Why Some Things Dissolve and Others Do Not
Dissolving separates solute particles and brings them into contact with a solvent. Attractions and the tendency of particles to mix affect solubility. Polarity helps explain why salt and sugar dissolve in water while oil forms a separate layer.
What a learner can do afterwards
- Explain dissolving as an exchange of interactions rather than a disappearance
- Predict whether a given solute will dissolve in water or in a non-polar solvent, and say why
- Use supplied solubility limits to calculate the mass crystallised when an initially saturated solution cools to equilibrium, with unchanged water mass and water-free crystals
1 · Read
You have compared attractions between molecules. Dissolving also involves a solvent, the liquid doing the dissolving. Sodium chloride contains ions held in a crystal. Polar water molecules attract these ions, separate them from the crystal and surround them. Attraction between an ion and a polar molecule is called ion-dipole attraction. The ions already existed in the solid. Water does not create them.
Essentially all dissolved sodium chloride separates into mobile ions, making it a strong electrolyte. Strong does not mean concentrated: conductivity also depends on how many ions are present. Sugar remains as neutral molecules and adds almost no conductivity to pure water. The solution's conductivity is not exactly zero. Compare similar concentrations with a supervised, battery-powered conductivity tester. Never use mains electricity or taste unknown solutions.
Dissolving changes contacts among solute and solvent particles. Attractions and the tendency of particles to mix both matter. Water often dissolves ionic or polar solutes, but some ionic solids barely dissolve. Sugar can hydrogen-bond with water. Oil-water contacts do not favour extensive mixing, although oil and water do attract. Nonpolar solutes often dissolve better in nonpolar solvents. Similar polarity is a guide, not a guarantee.
A saturated solution is at its equilibrium solubility limit. A supersaturated solution contains more dissolved solute than that limit at its temperature. Cooling can leave this temporary state without crystals forming. A seed crystal provides a surface for growth. In a supervised sodium acetate 'hot ice' demonstration, crystals can grow rapidly and release heat. This is an explanation to watch, not a home recipe. Hot liquids can burn; do not touch the liquid.
A solubility curve plots temperature horizontally and the equilibrium solubility limit vertically. Check units, often grams of solute per 100 g water. For cooling, keep the water mass fixed. Find the cold capacity for that mass. If the actual initial dissolved mass is at or below this capacity, no solute crystallises. Otherwise subtract the cold capacity from that mass. This gives the crystal mass at equilibrium if no water is lost or enters the crystals.
The solubility limits are 50 g at 60°C and 30 g at 20°C per 100 g water. A hot saturated solution contains 50 g solute and 100 g water. It cools without water loss. At equilibrium, 30 g remains dissolved and 50 - 30 = 20 g crystallises, with no water in the crystals. For an unsaturated start, compare the actual dissolved mass with the cold capacity first. If it does not exceed that capacity, all the solute stays dissolved.
Two readings from a solubility curve, not the whole curve. Both use 100 g of water.
Dissolving leaves solute particles in a solvent. Polarity guides predictions. Cooling calculations need stated units, a fixed solvent mass and a final equilibrium state.
Written for this dataset using chemistry and sodium acetate records as factual references. No source prose or preparation procedure was adapted.
The readings and cooling-calculation flashcard were written for this dataset. The readings use OpenStax Chemistry 2e and wikiHow's sodium acetate article as factual references. The other flashcards draw on imported facts-collection material.
More to read
Which dissolved particles carry charge?
An electrical current needs moving charge. In a salt solution, mobile ions carry charge through the liquid.
Sodium chloride contains ions before it enters water. Water molecules attract and surround these ions as they separate from the crystal. This contact between an ion and a polar molecule is called ion-dipole attraction. Dissolving frees the ions to move through the solution.
Essentially all the dissolved sodium chloride is present as separated ions, so sodium chloride is a strong electrolyte. Strong describes the dissolved substance's production of ions. It does not mean concentrated, nor does it guarantee the same conductivity at every concentration.
A weak electrolyte forms ions from only part of the dissolved substance. A nonelectrolyte such as sugar supplies no ions from its dissolved molecules. Pure water already conducts very poorly, so adding pure sugar does not make conductivity exactly zero.
Solubility is a separate issue. Even an ionic compound that dissolves only slightly can be a strong electrolyte if its dissolved portion separates into ions. To compare conductivity, control concentration and temperature and use suitable supervised low-voltage equipment.
When crystals begin to grow
A solubility limit applies at a stated temperature. At equilibrium, a saturated solution contains the limiting dissolved amount. An unsaturated solution contains less. A supersaturated solution temporarily contains more.
Cooling can produce supersaturation if crystals do not begin growing as the temperature falls. Crystallisation is therefore not always immediate. A suitable seed crystal offers a surface where growth can start.
Sodium acetate is used in supervised demonstrations called hot ice. It can be produced by reacting the acetic acid in vinegar with sodium bicarbonate. Carbon dioxide and water also form. This explains the source of the salt; it is not a preparation recipe. Heating and handling belong to the supervising adult, and the liquid should not be touched.
A prepared supersaturated sodium acetate solution can crystallise rapidly when seeded. The process releases energy and warms the mixture. The growing solid is not ordinary water ice.
For a cooling calculation, find the final equilibrium capacity for the stated water mass. If the actual initial dissolved mass is at or below that capacity, no solute crystallises. If it is higher, subtract the capacity from the initial dissolved mass to find the crystal mass at equilibrium. This assumes no water is lost and no water enters the crystals. It does not predict how soon crystal growth begins.
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