Solubility Data

Introduction

A solubility value is not an intrinsic number attached to a chemical name alone. It depends on solvent, temperature, pressure for gases, and often pH, ionic strength or complex formation. Tables may report grams per 100 g solvent, molarity or mole fraction, which cannot be compared until definitions and conditions are aligned.

Core explanation

Solubility is the maximum equilibrium amount of a specified solute form dissolved in a specified solvent under stated conditions. A saturated solution can contain undissolved solid in equilibrium; stirring changes how quickly equilibrium is reached but not the equilibrium limit under unchanged conditions. Some solids become more soluble when heated, but not all do. Many gases become less soluble as temperature rises, while increased gas partial pressure often increases dissolved amount. OpenStax's solubility chapter discusses these conditional trends and Henry's law.

For a dilute gas in a suitable solvent range, one Henry-law convention is c = kH p, with c dissolved concentration and p gas partial pressure. Other sources define the reciprocal constant or use different pressure and concentration units. Always inspect the equation accompanying a tabulated kH. A gas that reacts significantly after dissolving, such as one that acidifies water, may have total apparent uptake beyond simple physical dissolution; speciation must be included.

Weak acids and bases can change total aqueous solubility with pH. A neutral weak acid HA may dissolve to a limited molecular concentration, while at high pH deprotonation to A⁻ increases total dissolved analytical concentration. Metal ions can form soluble complexes, changing apparent solubility of a salt despite the same intrinsic lattice equilibrium. A table labeled “solubility in water” without pH or complexing conditions may be unsuitable for a strongly buffered or ligand-rich solution.

Read units carefully. “10 g per 100 g water” describes solute mass relative to solvent mass, not solution mass. Molarity needs solution volume, which changes with temperature and mixing. A value in mol/kg solvent is molality, not mol/L solution. Distinguish kinetic supersaturation from equilibrium solubility: a temporarily clear solution may be metastable and crystallize later.

Step-by-step reasoning

1. Identify solute form, solvent, temperature and pressure or pH conditions. 2. Read the solubility unit and its denominator exactly. 3. Check whether dissolution creates ions, complexes or acid-base species. 4. Convert units only with enough density or molar-mass information. 5. Distinguish an equilibrium table value from transient supersaturation.

Visual explanation

Draw a closed vessel with undissolved solid beneath a saturated solution, and double arrows showing dissolution and precipitation. A second plot shows total weak-acid solubility rising with pH as ionized A⁻ becomes dominant. A third sketch shows dissolved gas concentration increasing with partial pressure under a stated Henry-law convention.

Real-world analogy

The number of people who can sit in a room depends on the room, chair arrangement and whether standing is allowed. Saying only “capacity = 50” without conditions can mislead. Solubility likewise needs solvent and chemical conditions, especially when the dissolved species can change form.

Real-world example

A water-treatment chemist evaluates whether a metal will precipitate. A pure-water solubility table suggests low dissolved concentration, but the sample contains a ligand that forms soluble metal complexes. The total dissolved metal can be higher than the simple salt-table prediction. The chemist uses equilibrium speciation and measured pH instead of treating the tabulated number as a universal maximum.

Why?

Why can pH change total solubility of a weak acid? Removing dissolved neutral HA by converting it to A⁻ shifts dissolution toward more HA entering solution, provided the ionized form remains soluble. The intrinsic neutral-species solubility and the total analytical solubility are different quantities.

Common misconception

“All solids dissolve better when heated” has exceptions. “Stirring raises equilibrium solubility” confuses rate with equilibrium. “Henry constants have one universal unit convention” is false. “A clear solution is certainly below saturation” ignores metastable supersaturation.

Worked example

Suppose a gas has dissolved concentration 0.010 mol/L at partial pressure 50 kPa under a stated constant-temperature Henry form c = kH p. Then kH = 0.010/50 = 2.0 × 10⁻⁴ mol L⁻¹ kPa⁻¹. At 100 kPa, the idealized prediction is c = 0.020 mol/L. It assumes the same temperature, dilute behavior and no significant chemical reaction. If the gas reacts with water or pressure becomes high enough for nonideal behavior, the simple proportional estimate needs revision. The numerical example is illustrative, not a universal solubility value.

Quick check

1. Does “10 g solute per 100 g water” mean 10 g per 100 g of final solution? Answer: No. The denominator is solvent mass, so the final solution mass is at least 110 g if no material is lost.

Exam focus

Read solubility units and conditions precisely. Apply a clearly stated Henry-law convention and track its units. Explain how acid-base reaction or complexation changes total dissolved amount. Distinguish dissolving faster from a higher equilibrium solubility.

Advanced insight

At higher ionic strength, activities rather than bare concentrations govern thermodynamic equilibria. A solubility product based on activities can remain defined while concentration-based “apparent Ksp” shifts with medium. This is why a reference solution condition is essential when comparing literature solubilities of sparingly soluble salts.

Summary

Solubility data describe an equilibrium under named solvent, temperature and chemical conditions. Gas pressure, pH and complexation can change measured total dissolved concentration. Correct use requires reading units and distinguishing intrinsic dissolution from speciation and temporary supersaturation.

Practice questions

1. What does stirring mainly change when a solid dissolves? Answer: It usually changes the rate of reaching equilibrium, not the equilibrium solubility. 2. Under c = kH p, what happens to c when p doubles within the valid regime? Answer: It doubles at unchanged temperature and solvent conditions. 3. Why may a weak acid's total solubility rise at high pH? Answer: Deprotonation forms soluble A⁻ and pulls more neutral acid into solution. 4. Why may complexing ligands make a metal salt appear more soluble? Answer: They bind dissolved metal ions, reducing free-ion activity and shifting dissolution toward more total dissolved metal.