AP Chemistry 3.4: Solutions, Mixtures, and Solubility

• A mixture contains two or more substances physically combined rather than chemically bonded into a single pure substance.

• A homogeneous mixture has uniform composition throughout at the macroscopic level.

  • A solution is a homogeneous mixture.

• A heterogeneous mixture has a nonuniform composition or distinguishable regions/components.

• In a solution, the solvent is the medium doing the dissolving and is often the component present in greater amount.

• The solute is the substance dissolved in the solvent.

  • Example: in ordinary salt water, water is the solvent and NaCl is the solute.

  • Solutions can involve different phases; they are not limited to solids dissolved in liquids.

• Concentration describes how much solute is present relative to a certain quantity of solution.

• One of the most important concentration units in AP Chemistry is molarity :

  • M = moles of solute / liters of solution

  • Units are mol/L, commonly abbreviated M.

  • Example: 0.500 mol NaCl in enough solution to make 2.00 L:

  • M = 0.500/2.00

  • = 0.250 M

• Molarity uses the total solution volume , not simply the amount of solvent added.

• To find moles from molarity:

  • moles = M × V

  • where V is in liters.

  • Example: 250.0 mL of 0.400 M NaCl.

  • Convert 250.0 mL → 0.2500 L.

  • moles = 0.400 × 0.2500

  • = 0.100 mol

• Dilution occurs when additional solvent is added to decrease concentration.

  • During simple dilution, the number of moles of solute stays constant.

Therefore:

• M₁V₁ = M₂V₂

  • Example: 50.0 mL of 2.00 M solution is diluted to 200.0 mL.

  • (2.00)(50.0) = M₂(200.0)

  • M₂ = 0.500 M

  • The formula works because the amount of solute before dilution equals the amount after dilution.

  • Do not confuse dilution with a chemical reaction. Adding water to dilute a solution does not by itself remove moles of solute.

  • Khan Academy's current solutions lesson covers types of mixtures, molarity, and dilution.

• When solving solution calculations, always check whether volume must be converted from mL to L.

  • Example: 35.0 mL = 0.0350 L.

• Molarity can be used to connect solution volume to stoichiometry in later units because it converts liters of solution into moles of solute.

• AP Chemistry often expects students to connect symbolic equations, numerical concentration, and particulate diagrams .

  • A particulate representation shows individual molecules, atoms, or ions in a solution.

  • The diagram should reflect both the identity of the dissolved particles and their correct ratios .

• A soluble ionic compound can dissociate into ions when dissolved in water.

  • For example:

  • NaCl(aq) → Na⁺(aq) + Cl⁻(aq)

  • A particle diagram of dissolved NaCl should therefore show separated Na⁺ and Cl⁻ particles rather than intact NaCl "molecules" if complete dissociation is being represented.

• The ratio must match the chemical formula.

  • NaCl produces a 1:1 ratio of Na⁺ to Cl⁻.

  • CaCl₂ produces:

  • CaCl₂ → Ca²⁺ + 2Cl⁻

  • so the diagram must contain twice as many chloride ions as calcium ions .

  • AlCl₃ produces a 1:3 ratio of Al³⁺ to Cl⁻ if complete dissociation is being represented.

  • Charge must still be balanced overall.

• For equal solution volumes, a diagram with more dissolved solute particles generally represents a higher particle concentration.

  • However, you cannot compare concentration only from the number of particles unless the represented volumes are also considered.

  • Twenty particles in twice the volume can have the same concentration as ten particles in half the volume.

• Water orientation may also appear in particulate diagrams.

  • Around a cation, the oxygen side of water points toward the positive ion.

  • Around an anion, the hydrogen side points toward the negative ion.

  • This represents ion-dipole attractions and solvation.

• Solvation refers to solvent particles surrounding and stabilizing dissolved solute particles.

  • When water is the solvent, the process is often called hydration .

  • Khan Academy's current Unit 3 specifically includes representing solutions using particulate models.

• Mixtures can often be separated because their components have different physical properties .

  • Khan Academy's current Unit 3 covers distillation, distillation curves, thin-layer chromatography, retention factors, and column chromatography in this topic.

• Distillation separates substances mainly based on differences in volatility or boiling point .

  • A liquid mixture is heated.

  • The more volatile component, usually the one with the lower boiling point, enters the vapor phase more readily.

  • The vapor can then be cooled in a condenser.

  • It condenses back into liquid and is collected separately.

  • Distillation therefore involves:

  • vaporization,

  • separation in the vapor phase,

  • condensation.

  • A substance with weaker intermolecular forces often has a lower boiling point and greater volatility than a comparable substance with stronger IMFs.

  • Therefore IMF concepts connect directly with distillation.

• Simple distillation can work well when boiling points differ substantially or when separating a volatile liquid from a nonvolatile material.

  • If two liquids have similar boiling points, separation by ordinary simple distillation is less effective.

• A distillation curve commonly plots temperature against the amount/time of distillation.

  • Regions where temperature remains near a component's boiling point can provide information about the substances being collected.

  • Changes in temperature can indicate changes in vapor composition.

• Chromatography separates mixture components based on their different relative attractions for a stationary phase and a mobile phase .

  • The stationary phase stays in place.

  • The mobile phase moves through or across the stationary phase.

  • A component that interacts more strongly with the mobile phase generally travels farther or moves faster.

  • A component that interacts more strongly with the stationary phase tends to travel less or move more slowly.

  • The key idea is competition between solute-stationary and solute-mobile interactions .

• Thin-layer chromatography (TLC) uses a thin stationary layer, commonly on a plate, while a solvent travels upward through the material.

  • Different substances travel different distances because they have different affinities for the stationary and mobile phases.

• The retention factor , Rf, can be calculated:

  • Rf = distance traveled by solute / distance traveled by solvent front

  • The distances must be measured from the same starting line.

  • Example: solute travels 4.0 cm, solvent front travels 8.0 cm.

  • Rf = 4.0/8.0 = 0.50

  • Under ordinary TLC conditions, Rf values normally lie between 0 and 1 because the solute cannot travel farther than the solvent front.

  • A higher Rf means the substance traveled farther relative to the solvent front under those experimental conditions.

  • Rf values depend on the solvent, stationary phase, temperature, and other experimental conditions, so an Rf is not a permanent universal constant for a compound.

• Column chromatography uses a packed stationary phase inside a column.

  • The mobile phase travels through the column.

  • Components that interact weakly with the stationary phase move through more rapidly and elute earlier.

  • Components that interact more strongly with the stationary phase move more slowly and elute later.

• Polarity is often important because polar substances interact differently with polar and nonpolar stationary/mobile phases.

  • Never reduce chromatography to "polar travels farther" without knowing which phase is polar. The result depends on relative attraction to both phases .

• Solubility describes the amount of a substance that can dissolve in a given solvent under specified conditions.

• Dissolving requires interactions among:

  • solute particles,

  • solvent particles,

  • and newly formed solute-solvent interactions.

  • To dissolve a solute, some existing solute-solute attractions and solvent-solvent attractions must be disrupted.

  • New solute-solvent attractions then form.

  • Whether dissolving is favorable depends on the balance of these interactions and additional thermodynamic factors.

• A useful introductory rule is "like dissolves like."

  • Polar substances often dissolve better in polar solvents because they can form favorable dipole-dipole, hydrogen-bond, or ion-dipole interactions.

  • Nonpolar substances often dissolve better in nonpolar solvents because both rely strongly on dispersion interactions.

  • Ionic compounds may dissolve well in strongly polar solvents such as water when ion-dipole attractions between solvent molecules and ions are strong enough to stabilize the separated ions.

  • Example: water surrounding Na⁺ and Cl⁻ can form strong ion-dipole attractions.

  • However, not every ionic compound is highly soluble in water. Strong crystal-lattice attractions can sometimes outweigh the benefits of hydration.

  • "Like dissolves like" is therefore a prediction tool , not a perfect universal rule.

  • Khan Academy's current solubility section specifically focuses on solubility and intermolecular forces.

• Molecules that can hydrogen-bond strongly with water are often more water-soluble than similar molecules that cannot.

  • But molecular size matters.

  • As the nonpolar hydrocarbon portion of a molecule becomes larger, its water solubility can decrease even if the molecule still contains one polar functional group.

  • In comparing two molecules, consider both:

  • the polar/hydrogen-bonding regions,

  • and the size of the nonpolar regions.

• Increasing temperature can alter solubility, but the direction depends on the system.

  • Many solids become more soluble in water as temperature rises, but this is not universally true .

  • Gas solubility in liquids often decreases as temperature rises because higher kinetic energy makes it easier for dissolved gas particles to escape.

• Gas solubility generally increases when the partial pressure of that gas above the solution increases; this relationship is formalized by Henry's law, though the exact emphasis depends on the course question.

• For AP-style explanations, do not simply say "because it is polar." State what intermolecular attraction forms and why it is favorable.

Example:

• Instead of:

  • "ethanol dissolves because it is polar."

  • Better:

  • "ethanol can form hydrogen-bonding interactions between its O—H group and water molecules, making favorable solute-solvent attractions possible."