Chapter 6: Solutions

A solution is a homogeneous mixture of two or more substances. Its composition is the same throughout. A phase is a part of matter with uniform properties.

The substance dissolved in a solution is the solute. In simple examples, it is usually present in a smaller amount.

The substance doing the dissolving is the solvent. It is often present in a larger amount and generally determines the solution's state.

If water is the solvent, the solution is aqueous, shown by (aq); e.g. NaCl(aq).

Water dissolves many substances, so it is often called a universal solvent. It does not dissolve everything. Most oils do not dissolve in water.

A solution is a mixture. Its substances keep their identities, and their amounts can vary.

An unsaturated solution can dissolve more solute at the given temperature.

A saturated solution contains the maximum amount of dissolved solute at the given temperature. Any extra solute remains undissolved.

A supersaturated solution contains more dissolved solute than a saturated solution at the same temperature. It is not fully stable. Disturbing it or adding a small crystal may cause the extra solute to form crystals.

Dissolve a suitable solute in hot solvent until the solution is saturated.

Let it cool slowly. Avoid shaking it or adding seed crystals.

If more solute stays dissolved than normally allowed at the lower temperature, the solution is supersaturated.

Heating alone does not prove supersaturation. Neither do crystals at the bottom. Supersaturation means extra solute is still dissolved at the stated temperature.

A dilute solution has relatively little dissolved solute; a concentrated one has relatively more. Either can be unsaturated or saturated depending on the solute.

When only solvent is added, concentration decreases while the amount of dissolved solute stays unchanged.

M₁V₁ = M₂V₂.

1 means the starting solution. 2 means the diluted solution. V₂ is the final volume after dilution.

Fog and smoke contain tiny droplets or particles in air. They are not true solutions. Butter and cheese are also complex mixtures, not simple solid solutions.

Concentration in g/dm³ = mass of solute in grams ÷ volume of FINAL solution in dm³.

Use the final volume of the whole solution. Do not simply use the volume of water added. Dissolving and mixing can change the volume.

Concentration tells us how much is dissolved now. Solubility tells us how much can dissolve at the stated temperature and pressure.

% (m/m) = (mass of solute / mass of solution) × 100. Example: 5 g sugar in 100 g solution = 5% m/m.

% (m/v) = mass of solute (g) / FINAL solution volume (mL) × 100. Example: 5 g solute per 100 mL solution = 5% m/v.

% (v/v) = volume of solute / FINAL volume of solution × 100. Example: 5 mL ethanol in 100 mL solution = 5% v/v.

Volume by mass expresses the volume of solute in mL per 100 g of solution. Write these units clearly because volume and mass are different quantities.

Measure about 58.5 g NaCl: approximately 1 mole. Dissolve it in some water. Transfer it to a volumetric flask. Add water until the final solution volume is 1.000 L.

Concentration is the amount dissolved now. Solubility is the maximum amount that can stay dissolved at equilibrium.

As a useful guide, polar and ionic solutes tend to dissolve in polar solvents, while nonpolar solutes tend to dissolve in nonpolar solvents.

Some attractions between solute particles must be overcome so they can separate.

Some solvent particles must move apart too.

New attractions form between solute and solvent particles.

Dissolving depends on energy changes and how particles spread out. Attractions matter, but comparing attractions alone does not always tell us whether a substance will dissolve.

Water has slightly charged ends. Its oxygen end, δ−, attracts Na⁺. Its hydrogen ends, δ+, attract Cl⁻. Water surrounds the separated ions and helps keep them in solution.

Gasoline is mostly nonpolar. Water is polar and forms strong attractions with other water molecules. The two liquids mostly stay in separate layers.

Ethanol’s –OH group forms hydrogen bonds with water. This helps ethanol mix with water.

Heating lets many solid solutes dissolve in larger amounts. But some do not follow this trend. Dissolving faster and dissolving more are different ideas.

At fixed gas pressure, most gases become less soluble in water as temperature rises.

At fixed temperature, higher pressure of a gas above a liquid usually makes more of that gas dissolve. This is Henry’s law. The pressure due to that gas alone is its partial pressure.

A sealed fizzy drink holds CO₂ under pressure. Opening it lowers the pressure. Some CO₂ then escapes as bubbles.

A homogeneous mixture of molecules or ions, with solute particle size generally below about 1 nm. It is normally transparent without visible Tyndall scattering.

A colloid has particles roughly 1–1000 nm in size spread through another substance. Its particles are not mixed at the molecular level, so it is heterogeneous, though it may look uniform. It can remain stable for a long time and often scatters light.

A heterogeneous dispersion with larger particles, frequently above 1000 nm. Particles may settle on standing and are often trapped by ordinary filter paper.

Particle sizes are useful guides, not exact limits for every mixture. Some colloids look transparent. Filter behaviour can also vary.

Colloid particles scatter light sideways. Some light reaches our eyes, so we see the beam’s path. Salt and sugar solutions usually scatter too little light to show a bright path.

Milk has tiny fat droplets spread through water. This type of colloid is an emulsion.

Fog has tiny water droplets spread through air. This is an aerosol.

Smoke often has tiny solid particles or liquid droplets spread through air.

Paint particles may settle while the paint is stored.

Stirring spreads the separated particles through the paint again.

Glucose in water is a true solution. Starch paste often behaves as a colloid.

A solution is homogeneous because its components mix uniformly at the molecular or ionic scale.

Nitrogen is considered the solvent component of air because it is the most abundant gas (~78% of dry air).

An aqueous solution has water as solvent.

2 M solution is more dilute than 3 M solution of the same solute.

Unsaturated: can dissolve more at current temperature. Saturated: equilibrium limit reached. Supersaturated: more dissolved than equilibrium limit at same temperature.

10 g sugar + 140 g water gives 10/150 × 100 = 6.67% by mass.

A 10% m/m solution can be made from 10 g sugar + 90 g water (100 g final mixture).

1.2 M NaOH contains 1.2 × 40 = 48 g NaOH per dm³.

Milk and fog are colloids, not true homogeneous solutions.

Solubility depends on temperature and solvent; do not assume every solid solute increases in solubility with heating.

A saturated solution may contain undissolved solid, but this is not the definition of a supersaturated solution.

Adding more salt does not by itself prove that a solution is saturated. Check the temperature, solubility and whether extra solid remains after enough time for dissolving.

Water molecules attract Na⁺ and Cl⁻ ions and surround them. The oxygen side of water faces Na⁺, while the hydrogen side faces Cl⁻. Attraction between an ion and a polar molecule is called an ion–dipole attraction.

Saturated reaches the equilibrium solubility limit at specified temperature; unsaturated has room to dissolve more.

A solution is a homogeneous mixture made from solute (dissolved material) and solvent (dissolving medium).

Grams of solute per 100 mL of FINAL solution: % m/v = (mass g / solution mL) × 100.

Colloidal particles scatter visible light.

Dilution: lower concentration by adding solvent. Concentration: amount of solute per defined amount. Solubility: maximum equilibrium dissolvable amount. Molarity: moles per litre of solution.

Because polar solvents can stabilize solute ions or molecules through ion–dipole attractions, dipole interactions and hydrogen bonding.

The nonpolar solvent may not provide enough favourable interactions to compensate for disrupting those in the original solute.

Solutions are used in medicines, intravenous fluids, drinks, cleaning, plant nutrients and industry.

Polar water molecules surround Na⁺ and Cl⁻ ions and help separate them from the crystal. Surrounding dissolved particles with water molecules is called hydration.

N₂ is the major component of dry air (about 78% by volume).

Gasoline is mainly nonpolar and water polar; their interactions do not favour mixing.

Calculate stock volume using M₁V₁ = M₂V₂, measure carefully, transfer to a volumetric flask and add water until the FINAL target volume is reached. Mix well. For corrosive acids, follow proper laboratory precautions under trained supervision.

Solubility is the maximum equilibrium amount dissolved at a given temperature and pressure. Dissolving requires disrupting solute–solute and solvent–solvent forces and creating new solute–solvent interactions. Both energy changes and the spreading of particles help determine whether mixing occurs.

For many salts, heating favours greater equilibrium dissolution; the effect depends on the energy changes involved, so it does not occur for every salt. Increased dissolving SPEED is distinct from increased equilibrium solubility.

The δ− oxygen side of water faces Na⁺; the δ+ hydrogen side faces Cl⁻. The ions become hydrated and disperse in solution.

Ionic and polar substances often dissolve in polar solvents; nonpolar substances often mix with nonpolar solvents. This is a helpful guideline with exceptions.

A true solution is homogeneous with molecular/ionic particles. A colloid is a heterogeneous dispersion with particles often showing Tyndall effect. A suspension has relatively large particles that tend to settle and can often be filtered.