion bears a –2 charge, so the electrostatic attraction between the ions is intrinsically strong. Worth adding, the ionic radii of Ba²⁺ and SO₄²⁻ are relatively large, which forces the ions to approach each other closely in the crystal lattice, thereby increasing the magnitude of the Coulombic forces that must be overcome for dissolution It's one of those things that adds up..

The second factor is the limited ability of BaSO₄ to attract water molecules. The sulfate group is a relatively weak hydrogen‑bond acceptor, and the barium cation, being large and highly polarizable, does not polarize water molecules efficiently. This means the hydration (solvation) energy released when BaSO₄ ions become surrounded by water is modest compared with the lattice energy that holds the solid together. The balance between these two energies lies far on the side of the lattice, leaving only a trace amount of solid in solution.

Some disagree here. Fair enough.

Because of this combination of high lattice energy and low hydration energy, BaSO₄ precipitates instantly when introduced to an aqueous medium, a characteristic that is exploited in qualitative analysis to confirm the presence of sulfate ions and in radiology as a contrast agent that remains largely undissolved in the body. Its practical insolubility also means that it does not contribute to water hardness or interfere with most chemical processes, making it a valuable inert material in industrial applications.

Simply put, among the compounds listed, barium sulfate (BaSO₄) demonstrates the lowest solubility in water due to its pronounced lattice energy and weak hydration energy, which together prevent the dissolution of appreciable quantities of the solid. This extreme insolubility defines its role in both analytical chemistry and various applied fields, cementing its status as the least soluble substance in the given set.

It sounds simple, but the gap is usually here.

Real‑World Consequences of BaSO₄’s Near‑Insolubility

These examples illustrate that the very property that makes BaSO₄ a “problem” in some contexts (its refusal to go into solution) is precisely why it is prized in others The details matter here. Turns out it matters..

Thermodynamic Perspective

The solubility product (Kₛₚ) of barium sulfate quantifies its insolubility:

[ \text{BaSO}{4(s)} \rightleftharpoons \text{Ba}^{2+}{(aq)} + \text{SO}{4}^{2-}{(aq)} \qquad K_{sp}=1.1\times10^{-10};(\text{at }25^{\circ}\text{C}) ]

A Kₛₚ of 10⁻¹⁰ means that, at equilibrium, the product of the molar concentrations of Ba²⁺ and SO₄²⁻ is only one part in ten billion. 8 × 10⁻¹⁰, CaCO₃ ≈ 3.Day to day, g. 3 × 10⁻⁹). On the flip side, , AgCl Kₛₚ ≈ 1. By contrast, the Kₛₚ values for the other compounds in the list are orders of magnitude larger (e.Worth adding: the extremely low Kₛₚ aligns perfectly with the experimentally observed solubility of 0. 0002 g / 100 mL Simple, but easy to overlook. No workaround needed..

The underlying thermodynamic balance can be expressed as:

[ \Delta G_{\text{dissolution}} = \Delta H_{\text{lattice}} + \Delta H_{\text{hydration}} - T\Delta S_{\text{hydration}} ]

For BaSO₄, (\Delta H_{\text{lattice}}) is large and positive (energy required to break the crystal), while (\Delta H_{\text{hydration}}) is comparatively small and only partially offsets the lattice term. Now, the entropy gain from dispersing ions ((\Delta S_{\text{hydration}})) is modest because only a few ions are released. The net (\Delta G_{\text{dissolution}}) remains positive, indicating a non‑spontaneous process under standard conditions Simple as that..

Common Misconceptions

1. “All sulfates are soluble.”
   Textbook generalizations often state that most sulfates dissolve well, but BaSO₄, PbSO₄, and CaSO₄ are notable exceptions. The rule‑of‑thumb fails because it neglects the interplay of ionic charge, size, and lattice geometry that can tip the balance toward insolubility Which is the point..

2. “Barium ions are always toxic because they dissolve readily.”
   The toxicity of barium stems from soluble salts such as BaCl₂, where the lattice energy is low enough for appreciable dissolution. In BaSO₄, the ions are essentially locked in the solid, so the compound is considered biologically inert when used as a contrast medium—provided it is not inhaled as a fine dust Simple, but easy to overlook..

3. “A white precipitate always means BaSO₄.”
   While BaSO₄ forms a characteristic dense white precipitate, other barium salts (e.g., BaCO₃) and even some non‑barium compounds (e.g., AgCl) can produce similar appearances. Confirmatory tests—such as adding dilute HCl to see if the precipitate dissolves (BaSO₄ remains undissolved, AgCl dissolves)—are essential for accurate identification.

How to Predict Insolubility Without Tables

A quick mental check for a compound’s likely solubility involves three steps:

1. Identify the cation and anion charges. Higher absolute charges increase lattice energy.
2. Consider ionic radii. Smaller ions pack more tightly, raising lattice energy; larger ions can sometimes lower it, but the effect is secondary to charge.
3. Assess hydration potential. Highly charged, small ions (e.g., Mg²⁺, Al³⁺) are strongly hydrated, which can offset a high lattice energy. Large, low‑charge ions (e.g., Ba²⁺, Sr²⁺) gain little from hydration.

Applying this to BaSO₄: a +2 cation paired with a –2 anion (high charge product) and relatively large ionic sizes yield a high lattice energy, while the hydration of Ba²⁺ is weak. The result is a very low Kₛₚ and thus the least solubility among the listed compounds Simple, but easy to overlook..

Conclusion

When asked to identify the least water‑soluble substance among CaCO₃, AgCl, and BaSO₄, the data and thermodynamic reasoning unequivocally point to barium sulfate (BaSO₄). In practice, its exceptionally low solubility—0. 0002 g per 100 mL of water—stems from a potent combination of high lattice energy and modest hydration energy. This physicochemical profile translates into practical advantages: BaSO₄ serves as a reliable precipitating reagent in analytical chemistry, an inert white pigment in industry, and a safe, non‑absorbing contrast agent in medical imaging. Understanding why BaSO₄ behaves the way it does also clarifies broader solubility trends and helps avoid common misconceptions about sulfate salts. In short, BaSO₄’s near‑invisibility to water not only answers the quiz question but also underlies a suite of valuable applications that hinge on its stubborn refusal to dissolve.