Source: Comprehensive Chemistry Study Guide, General Chemistry (Purdue University)
Tags: phase changes, evaporation, condensation, sublimation, deposition, melting, freezing, heating curve, cooling curve, enthalpy of vaporisation, enthalpy of fusion, solutions, solubility, saturated, unsaturated, supersaturated, molarity, ppm, mass percent, concentration units
Difficulty: Intermediate Prerequisites: Intermolecular forces, basic thermochemistry (q = mcΔT), mole concept, and unit conversions.
Phase changes and solutions sit on either side of the same coin: both are governed by the intermolecular forces you studied in the previous section. Understanding how substances move between solid, liquid, and gas phases tells you about energy flow in physical processes, while understanding solutions tells you how substances mix at the molecular level and how we quantify that mixing. Together, these topics set the stage for colligative properties and chemical equilibrium. If you are comfortable with heating curves and concentration calculations, the later chapters will feel far more manageable.
Every phase change involves breaking or forming intermolecular attractions, which means absorbing or releasing energy. Solutions form when solute-solvent interactions are strong enough to overcome solute-solute and solvent-solvent forces, and we express "how much solute" using units like molarity, ppm, and mass percentage, each suited to different practical contexts.
Evaporation (vaporisation)
The transition of a substance from liquid to gas, occurring at the surface of the liquid at any temperature. An endothermic process: molecules must absorb energy to overcome IMFs and escape into the vapour phase.
Think of it as molecules at the surface getting lucky with enough kinetic energy to break free.
Condensation
The transition from gas to liquid. An exothermic process: gas molecules release energy as they form intermolecular attractions with the liquid.
Melting (fusion)
The transition from solid to liquid. Endothermic: energy is absorbed to partially overcome the ordered intermolecular interactions in the solid.
Freezing (solidification)
The transition from liquid to solid. Exothermic: energy is released as molecules settle into an ordered arrangement.
Sublimation
The direct transition from solid to gas, bypassing the liquid phase. Endothermic. Dry ice (solid CO₂) is the classic example.
Deposition
The direct transition from gas to solid. Exothermic. Frost forming on a cold surface is an everyday example.
Enthalpy of vaporisation (ΔH_vap)
The energy required to convert one mole of a substance from liquid to gas at its boiling point. Larger ΔH_vap means stronger IMFs in the liquid.
Enthalpy of fusion (ΔH_fus)
The energy required to convert one mole of a substance from solid to liquid at its melting point. Always smaller than ΔH_vap for the same substance, because melting only partially disrupts IMFs.
Heating curve
A graph of temperature versus heat added for a substance. Flat (horizontal) segments correspond to phase changes, where added heat breaks IMFs rather than raising temperature. Sloped segments correspond to heating within a single phase.
Solubility
The maximum amount of solute that dissolves in a given quantity of solvent at a specified temperature, producing a saturated solution.
In simple terms, it is the "capacity" of a solvent for a particular solute under defined conditions.
Saturated solution
A solution that contains exactly the maximum amount of dissolved solute at a given temperature. Any additional solute will not dissolve.
Unsaturated solution
A solution containing less solute than the solvent can dissolve at that temperature. More solute can still be added.
Supersaturated solution
A solution that temporarily holds more dissolved solute than its saturation limit. Supersaturated solutions are unstable: a seed crystal or disturbance can trigger rapid crystallisation of the excess solute.
Think of it as a solution that is "over capacity" and just waiting for an excuse to dump the extra solute out.
Molarity (M)
Moles of solute per litre of solution. The most commonly used concentration unit in general chemistry.
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, such as trace contaminants in water.
Mass percentage (mass %)
Mass of solute divided by total mass of solution, multiplied by 100. Temperature-independent, which makes it useful for industrial and commercial labelling.
Volume percentage (volume %)
Volume of solute divided by total volume of solution, multiplied by 100. Common for liquid-liquid mixtures such as ethanol in beverages.
Six phase transitions exist, grouped into three endothermic/exothermic pairs:
Melting ↔ Freezing
Evaporation ↔ Condensation
Sublimation ↔ Deposition
Endothermic transitions (melting, evaporation, sublimation) absorb heat from the surroundings.
Exothermic transitions (freezing, condensation, deposition) release heat to the surroundings.
Sloped regions: Temperature changes as heat is added or removed. Use q = mcΔT, where m is mass, c is specific heat capacity for that phase, and ΔT is the temperature change.
Flat (plateau) regions: Temperature stays constant while phase change occurs. Use q = nΔH (n = moles, ΔH = enthalpy of fusion or vaporisation).
The plateau at the melting point is shorter than the plateau at the boiling point for the same substance, because ΔH_vap > ΔH_fus.
On a cooling curve the logic reverses: the substance releases heat at each plateau.
During melting, particles gain enough energy to move out of fixed lattice positions but remain close together.
During evaporation, particles gain enough energy to completely overcome IMFs and separate into the gas phase.
During a phase change, kinetic energy (and therefore temperature) does not increase, because the added energy goes into overcoming potential energy (IMFs) rather than speeding up particles.
Dissolution involves three energy steps:
Separating solute particles (endothermic, overcoming solute-solute attractions).
Separating solvent particles (endothermic, overcoming solvent-solvent attractions).
Forming solute-solvent interactions (exothermic).
If the energy released in step 3 is comparable to or greater than the energy consumed in steps 1 and 2, the substance dissolves readily.
"Like dissolves like": Polar and ionic solutes tend to dissolve in polar solvents. Nonpolar solutes dissolve in nonpolar solvents.
When an ionic compound dissolves in water, ion-dipole forces pull ions away from the crystal lattice.
Each cation gets surrounded by water molecules with the oxygen (partial negative) pointing inward.
Each anion gets surrounded by water molecules with the hydrogens (partial positive) pointing inward.
This shell of oriented water molecules around an ion is called the hydration shell (or solvation shell).
Molarity: M = moles of solute / litres of solution.
To find moles from mass: n = mass / molar mass.
Dilution shortcut: M₁V₁ = M₂V₂.
Parts per million: ppm = (mass of solute / mass of solution) × 10⁶.
For dilute aqueous solutions, density ≈ 1 g/mL, so ppm ≈ mg/L.
Mass percentage: mass % = (mass of solute / mass of solution) × 100.
Volume percentage: vol % = (volume of solute / volume of solution) × 100.
Mass/volume percentage: (mass of solute in g / volume of solution in mL) × 100.
Converting between units often requires knowing the density of the solution and the molar mass of the solute.
Temperature:
For most solid solutes in liquid solvents, solubility increases with temperature.
For gases dissolved in liquids, solubility decreases with temperature (which is why a warm fizzy drink goes flat faster).
Intermolecular forces: Stronger solute-solvent attractions relative to solute-solute and solvent-solvent attractions favour dissolving.
Nature of solute and solvent: Polar solvents dissolve polar/ionic solutes. Nonpolar solvents dissolve nonpolar solutes.
Pressure (for gases): Henry's law states that gas solubility is directly proportional to the partial pressure of the gas above the solution. S = k_H × P.
Heat calculations within a single phase:
q = mcΔT
where q = heat (J), m = mass (g), c = specific heat capacity (J/g·K), ΔT = temperature change (K or °C).
Heat calculations during a phase change:
q = nΔH_fus (melting/freezing) or q = nΔH_vap (boiling/condensation)
where n = moles.
Molarity:
M = mol solute / L solution
Dilution:
M₁V₁ = M₂V₂
Henry's law (gas solubility):
S = k_H × P
where S = solubility, k_H = Henry's law constant, P = partial pressure of the gas.
Heating curves are exactly what engineers use to size industrial boilers and refrigeration systems: knowing how much energy each phase transition requires determines the capacity of the equipment. Concentration units show up everywhere from medicine (IV drip concentrations in mass/volume %) to environmental science (pollutant levels in ppm) to brewing (alcohol content in volume %). Henry's law explains why divers must ascend slowly: dissolved nitrogen comes out of solution as pressure drops, and rising too fast produces dangerous gas bubbles in the blood (decompression sickness).
"Temperature always rises when you add heat." During a phase change, temperature remains constant even as heat is added. The energy goes into breaking intermolecular forces, not increasing kinetic energy.
"Supersaturated solutions are very concentrated, so they must be stable." They are the opposite of stable. They hold more solute than equilibrium allows and will crystallise readily when disturbed.
"Molarity and molality are the same thing." Molarity (M) is moles per litre of solution; molality (m) is moles per kilogram of solvent. They differ whenever the solution's density is not exactly 1 kg/L, and molality is temperature-independent because it uses mass rather than volume.
"All solutes become more soluble as temperature rises." This is true for most solids in water, but gas solubility in liquids decreases with temperature.
⚠️ Heating curve problems are a staple of general chemistry exams. Be ready to calculate the total energy needed to take a substance from one temperature and phase to another, combining q = mcΔT segments with q = nΔH segments.
⚠️ Know how to convert between concentration units. A common exam question gives you mass %, density, and molar mass and asks for molarity.
⚠️ Distinguish saturated, unsaturated, and supersaturated solutions. If a problem says "excess solid remains at the bottom," the solution is saturated.
⚠️ Ion concentrations in solutions of ionic compounds: dissolving 0.10 M CaCl₂ gives 0.10 M Ca²⁺ and 0.20 M Cl⁻. Do not forget to multiply by the subscript.
True or false: Sublimation is an exothermic process.
Fill in the blank: On a heating curve, the flat regions correspond to __________.
True or false: A supersaturated solution contains less solute than its saturation limit.
Fill in the blank: Molarity is defined as moles of solute per __________ of solution.
True or false: Gas solubility in water generally increases as temperature increases.
Answers: 1. False (sublimation is endothermic). 2. Phase changes (where temperature is constant). 3. False (it contains more). 4. Litre (L). 5. False (gas solubility decreases with temperature).
Q: Calculate the total energy required to convert 36.0 g of ice at 0 °C to steam at 100 °C. (ΔH_fus = 6.02 kJ/mol, ΔH_vap = 40.7 kJ/mol, c_water = 4.184 J/g·°C.)
A: First, find moles: 36.0 g / 18.0 g/mol = 2.00 mol. Melting: q₁ = 2.00 mol × 6.02 kJ/mol = 12.0 kJ. Heating liquid water from 0 °C to 100 °C: q₂ = 36.0 g × 4.184 J/g·°C × 100 °C = 15,062 J ≈ 15.1 kJ. Vaporisation: q₃ = 2.00 mol × 40.7 kJ/mol = 81.4 kJ. Total = 12.0 + 15.1 + 81.4 = 108.5 kJ.
Q: A solution is prepared by dissolving 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 Na⁺ ions?
A: Moles of NaCl = 5.85 / 58.44 = 0.1001 mol. Molarity of NaCl = 0.1001 mol / 0.5000 L = 0.200 M. NaCl dissociates completely: NaCl → Na⁺ + Cl⁻. So [Na⁺] = 0.200 M.
Q: Explain at the molecular level why temperature remains constant during melting even though heat is being added.
A: The added energy is used to overcome the intermolecular forces holding particles in the rigid solid lattice. It increases the potential energy of the system rather than the kinetic energy, so the temperature (which reflects average kinetic energy) stays constant until the entire solid has melted.
Q: A student opens a bottle of fizzy water and leaves it on a warm windowsill. The next day, the water is completely flat. Explain using two relevant principles.
A: Once the bottle is opened, the partial pressure of CO₂ above the liquid drops (Henry's law: lower pressure means lower gas solubility). The warm temperature further decreases gas solubility. Both effects drive dissolved CO₂ out of solution.
Phase changes connect back to intermolecular forces (ΔH_vap and ΔH_fus reflect IMF strength) and forward to colligative properties (adding solute disrupts the solvent's phase-change behaviour, raising the boiling point and lowering the freezing point). Concentration units are essential for stoichiometry in solution, colligative property calculations, and equilibrium expressions (where molarity appears in K expressions).
phase change, phase transition, evaporation, vaporisation, condensation, melting, fusion, freezing, sublimation, deposition, heating curve, cooling curve, enthalpy of vaporisation, enthalpy of fusion, heat of vaporisation, heat of fusion, specific heat, solution, solute, solvent, saturated, unsaturated, supersaturated, molarity, parts per million, ppm, mass percent, volume percent, concentration, dilution, Henry's law, like dissolves like, hydration shell, solvation, ion-dipole