A structured visual guide to the chemistry of solutions — from concentration and solubility to Henry's law, vapour pressure, ideal and non-ideal behaviour, colligative properties, osmotic pressure and abnormal molar mass.
Click any concept card for a focused explanation.
Mass %, volume %, ppm, mole fraction, molarity and molality.
Saturation, temperature effects, pressure and dissolution behaviour.
Gas pressure, mole fraction and gas solubility in liquids.
Raoult's law, deviations, azeotropes and molecular interactions.
Vapour pressure, boiling point, freezing point and osmotic pressure.
Association, dissociation and van't Hoff factor.
Different numerical ways to describe composition.
Percentage concentration compares the amount of a chosen component with the total quantity of the solution.
Mole fraction expresses the share of one component in terms of total moles present. It is especially useful in vapour-pressure and phase-equilibrium calculations.
Molarity is based on the volume of the complete solution. Because liquid volume can change with temperature, molarity is temperature dependent.
Molality relates moles of solute to the mass of solvent. Since mass does not change simply because temperature changes, molality is preferred in many temperature-sensitive calculations.
ppm is designed for very small concentrations. It is commonly encountered in environmental chemistry, water analysis and trace-component measurements.
The most useful concentration unit depends on what is being measured and what kind of calculation is required. Mass-based units are generally less sensitive to temperature than volume-based units.
Why substances dissolve — and what controls the limit.
Solubility describes the maximum amount of a substance that can dissolve in a specified amount of solvent under specified conditions, especially temperature and pressure.
The solution contains less dissolved solute than its equilibrium capacity. Additional solute can therefore dissolve under the same conditions.
At equilibrium, the solvent has reached its dissolution capacity. Adding more solute does not increase the dissolved amount under the same conditions.
For many solids, dissolution becomes more favourable as temperature rises when the process is endothermic. If dissolution releases heat, increasing temperature can have the opposite effect. Gas solubility in liquids generally decreases as temperature increases.
Pressure–solubility relationship for gases.
Pressurising a beverage encourages carbon dioxide to remain dissolved. Opening the container reduces pressure and allows dissolved gas to escape.
Aquatic ecosystems depend on dissolved oxygen. Since warmer water generally holds less oxygen, rising temperature can reduce oxygen availability.
Increased pressure can increase gas dissolution in body fluids. Controlled ascent is therefore essential when pressure changes rapidly.
How molecular interactions change solution behaviour.
An ideal solution follows Raoult's law throughout its composition range. Interactions between unlike molecules are approximately comparable to those between like molecules.
If unlike-molecule attraction is weaker than the original interactions, molecules escape more readily into the vapour phase.
Stronger unlike-molecule interactions hold molecules more tightly in the liquid phase. The vapour pressure consequently becomes lower than the ideal prediction.
An azeotropic mixture has a vapour composition that matches the liquid composition at a particular composition and boils at a constant temperature. Ordinary fractional distillation cannot completely separate such a mixture.
A useful conceptual question is: “Are unlike molecules attracting each other more strongly or less strongly than they did in the original liquids?” The answer predicts the direction of deviation.
Properties governed primarily by the number of solute particles.
Adding a non-volatile solute reduces the fraction of solvent molecules available to escape into the vapour phase.
A solution containing a non-volatile solute generally requires a higher temperature to reach the external pressure needed for boiling.
Dissolved particles lower the temperature at which the solvent's solid and liquid phases reach equilibrium.
Osmotic pressure is the pressure that must be applied to a solution to stop the net movement of solvent through a semipermeable membrane.
Exam-focused equations in a separate visual layer.
Drawn entirely with SVG — no image files required.
Ideal binary liquid solution — schematic representation.
Addition of non-volatile solute shifts the boiling point upward.
A solution freezes at a lower temperature than the corresponding pure solvent.
A quick revision comparison.
| Quantity | Based On | Main Advantage | Temperature Sensitivity |
|---|---|---|---|
| Mass % | Mass of solution | Simple composition reporting | Very low |
| Volume % | Volume of solution | Convenient for liquid mixtures | Relatively high |
| ppm | Trace quantities | Extremely dilute systems | Depends on definition |
| Mole fraction | Number of moles | Phase-equilibrium calculations | Low |
| Molarity | Solution volume | Common laboratory calculations | Yes |
| Molality | Solvent mass | Temperature-independent mass basis | Very low |
Use this sequence before solving questions.
Decide which substance is the solvent and which is the solute.
Look at whether the question gives mass, volume, moles, solvent mass or trace concentration.
Be careful whenever solution volume or molarity appears.
Vapour pressure → Raoult; gas solubility → Henry; particle effects → colligative relations.
If experimental and theoretical values disagree, consider the van't Hoff factor.
Convert litres, kilograms, grams and temperature units before substituting into equations.
The central idea of the chapter is that adding one component to another changes measurable physical properties according to composition, molecular interactions and the number of dissolved particles. Once concentration units, solubility behaviour, vapour-pressure laws and colligative relationships are connected, most of the chapter becomes a single logical system rather than a collection of unrelated formulas.