Difficulty: Intermediate | Prerequisites: Chapter 11 (intermolecular forces), basic stoichiometry, mole concept
This chapter builds directly on intermolecular forces by asking: what happens when you mix two substances? It covers why things dissolve, how to express concentration in several different units, what controls solubility, and then introduces colligative properties, which depend on how many solute particles are present rather than what those particles are. It closes with colloids and micelles. If you are not yet comfortable identifying IMFs from Lewis structures, go back to Chapter 11 before starting here.
Solutions form when solute-solvent interactions are strong enough to overcome solute-solute and solvent-solvent attractions. Concentration can be expressed as molarity, mass %, volume %, ppm, or mass/volume %. Colligative properties (vapour pressure lowering, boiling point elevation, freezing point depression, osmotic pressure) depend on the number of dissolved particles, not their identity. Colloids are mixtures with particle sizes between those of true solutions and suspensions, and micelles are a key colloidal structure.
Solubility
The maximum amount of solute that can dissolve in a given amount of solvent at a specific temperature. Think of it as the capacity of the solvent to hold that solute.
Saturated solution
A solution that contains the maximum amount of dissolved solute at a given temperature. Any additional solute will remain undissolved.
Unsaturated solution
A solution that contains less solute than the maximum the solvent can dissolve. More solute could still be added and it would dissolve.
Supersaturated solution
A solution that contains more dissolved solute than a saturated solution would at the same temperature. These are unstable: a disturbance (a seed crystal, a scratch) can cause the excess solute to crystallise out rapidly.
Molarity (M)
Moles of solute per litre of solution. The most commonly used concentration unit in general chemistry.
Mass percent (mass %)
Mass of solute divided by total mass of solution, multiplied by 100.
Volume percent (volume %)
Volume of solute divided by total volume of solution, multiplied by 100.
Parts per million (ppm)
Mass of solute per million parts of solution (mg solute per kg solution, or mg/L for dilute aqueous solutions). Used for very low concentrations.
Mass/volume percent (m/v %)
Mass of solute (in grams) divided by volume of solution (in mL), multiplied by 100.
Colligative properties
Properties of solutions that depend on the number of dissolved solute particles, not the identity of those particles. The four main colligative properties are vapour pressure lowering, boiling point elevation, freezing point depression, and osmotic pressure.
Isotonic
Two solutions with equal osmotic pressure (equal concentration of dissolved particles). A red blood cell placed in an isotonic solution maintains its shape.
Hypertonic
A solution with higher osmotic pressure (more dissolved particles) relative to another. Cells in a hypertonic environment lose water and shrink (crenation).
Hypotonic
A solution with lower osmotic pressure (fewer dissolved particles) relative to another. Cells in a hypotonic environment gain water and may swell or burst (lysis).
Colloid
A mixture in which dispersed particles are larger than those in a true solution but small enough not to settle out. Examples include milk, fog, and paint.
Micelle
A spherical arrangement of surfactant molecules in water. The hydrophobic tails point inward (away from water), and the hydrophilic heads point outward (toward water). The interior is nonpolar; the exterior is polar.
A solution is a homogeneous mixture of a solute (the substance dissolved) and a solvent (the substance doing the dissolving).
Saturated: maximum solute is dissolved. Unsaturated: less than the maximum. Supersaturated: more than the maximum is temporarily dissolved (unstable, will crystallise when disturbed).
Three types of interactions govern dissolution
Solute-solute interactions: forces holding solute particles together. These must be overcome (energy input).
Solvent-solvent interactions: forces holding solvent molecules together. These must be partially overcome (energy input).
Solute-solvent interactions: new attractions between solute and solvent. These form and release energy.
If solute-solvent interactions are comparable to or stronger than the sum of the other two, the solute dissolves readily. This is the basis of "like dissolves like": polar solutes dissolve in polar solvents, nonpolar solutes in nonpolar solvents.
Formula units and ions
An ionic compound's formula unit tells you the identity and number of each cation and anion produced on dissolution. For example, CaCl₂ produces one Ca²⁺ and two Cl⁻ per formula unit, giving three particles total.
This particle count matters for colligative properties later.
Molarity (M)
M = moles of solute / litres of solution
To calculate: convert mass of solute to moles (using molar mass), then divide by volume of solution in litres.
Mass percent
mass % = (mass of solute / mass of solution) × 100
Volume percent
volume % = (volume of solute / volume of solution) × 100
Parts per million
ppm = (mass of solute / mass of solution) × 10⁶
For dilute aqueous solutions, ppm is approximately mg solute per litre of solution.
Mass/volume percent
m/v % = (mass of solute in g / volume of solution in mL) × 100
Interconverting units
You may be asked to convert between these. The key is tracking what is in the numerator and denominator and using density if needed to bridge mass and volume.
Ion concentrations from molarity
If a 0.10 M solution of Na₂SO₄ fully dissociates, the Na⁺ concentration is 0.20 M (2 ions per formula unit) and the SO₄²⁻ concentration is 0.10 M.
Temperature: for most solid solutes, solubility increases with temperature. For gases dissolved in liquids, solubility decreases with increasing temperature.
Intermolecular forces: "like dissolves like." Polar/ionic solutes tend to dissolve in polar solvents. Nonpolar solutes dissolve in nonpolar solvents. The stronger the solute-solvent IMF relative to solute-solute and solvent-solvent IMFs, the greater the solubility.
Pressure: primarily affects gas solubility (Henry's law). Higher pressure forces more gas into solution.
Colligative properties depend on the number of dissolved solute particles, not their chemical identity. For this exam, understand the qualitative relationships (no calculations required).
Vapour pressure lowering
Adding a nonvolatile solute lowers the vapour pressure of the solvent. More particles dissolved means a greater lowering.
Boiling point elevation
A solution boils at a higher temperature than the pure solvent. More dissolved particles means a greater elevation.
Freezing point depression
A solution freezes at a lower temperature than the pure solvent. More dissolved particles means a greater depression. This is why salt is spread on icy roads.
Osmotic pressure
Osmotic pressure is the pressure needed to stop osmosis (the net flow of solvent through a semipermeable membrane from a dilute to a concentrated solution). More dissolved particles means higher osmotic pressure.
Particle count matters
An ionic compound that dissociates into more ions has a greater colligative effect per mole than a molecular (nonionic) solute. For example, 1 mol NaCl gives 2 mol particles; 1 mol CaCl₂ gives 3 mol particles; 1 mol glucose gives 1 mol particles.
Tonicity
Isotonic: same particle concentration on both sides of a membrane. Cells stay the same size.
Hypertonic: higher particle concentration outside the cell. Water flows out; cells shrink (crenation).
Hypotonic: lower particle concentration outside the cell. Water flows in; cells swell and may burst (lysis).
Colloids are mixtures with dispersed particle sizes roughly 1 to 1000 nm, larger than dissolved solute particles but small enough not to settle out under gravity.
Examples: milk (fat droplets in water), fog (water droplets in air), paint (pigment particles in liquid).
The Tyndall effect: colloids scatter light, making a beam visible as it passes through. True solutions do not scatter light this way.
Micelles
Formed by surfactant molecules (e.g., soap) in water.
Structure: hydrophobic tails cluster inward, away from water. Hydrophilic heads face outward, interacting with water.
The nonpolar interior can trap grease and oil, which is how soap cleans.
Applications in agriculture: delivery of hydrophobic pesticides. In health sciences: drug delivery systems, bile salts in digestion.
Molarity
M = \frac{\text{moles of solute}}{\text{litres of solution}}Mass percent
\text{mass \%} = \frac{\text{mass of solute}}{\text{mass of solution}} \times 100Volume percent
\text{volume \%} = \frac{\text{volume of solute}}{\text{volume of solution}} \times 100Parts per million
\text{ppm} = \frac{\text{mass of solute}}{\text{mass of solution}} \times 10^6Mass/volume percent
\text{m/v \%} = \frac{\text{mass of solute (g)}}{\text{volume of solution (mL)}} \times 100Micelle diagram (know how to draw)
Draw a circle. Place surfactant molecules around it with hydrophobic tails pointing inward (toward the centre) and hydrophilic heads pointing outward (toward the surrounding water). Label the nonpolar interior and the polar exterior. Any trapped grease or oil sits inside the nonpolar core.
IV fluids and tonicity. Hospitals use isotonic saline (0.9% NaCl) for intravenous drips because it matches the osmotic pressure of blood. A hypertonic IV would cause red blood cells to shrink; a hypotonic IV would cause them to swell and potentially burst.
Salting roads in winter. Spreading salt (NaCl or CaCl₂) on icy roads exploits freezing point depression. The dissolved salt lowers the freezing point of water, causing ice to melt at temperatures below 0 °C. CaCl₂ is more effective per mole than NaCl because it dissociates into three ions rather than two.
Colloids in agriculture. Pesticide delivery often uses colloidal suspensions or micelle-based formulations to carry hydrophobic active ingredients through water-based spray systems. The micelle's nonpolar interior dissolves the pesticide; the polar exterior keeps the whole assembly suspended in water.
Bile salts in digestion. Bile salts form micelles in the small intestine, emulsifying dietary fats into smaller droplets so that lipase enzymes can access and break them down more efficiently.
Students often think "like dissolves like" means identical molecules. It means matching polarity: polar dissolves in polar, nonpolar in nonpolar. Ethanol (polar) dissolves in water (polar) even though they are different molecules.
A frequent mistake is forgetting to count all the ions when determining colligative effects. NaCl produces 2 particles per formula unit, not 1. CaCl₂ produces 3. Glucose produces 1 (it does not dissociate).
Students sometimes confuse a saturated solution with a concentrated one. A saturated solution of a poorly soluble substance can still be very dilute. Saturated refers to reaching the solubility limit, not to having a lot of solute.
Many students assume colligative properties depend on what the solute is. They do not. A mole of NaCl particles and a mole of KBr particles produce the same freezing point depression if they generate the same number of dissolved particles.
⚠️ Concentration calculations (molarity from mass, volume, and molar mass) are heavily tested. Practise unit conversions until they are automatic.
⚠️ Be ready to calculate the concentration of individual ions from the molarity of an ionic compound. Know how many ions each formula unit produces.
⚠️ Qualitative colligative property questions: given two solutions, determine which has the lower freezing point or higher boiling point based on particle count.
⚠️ Isotonic, hypertonic, and hypotonic: know what happens to a cell in each type of solution.
⚠️ Be able to draw and describe a micelle, including the orientation of hydrophobic and hydrophilic parts and what the interior and exterior environments are like.
True or False: A supersaturated solution contains less solute than a saturated solution at the same temperature. Answer: False. It contains more.
Fill in the blank: Molarity is defined as moles of solute per ______ of solution. Answer: Litre (L).
True or False: Adding a nonvolatile solute to water raises its vapour pressure. Answer: False. It lowers the vapour pressure.
Fill in the blank: CaCl₂ dissociates into ______ particles per formula unit in water. Answer: 3 (one Ca²⁺ and two Cl⁻).
True or False: In a hypertonic solution, a red blood cell will swell. Answer: False. It will shrink (crenation). Cells swell in a hypotonic solution.
Q: You dissolve 5.85 g of NaCl (molar mass 58.44 g/mol) in enough water to make 500.0 mL of solution. What is the molarity of NaCl? What is the concentration of Na⁺ ions? Of Cl⁻ ions?
A: Moles of NaCl = 5.85 / 58.44 = 0.100 mol. Molarity = 0.100 mol / 0.5000 L = 0.200 M. NaCl dissociates into one Na⁺ and one Cl⁻, so [Na⁺] = 0.200 M and [Cl⁻] = 0.200 M.
Q: Which solution has the lowest freezing point: 0.10 M glucose, 0.10 M NaCl, or 0.10 M CaCl₂? Explain.
A: 0.10 M CaCl₂. Glucose does not dissociate (0.10 mol particles). NaCl gives 0.20 mol particles. CaCl₂ gives 0.30 mol particles. More particles means a greater freezing point depression.
Q: A student adds a small crystal to a clear solution and observes rapid crystallisation. Was the original solution saturated, unsaturated, or supersaturated?
A: Supersaturated. The seed crystal triggered the excess dissolved solute to come out of solution.
Q: Draw a simple micelle in water. Label the hydrophobic tails, hydrophilic heads, nonpolar interior, and polar exterior.
A: Draw a roughly spherical arrangement. Wavy lines (tails) point inward toward the centre, labelled "hydrophobic tails" and "nonpolar interior." Circles (heads) line the outside, labelled "hydrophilic heads." The surrounding area is labelled "water / polar exterior."
Q: Explain, in terms of the three types of interactions, why NaCl dissolves in water but not in hexane.
A: NaCl has strong ion-ion interactions (solute-solute). Water has strong hydrogen bonding (solvent-solvent). The ion-dipole interactions between Na⁺/Cl⁻ and water (solute-solvent) are strong enough to compensate. Hexane is nonpolar, so ion-induced dipole interactions are far too weak to overcome the ionic lattice energy of NaCl.
This connects to Chapter 11 because the "like dissolves like" principle is a direct consequence of intermolecular force matching. Ion-dipole forces, introduced in Ch. 11, are the primary driver of ionic compound dissolution covered here.
Colligative properties connect forward to equilibrium concepts (Ch. 14): vapour pressure lowering is an equilibrium phenomenon, and osmotic pressure calculations share a structural similarity with the ideal gas law.
Micelles and colloids connect to biochemistry and health sciences: bile salt micelles in digestion, liposomal drug delivery, and emulsion chemistry in food science.
Solutions, solubility, saturated, unsaturated, supersaturated, molarity, moles per litre, mass percent, volume percent, ppm, parts per million, mass/volume percent, concentration units, like dissolves like, solute-solvent interactions, ion-dipole, dissolution, colligative properties, vapour pressure lowering, boiling point elevation, freezing point depression, osmotic pressure, osmosis, isotonic, hypertonic, hypotonic, crenation, lysis, colloid, Tyndall effect, micelle, surfactant, hydrophobic, hydrophilic, formula unit, ion concentration, CHM 11200, Purdue general chemistry, Chapter 12